Chemistry 1A: General Chemistry

Transcription

Laboratory Manual
Prepared by
Las Positas College
Chemistry Faculty and Staff
Past and Present
Fall 2012 Edition
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Las Positas College, Chemistry 1A Lab Manual Fall 2012
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Table of Contents
Safety
Laboratory Safety: Laboratory Rules, Self Protection, and Handling of Chemicals and
Glassware………………………………………………………………………………….6
Laboratory Assignments
Math Review
Writing Formulas and Nomenclature Worksheets
Experiment 1
Significant Figures in Data Collection and Calculation
Experiment 2
Library Assignment
Experiment 3
Composition and Formula of a Hydrate
Experiment 4
Mixing Alcohol and Water—A Thumbsucking Exercise
Experiment 5
Identification of Reaction Products
Worksheet
Stoichiometry Problem Set
Experiment 6
Ions in Solution: Electrolyte Strength and Electrical Conductivity
Experiment 7
Net Ionic Equations and Reactions in Aqueous Solution
Experiment 8
Determination of Copper in a Coin
Experiment 9
Oxidation –Reduction Reactions: Predictions and Equations
Experiment 10
Determination of the Gas Constant, R
Experiment 11
Determination of Sodium Bicarbonate in Alka-Seltzer
Experiment 13
Determination of Heat of Reaction
Experiment 14
Determination of Crystal Violet by Spectrometry
Experiment AA
Measurement of Iron by Atomic Absorption (AA) Spectrometry
Experiment 15
Group Relationships and Periodic Properties
Experiment 16
Model Making and Geometry
Experiment 17
Metallic and Ionic Crystal Lattices
Experiment 18
Evaporation and Intermolecular Attractions
Experiment 19
Using Freezing-Point Depression to Find Molecular Weight
Experiment 20
Acid Rain
Experiment 21
Chemical Equilibrium: Finding a Constant, Kc
Experiment 24
Measuring Sulfur Dioxide in Wine
Reference Material
Lab Report Format
Periodic Table
Errors, Precision and Accuracy
Treatment of Experimental Data
Statistics and Uncertainty in the Laboratory
Names, Formulas and Oxidation Numbers of Some Common Ions
Net Ionic Equations
Solubility Tables (4 versions)
Colors of Ions in Aqueous Solutions
Common Oxidation States of Six Elements Important in Redox Chemistry
Activity Series of Metals and Nonmetals
Acids and Bases
Properties of Water: Density and Vapor Pressure
11
13
17
21
25
33
35
43
45
53
61
73
81
89
97
109
115
123
145
157
171
177
181
189
195
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Safety
Las Positas College
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Laboratory Safety
Read this section on Laboratory Safety. The material contained in this section will be
discussed and augmented by your instructor. All students must also complete “Your
Safety in the Laboratory” on the following pages. Each student must turn in the
completed answer key and the signed “Safety Rules Agreement” before beginning
laboratory work. You should have no trouble if you have done your reading and
participated in the discussion.
A. Laboratory Rules
1. Never work alone! You will work individually, but you should never work in the lab if
the instructor is not present.
2. Report any injury to your instructor at once, no matter how slight it may appear to be.
3. Pregnant students must be especially careful around hazardous chemicals. Pregnant
students must see their instructor for additional safety considerations.
4. Maintain an orderly and clean laboratory desk. Immediately clean up anything you
spill or break. Use a dustpan and brush for broken glass and dispose of broken glass in
the specially labeled container. Keep drawers closed while you are working and keep
stools and backpacks from obstructing the aisles. You should find a safe place to store
your bags and backpacks while you are working.
5. Do not perform any experiments other than those authorized for use that day, unless
you first secure permission from your instructor. Horseplay or gags of any kind are
strictly prohibited.
6. No smoking, drinking, eating or chewing is permitted in the laboratory at any time.
Smoking is only allowed on campus in parking lots or in designated smoking areas, NOT
OUTSIDE LABS! The fume hoods draw smoke directly back into the labs.
7. Do not put equipment in your drawer except for the equipment you were originally
issued. You are encouraged to store your goggles, gloves, and hair bands in your drawer.
8. Do not leave a heat source unattended.
9. Never take a strong whiff of any chemical. If you are instructed to smell a substance,
always waft a small sample of the vapor toward your nose and smell it cautiously.
10. Know the location of and how to use the emergency equipment in your area. Be
familiar with emergency procedures.
11. Avoid distracting or startling any other person. Practical jokes or horseplay cannot be
tolerated.
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12. Most experiments are to be performed individually. Experiments in which partners
are allowed or recommended will be specifically identified by your instructor. In any
case, lab reports must be made and submitted individually.
13. Do not leave the laboratory before the period is over, unless you have completed all
calculations and write-up for the experiment. You may find that you have questions or
need to repeat part of the experiment.
14. When you finish your experiment, clean and return equipment to its proper storage
area. Clean the lab benches with a damp towel, clean the sink of all debris, and discard
all chemically contaminated paper and chemical waste in the designated containers. Ask
your instructor if you are unsure about where something should go.
B. Self Protection
1. Safety goggles, of the type approved for chemistry, must be worn in the laboratory.
These goggles seal to the face and have only indirect air vents to prevent chemicals from
dripping or splashing into your eyes.
2. Do not wear contact lenses in the laboratory.
3. If any substance should get in your eye, flush the eye (or eyes) thoroughly (at least 15
minutes) in the eyewash fountain.
4. It is often advisable to wear a lab coat or plastic apron in the lab to protect your
clothing.
5. Wear only fully enclosed shoes; sandals are not permitted.
6. Tie back long hair so that it doesn’t contact flame or chemical solutions.
7. If chemicals are spilled over a large part of the body, use the safety shower at once.
Remove the contaminated clothing. Flood the chemically burned area with water for 15
minutes. Notify your instructor immediately.
8. Do not taste chemicals or anything in the lab. Wash your hands before leaving the lab
so that you do not accidentally ingest chemicals.
9. If you feel faint or dizzy, sit down on the floor. Don’t walk for help, let your neighbor
do that.
10. Don’t panic, whatever the emergency.
C. Handling Chemicals and Glassware
1. Always wash your hands before leaving the laboratory.
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2. Do not take reagent bottles to your desk or into the balance room. Instead, use them at
the side shelves and bring your own container to take what you need.
3. Do not lay the stopper from a bottle down on the table; learn to hold it between your
fingers. Be careful not to exchange stoppers from one bottle to another.
4. If the outside of a bottle becomes wet with liquid, wipe the bottle with a damp towel.
If the bottle has left a ring of liquid on the countertop, wipe up that contamination, too.
5. Whenever pouring liquids from a bottle, “guide” the liquid down a glass stirring rod
held against the lip of the bottle. Done properly, this will practically eliminate the
dribbling problem.
6. Check the label on the bottle both before and after removing the reagent. Check not
only name and formula, but the concentration.
7. Never pipet by mouth or put anything from the lab in your mouth!
8. Never return any excess chemical to a reagent bottle. Dispose of it as directed in the
experiment or ask your instructor.
9. Never use more of a chemical than is called for in the experiment.
10. Always make dilutions by pouring concentrated solutions slowly into water, not the
reverse. Much heat may be evolved, especially in diluting acids or bases. Remember,
add acid to water, not the other way around.
11. Do not grasp recently heated glassware or iron rings, etc. These objects may still be
hot enough to burn you. If you should burn yourself, notify your instructor immediately
and run cool tap water on the burn.
12. Handle glass thermometers carefully. Thermometers with a metallic liquid contain
mercury. Mercury spills require immediate attention. A technician will clean up the spill
immediately.
13. Make sure glassware and equipment are clean before you start your experiment.
14. Disposal of hazardous chemical wastes will be spelled out in the lab manual and/or
by your instructor. Always ask if you aren’t sure where to dispose of something.
15. Before use, inspect all glassware for damaged edges and cracks. Start heating test
tubes and other glassware slowly. When heating a test tube, make sure you are not
pointing it at anyone.
16. Beware of table edges when setting glassware down. Do not place items near the
edge where they can roll or be knocked off.
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Laboratory Assignments
Las Positas College
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Name_______________________________
Las Positas College
Chemistry 1A
Math Review
Answer the following problems in the space provided. Show units and significant
figures.
23. Express 4521.3 in scientific notation
6
5
1. 1.61 x 10 + 1.9 x 10 =
=
2. 1.61 x 106 - 1.9 x 105 =
24. Express 0.0000456700 in scientific
3. 1.91 x 10-5 + 1.6 x 10-6 =
notation =
4. 1.91 x 10-5 – 1.6 x 10-6 =
25. Express 6.72123 x 104 without
-8
23
5. (2.6 x 10 )(6.02 x 10 ) =
exponents =
-8
-9
6. (2.6 x 10 )/(0.52 x 10 ) =
26. Express 78.7654 x 10-3 without
7. (2.6 x 10-8)(0.25 x 1017)/(4.6 x 10-9)
exponents =
=
27. Find A where 2A + 2 = 25
________________
A=
8. (2.46)(1.98)/(0.82)(273) =
28. Find X where XY = 16 and Y2 = 225
________________
X=
1/2
9. (3.2) =
29. Find Y where Y2/(0.1) = 4.0 x 10-9,
10. (4.0 x 10-6)1/2 =
Y=
11. (4.0 x 107)1/2 =
30. (6.4 x 10-14)1/3 =
12. (4.0 x 10-7)1/2 =
31. Find Z where Z2 + 3Z – 10 = 0, Z =
-7 2
13. (4.0 x 10 ) =
32. Find Y where Y2 + 2Y = -0.3 Y =
-4 3
14. (3.0 x 10 ) =
33. (0.070)(0.6023 x 1023)/(22)=
3 4
15. (3.0 x 10 ) =
34. 77.777 – 44 =
16. (12.35)-2 =
35. How many significant figures
17. (1.5)-3/(2.5)-2 =
in 6.040 x 108?
3
-2
18. (2.5) (1.5) =
36. (-2.2 x 104)2 =
3
19. (5.0 cm) /(2.0 cm) =
37. (-1.3 x 103)3=
2
3
4
20. (4.0 cm) / (3.0 cm) (2.0 cm)
38. (-2.0 x 10-2)4=
21. (4.0 cm2)2 =
39. (4.5)(-2.0) =
22. (7.5 cm)/(1.5 cm-2)-3 =
40. –(-2.5)(-3.0) =
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Writing Formulas Worksheet
Fill in the blanks with the correct formulas.
Anions 
Chloride
Sulfate
(Cl-)
Cations ↓
Ammonium
(NH4Cl)
(NH4+)
Silver
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Name: _________________________
Carbonate
Nitrate
Phosphate
Sulfide
Copper (II)
Calcium
Potassium
Mercury (II)
Anions 
Cations ↓
Sodium
Magnesium
Iron (II)
Copper (II)
Zinc
Aluminum
Hydroxide
Oxide
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Nomenclature Worksheet
Name __________________________________
I. Name the following binary compounds, all of which are composed of non-metals:
NO2___________________________ NO___________________________
CO2 ___________________________ CO___________________________
CS2___________________________ CBr4___________________________
PCl3___________________________ PCl5___________________________
N2O3___________________________ H2S___________________________
II. Name first by the –ous/-ic system and secondly by the IUPAC System
Compound
CuCl
-ous/-ic name
IUPAC System name
Cuprous chloride
Copper(I) chloride
Hg2O
SnF4
Hg(NO3)2
Fe2O3
III. Name each of the following compounds.
KBr
Be(NO3)2
(NH4)2S
Ag3PO4
Ca(HCO3)2
Li2CO3
BaCrO4
KH2PO4
CdSO4
NaH
MgSO3
Mg(HSO4)2
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IV. Write formulas for the following:
Zinc hypochlorite
Silver oxide
Sodium
permanganate
Sulfurous acid
Carbonic acid
Zinc phosphate
Ammonium iodide
Stannous fluoride
Hydroiodic acid
Copper(I) oxide
Mercuric oxide
Calcium carbonate
Iron (III) sulfate
Zinc oxide
Nickel (II)
carbonate
Ammonium sulfate
V. Adjacent to the formula, write the name of the salt, then give the formula and the
name of the acid from which each salt may be derived.
FeSO4
Iron (II) sulfate
H2SO4
Sulfuric acid
Fe2(SO4)3
Mg(NO3)2
Na2HPO4
KHSO4
NiCl2
SnS2
CaF2
VI. What is the formula (not the symbol) for each of the following elementary substances
when it occurs free in nature?
Hydrogen
Chlorine
H2
Nitrogen
Oxygen
Fluorine
Bromine
Iodine
Helium(!)
VII. Name the following:
H2O ___________________ NH3 ____________________ CH4 _________________
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Experiment 1
Significant Figures in Data Collection and Calculation
PURPOSE
The objective is to show how the number of significant figures is limited by the precision of
the measuring instruments. Here the measurement will be volume, and the measuring
instruments will be beakers, graduated cylinders, pipets and burets.
Method
Various measurements are to be made and reported to the precision allowed by the
measuring instrument. In data collection, always include all certain digits and the first
uncertain (estimated) digit beyond these. This set of digits constitutes the significant figures
appropriate to that measurement.
How the number of significant figures depends upon the measuring instrument is illustrated
below. The same linear dimension is measured with rulers of different degree of precision.
Enter your estimate of the dimension in both situations; you should have two significant
digits for Figure A and three for Figure B.
(A) ____________________ cm*
(B) ____________________ cm*
*Note: enlarged view, not the true size of
a centimeter.
If the instrument (such as an electronic balance) has a digital readout, the last digit may
fluctuate. Take what seems to be a middle value.
PROCEDURE AND DATA SECTION: Write up the procedure and data section below
in your notebook before coming to lab.
Throughout the semester, record all measurements in INK, including proper units,
immediately in your laboratory notebook! The spaces given in this manual are an
indication that you should write something in your notebook, not a proper place to
record data!
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Please read other the reference material for information on errors, uncertainty, statistics and
calculations in the laboratory.
1. BEAKERS AND ERLENMEYER FLASKS
Take a 600 mL beaker and a 500 mL Erlenmeyer flask from your locker. Note the
graduations stenciled on the side of each. Note also that the precision of calibration claimed
is only ±5%. If you were to fill the beaker to the 500 mL mark, how should you express this
volume to show the absolute uncertainty? Fill in the blank below.
Volume = 500 mL ± ______ mL
Now, actually fill the beaker as carefully as you can to the 500 mL mark with tap water.
Then, carefully transfer the water from the beaker to the Erlenmeyer flask.
Questions (answer in your notebook):
1. How well does the new volume reading agree with the old?
2. Although both pieces of glassware have the same estimate of accuracy, why might you
find the Erlenmeyer flask more reliable?
2. GRADUATED CYLINDERS
Visit Station 2 in the lab, where two graduated cylinders partially filled with water are on
display. As precisely as you can, read the volume level in each. Enter your readings in your
notebook in the format shown below. Note that a ± uncertainty is asked for also.
10 mL graduated cylinder ___________ ± ______
100 mL graduated cylinder ___________ ± ______
3. BURETS
Burets are carefully calibrated to be very precise, usually on the order of ±0.01 or ±0.02 mL.
Visit Station 3 in the lab, where two burets are set up as if involved in a titration. The buret
on the left represents the initial reading, and the one on the right represents the final reading
(after you have drained some liquid out of the buret). Record each buret reading to the
proper number of significant figures and uncertainty.
Final buret reading* ___________ ± ______
Initial buret reading* ___________ ± ______
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Now calculate the volume of liquid delivered. Report your answer to the correct number of
significant figures and include the total uncertainty. (The total uncertainty will be the sum
of the absolute uncertainties.)
___________ ± ______
Question 3: Why do we say "reading" here, and not "volume"?
4. PIPETS
NEVER use your mouth for suction. An essential part of the exercise is to learn to use a
rubber bulb or other device for filling the pipet.
You will be asked to weigh the water (to ±0.001 g) after transfer to a small flask, and to
repeat this operation two times. The reproducibility (precision) of your work will be a good
indicator of the quality of your technique. One measure of precision, as used here, is to
calculate the average deviation of the data. First, find the average mass of water, then the
deviation of each trial from the average, and finally the average of the absolute values of
these deviations. (Absolute values must be used. Otherwise the sum of the deviations will
equal zero.) Even more significant than average deviation is relative average deviation. To
calculate this, divide the average deviation by the average value of the three trials, and
multiply by 100 to get relative average deviation in parts per hundred (pph or %).
Description of pipet:
Size: ______ Accuracy (from label or manufacturer's catalog): ______
(Also note representative values in other handouts.)
Volume (expressed to proper number of significant figures): _________
Description of water used (This must be deionized water kept at room temperature for some
time.)
Temperature of the water: _________
Density of the water at this temperature : ____________
(from the CRC Handbook of Chemistry and Physics)
Trial
#1
#2
#3
mass of water + flask
__________ __________ __________
mass of flask empty
__________ __________ __________
mass of water
__________ __________ __________
Average mass of water
__________
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Deviation from average
__________ __________ __________
Average deviation
__________(Show calc.)
Relative average deviation
( = average deviation/average mass)
__________
Calculated volume of average mass of water
(Show calculation in notebook.)
__________
Percent error for the volume of water: (Show calculations in notebook.)
experimental value (label) - true value (calculated)
x 100% =
true value (calculated)
5. BAROMETER
Anticipating its relevance to future experiments, we will ask you to practice reading the
barometer and then applying corrections to that reading.
The major correction is that for temperature. As the temperature rises, mercury expands
inside the glass tube and stands higher even though the atmospheric pressure has not
changed. The other correction is for the fact that the gravitational pull of the earth varies
with latitude.
Record the "uncorrected" reading of the barometer:
_______________
Record the temperature of the barometer:
_______________
Record the temperature correction:
_______________
Record the gravitational (latitude) correction:
_______________
Calculate the corrected barometric pressure:
_______________
Next: Find the atmospheric pressure from the Official LPC Weather site at
http://www.aws.com/AWS/wx.asp?id=LPCAL
Pressure:
________________ inches of Hg.
Convert to mmHg:
_________________mmHg
Experiment 2
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Name ______________________
Library Assignment
**You can do this assignment here in the manual, instead of your lab notebook.
The objective is to give you experience in using our Learning Resource Center (the Library)
to look up scientific and especially chemical information. The three main areas will be (1)
browsing through the stacks of Non-fiction, (2) using the Reference section, with its various
chemical handbooks; and (3) scanning the various scientific magazines to find some
chemical articles that interest you.
Part I Browse through the open stacks of the nonfiction section of the Library. Find one
book in each of the following Library of Congress classifications that both interests you and
would seem useful in your study of chemistry. Note that two letters of the classification
have been given (except for the one which is just Q). Give the rest of the classification
(numbers and letters) and complete the rest of the table.
First two Rest of the
call letters Classification
Q
Author(s)
Title
QC
QD
QH
QP
RS
TA
TD
As much as you can generalize, what subject matter is treated by books whose call letters
start with:
Q
_________________
QC
___________________
QD _________________
QH
___________________
QP
_________________
RS
___________________
TA __________________
TD
___________________
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Part II Reference Section
Before visiting the Reference Section of the Library, use the library catalog to find the call
letters and numbers for the following works and enter them below.
CRC Handbook of Chemistry and Physics
(___________________)
Merck Index.
(___________________)
McGraw-Hill Encyclopedia of Science & Technology (___________________)
Now choose one of the chemical elements to look up in the Encyclopedia; choose an
element that starts with the same letter as your last name. [If your name begins with J,
choose iodine ("Jod" in German); if Q, choose mercury ("quicksilver"); if W, choose
tungsten ("Wolfram" in German)].
Element _______________
volume & page from Encyclopedia ____________
1) How is this element obtained in the free state?
2) List some commercial applications of this element.
3) Next refer to the Handbook of Chemistry & Physics. Use the index to find the Table of
Isotopes. Find the element you have chosen. How many different isotopes are listed in the
table for this element? Fill in the three blanks below.
Handbook: (______) Edition;
page reference (______);
Number of isotopes (____)
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4) Before this next assignment in the Reference Section, find the chemical name of an
ingredient listed on some prescription or over-the-counter drug found in the medicine
cabinet at home. Then look it up in the Merck Index. (Merck is one of the principal
pharmaceutical companies in the world.) Look up the side effects associated with this drug.
Fill in the four blanks below:
Name(s) of drug: _______________________
Merck Index: (_____) Edition; page (_____);
side effects:
4b. Now, search for the same medicine on the Web.
What are the uses of this medicine?
URL (web address): _________________________________
Part III Browse through the Periodicals Section of the Library. List below the names of
five magazines (or "journals", the more professional-sounding term) that regularly seem to
contain articles on chemistry.
1)
3)
5)
2) ___________________________________
4) ___________________________________
Now, from one of these magazines, choose an article on chemistry that interests you. In the
space below write a paragraph telling a little about the content of this article and why it
interests you. Preface the paragraph with title of the article, name of author, name of the
magazine, date of issue and page.
_______________________________________________
_______________________________________________
_______________________________________________
_______________________________________________
_______________________________________________
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Experiment 3
Prelab – Complete the Prelab on page 31 before coming to lab.
PURPOSE
To determine the percent, by mass, of water in a hydrate compound and to determine the
number of moles of water per mole of hydrate compound – the empirical formula of the
hydrate compound.
Introduction
What is a hydrate? A hydrate is a solid ionic compound that contains water molecules within
its crystalline lattice structure. Frequently the water is chemically bonded less tightly than
other bonds within the compound. Subjecting a sample of hydrate to the heat of a laboratory
Bunsen burner for 5-10 minutes will often drive off all of the water and leave an anhydrous
salt compound residue.
CuSO45H2O(s) + heat  CuSO4(s) + 5 H2O(g)
(hydrate)
(anhydride)
From the masses of the hydrate, anhydride, eliminated water, and the formula of the
anhydrous salt the composition and the empirical formula of the hydrate can be
determined. Proper technique and careful heating and cooling of the sample should give
precise and accurate results.
PROCEDURE
A. Verify the formula of a known hydrate and perfect your technique with BaCl22H2O.
1. Thoroughly clean a crucible and crucible lid. Place them in a clay triangle suspended
on a ring on a ringstand. Heat them for about 3-5 minutes. Using crucible tongs,
remove the crucible and the lid and let them cool to approximately room temperature.
Carefully weigh the crucible and lid on the analytical balance. Don’t touch the
crucible, even when it’s cool, because fingerprints can affect your results! Use your
time efficiently by cleaning and heating the second sample while the first one is
cooling.
2. When the crucible mass is obtained, add 1-2 grams of solid barium chloride hydrate
to crucible and weigh it – all weights should measured to the nearest milligram. Then
heat the crucible and contents strongly for 4-8 minutes. The crucible should be
cherry red during the last three minutes of heating. The lid of the crucible should be
in place during the heating. Cool and weigh.
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3. Repeat with a second sample.
4. Calculate the percent water in the hydrate compound.
5. If your two runs do not agree within 2% relative range, consult with your instructor
before going to part B. If time allows, you probably need to do a third trial.
B. Unknown Hydrate
1. Clean your crucibles and lids after Part A is completed and heat them carefully as
before. Add 0.5-0.6 g of unknown hydrate to your crucible and heat as in Part A. Be
careful to avoid splattering. Cool and weigh.
2. Repeat with a second sample – a third if time permits.
3. Calculate the percent by mass of water in the unknown hydrate.
4. When you have completed the calculations in part 3 above, see your instructor for
the formula of the anhydrous compound and then compute the formula for the hydrate
(i.e. how many moles of water are there per mole of the anhydrous compound).
POST-LAB QUESTIONS: Include the answers to the following questions with your
laboratory report.
1. How can an analysis have good precision and poor accuracy?
2. List possible errors that will cause your results to be too low or too high. Indicate if
each error will have a large effect on the results or a minor effect.
3. In a similar analysis, a student determined that the percent of water in the hydrate was
25.3%. The instructor informed the student that the formula of the anhydrous
compound was CuSO4. Calculate the formula of the hydrated compound.
DATA
Enter the sample data report form below in your notebook before coming to lab. During
the experiment, record all data (with proper units and number of significant digits)
directly into your notebook in ink. Show sample calculations where an asterisk, *,
appears at the beginning of the entry description.
Part A.
Run 1
Run 2
1. mass clean empty crucible and lid
______
______
2. mass crucible, lid, and BaCl2 . 2H2O
______
______
3. mass BaCl2 . 2H2O
(hydrate only)
Page 27
______
______
4. mass crucible, lid & anhydride after 1st heating
______
______
5. mass crucible, lid & anhydride after 2nd heating
______
______
6. mass crucible, lid & anhydride after 3rd heating
(if necessary)
7. mass crucible, lid & anhydride after 4th heating
(if necessary)
______
______
______
______
8. mass BaCl2
______
______
9. mass water, H2O, driven off
______
______
10. * % water by mass in hydrate, BaCl2 . 2H2O
______
______
(anhydrous salt only)
11. * average % water by mass in hydrate, BaCl2 . 2H2O
______
12. * % precision or % relative range of % water by mass
in hydrate, BaCl2 . 2H2O
(see error references)
.
[(high-low)/average] 100%
______
13. * theoretical % water by mass in hydrate, BaCl2 . 2H2O
______
14. * percent accuracy or relative error of % water by mass
in hydrate, BaCl2 . 2H2O
[(expt'l ave. - theo.)/theo.] . 100%
______
Page 28
Part B.
Unknown Number ______ (required)
Run 1
Run 2
1. mass clean empty crucible and lid
______
______
2. mass crucible, lid, and unknown
______
______
3. mass unknown
______
______
4. mass crucible, lid & anhydride after 1st heating
______
______
5. mass crucible, lid & anhydride after 2nd heating
______
______
6. mass crucible, lid & anhydride after 3rd heating
(if necessary)
7. mass crucible, lid & anhydride after 4th heating
(if necessary)
8. mass unknown anhydride only
______
______
______
______
______
______
9. mass water, H2O, driven off
______
______
(hydrate only)
10. * % water by mass in unknown hydrate
______
11. * average % water by mass in unknown hydrate
______
12. * % precision or % relative range of % water by mass
in unknown hydrate
.
[(high-low)/average] 100%
______
______
Page 29
To calculate formula of unknown hydrate:
13. formula of your unknown anhydride (given by instructor) ___________________
14. * number of moles of water driven off
from unknown hydrate
______
______
15. * number of moles of anhydride after heating
______
______
16. * ratio of moles water to moles anhydride
______
______
17. * average ratio of moles water to moles anhydride
18. * formula of unknown hydrate ______________________________
_____
Page 30
Page 31
PRE-LAB
Name_________________
(To be completed before coming to lab)
Experiment 3
1. If you start with 0.572 grams of unknown hydrate and end with 0.498 grams of
anhydrous compound, what is the percentage of water by mass in the sample?
2. How could you determine if you have heated the sample long enough to drive off all of
the water?
.
3. If you begin with 1.534 grams of BaCl2 2H2O and heat it for 4 to 8 minutes, what will
be the mass of BaCl2(s) that remains?
Page 32
Page 33
Experiment 4
Mixing Alcohol and Water--a Thumbsucking Exercise
Despite its frivolous-sounding title, this exercise will provide you with an opportunity to
speculate about what must be happening on the molecular level when these two simple
substances are mixed.
PURPOSE
Very little introduction will be given to this exercise. You are to perform two rather simple
operations, make close observations of the phenomena observed, and then speculate about
an explanation.
PROCEDURE
SAFETY FIRST
The hazard level is low in this exercise. However, there is a risk of splashing alcohol in the
eyes, and, as is usual, goggles are required. Also, the use of your thumb to close off the top
of a test tube, as called for here, is not allowed in any future work.
1. Remove two small (10-cm) test tubes from your locker along with a small beaker to
hold them upright. Locate two plastic wash bottles: one, containing deionized water,
will be unlabeled; and the other, containing ethyl alcohol, will be labeled with that name,
along with the warning "Flammable!” There won’t be a flame in today's assignment, but
you should be aware that alcohol, like many non-aqueous solvents, will burn.
2. Fill one test tube about half full with water from the unlabeled wash bottle.
3. Now fill it to the top with alcohol from the labeled wash bottle. Do this carefully. Try
not to mix the two liquids unnecessarily. (They will dissolve in each other, given the
chance, but it is possible to layer the alcohol above the water.)
4. When the tube is "brimful", slide your thumb over the top of the tube carefully so that no
bubbles of air are trapped. (Again, we will not use our thumbs this way in any further
chemical work. It is not good for the thumb or the chemical study. Here, there is no
hazard, and direct contact is necessary to the study.)
5. Be sure your thumb is squarely and firmly positioned on top. Now invert the tube once
or twice and return to the original position, keeping your thumb on top. Observe any
changes in appearance inside the tube and any sensations in your thumb. (Some people
find that they can suspend the tube from their thumb without supporting it! But keep
your free hand underneath, just in case!)
Page 34
6. Discard the solution into the proper waste container.
7. Try the experiment again, but now this time add the alcohol first, and then fill with
water. Do you get the same effect? If not, why not?
POSTLAB QUESTIONS: Answer the following questions in your notebook.
1. In your lab write-up, describe what you saw and experienced in the two different
exercises. Please use complete sentences.
2. What is your explanation for what happened in A, and for the difference in behavior
between A and B? You are not expected to know the whole story. Any hypothesis,
however "wild" is invited.
Page 35
Experiment 5
Prelab – Complete the Prelab on page 41 before lab.
PURPOSE
The purpose of this experiment is to recognize evidence of chemical change, and to write
proper equations, both complete and net ionic, for reactions observed to occur.
Discussion (Read the discussion below but you do not need to enter this section into your
notebook.)
For reactions in this experiment, we will limit our study to the category variously known
as "metathesis," "double displacement," or "partner exchange." Partner exchange can be
illustrated by the generalized equation
AB
+
CD

AD
+
CB
It can be assumed that the reaction is conducted in water solution, and that, in this
experiment, reactants AB and CD are not only soluble but are dissociated into their
ions.
Question: Why are the products not AC and BD? And why are they not DA and BC?
One can write on paper many chemical equations of this form, with proper
formulas and properly balanced equations, but whether or not they really represent
true reactions depends on experimental investigation.
1) One of the products (AD or CB) is not soluble; that is, it precipitates. (See solubility
rules in the reference section.)
2) One of the products (AD or CB) is a weak acid or base formed from a strong acid
or base (AB or CD). Examples are the formation of the weak acid acetic acid,
HC2H3O2, from the strong acid HCl, or the formation of the weak base NH3 from
the strong base NaOH. (See listing of common weak and strong acids in the
reference section.)
3) One of the products (AD or CB) is a water molecule formed from the
reaction of an acid and base (AB and CD). The reaction, commonly called
"neutralization," has the form "acid + base --> salt + water."
You will be able to detect these chemical changes by careful observation.
1) Precipitates usually form first as cloudy suspensions which then settle out
as visible particles. Some are white; others are colored.
Page 36
2) Strong acid  weak acid reactions often show very little sign of chemical change.
In the reactions performed here, the weak acid is unstable and decomposes into a
gaseous component which escapes. The two most common examples are the
formation of CO2 (from H2CO3) and SO2 (from H2SO3). Strong base --> weak base
reactions are similarly hard to detect. However, the most common weak base, NH3,
has a characteristic odor.
3) For neutralization reactions heat evolution may be the only obvious evidence of
reaction. This criterion for reaction is tricky. For instance when concentrated
sulfuric acid is added to water, much heat is evolved but the chemical change
involves just the ionization of the acid, actually:
-
H2SO4(l) + H2O(l)  H3O+(aq) + HSO4 (aq)
PROCEDURE
SAFETY FIRST
Wear safety goggles throughout the experiment. Exercise care in handling acids and
bases. Note the caution below on handling AgNO3. Compounds of heavy metals (Ni,
Cu, Ag, etc.) are poisonous if ingested. Wash hands frequently.
In Table I, some ten or so pairings of possible reactants are listed for you to investigate.
Work independently! Aqueous solutions of these compounds in the recommended
concentrations have been prepared for you. Volumes suggested (often just 2 mL) can be
estimated rather than using the graduated cylinder each time. (You can use your
graduated cylinder to measure 2 mL of water into a test tube and keep that tube for
comparison.)
Use your large size test tubes; they will allow you more freedom in mixing. Observe
and record original colors. Add the second reactant to the first gradually, swirling the
tube between additions.
Do not be too quick to report "no reaction". Some precipitates are slow in forming.
Rubbing the inside wall of the test tube with a glass rod sometimes initiates
precipitation in the solution that wets it.
Some precipitates are so finely divided they stay suspended and look like an emulsion-they look "milky". With time, however, they usually coagulate and settle out. To speed
this process, we often use a centrifuge.
To use a centrifuge:



always insert pairs of tubes (either two or four of your small test tubes, use only
small test tubes);
be sure each tube contains approximately the same volume of liquid; fill one with
plain water if necessary;
insert each pair of tubes in the tube holders at 180° (i.e. opposite) to each other (to
keep the centrifuge balanced);

Page 37
keep the lid down until the centrifuge has completely stopped; usually one minute of
centrifuging at full speed is sufficient.
One reason for using the centrifuge is that a clean separation of precipitate is often
necessary before the true color of the precipitate can be distinguished from that of the
supernatant solution. If the color of the precipitate is still not evident, pour off the
solution, add about 1 mL of deionized water, stir the mixture well, and centrifuge it
again.
After the reaction has been studied, dispose of the contents of the test tube into the
proper waste container. Rinse tubes but it is not necessary to dry them before going to
the next test pair.
T a b l e I : Partner Exchange: Mixing Directions
Set
Reactants
Mixing Directions
1
FeC13 and NaOH
0.1 M FeC13 (2 mL)
6 M NaOH, added dropwise; 10 drops
2
FeC13 and NH3
0.1 M FeC13 (2 mL)
6 M NH3, added dropwise; 10 drops
3
FeC13 and H2SO4
0.1 M FeC13 (2 mL)
3 M H2SO4, added dropwise; 10 drops
4
NiCl2 and Na2CO3
0.1 M NiC12 (2 mL)
0.3 M Na2CO3 (2 mL); centrifuge
5
NiC12 and AgNO3
0.1 M NiC12 (2 mL)
0.1 M AgNO3 (2 mL)
AgNO3 stains the skin; wash hands after contact.
6
Na2CO3 and HC1
1.0 M Na2CO3 (2 mL)
3.0 M HC1, added dropwise; 10 drops
7
NH4NO3 and NaOH
0.2 M NH4NO3 (2 mL)
3 M NaOH (1 mL)
Heat mixture by placing test tube in small beaker of
hot water; hold moist red litmus in vapors; waft
vapors to nose.
8
CuSO4 and Na3PO4
0.1 M CuSO4 (2 mL)
0.1 M Na3PO4 (2 mL); centrifuge
Page 38
9
ZnSO4 and(NH4)2S
0.1 M ZnSO4 (2 mL)
0.1 M (NH4)2S (2 mL)
Mix in HOOD!
10
NiCl2 and(NH4)2S
0.1 M NiCl2 (2 mL)
0.1 M (NH4)2S (2 mL)
Mix in HOOD!
11
H2SO4 and NaOH
3 M H2SO4 (2 mL)
6 M NaOH (2 mL)
Mix all at once!
DATA
Enter a data table for all 11 sets of reagents in the style of the following Sample Data
Table in your notebook before class, but leave enough space for observations. Remember
that the reader might not have performed this experiment, so write sufficient detail to be
informative. Use the format below to record data. For each set of reagents, you should
have five lines blank; A, B, C, D, & E, intended to receive the following information.
Report your observations in ink directly into your lab notebook as you go along.
#1
FeC13 and NaOH
A:
B:
C:
D:
E:
____________________________________________________________
A: Give the complete or "molecular" equation for the predicted partner
exchange reaction.
B: Describe original colors. Describe changes (if any) upon mixing the proposed
reactants. Report any color change; separation of a solid (and its color);
evolution of a gas (and its odor--but waft only a very little to your nose!); and
evolution of heat.
C: If there was a reaction, what is the likely identity of the new substances
formed?
D: Write the “Total Ionic Equation,” that is write each aqueous compound from
part A as separate ions (labeled with the correct charges). Don’t remove
spectator ions yet!
Page 39
E: Write the net ionic equation for the reaction. If there was no reaction, say so,
"no rxn". (Hint: at least one is a "No Reaction" pair.)
Example of Entry into Report
Set:
BaCl2 and CuSO4

A:
BaCl2(aq) + CuSO4(aq)
BaSO4(aq) + CuCl2(aq)
B:
CuSO4 sol'n blue, the other colorless. When mixed, white
precipitate gradually settled out from a blue solution.
C:
There was a precipitation reaction. The copper ion remained in
solution. The precipitate was probably BaSO4.
D:
Ba2+(aq) + 2Cl-(aq) + Cu2+(aq) + SO42-(aq)  BaSO4(s) + 2Cl-(aq) + Cu2+(aq)
E:
Ba2+(aq) + SO42-(aq)  BaSO4(s)
Page 40
Page 41
PRE-LAB
Name_________________
Experiment 5
1. Find all the pairs of ions in the following list which cannot exist together (in
appreciable concentration) in the same aqueous solution without precipitating. That is,
write any combination of the following ions which would form a precipitate, a weak acid,
a weak base, or water. [Note: PbCl2 is slightly soluble.]
Fe 3+
Answer:
H+
OH -
Na +
Cl
-
Pb 2+
NO 3
-
Mg2+
__________
2. Convert the following (a) complete or "molecular" equation to (b) the total ionic, and
then to (c) the net ionic equation.
(a)
Na 2 C rO 4 ( a q )
+
2AgN O 3 ( a q )
 2 NaNO3(aq)
Ag2CrO4(s)
(b)____________ __ _____________________________________
(c)______________ ______________________________________
3. Write all three balanced equations (a, b, and c, as above), given this word equation
for a chemical reaction:
barium hydroxide + sulfuric acid  barium sulfate + water
solution
solution
solid
(a)
__________________________________________________________
(b)
__________________________________________________________
(c)
___________________________________________________________
Page 42
Stoichiometry Problem Set
1.
Page 43
Chem 1A (LPC)
a.) How many moles of H3PO4 are represented by 294 grams of H3PO4?
b.) How many atoms of hydrogen are in 196 grams of H3PO4?
c.) What is the % composition of each element in phosphoric acid?
d.) How many grams of phosphorus are in 294.1 grams of H3PO4?
2. a.) Lactic acid is a metabolite formed in the body during muscular activity. It is
composed of 40.00% carbon, 6.71% hydrogen and 53.9% oxygen by weight. What is the
empirical formula of lactic acid?
b.) The molecular weight of lactic acid is 90.08 g/mole. What is the molecular formula?
3. Some commercial baking powders contain a mixture of sodium bicarbonate (baking
soda) and calcium dihydrogen phosphate. When the powder is moistened, carbon dioxide
gas is liberated and makes the batter or dough rise. Balance the equation for the
following reaction.
NaHCO3 + Ca(H2PO4)2
Na2HPO4 + CaHPO4 + CO2 + H2O
4. Chlorine for use in water purification systems may be obtained from the electrolytic
decomposition of seawater. The unbalanced chemical equation for this reaction is:
NaCl(aq) +
H2O(l)
NaOH(aq) +
H2(g) +
Cl2(g)
a.) Balance the equation.
b.) What weight of sodium chloride would be consumed in the production of 25 metric
tons of chlorine (1 metric ton = 1000 kg).
5. Iron oxides found in iron ores can be reduced to metallic iron when reacted with
carbon monoxide. The equation for this reaction is:
Fe2O3 + 3CO
2Fe + 3CO2
a.) How many kilograms of elemental iron can be formed if 16.0 kg Fe2O3 is reacted
with 10.0 kg CO?
b.) How many kg of CO2 will be produced in the reaction described in 5a?
c.) 16.0 kg Fe2O3 + 10.0 kg CO = 26.0 kg of reactants
11.2 kg Fe + 13.2 kg CO2 = 24.4 kg products
Explain this apparent inconsistency in the conservation of mass.
6. Xylocaine, a local anaesthetic which has largely replaced novocaine in dentistry, is a
compound of carbon, hydrogen, nitrogen and oxygen. Combustion of a 0.4817 g sample
of xylocaine yielded 1.2665 g of CO2 and 0.4073 g of H2O. A separate nitrogen assay,
using another 0.4817 g sample of xylocaine formed 0.07006 g NH3. What is the
empirical formula of xylocaine?
Page 44
Page 45
Experiment 6
Ions in Solution: Electrolyte Strength and Electrical Conductivity
Prelab – Please complete the pre-lab on page 51 before lab.
PURPOSE
 to relate electrical conductivity to the presence of ions in aqueous solution
 to classify substances as strong, weak, and non-electrolytes
 to relate electrical conductivity and chemical reactivity to the presence of
hydronium ions in solutions of acids and hydroxide ions in solutions of bases
 to classify acids and bases as weak or strong
notebook.)
This experiment addresses the question “Do ions form when a given substance is placed
in solution?” The test we apply is the electrical conductivity of the solution. Whereas
electrical conductivity in metals is based on the movement of electrons, that in solution is
based on the movement of + and – ions. On the basis of the observed conductivity of an
aqueous solution of a substance, we will classify that substance as one of the following:
1. strong electrolytes: substances whose aqueous solutions conduct electricity very
well because many ions are formed
2. weak electrolytes: substances whose aqueous solutions conduct only slightly
because only a few ions are formed
3. non-electrolytes: substances whose conductivity in solution is less than or equal to
that of water because they form hardly any ions or no ions at all
In this experiment you will test the electrical conductivity of aqueous solutions of a
variety of substances that differ in their bond types. These will vary from ionic
substances (salts with a difference in electronegativity, ∆EN, greater than about 1.8 to
2.0, to polar covalent molecules with ∆EN between about 0.4 and 1.8. You will not test
any non-polar covalent molecules (those with ∆EN of 0.4 or less) because they are
generally not soluble in water! Your observations will correlate well with the bond type
of the substances. A few generalizations will help.
1.
Solid salts do not conduct electricity when not in aqueous solution. Even ionic
compounds such as NaCl, which exist in the solid state primarily as distinct Na+
and Cl– ions, do not conduct electricity. Because of the solid lattice, ions are
prevented from moving and thus cannot pass charge through the solid.
2.
Molten salts do conduct electricity quite well. Melting the solid frees the ions to
move around in the liquid and thus to move charge through the liquid.
3.
Salts of high solubility show high conductivity in aqueous solution and thus are
strong electrolytes. Dissolving a salt in water separates the ions, primarily due to
Page 46
the ion/dipole attractions between the ions and the water molecules. The hydrated
ions are free to move through the solution and so conduct electricity quite well.
An example is sodium chloride:
+
NaCl(s)
4.
Na (aq) + Cl
–
Salts of low solubility do not dissolve sufficiently to free many of their ions into
solution. Because so few ions go into solution, no increase in conductivity is seen.
An example is calcium carbonate:
2+
CaCO3(s)
2–
Ca (aq) + CO 3 (aq)
A substance must dissolve in water before it can be classified as a strong, weak,
or non-electrolyte. When asked to classify such a substance that does not dissolve
in water, reply “not soluble”.
5.
Many polar molecules show very low electrical conductivity in aqueous solution
and are non-electrolytes. (The major exceptions are acids and bases, discussed
below.) Water itself ionizes slightly (a process called autoionization).
H2 O(l) + H 2 O(l)
+
H3O + OH
–
However, the concentration of ions is very low, about 10–7 M in pure water at
25°C, and so the conductivity is very low. Most other polar molecules show even
less tendency to form ions. Even some metal/nonmetal compounds with low ∆EN
values show very little ion formation. Mercury(II) chloride, HgCl2, ∆EN = 1.2, is
a water soluble solid, but it is a non-electrolyte. HgCl2 dissociates to form fewer
ions than does water!
HgCl 2(s)
6.
+
HgCl2 (aq)
–
HgCl (aq) + Cl (aq)
The conductivity of acids of general formula HX is due to their reaction with
water to form H3O+ and X– ions. Strong acids, such as HCl, react extensively.
Weak acids, such as acetic acid, HC2H3O2, react only slightly.
+
–
HCl(aq) + H 2 O(l)
H3 O (aq) + Cl (aq)
HC2 H3 O2 (aq) + H 2 O(l)
H3 O (aq) + C 2 H3 O2 (aq)
+
–
Strong acids are strong electrolytes and form almost exclusively H3O+ and X– ions
in solution. Weak acids are weak electrolytes. They exist primarily as
undissociated HX(aq) molecules in solution, with only minor amounts of H3O+
and X– ions.
The common strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO4, and HClO3.
All other common acids are weak. Only the first hydrogen of sulfuric acid is
strong, essentially 100% ionized. The second is weak, only about 20% ionized at
Page 47
typical concentrations. Note the positions of the equilibrium arrows in the
ionizations of the two H’s of sulfuric acid:
+
H2 SO 4 (aq) + H 2 O(l)
–
+
HSO 4 (aq) + H 2 O(l)
7.
–
H 3 O (aq) + HSO 4 (aq)
2–
H 3 O (aq) + SO 4 (aq)
The conductivity of bases is due to the presence of OH– ions. The strong bases are
the soluble hydroxides such as NaOH. For these ionic substances, the ions are
already present and are freed into solution when the substance dissolves.
+
–
NaOH(s)
Na + OH
Weak bases such as ammonia, NH3, react only slightly with water. (Most neutral
molecules that are weak bases are amines, organic derivatives of ammonia. Their
general formulas can be written NR3, where R represents either a hydrogen and an
organic carbon group.)
NH3 (aq) + H 2 O(l)
+
–
NH4 (aq) + OH (aq)
Weak acids and bases are weak electrolytes and form only a few ions in aqueous
solution. Strong acids and bases are strong electrolytes and form lots of ions in
aqueous solution. This ion formation will be detected in this experiment by
observing the conductivity of various solutions.
Chemicals
NaCl(s)
1 M NaCl(aq)
CaCO3(s)
CH3CH2OH(l) and (aq, 3 M)
C12H22O11 (aq, 3 M)
1.0 M HCl(aq)
1.0 M NaOH(aq)
1.0 M NH3(aq)
15 M NH3(aq)
HC2H3O2(aq, 1 M)
HC2H3O2(l) [glacial, 17 M]
saturated HCl in toluene
Equipment
conductivity apparatus
50-mL beakers for each substance tested
instructor demo:
crucible
clay triangle
KC2H3O2(s)
Page 48
PROCEDURE
Wear safety goggles. Substance 14 should be worked with under a fume hood. Several
workstations will be set up around the lab. At each station you will measure the
conductivity of one or two substances or solutions. Classify each observation using the
level of conductivity (see instructions on the next page on how to use the conductivity
meter). The instructor will begin by demonstrating some of the conductivities. The
following conductivity tests should be performed:
1.
2.
3.
4.
5.
6.
7.
8.
9.
deionized water
tap water
NaCl(s)
NaCl(aq)
CaCO3(s) + H2O(l)
KC2H3O2(s) (demonstrated)
KC2H3O2(l) (demonstrated)
KC2H3O2 (aq) (demonstrated)
CH3CH2OH(l)
10.
11.
12.
13.
14.
15.
16.
17.
18.
CH3CH2OH(aq)
C12H22O11(aq)
HC2H3O2(l)
HC2H3O2(aq)
HCl in toluene
HCl(aq)
NaOH(aq)
1.0 M NH3(aq)
15 M NH3(aq)
(17 M)
DATA
Enter a complete data table for all the substances you will test (see sample below) in your
notebook before coming to lab. During the experiment, record all data directly into
your notebook in ink.
Substance
1. H2O
deionized
Conductivity
Type
(high, medium,
low…)
Electrolyte
Type
(strong, weak,
non)
Major Species
(sol’ns only)
Minor Species
(sol’ns only)
Page 49
How to Use the Double-Diode Conductivity Meter
The conductivity meter will enable you to learn whether charged particles — electrons or
ions — are free to move in the substance you are examining. This meter detects five or
six levels of conductivity.
High conductivity means that electrons are free to move easily in the substance (as in a
metal), or that ions are present in large numbers and are free to move.
Low conductivity means that electrons in the substance are not free to move (as in a
nonmetal), or that there are few free ions in the substance or solution.
Before You Use the Meter
1) Turn the switch ON.
2) Pinch the electrodes together. If both red and green diodes glow brightly, the meter is
working and the battery is probably strong enough. The electrodes are best viewed
from the side rather than straight on. (If they don’t light up, or don’t glow brightly,
tell the technician.)
When You Use the Meter
1) Turn the switch to ON.
2) Test the meter frequently, as described above, to be sure it is working.
3) Touch both electrodes simultaneously to the solution or surface you are
testing. On a solid surface, scratch them back and forth several times to
improve contact.
GREEN diode RED diode CONDUCTIVITY
Bright
Bright
High
Medium
Bright
Medium
Dim
Medium
Low
Off
Dim
Very low
Off
Very dim
Extremely low
Off
Off
None
4) When testing a solution, it is sometimes helpful to pinch the electrodes
together and release them a few times in succession while they are
immersed in the solution in order to distinguish between bright and
medium conductivity levels.
5) Clean the electrodes frequently with a tissue to remove traces of the substance you
tested.
6) Turn the switch to OFF when you finish using the meter.
Las Positas College
Livermore, California
JHA110-8-96
Page 50
PRE-LAB:
Experiment 6
1.
2.
3.
NAME______________________________
Ions in Solution: Electrolyte Strength and Electrical
Conductivity
For each of these substances identify the bond types present (covalent, ionic, or
both). Hint: Check EN (difference in electronegativities).
A.
BaC12
B.
SO2
C.
HBr
D.
KNO3
E.
PbI2
F.
H2C2O4
Classify each of these ionic compounds as soluble, not soluble, or slightly soluble.
A. Mg (NO3)2
B.
PbCl2
C. Fe(OH)3
D.
K2SO4
Classify each of these as a strong acid, weak acid, strong base, or weak base.
A.
C.
E.
4.
Page 51
H2SO4
CH3CH2CO2H
NH3
B.
Ba(OH)2
D.
HClO3
F.
HF
a. Write balanced molecular (not ionic) equations for each of the following.
b. Identify each as PRECIPITATION (forming a solid which is not soluble),
ACID-BASE (formation of a weak acid or base), NEUTRALIZATION (acid +
base forms a salt and water), or DECOMPOSITION (one substance breaks up
into two).
If no reaction occurs, write NO OBSERVED REACTION. (Some
reactions may involve more than one type.)
A.
HNO3 + KC2H3O2 
B.
Mg(ClO3)2 + Fe2(SO4)3 
C.
H3PO4 + NH3
D.
Ba(OH)2 +

H2SO4

Page 52
Page 53
Experiment 7
PURPOSE
 to predict products for reactions among salts, acids, and bases
 to write balanced net ionic equations for those reactions
notebook.)
This experiment presents a brief introduction to reactions involving ions in aqueous
solution and how to write correct balanced net ionic equations for those reactions. The
reactions that will be demonstrated fall into three major types:
1.
Precipitation reactions: aqueous solutions of two substances are mixed and form a solid precipitate.
Equations for these reactions may be written by:
a.
predicting products by double replacement (partner exchange)—remember that the products
of ammonia reactions can be predicted by thinking of aqueous ammonia as if it existed in
solution as ammonium ions and hydroxide ions
b.
applying solubility rules to products to determine whether either of the products (or one of the
reactants!) is insoluble—if no precipitate will form, look for acid-base reactions
c.
balancing the equation in molecular form
d.
writing the total ionic equation by separating strong electrolytes into ions but keeping weak
and non-electrolytes as molecules—remember to write “NH4OH” as NH3(aq) + H2O(l)
e.
canceling unreacted spectator ions, and simplifying coefficients to obtain the net ionic
equation
2.
Acid-base reactions: the transfer of H+ between the reactant acid and base. Equations for these
reactions may be written by:
a.
predicting products by double replacement (partner exchange)—remember that the products
of ammonia reactions can be predicted by thinking of aqueous ammonia as if it existed in
solution as ammonium ions and hydroxide ions
b.
looking for reaction as strong acid or base  weak acid or base or as neutralization: acid +
base  salt + water—without the formation of a precipitate, a weak acid or base, or water, no
reaction will occur, so write no observed reaction/no expected reaction.
c.
balancing the equation in molecular form
d.
writing the total ionic equation as above
e.
writing the net ionic equation as above
3.
Complex formation reactions: a Lewis acid and a Lewis base combine to form a complex. One
type of complex formation reaction will be performed. (See note in the procedure.) Students are not
responsible for writing net ionic equations for complexation reactions at this time.
PROCEDURE
Safety
Many of the chemicals used are toxic, corrosive, or both. Handle them with care. Wear
your goggles at all times. Wear gloves. Wear a lab coat, or else wear old clothes to the
lab. Dispose of the chemicals in the waste jars provided.
Page 54
To perform each reaction, mix about 3 mL (estimated, do not waste time carefully
measuring exactly 3 mL) of each reactant. Record observations of colors, odors, and so
forth before mixing.
Mix the reactants well (the best way to do this is to cover the test tube with parafilm and
shake thoroughly), and then observe the result. Record changes in color, state (formation
of a precipitate or a gas), and temperature (note that temperature changes are often small
and not detectable). Note any odors produced. Remember the technique for detecting
odors: Hold the test tube about 6 inches from your nose. With your free hand, waft the
vapors toward your face. If no odor is detected, gradually move the test tube toward you
nose until an odor is detected (or until the test tube reaches your nose!) Another way to
achieve mixing is to pour the contents of the test tube back and forth between 2 test tubes.
Note also that the test tubes need not be dry if you are going to place aqueous solutions in
them. Just clean, rinse, give a final rinse with deionized water, and shake out the water.
For two of the reactions, you will be asked to use a precipitate from the previous reaction.
To do this, centrifuge the mixture, pour off most of the solution, add about 10 mL of
deionized water, stir the mixture well, centrifuge it, and again pour off most of the
solution. Now proceed with the reaction of the solid. The following reactions should be
performed:
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
Pb(C2H3O2)2(aq) + Na2CO3(aq)
The precipitate from reaction 1 plus excess HNO3(aq)
Pb(C2H3O2)2(aq) + H2SO4(aq)
The precipitate from reaction 3 plus excess HNO3(aq)
Pb(C2H3O2)2(aq) + NaOH(aq), add NaOH dropwise until precipitate forms
Continue adding NaOH(aq) until the precipitate in reaction 5 dissolves (*)
Pb(C2H3O2)2(aq) + 1M NH3(aq)
Co(NO3)2(aq) + Na2CO3(aq)
Co(NO3)2(aq) + H2SO4(aq)
Co(NO3)2(aq) + NaOH(aq)
Co(NO3)2(aq) + 1 M NH3(aq)
Co(NO3)2(aq) + 6 M NH3(aq) (**)
Co(NO3)2(aq) + 15 M NH3(aq) (**)
NH4C2H3O2(aq) + Na2CO3(aq) (***)
NH4C2H3O2(aq) + H2SO4(aq) (***)
NH4C2H3O2(aq) + NaOH(aq) (***)
H2SO4(aq) + Na2CO3(aq)
H2SO4(aq) + NaOH(aq) (****)
H2SO4(aq) + 1M NH3(aq) (****)
(*)
complex forms:
–
Pb(OH)2 (s) + OH (aq)
–
Pb(OH)3 (aq)
(**)
Page 55
complex forms:
Co(OH)2 (s) + 6NH 3(aq)
Co(NH3 )6
2+
–
+ 2OH (aq)
(***) The only observation possible for these three is an odor.
(****) The only observation possible for these two is a temperature change.
DATA
Enter the sample observations table below in your notebook before coming to lab. Make
sure to leave enough space. During the experiment, record all data/observations
directly into your notebook in ink.
Reaction
1. Pb(C2H3O2)2(aq) + Na2CO3(aq)
2. ppt #1 + HNO3
3. Pb(C2H3O2)2(aq) + H2SO4(aq)
4. ppt #3 + HNO3
5. Pb(C2H3O2)2(aq) + NaOH(aq)
6. ppt #5 + excess NaOH
7. Pb(C2H3O2)2(aq) + NH3(aq) (1 M)
8. Co(NO3)2(aq) + Na2CO3(aq)
9. Co(NO3)2(aq) + H2SO4(aq)
10. Co(NO3)2(aq) + NaOH(aq)
11. Co(NO3)2(aq) + 1 M NH3(aq)
12. Co(NO3)2(aq) + 6 M NH3(aq)
13. Co(NO3)2(aq) + 15 M NH3(aq)
14. NH4C2H3O2(aq) + Na2CO3(aq)
15. NH4C2H3O2(aq) + H2SO4(aq)
16. NH4C2H3O2(aq) + NaOH(aq)
17. H2SO4(aq) + Na2CO3(aq)
18. H2SO4(aq) + NaOH(aq)
19. H2SO4(aq) + NH3(aq)
Observations
Page 56
Balanced Formula, Complete Ionic, and Net Ionic Equations
For each of the mixtures above for which there was a reaction, write the balanced
molecular, total ionic, and net ionic equations. If no reaction occurred, write the phrase
“No observed reaction”. The numbers correspond to the reaction list. Equations may be
written in pencil. See sample data table below.
1. molecular
total ionic
net ionic
2. molecular
total ionic
net ionic
3. molecular
total ionic
net ionic
4. molecular
total ionic
net ionic
5. molecular
total ionic
net ionic
6. molecular
total ionic
net ionic
7. molecular
total ionic
net ionic
8. molecular
total ionic
net ionic
9. molecular
total ionic
net ionic
10. molecular
total ionic
net ionic
11. molecular
total ionic
net ionic
12. molecular
total ionic
net ionic
13. molecular
total ionic
net ionic
14. molecular
total ionic
net ionic
15. molecular
total ionic
net ionic
Pb(OH)2(s) + NaOH(aq)  Na[Pb(OH)3](aq)
Pb(OH)2(s) + Na+ + OH–  Na+ + Pb(OH)3–
Pb(OH)2(s) + OH–  Pb(OH)3–
Co(NO3)2(aq) + 6NH3(aq)  [Co(NH3)6](NO3)2(aq)
Co2+ + 2NO3– + 6NH3(aq)  Co(NH3)62+ + 2NO3–
Co2+ + 6NH3(aq)  Co(NH3)62+
Co(NO3)2(aq) + 6NH3(aq)  [Co(NH3)6](NO3)2(aq)
Co2+ + 2NO3– + 6NH3(aq)  Co(NH3)62+ + 2NO3–
Co2+ + 6NH3(aq)  Co(NH3)62+
16. molecular
total ionic
net ionic
17. molecular
total ionic
net ionic
18. molecular
total ionic
net ionic
19. molecular
total ionic
net ionic
Page 57
Page 58
Page 59
POST-LAB
Name_________________
(Problems and discussion to be turned in with the lab report)
Experiment 7
Write correct balanced molecular, total ionic, and net ionic equations for the reactions
that occur when the following substances are mixed. All are in aqueous solution except as
noted. Answers may be written in pencil.
1.
magnesium chloride and sodium carbonate
2.
aqueous ammonia and acetic acid
3.
nitric acid and magnesium acetate
4.
ammonium chloride and sodium hydroxide
5.
barium chloride and calcium nitrate
Page 60
6.
sulfuric acid and excess potassium hydrogen carbonate
7.
nitric acid and solid silver chloride
8.
solid aluminum hydroxide and nitric acid
9.
excess aqueous ammonia and sulfuric acid
10.
magnesium nitrate and aqueous ammonia
Page 61
Experiment 8
Prelab: Complete the Prelab questions on page 71 before lab.
PURPOSE
The purpose of this experiment is to determine the mass percent copper in a nickel coin.
Method
A nickel coin is an alloy of two metals, copper and nickel. The alloy dissolves readily in
nitric acid, yielding a solution of Cu2+ ions and Ni2+ ions. Excess nitric acid, a volatile
acid, is removed by adding the less volatile sulfuric acid, and heating.
The acidity of the solution is reduced somewhat by adding a controlled amount of
ammonia water. Under these conditions, the Cu2+ ion can be reduced to Cu1+ by the
hydrogen sulfite ion, HSO3-, and then precipitated as CuSCN by the thiocyanate ion,
SCN-. The precipitate, copper(I) thiocyanate, is collected on a filter paper, washed, dried
and weighed. From the two weights--that of the original alloy sample and of the final
precipitated compound--the mass percent copper in the original alloy can be calculated.
PROCEDURE
Safety First
As is usual, goggles must be worn throughout this experiment. Hot solutions of nitric
and sulfuric acids are involved. Wash with water at once after any accidental contact of
skin with acid. At the beginning of the experiment, evolution of noxious gases (oxides of
nitrogen) will require use of the fume hoods.
Some more safety notes: NH4SCN: toxic by ingestion; CuSCN: moderately toxic;
DMG: no particular hazards; Ni(NO3)2: strong oxidizing agent, tolerance (as dust) = 1
mg/m3, dangerous fire risk for solid; NaHSO3: contact with solid causes burns to
skin/eyes, strong irritant to skin and tissue, tolerance (as dust) = 5 mg/m 3;
SO3:
poisonous if inhaled, highly toxic, strong irritant to tissue, oxidizing agent, fire risk with
organics, forms H2SO4 with water producing heat
This procedure assumes a sample size of about 1.25 g, or one-fourth of a nickel. A whole
nickel can be used if the amounts or reagents are scaled up (four times), but the problem
with using the larger size sample is that longer times are required for filtering, washing
and drying the precipitate at the end.
The nickel will have been cut for you. You will be provided with two pieces (for two
runs) along with the year and mint of issue.
Page 62
Plan on making two trials or runs. Read through the directions that follow and see how a
second trial can be carried along at the same time as the first. Be sure to label weighing
dishes, flasks and filter papers to distinguish one sample from the other. Do not get the
two runs mixed up!
Work independently. Steps 1 through 15 must be completed during the first laboratory
period. Read labels carefully; be certain to use the specified concentration. Several
reagents are used with different concentrations at different times during this experiment.
Part I Separation of copper and precipitation as copper(I) thiocyanate.
1. Rub one of the pieces of coin with a paper towel and avoid touching it with your
fingers thereafter. Weigh it to ±0.001 g in a preweighed plastic weighing dish.
2. Transfer the sample to a 250 mL Erlenmeyer flask. Add 15 mL of 6 M HNO3 (dilute
nitric acid, from the reagent shelf). In the fume hood, heat gently (hot plate or low
flame) to get (react) the nickel into solution. Note color changes and fumes emitted.
Do not breathe the fumes. CAUTION! DO NOT LET THE FLASK DRY OUT
OR THE SOLID WILL BEGIN TO "BUMP!" Have some extra 6 M HNO3(aq)
readily available in order to add more 6 M HNO3 as needed to always keep some
liquid in the flask. If someone else cannot use your extra nitric acid, dispose of any
excess nitric acid in the appropriate waste container.
3. When the coin is completely dissolved (10 to 20 minutes), add 10 mL of 3 M H2SO4
(dilute sulfuric acid from the reagent shelf). Heat the solution to boiling. If a solid
forms, add 2 mL more of the 3 M H2SO4. White fumes should appear, at least faintly,
above the mouth of the flask when moist breath is blown across it. This would be
gaseous SO3 (from decomposition of H2SO4) reacting with water vapor to re-form a
fog of H2SO4. This usually takes 10 to 20 minutes also.
Make sure you start the second trial while you are waiting for the first to react.
4. Allow the flask to cool (Never put a very hot flask in cold water, but you may speed
the cooling by setting the flask in cool water) and move it to your desk.
5. Add about 25 mL of deionized water.
6. Measure out into your graduated cylinder 15 mL of 6 M NH3 (dilute aqueous
ammonia, sometimes labeled "ammonium hydroxide", from the reagent shelf). Start
adding the ammonia in small increments to the solution in the flask. You will soon
see a milky blue precipitate, Cu(OH)2, form where the ammonia hits the acidic
solution, and then go away as you swirl the flask. The object is to continue adding
ammonia only until that light blue precipitate just persists after swirling. (It may be
necessary to use more than 15 mL of the ammonia (possibly as much as 25 to 30 mL
more, depending on the amounts nitric and sulfuric acids used earlier) to get the light
blue precipitate. If a dark blue solution containing NO precipitate forms, you have
Page 63
added too much NH3. See the instructor.) This step takes a lot of ammonia. After
you have slowly added the 15 mL of ammonia, add more using small increments. It
is critical not to overshoot. Be patient! Discard the unused ammonia solution in the
appropriate waste container.
7. Add 15 mL of 1.0 M H2SO4. This solution will be prepared especially for this
experiment, and is not the reagent shelf sulfuric acid used earlier. The precipitate
should dissolve leaving a clear blue or blue-green solution. The solution is now at the
proper acidity, ready for the reducing agent.
8. Add 15 mL of 5% NaHSO3 (sodium hydrogen sulfite, or sodium bisulfite solution).
Note, very cautiously, the characteristic choking odor of SO2 that emanates. Very
little additional change will be observed at this point.
9. Promptly add 15 mL of 10% NH4SCN, ammonium thiocyanate solution, the
precipitating agent. Swirl the flask as the precipitate forms. Continue to swirl the
flask from time to time as the precipitate coagulates; that is, changes from a very fine
grain to a larger grain size. The larger grain size will be more filterable. Warming
will help this process, but do not boil the mixture. If you do not coagulate the
precipitate, there is a good chance that much will pass through the filter paper and be
lost. Contact instructor in this case before disposing of the filtrate.
10. While the precipitate is coagulating, prepare the funnel and filter paper for filtration.
Obtain a piece of 15 cm filter paper (Whatman #40 or equivalent). Fold the paper as
directed by your instructor. Briefly, the directions are to fold the paper in half, then
fold again but stop short of a 90o fold. Weigh the folded portion (after the corner is
torn off). Do not wet the paper while folding; it must be weighed dry.
11. Open up the larger quadrant of the weighed filter paper and fit it snugly into the
funnel. With your wash bottle, moisten the filter paper and press its top edge against
the glass to make a tight seal.
12. Start filtering the reaction mixture. Make a quantitative (without loss) transfer of
precipitate to the filter cone, using squirts of water from the wash bottle to assist you.
The filtrate should be collected in a beaker or flask and saved for proper disposal
later.
13. After all the precipitate has been collected on the filter paper, it must be washed until
free of Ni2+ ions. A suitable test for traces of nickel ion in the wash water is provided
by the reaction of Ni2+ with DMG in the presence of ammonia. DMG is
dimethylglyoxime, an organic compound that forms a bright red precipitate with the
nickel ion. Of course, what we want here is the absence of a red color, proving the
absence of nickel ion in the wash water.
a. Squirt about 1 mL of deionized water onto the solid and let it drain through. Repeat
this twice more.
Page 64
b. Collect the last mL of wash water directly from the funnel stem in a test tube. Be
very careful to avoid ANY contamination from previous washings as you collect
this mL sample of wash.
c. Then add about 1 mL of 6 M NH3 and about 1 mL of 1% DMG solution to the test
tube containing the last mL of wash.
d. Repeat the washing process until the test sample does not show the presence of Ni2+
ion. Collect a fresh sample of the latest washing, that is NOT contaminated with
any of the earlier wash solutions, for each repeated test to determine whether the
precipitate is now free of Ni2+ ion.
14. When washed free of Ni2+ ions, the precipitate is allowed to drain in the funnel. With
a spatula, open up a channel between the paper and the glass to allow the stem to
drain. Later, carefully lift the paper plus precipitate and rest it on a watch glass. Still
later, as it dries, open up the filter paper with your spatula, being careful not to lose
precipitate. In the meantime, dispose of the filtrate in the large jar labeled for this
purpose.
15. Place the watch glass in a safe place in your drawer to dry for two days or more. If
time remains in this lab period, you may start drying your sample in an oven set at
95oC for one hour and you may start the test tube experiments described later in the
section entitled "Some Reactions of Nickel and Copper Ions." Otherwise, do this part
next lab period. It is safe to stop the experiment at this step.
16. On the next lab day, complete the drying of the paper and precipitate by placing the
watch glass either (a) in an oven set at 95oC for one hour, or, (b) under a heat lamp
for one hour. While the sample is drying, perform Part II.
17. Allow the glass to cool, and then weigh just the paper and precipitate. Protect the
balance pan with a piece of paper which has been previously weighed, so that you can
quantitatively determine the weight of only your precipitate and filter paper.
18. If time permits, continue the drying operation for a half hour more, or so. Cool and
reweigh. If the second weighing is within 0.01 g of the first, the sample may be
considered dry. If there is a loss in weight greater than 0.01 g, you would ideally heat
it again until there was no significant change. Such a routine is known as "drying to
constant weight". Dispose of the precipitate in the disposal jar provided.
19. Calculate the mass percent of copper in the nickel coin, assuming that all of the
copper in your alloy sample has ended up as CuSCN.
Page 65
DATA
directly into your notebook in ink. Show sample calculations for all calculated data
for at least one trial if multiple runs were made.
Part I
Description of coin:
Date: _______ Mint: _______
Mass of weighing paper
Mass of weighing paper and nickel
Mass of sample of nickel
Mass of dry filter paper
Mass of protective paper (optional)
Mass of papers + precipitate (1st)
Mass of papers + precipitate (2nd)
Mass of papers + precipitate (3rd)
(if necessary)
Mass of precipitate
Trial #1
____________
____________
____________
____________
____________
____________
____________
____________
Trial #2
____________
____________
____________
____________
____________
____________
____________
____________
____________ ____________
Calculation of copper content Preserve proper precision, significant digits, in all
calculations!
Mass of copper present .............
(Show a clear sample calculation.)
____________ ____________
Mass percent copper in coin ........
(Show a clear sample calculation.)
____________ ____________
Average..................................
____________
Precision
Deviation from average..............
____________ ____________
Average deviation ............................
____________
Relative average deviation, % ...............
____________
(Average deviation/average mass percent)*100%
Accuracy (to be filled in by instructor)
____________% Error
Page 66
Part II - Reactions of nickel and copper ions
PURPOSE
In this second part of the experiment we will investigate some reactions of nickel(II) and
copper(II) ions which will allow us to tell them apart.
Method
A solution containing Ni2+ ions and a separate solution containing Cu2+ ions will be
subjected to the following test reactions (which are run independent of one another):
1) add aqueous ammonia solution
2) add aqueous sodium hydroxide solution
3) add ammonia followed by DMG (dimethylglyoxime) solution
PROCEDURE
For this part of the experiment you will need seven clean test tubes in all. You may work
with a partner, provided that each of you participates in all tests and that each of you
records your own observations. NO more than two students should be working together
on these test reactions!
1. Place 1 mL (20 drops) of 0.1 M Ni(NO3)2 in each of three test tubes. Now in three
other clean test tubes place 1 mL (20 drops) of 0.1 M Cu(NO3)2.
2. Ammonia test.
a. To one of the test tubes containing Ni(NO3)2 add 6 M NH3 drop by drop
until 10 drops have been added. Record your observations after the first
drop has been added and after all 10 drops have been added.
b. To one of the test tubes containing Cu(NO3)2 add 6 M NH3 dropwise as
before until 10 drops have been added. Record your observations.
3. Sodium hydroxide test.
a. To another portion of Ni(NO3)2 add 6 M NaOH dropwise until 10 drops have
been added. Record your observations.
b. To a sample of Cu(NO3)2 add 6 M NaOH dropwise as above.
4. Dimethylglyoxime test.
a. Place 1 drop of Ni(NO3)2 solution in the seventh clean test tube.
Page 67
b. To the tube with 1 drop of Ni2+ solution add 20 drops of deionized water.
Observe the effect of dilution on the color of the Ni2+ by comparing the color
in this test tube to the color in the one remaining tube containing the 20 drops
of Ni(NO3)2 solution that you prepared above.
c.
To both of these solutions in b. immediately above (the test tube with 1 drop
of Ni2+ and 20 drops of water AND the remaining tube with the 20 drops of
Ni2+) add 1 drop of 6 M NH3 followed by 10 drops of DMG solution. Do you
get a positive test for Ni2+ in both?
d. Perform this test on the remaining Cu(NO3)2 solution from above. Record
your
observations.
DATA
Report in ink all observations on mixing reactants directly into your notebook. See
sample observations table below.
Solutions
Test tube with 1 mL 0.1 M
Ni(NO3)2
Test tube with 1 mL 0.1 M
Cu(NO3)2
6 M NH3
(1st drop)
6 M NH3
(10 drops)
6 M NaOH
(10 drops)
Test tube with 1 drop
Ni(NO3)2 + 20 drops
H2O (blank)
Test tube with 1
mL Ni(NO3)2
6 M NH3 +1%
DMG
Answer the post-lab questions on the following page.
Test tube with 1 mL
0.1 M Cu(NO3)2
Page 68
Page 69
POST-LAB
Name_________________
Experiment 8
Part II: Reactions of nickel and copper ions
1. How do Ni2+ and Cu2+ differ in their reactions with NH3? with NaOH?
2. Calculate the concentration after 1 drop of 0.1 M Ni2+ solution was diluted with 20 drops
of water. (HINT: M1V1 = M2V2 where M1 and M2 are the molarities before and after
dilution and V1 and V2 are the volumes before and after dilution.)
3. Could you still see the green color of Ni2+ even after dilution with water?
4. Could you still detect the characteristic color of the DMG test for Ni2+ in the diluted
solution?
Equations - Complete and balance these ionic equations:
Ni2+(aq) + NH3(aq) + H2O(l)  Ni(OH)2(s) + NH4+(aq)
Cu2+(aq) + NH3(aq) + H2O(l)  Cu(OH)2(s) + NH4+(aq)
Ni(OH)2(s) + NH3(aq)  Ni(NH3)62+(aq) + OH-(aq)
Cu(OH)2(s) + NH3(aq)  Cu(NH3)62+(aq) + OH-(aq)
Ni2+(aq) + OH-(aq) 
Cu2+(aq) + OH-(aq) 
Page 70
PRE-LAB
Experiment 8
Page 71
NAME___________________
1. Why does nitric acid dissolve (react with) so many more metals than does either
sulfuric or hydrochloric acid? They are all strong acids. (Hint: Dissolving (reacting) a
metal in an acid is an oxidation-reduction reaction.)
2. What is a volatile acid? (See a dictionary).
Which of the following are volatile? HCl HNO3 H2SO4 H3PO4
(Circle the volatile ones. Hint: look up the boiling points and characteristics of each acid,
possibly in the Merck Index).
3. Balance the following net ionic equations:
a. For dissolving metals into solution by reaction with nitric acid
-
Cu2+(aq) +
-
Ni2+(aq) +
Cu(s) +
H+(aq) +
NO3 (aq) 
Ni(s) +
H+(aq) +
NO3 (aq) 
NO2(g) +
NO2(g) +
H2O(l)
H2O(l)
b. For reducing Cu(II) to Cu(I):
-
-
Cu2+(aq) + HSO3 (aq) + H2O(l)  Cu+(aq) + HSO4 (aq) + H+(aq)
c. To precipitate Cu(I) as copper(I) thiocyanate:
Cu+(aq) + SCN-(aq) 
4. What is the mass percent copper in copper(I) thiocyanate?
Show clear calculations.
Page 72
Page 73
Experiment 9
PURPOSE
 To use patterns of reactions involved in oxidation-reduction in predicting redox
behavior.
 To confirm these predictions from observations of some redox reactions.
 To represent these redox reactions using balanced chemical equations.
Introduction
"Redox" is a convenient term for oxidation-reduction reactions. These reactions often
involve a transfer of electrons from one reactant to another. They can be recognized by
noting a change in oxidation number (charge) in one or more elements as they move
from being part of a reactant to being part of a product.
Other references are extremely helpful for writing correct chemical equations which
describe the reactions that you observe. See other sources for discussion and drill on:
1)
2)
3)
4)
5)
6)
ionic equations
solubility
assignment of oxidation numbers
common compounds to illustrate various oxidations states of N, Mn, O, Cl, S, etc,
methods of balancing redox equations: the half-reaction method
acids and bases
Method
Redox reactions cannot be predicted as easily as simple partner exchange reactions, but it
can be done.
Transfer of electrons depends on two things:
1) the ease with which the reducing agent parts with electrons, and
2) the strength with which the oxidizing agent attracts the electrons.
You will not be required to make predictions of reaction products at this time. The factors
involved in such predictions will be discussed later. However, for this lab, we will
introduce the patterns of reactions involved in oxidation-reduction. For this purpose we
can make use of additional information that can be found at the back of this lab manual in
the Reference section:


The Activity Series of Metals
Nonmetals Common Oxidation States of Various Elements
An example of how the second of these tables can be used follows.
Page 74
Example: Predict the products of reaction between K2Cr2O7 and H2O2 in acidic
medium. (It makes a big difference in many redox reactions whether the solution is
acidic or basic.)
Answer to above example:
System:
KCr2O7 and H2O2, acidic solution
Prediction:
From information on oxidation states, we see that Cr in Cr2O72- is in its highest
oxidation state (+6). Therefore, Cr2O72- cannot be a reducing agent, but just might be a
good oxidizing agent. We also see that oxygen in H2O2 is in the -1 state, and can go
either up to zero (as in O2) or down to -2 .(as in H2O) . However, if there is to be a
reaction with KCr2O7 H2O2 must act as a reducing agent, and yield O2. The fate of
chromium in Cr2O72- may be revealed by a change in color.
Experiment:
A dilute solution of K2Cr2O7, was acidified with dilute hydrochloric acid, and H2O2
was added dropwise. The characteristic orange color of the Cr2O72- ion changed to a
greenish color, and bubbles of a gas were evolved.
Because Cr3+ is greenish in color in this case, and because a gas (O2) is
expected from the reaction of H2O2, we feel justified in proposing the skeleton
equation:
(Acidic)
Cr2O72-(aq) + H2O2 (aq)

Cr3+(aq) + O2(g)
To balance this equation, we use the technique of balancing redox equations discussed
in lecture and in your textbook. In the course of this process we will arrive at the two
half reactions:
6 e- + Cr2O72-(aq) + 14 H+(aq)  2 Cr3+(aq) + 7 H2O(l)
and
H2O2(aq)  O2(aq) + 2 H+(aq) + 2 e-
Continuing the process, these two half reactions, properly multiplied and added, yield
the net ionic redox reaction:
Cr2O72-(aq) + 3 H2O2(aq) + 8 H+(aq)  2 Cr3+(aq) + 3 O2(g) + 7 H2O2(aq)
Page 75
PROCEDURE
Four possible redox systems will be investigated. Work individually! Specific directions
are given.
Redox Reaction Mixing Directions
SAFETY FIRST: Wear goggles throughout the experiment. Exercise care in handling
acids and bases. Brown fumes from reaction of nitric acid are toxic. Perform such
reactions in the hood. Do not breathe or smell the fumes. To quench (stop) a reaction,
add several volumes of water.
System 1: Cu and HNO3(aq)
Perform in HOOD!
Place 2 or 3 copper shot (about the size of BB's) in a small test tube in the hood. Add
about 5 mL dilute HNO3 (6 M). Look for a reaction over a period of 5 minutes or so.
Note the color of the fumes above the test tube. Covering mouth of test tube with
parafilm, waiting, and holding tube in front of a white background may help in seeing
the color of gas produced.
System 2: Cu and HCl(aq)
Repeat the procedure of System 1, but substitute 5 mL 6 M HCl for HNO3.
System 3: NaHSO3 and KMnO4
Do three variations: [3A], [3B] & [3C]
[3A] Acidic solution
In large test tube, 0.1 M NaHSO3 (2 mL) + 3 drops of 0.5 M H2SO4; then add 10
drops of 0.1 M KMnO4, drop by drop. The sign of reaction will be the loss of the
permanganate color.
[3B] Slightly basic solution
0.1 M NaHSO3 (2 mL) + 2 drops 1 M NaOH; then add 0.1 M KMnO4 dropwise
(about 10 drops).
[3C] Strongly basic solution
0.1 M NaHSO3 (2 mL) + 1 mL 6 M NaOH; then add 0.1 M KMnO4
dropwise (about 10 drops). Note two reactions: the first, a color change, is
unique to system [3C]; the second, a precipitate, is like that of [3B].
System 4: KI and FeCl3
In a large test tube place 0.1 M KI (2 mL) + 0.1 M FeC13 (2 mL)
Note the color and form of the precipitate of I2(s) formed. This precipitate
may be very slow to form.
Page 76
DATA
Enter the sample report form below in your notebook before coming to lab. Record
your observations, as you go along, directly into tables in your lab notebook in ink,
not into the sample tables provided here. Remember that the other references on
oxidation and reduction are extremely helpful in writing the correct equations to describe
these redox reactions. Answer as many of these questions as possible before coming to
lab.
SYSTEM
Prediction: Do you expect
H2(g) to be evolved? Why
or why not?
What gas(es) might be
emitted if the nitrate ion
reacts?
Observations:
#1: Cu and HNO3(aq)
#2: Cu and HCl(aq)
Net ionic equation for reaction, if any.
System #1: Cu and HNO3(aq)
_____________________________________________________________________
System #2: Cu and HCl(aq)
_____________________________________________________________________
System 3: NaHSO3 and KMnO4
Prediction:
Is the HSO3- ion a possible oxidizing agent?
_________________________
If 'yes', to what species could it change?
_________________________
Is it a possible reducing agent?
_________________________
If 'yes', to what species could it change?
_________________________
Is the MnO4- ion a possible oxidizing agent?
_________________________
Page 77
If 'yes', list the species to which it could change, the color, and conditions that favor each
change:
species
color
conditions
Why is the MnO4- ion not a possible reducing agent? ________________________
Observations:
[3A] acidic:
____________________________________________________
[3B] slightly basic:
____________________________________________________
[3C] basic:
___________________________________________________
Net Ionic Equations: Write balanced redox equations for the three reactions in System 3
using the half reaction method. (Hint: one of the products for [3A] is HSO4-(aq))
[3A]
__________________________________________________________
[3B]
__________________________________________________________
[3C]
__________________________________________________________
System 4: KI and FeCl3
Prediction:
What possible redox role(s) can the iodide ion assume?
____________
To what species could it be converted?
____________
What possible redox role(s) can the Fe3+ ion assume?
____________
To what species could it be converted?
____________
Observations: ___________________________________________
Net Ionic Equation (Use half reaction method):
_______________________________________________________________________
Page 78
PRE-LAB
Page 79
NAME___________________
Experiment 9
See Reference: Common Oxidation States of Various Elements with Representative
Compounds
1. Determine the oxidation number of each element in the following ions or compounds:
(a) OH-
(c) O22-
(b) H2O2
(d) S 2 O 7 2 -
(e) Cr(OH)4-
2. For each of the following reactions, tell whether or not it is a redox reaction. If it is,
tell which reactant is the oxidizing agent (OA) and which is the reducing agent
(RA).
Yes/No
OA
RA
(a)
Cu(s) + 2 AgNO3(aq) 
(b)
NaHCO3(s) + HC1(aq)  NaCl(aq) + H2O(l) + CO2(g)
(c)
Fe(s) + Cl2(g) 
Cu(NO3)2(aq) + 2 Ag(s)
FeCl3(s)
3. In the following problem, "o.s." means "oxidation state". Write the formula of any salt
you can think of which contains:
S in +6 o.s.
S in +4 o. s.
S in -2 o. s.
N in +5 o.s.
N in +3 o.s.
N in -3 o.s.
Cl in +7 o.s.
C1 in +5 o.s.
Cl in -1 o.s.
4. Balance the following half reaction, and tell whether it is a reduction or an oxidation half
reaction.
NH4+ (aq) 
NO3-(aq)
(Acidic)
Which is it?
Page 80
Page 81
Experiment 10
PURPOSE
In this experiment we will evaluate the magnitude of the gas constant R in the ideal gas law
equation, PV = nRT. We will assume that we don’t know the value of R and we will
measure P, V, n, and T.
Method
To solve the ideal gas law, PV = nRT, for R, we need to have values for all the others: P, V,
T and n, describing an “ideal gas.” A direct approach would be to weigh a known volume
of a known gas and convert it to moles using molar mass. However, because of their low
densities, gas weights are difficult to measure directly. In this experiment the gas (H2) will
not be weighed. Instead, the number of moles present will be calculated from the
stoichiometry of the reaction which yielded it, namely
Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
The mass of one of the reactants, magnesium, will be measured carefully and then the molar
amount can be calculated. The other reactant, hydrochloric acid, will be taken in excess.
And, of course, the volume, temperature and pressure of the gas will be recorded.
The volume and temperature of the gas present no problems because they are direct
observations. The volume will be measured with a gas buret. (See the figure on the next
page.) The temperature will be measured with a thermometer. The number of moles of gas
is derived from the moles of magnesium reacted and the stoichiometry of the reaction
equation.
Measuring the pressure of the gas is complicated by two factors:
1. The gas pressure inside the buret is not the same as the barometric pressure.
2. The gas inside the buret is really two gases: H2 and H2O(g) (water vapor). Each gas
occupies the same volume but each at its own partial pressure.
First, we will find the total pressure of the gases, Pgas. It is the corrected barometric pressure
of the air, Pair, minus the pressure difference between the outside air and the trapped gases,
H2(g) and H2O(g). This pressure difference, ΔP, is due to the height of the column of water, h.
(Again, refer to the figure on the next page.) How to make this correction as well as the
reasoning behind it is given elsewhere. The relationship is:
Pgas = Pair - ΔP
Page 82
Second, we will find the vapor pressure of water at the temperature of the water, PH2O, (from
Appendix) and subtract it from the total pressure of the gases, Pgas, to yield the partial
pressure of the hydrogen, PH2.
PH2 = Pgas - PH2O
It is this partial pressure of hydrogen that we will use in the ideal gas law
PH2V = nH2RT .
PROCEDURE
SAFETY
Wear your safety goggles throughout this experiment. It involves an acid solution from
which a gas is evolved. The gas, hydrogen, is explosive. No flames or hot plates are
allowed in the vicinity.
1. Assuming you are using a gas buret of 100 mL capacity, your magnesium sample must
not exceed 0.080 g (about 8 cm). However, the oxide coating usually present must first be
removed before the sample is weighed. Place the ribbon on a piece of cardboard, and polish
each side with a piece of emery paper until the metal is
shiny. Wipe the ribbon free of dust with a paper towel.
Until it is weighed, avoid touching the strip with your
fingers, which could leave oils and moisture on the
surface.
2. Carefully weigh the magnesium to
±0.001 g.
Compress the weighed
sample into a small bundle.
(Fingers are okay now.) You can coil it loosely around a
pencil and then squash it. Next take about 20 cm of fine
(#24) copper wire and wrap it all around the bundle like a
cocoon. Leave only uniformly small holes. The idea is to
make a cage about the magnesium to allow the acid to
move through the holes to react with the magnesium and
to also hold the magnesium in place as it disintegrates
during reaction. Check to see that the wire cage will slip
inside a gas buret tube and then set it aside.
3. Along with the 100 mL gas buret, obtain a ring stand
with buret clamp attached. Rinse the buret with tap
water. If beads of water remain on the inside wall when
it drains, the buret is not clean. Scrub it with a buret
brush and detergent solution, and rinse it thoroughly.
Page 83
4. Into the gas buret pour about 10 mL of 6 M HCl (dilute hydrochloric acid). Use a funnel
in this operation to avoid getting acid on your hands. If you do, wash your hands promptly.
5. Remove the funnel and start adding deionized water to the buret from a small beaker.
Hold the buret at an angle from the vertical, and pour the water down the side slowly. The
idea is to layer the water over the acid with the least possible mixing. Fill it to the very top,
brimful. Clamp the buret in the vertical position while you prepare the rest of the assembly.
6. Have ready a 400 mL beaker, half full of water. (Tap water is okay here.) Also have
ready a one- or two-hole rubber stopper that will fit the buret.
7. Insert the magnesium sample about 4 cm into the buret. Lock it in place by inserting the
rubber stopper, pinching the wire stem. Be sure your stopper has a hole (or two) in it!
There should be no air bubbles below the stopper. See figure (a).
8. Quickly cover the stopper hole with your finger, and invert the buret into the beaker of
water. Remove your finger when the stopper is immersed. Clamp the buret in this position.
See figure (b).
9. The more dense acid, now at the top, will sink to the bottom and start to react with the
magnesium. As hydrogen is generated, it rises to the top, displacing water, which leaves
through the hole in the stopper. Toward the end of the reaction, small bits of magnesium
may break free and float to the top. (How can magnesium metal float on water?) Do not let
these pieces get stranded on the walls. Cautiously tilt the buret, if necessary, to release
them, but be sure to keep the rubber stopper submerged at all times.
10. When all the magnesium has disappeared and no more bubbles of gas arise, the
reaction is over. However, do not be too quick to disassemble the apparatus. You first must
make three measurements.
a. Read the volume of gas collected.
b. Measure, with a millimeter ruler, the height, h, in figure b, the height the water
stands in the buret above the surface of water in the beaker.
c. Measure the temperature of the water, which we will take to be the temperature of
the gas overhead.
11. Read and record the room temperature and uncorrected barometric pressure.
12. Empty the buret and beaker into the waste container. Rinse the buret, and repeat the
determination at least two more times.
DATA
Page 84
Room temp._________
Barometer
temp.
latitude
corrected
reading...__________corr'n________corr'n.._______bar. reading_____________
Trial 1
Trial 2
Trial 3
1. Mass of magnesium
__________ __________ __________
2. Volume of gas collected
__________ __________ __________
3. Temperature of water (& gas)
__________ __________ __________
4. Height, h, elevation of water
level in buret (in mm)
__________ __________ __________
5. Vapor pressure of water, PH2O
__________ __________ __________
CALCULATIONS
Show sample calculations where an asterisk, *, appears at the beginning of the entry
description.
1. * Partial pressure of H2, PH2.
__________ __________ __________
(corrected for vapor pressure and height difference of water)
2. * Moles of hydrogen, nH2
__________ __________ __________
3. * Calculated value of R
(from above exp'tl data)
__________ __________ __________
4. Average
5. Deviation from average
__________
__________ __________ __________
6. Average deviation
__________
7. Relative average deviation
__________
8. * % error (line 4 value compared to
accepted value, 0.08206 L.atm/mole.K)
______________
Page 85
POST-LAB
Name_________________
Experiment 10
Another way of expressing the results of this experiment is to calculate the molar volume,
V/n, of the gas at 0oC and 1.000 atm pressure, conditions known as STP, standard
temperature and pressure. Calculate from your data the value of molar volume, V/n, in units
of liters per mole at STP. Note that you are to use your experimental values for pressure,
volume and temperature in this calculation and NOT use either the experimental or
theoretical value of the gas constant, R, anywhere during this calculation.
Ans.
Would you expect to get the same % error as you got for your determination of R? Try it
and see. The accepted value for molar volume is 22.41 L/mol at 273.15 K and 1.000 atm
pressure. Show your supporting calculations.
% error __________
Is it the same yes or no ?
Page 86
Page 87
PRE-LAB
Name_________________
Experiment 10
Determination of the Gas Constant R
1. Write two balanced equations:
a. the reaction of magnesium with hydrochloric acid:
b. the reaction of magnesium oxide with hydrochloric acid:
2. In this experiment, care is taken to remove any "oxide coating" from the magnesium
metal before it is weighed. If a student failed to do this, how would this error affect the
calculation of R?
Too high______?
Show your reasoning:
Too low_______?
No effect_____?
Hint:
Solve the ideal gas law equation PV = Nrt for R. Then see how the error would affect
a. the numerator
b. the denominator
3. The magnesium sample is wrapped in fine copper wire before insertion into the acid
solution. Why isn't fine iron wire used instead? It's cheaper!
4. The experiment calls for about 0.08 g of Mg and 10 mL of 6 M HCl. What volume of
6.0 M HCl is theoretically required to react with 0.080 g of Mg? Show calculations:
Page 88
Page 89
Experiment 11
PURPOSE
The purpose of this experiment is to determine the mass percent of sodium hydrogen
carbonate (sodium bicarbonate) in an Alka-Seltzer tablet.
Method
The carbon dioxide evolved when Alka-Seltzer is added to water will be collected. Its
volume will be determined by measurement of the volume of water it displaces.
Application of the ideal gas law will allow calculation of moles of CO2. The number of
moles of CO2 is related to moles of NaHCO3 by the net ionic equation:
HCO3
(aq)
+ H+(aq)  H2O(l) + CO2(g)
Ordinarily, the acidity necessary for the reaction comes in the tablet itself in the form of a
solid organic acid, citric acid. However, in this experiment an inorganic acid, HCl, will be
added to the water as well to speed the reaction to completion.
One problem that must be faced is the appreciable solubility of CO2 in water. To reduce this
source of error, the water in the collection system will be pre-saturated with the gas. Then,
in calculating the final result, we will correct for the solubility of CO2 in the reaction
solution.
Another problem is that the Alka-Seltzer tablet must be kept dry until the controlled reaction
starts. Direct contact with the fingers and other sources of moisture must be avoided.
PROCEDURE
SAFETY
Wear your safety goggles throughout this experiment. It involves handling an acid solution
from which a gas is evolved. The gas, CO2, is harmless, but could expel acid if released
suddenly. Note: instructor's approval of apparatus is required before first run.
1. Handling the Tablet
a. Work individually. Each student will receive a pair of tablets sealed in foil, and
several small plastic weighing dishes. Use a pencil to make identifying marks on the
weighing dishes.
Page 90
b. Accurately pre-weigh (± .001 g) one of the dishes and record the mass. Open the foil
packet and, using the pre-weighed dish, accurately weigh one of the tablets to ±0.001 g.
Record the mass. Avoid wasteful opening of any more foil packets than absolutely
necessary.
c. Pick up the tablet using a piece of the foil or a piece of Parafilm to cover your
fingers. Place a ruler as a guide across a diameter of the tablet and score (scratch) the
tablet with a knife point or large pin. Still using the foil or Parafilm snap the tablet into
two halves. Shake off any crumbs or powder.
d. Using a pre-weighed dish, accurately weigh one of the halves (±.001 g). Then affix
one end of a fine 10 inch thread to the half tablet. To do this, use a very small piece of
cellophane tape (about one-fourth inch). Save the other half-tablet for saturating the
water in the collection system with CO2.
2. Saturating the Water with CO2
a. Fill a 500 mL Florence flask about half full with tap water. Pour in about 10 mL of 6
M HCl. Swirl to mix the solution, and add more tap water until the level is about threequarters up the neck of the flask.
B.Into the flask, drop the unweighed half-tablet, and set the flask aside as the solution
effervesces.
3. Establishing the Siphon
a. The siphon system, shown in Figure 1 to the
right, has been preassembled for you. Insert it
into the Florence flask after effervescence
(fizzing) has stopped. Have available a 600
mL beaker and a pinch clamp.
B.Using your pipet bulb, squeeze air into
stopper D to start a flow of water into the large
beaker, C, in Figure 2. When you see no more
air bubbles passing through the siphon, apply
the pinch clamp E to the rubber tubing to stop
the flow.
c. If the level of water in B is now appreciably below the neck of the flask, lift beaker C
above the level of water in flask B, loosen the clamp E and allow water to return from
beaker C to flask B, while you always keep the end of the glass tube below the surface of
the water in beaker, C, but retain some water in the beaker C for pressure equalization
(discussed later). Note that you must keep the end of the glass tube submersed under the
water level in beaker C as you run the water from beaker C back to flask B. Be sure the
stopper makes a very tight seal in flask B.
Page 91
4. Preparing the reaction flask, A
a. Insert a funnel into a 250 mL Erlenmeyer flask. Pour in 25 mL of 6 M HCl through
the funnel so as to not wet the sides of the flask. Remove the funnel and clamp the flask
to a ring stand.
b. Pick up the thread with the half tablet attached, and lower it inside the flask (very
carefully!) until it is about 2 to 3 cm above the acid solution. Be sure the tablet stays dry
on the way down. Immediately insert stopper D to secure the thread and to make an
airtight seal in flask A.
Stoppers must be tightly pushed into flasks B & A. Check!!!
c. Remove the pinch clamp E. There should be just a very small, brief flow of water
from flask B to beaker C. If water continues to flow, you have a leak and all connections
should be checked.
5. Pressure Equalization
a. The pressure inside flask A will not be the same as that in the room. To equalize the
pressures, keep the end of the glass tube under the surface of the water in beaker C and
raise beaker C until the level of water in beaker C is the same as in flask B. While
holding the water levels in beaker C at the same level as in flask B, put the pinch clamp
back on the tubing. If you need more hands to accomplish this task, ask someone around
you for help.
b. Now carefully hold the glass tube so no water drips out and pour the water from the
beaker C into another container to save as carbonated water for future trials. Drain the
beaker C well but do not dry it. Return the beaker to its normal position C with the glass
tube inside again. (This beaker C will collect the water displaced by the gas evolved
during the reaction.)
c. Remove the pinch clamp. Like before, you may see a very small, brief flow of water
from flask B to beaker C. You are now ready to start the reaction, but only after you get
Page 92
your instructor's approval and initials on your data sheet. Failure to remove the pinch
clamp would cause pressure to build up and the system would blow apart somewhere. So
be sure to get your instructor's approval before the first trial. On subsequent trials, ask a
person near you to provide an extra check to verify that the system is ready.
6. Starting the Reaction
a. Tilt flask A until the tablet sits in the acid solution. Clamp the flask A in this position,
but swirl it from time to time to dissolve the entire tablet. When no more effervescence is
noted, the reaction is over (about 10 minutes).
b. Do not measure the water in the beaker just yet; the gas pressure must first be
equalized as it was originally. Raise or lower the beaker C until the water levels are the
same between beaker C and flask B, just as you did at the start of this run keeping the end
of the glass tube submerged) and replace clamp E on the tubing. You have now captured
a volume of water in the beaker equal to the volume of gas evolved (CO2 and water
vapor) measured at the barometric pressure of the laboratory.
c. Measure the water in beaker C carefully in a graduated cylinder. Do not discard this
carbonated water. Save it for future trials.
d. Open the system and measure the temperature of water in flask B. We will take this to
be the temperature of the gas as well.
e. Empty flask A into the waste container, rinse it, and wipe dry the inside of the neck.
You are now ready for another trial. Using a preweighed dish accurately weigh the
second tablet and record the mass. Score it and snap it into two pieces. Discard small
fragments. Using preweighed dishes, accurately weigh separately each of the two large
pieces and record the masses. Be sure to put identifying marks on the plastic weighing
dishes so that you know which mass belongs to which run.
7. Collecting Other Needed Data
a. Read the barometer, and note the temperature and latitude corrections.
b. Determine the vapor pressure of water at the temperature of the gas collected; consult a
table elsewhere.
c. The solubility of CO2 in the acid solution in reaction flask A is assumed to be 20 mL.
d. Record the manufacturer’s listing of the composition of Alka-Seltzer on the data sheet.
Page 93
DATA
Name of Alka-Seltzer preparation tested: _____________________
1st tablet
2nd tablet
1. Mass of weighing dish
_________
_________
2. Mass of weighing dish and whole tablet
_________
_________
3. Mass of whole tablet
_________
_________
Trial 1
Trial 2
_______
_______
_______
5. Mass of dish and half-tablet
_______
_______
_______
6. Mass of half-tablet
_______
_______
_______
Approval of apparatus
Trial 3
(instructor)
7. Volume of water collected
_______
_______
_______
8. Temperature, water (and gas)
_______
_______
_______
9. Vapor pressure of water
_______ _______
_______
10. Barometric pressure __________ - __________- __________ = __________
uncorrected temperature latitude
corrected
reading correction correction
reading
CALCULATIONS Show a sample calculation for each one below for at least one trial.
1. partial pressure of CO2
_______
2. solubility of CO2 in flask A
20 mL
_______
20 mL
_______
20 mL
3. corrected volume of CO2
_______
_______
_______
4. moles of CO2
_______
_______
_______
Page 94
5. moles of NaHCO3 reacted
_______
_______
_______
6. mass of NaHCO3 reacted
_______
_______
_______
7. % by mass NaHCO3 in this
Alka-Seltzer preparation
_______
_______
_______
8. average mass percent NaHCO3
9. deviation from average
_______
_______
_______
10. average deviation
_______
11. relative average deviation, %
_______
_______
Calculation of % by mass sodium bicarbonate claimed on label
Name of Alka-Seltzer preparation
____________________
1. from the label: mass of NaHCO3 per tablet
__________
2. average mass of a tablet (from your data, line 3)
__________
3. calculated mass % NaHCO3, based on label
Calculation:
__________
4. taking label for true value, calculate % error
Calculation:
__________
POST-LAB QUESTION
Answer the following question in your notebook.
What factors might explain why your determination of mass percent NaHCO3 may not agree
with the information on the label? Be specific, but do not say things like "mistake in reading
the balance" or "mistake in reading the graduated cylinder." We assume you do not make
blunders like that.
Page 95
PRE-LAB
Name_________________
Experiment 11
Use the following data obtained in the analysis of an Alka-Seltzer tablet to make the
calculations asked for. They parallel those in the experiment to be performed. See the
sequence of calculations given in the Data and Calculations sheet for helpful hints.
Data
Mass of whole tablet
3.478 g
Mass of half-tablet taken for reaction
1.973 g
Volume of CO2 + water vapor (corrected)
314 mL
Temperature of gas
22.0oC
Vapor pressure of water at 22.0oC
21 torr
Barometric pressure, corrected
757 torr
Calculations (Show your work.)
1. partial pressure of CO2 __________
2. moles of CO2 __________
3. moles of NaHCO3 __________
4. mass of NaHCO3 __________
5. mass % NaHCO3 in Alka-Seltzer __________
6. mg of NaHCO3 per one tablet, calculated __________
On its label, the company (Miles Laboratories) claims 1904 mg NaHCO3 per tablet. Taking
this to be the true value, what is the % error in the above experimental work?
__________
Page 96
Page 97
Experiment 13
Determination of Heat of Reaction (using Logger Pro®)
PURPOSE
The objects of this experiment are to:
a. Determine the total heat capacity of a simple Styrofoam cup calorimeter
b. Measure the heat (enthalpy) change of two related chemical reactions, and
c. Calculate the heat change of a third reaction by applying Hess' law.
Method
The calorimeter used here will be similar to the one used previously. You will use
LoggerPro® to collect and display data as a graph and table, analyze your experimental
data, and print a graph and data table.
Two chemical reactions will be studied, both of which are exothermic:
(1)
Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
ΔH1
(2)
MgO(s) + 2 HCl(aq)  MgCl2(aq) + H2O(l)
ΔH2
The amount of heat evolved will be calculated from the temperature rise of the aqueous
solution inside the calorimeter. From the amounts of Mg and MgO reacting, ΔH1 and
ΔH2 will be determined. From these results, ΔH for the following reaction, difficult to
determine experimentally, will be calculated by application of Hess's law:
(3)
Mg(s) + H2O(l)  MgO(s) + H2(g)
ΔH3
In this experiment, the heat absorbed by the calorimeter will be taken into consideration:
heat evolved by reaction = heat absorbed by H2O + heat absorbed by calorimeter
qrxn = -(mH2OcH2OΔT + CcalΔT)
(1)
where Ccal is the total heat capacity of the styrofoam cup part of the calorimeter, mH2O is
the mass of water, cH2O is the specific heat of water, and ΔT is the temperature change.
The value of Ccal will be determined by adding a measured mass, mH2O, of hot water at Th
to the same mass, mH2O, of cold water in the styrofoam cup. The cold water and cup are
both at Tc. Heat will transfer from the hot water to the cold water until they reach the
same final temperature, T2.
heat lost by hot water = heat gained by cold water and styrofoam cup
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-mH2OcH2O(T2 - Th) = +mH2OcH2O(T2 - Tc) + Ccal(T2 - Tc)
(2)
The value of Ccal can be calculated from this relationship.
PROCEDURE
SAFETY
Wear your safety goggles as usual. Be careful to keep the cables from the temperature
probes away from the hot plate or the flame of the Bunsen burner!!! They are quickly
destroyed by excessive heat! In Part B and Part C you will be working with acidic
solutions and in Part B a flammable gas, H2, is evolved. Avoid breathing this gas, and
keep it away from flames and heaters.
Carefully measure all volumes, masses and temperatures (read to the estimated limit on
all scales) throughout this experiment!
Part A Determination of the Calorimeter Constant
1. Assemble a coffee cup calorimeter. Use two nested 8 oz. clean dry foam plastic cups,
placed in a 400 mL glass beaker, with a cardboard cover lid which has a hole in the
center. Set up LoggerPro with the Vernier Interface and two stainless steel
temperature probes (using CH1 and CH2). (Double click on the Logger-Pro icon on
the desktop. Change the time interval for data collection by going to EXP  DATA
COLLECTION and changing the time to a longer interval, e.g. 800 seconds).
2. Using a graduated cylinder, carefully measure 50.0 mL of water into the calorimeter.
Place the temperature probe connected to CH1 through the hole in a cardboard cover
and into the water in the calorimeter.
3. In a drained wet 150 mL beaker, heat another carefully measured 50.0 mL of water to
about 70oC. Observe its temperature with the temperature probe connected to CH2.
BE CAREFUL TO KEEP THE CABLES CONNECTED TO THE
TEMPERATURE PROBES AND ALL OTHER CABLES AWAY FROM THE
HOT PLATE OR THE FLAME OF THE BUNSEN BURNER!!! (They are
quickly destroyed by high temperatures!)
4. Pour all of the hot water from the 150 mL beaker into a third drained wet Styrofoam
cup and also transfer the temperature probe (connected to CH2) with the hot water in
the Styrofoam cup. Beware, this cup with temperature probe easily tips over, so you
may find it helpful to place this cup in a holder such as a glass beaker to give a little
more stability.
5. Begin data collection by clicking on "Collect". Gently stir both the hot and cold
water in the Styrofoam cups with the temperature probes in their respective cups
(cold still connected to CH1 and hot still connected to CH2) and watch for the
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temperatures (both hot and cold) to be holding approximately constant (probably
within about 30 seconds). (The hot water temperature will actually be very gradually
decreasing.)
6. When the two temperatures are holding approximately constant (not the same
temperature, but each separately constant), then quickly pour all of the hot water from
the Styrofoam cup into the cold water calorimeter and gently swirl and stir the
hot/cold mixture with the temperature probe (still connected to CH1) through the hole
in the lid on the calorimeter. Collect data a minute or so after mixing the hot and cold
water with the continuous gentle stirring.
7. When the temperature of the hot/cold mixture becomes constant (actually very slowly
decreasing), you may then stop collecting data by clicking on "Stop".
8. Look at the data you have collected and record the constant temperature of the cold
water just before mixing. Similarly record the best almost constant temperature of the
hot water just before you started pouring. Excepting for a brief spike, which might
appear, record the highest constant (almost) temperature of the hot/cold water mixture
observed soon after pouring together.
9. As an option, you may store your data by clicking on the "Data" Menu (top left of screen) and
then click on "Store Latest Run" in the Data Menu.
10. Drain the calorimeter, and wipe the inside dry.
11. Make at least three repeat runs (at least four runs total)
12. Perform the calculations for the calorimeter constant. Note that if the heat given up
by the hot water is less than the heat gained by the cold water, the calorimeter
constant will appear to be negative! In this experiment, since the mass, mH2O, of the
hot water and the cold water are the same (almost) and we assume the specific heat,
cH2O, of the hot and cold water are the same, a quick check on the experimental
results can be made by comparing the temperature change of the hot water, (T2 - Th),
and the temperature change of the cold water, (T2 - Tc). If the temperature change of
the hot water, (T2 - Th), is less than the temperature change of the cold water, (T2 Tc), that is (T2 - Th) < (T2 - Tc), then you will determine a negative calorimeter
constant, Ccal. Calorimeters usually do not have negative calorimeter constants, Ccal.
13. If you get negative values for Ccal, you should rerun the calorimeter constant
determinations. If you are limited by time, it may be possible to make additional
calorimeter constant runs after you finish the other parts of this experiment on
another day.
Very carefully measure all volumes, masses and temperatures (read to the estimated
limit) throughout this experiment!
Page 100
Part B
Magnesium Plus Hydrochloric Acid
1. Carefully measure out 100.0 mL of 1.0 M HCl(aq) into the clean dry calorimeter. You
may disconnect the temperature probe connected to CH2. Place the temperature
probe connected to CH1 through the hole in the cardboard cover and into the solution.
The calorimeter should now be covered with the cardboard lid and the temperature
probe should be immersed in the solution. Set this calorimeter aside.
2. Obtain a piece of magnesium, Mg, metal, in ribbon form, about 30 cm long. Because
such ribbon is usually covered with an oxide coating, it should be polished before
weighing. This can be done by stretching the ribbon on top of a piece of cardboard
and rubbing it with a piece of emery cloth. Polish both sides of the ribbon, and wipe
it free of dust. With minimum handling, cut the ribbon into lengths shorter than 2 cm.
Weigh these strips by placing plastic weighing dish on the balance and tare the
balance. Then measure the mass of the magnesium (you should have about 0.25 to
0.30 g of magnesium).
3. Begin data collection by clicking on "Collect". Gently stir the solution in the
calorimeter with the temperature probe and watch for the temperature to be holding
about constant. When the temperature is constant (probably within about
15 seconds), add the Mg strips and replace the cover. Swirl and stir the calorimeter
gently with the temperature probe (still connected to CH1). The temperature will
start to rise. Avoid breathing the gas (H2) evolved.
4. Collect data as the temperature rises to a nearly constant maximum and then begins to
very gradually drop to lower temperatures. You may then stop collecting data by
clicking on "Stop".
5. Look at the data you have collected and record the constant temperature of the HCl(aq)
solution just before adding the magnesium metal. Similarly record the constant (almost)
highest temperature (ignore very brief spikes) of the solution after adding the Mg.
6. You may store your data by clicking on the "Data" Menu (top left of screen) and then click on
"Store Latest Run" in the Data Menu.
7. Discard the contents into the waste container. (There should be solution only and not
any unreacted Mg metal remaining in the calorimeter!). Rinse the calorimeter with
water and wipe it dry.
8. Repeat the complete determination, starting with another carefully measured 100.0
mL of 1.0 M HCl(aq) and weighed cleaned Mg.
Part C Magnesium Oxide and Hydrochloric Acid
1. Carefully measure out 100.0 mL of 1.0 M HCl(aq) into the clean dry calorimeter.
Place the temperature probe connected to CH1 through the hole in the cardboard
Page 101
cover and into the solution. The calorimeter should now be covered with the
cardboard lid and the temperature probe should be immersed in the solution. Set this
calorimeter aside.
2. Weigh a plastic weighing dish. Add about 0.5 g of MgO to the dish, and carefully
reweigh it to ±0.001g. Reseal the bottle of MgO promptly. (After pouring out the
MgO from the weighing dish, the dish will need to reweighed, because some of the
MgO may stick to it, Step #7 below.)
3. Begin data collection by clicking on "Collect". Gently stir the solution in the
calorimeter with the temperature probe and watch for the temperature to be holding
about constant. When the temperature is constant (probably within about 15 seconds),
dump in the MgO rapidly but carefully so that it does not stick to the sides of the
calorimeter. Set the emptied weighing dish aside to reweigh later. At once, swirl the
cup gently and stir the calorimeter with the temperature probe (still connected to CH1)
to keep the solid from caking on the bottom. The temperature will start to rise.
4. Collect data as the temperature rises to a nearly constant maximum and then begins to
very gradually drop to lower temperatures. You may then stop collecting data by
clicking on "Stop".
5. Look at the data you have collected and record the constant temperature of the HCl(aq)
solution just before adding the magnesium oxide. Similarly record the constant
(almost) highest temperature (ignore very brief spikes) of the solution after adding the
MgO.
6. You may store your data by clicking on the "Data" Menu (top left of screen) and then
click on "Store Latest Run" in the Data Menu.
7. Reweigh the emptied plastic weighing dish.
8. Before discarding the contents of the calorimeter, check to see that no solid is left. (If
there is, you may want to change your technique slightly on the next trial and not use
the results of this run.) All of the solid MgO should be reacted with the HCl(aq).
9. Rinse the calorimeter and dry the inside.
10. Repeat the complete determination, starting with another carefully measured 100.0
mL of 1.0 M HCl(aq) and weighed MgO.
Page 102
DATA
Part A: Calorimeter Constant
Run:
1
2
3
4
1. Temperature of 50.0 mL cool water (Tc)
______ ______ ______ _______
2. Temperature of 50.0 mL warm water (Th)
______ ______ ______ _______
3. Maximum temperature on mixing (T2)
______ ______ ______ _______
4. Temperature change of hot water,
(T2 - Th)
5. Temperature change of cold water,
(T2 - Tc)
6. Heat lost by hot water,
mH2OcH2O(T2 - Th)
______ ______ ______ _______
______ ______ ______ _______
______ ______ ______ _______
7. Heat gained by cold water,
+mH2OcH2O(T2 - Tc
______ ______ ______ _______
8. Heat transferred ("lost") to
calorimeter, qcal,
______ ______ ______ _______
9. Total heat capacity of the calorimeter,
qcal/ΔTcal = Ccal, in joules per degree,
______ ______ ______ _______
where ΔTcal = (T2 - Tc). Note: If you get negative values for Ccal, you should rerun
these calorimeter constant determinations after you finish the other parts of this
experiment.
10. Average:
__________
(Note: You should repeat runs for the calorimeter constant until you get positive values
for Ccal. If your results continue to be negative, consult your instructor. (In place of
continually negative results, you may be instructed to use Ccal = 10 J/oC.)
Part B:
Mg plus HCl(aq)
Trial 1
Trial 2
__________
__________
2. Mass of weighing dish + magnesium
__________
__________
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3. Mass of magnesium
__________
__________
4. Initial temperature of acid sol'n (T1)
__________
__________
5. Final temperature (maximum, T2)
__________
__________
6. Heat absorbed by the solution, msolncH2OΔT,
(assume 100.0 g, with c = 4.18 J/goC)
__________
__________
7. Heat absorbed by the calorimeter, CcalΔT
__________
__________
8. Total heat evolved by reaction, qrxn
__________
__________
9. Moles of magnesium
__________
__________
10. ΔH (kJ/mol Mg)
__________
__________
11. Average
Part C
__________
MgO plus HCl(aq)
Trial 1
__________
Trial 2
__________
2. Mass of weighing dish plus MgO
__________
__________
3. Mass of weighing dish, after using MgO
__________
__________
4. Net mass of MgO taken for reaction
__________
__________
5. Initial temperature of acid solution (T1)
__________
__________
6. Final temperature (maximum, T2)
__________
__________
7. Heat absorbed by the solution, msolncH2OΔT,
(assume 100.0 g, with c = 4.18 J/goC)
__________
__________
8. Heat absorbed by the calorimeter, CcalΔT
__________
__________
9. Total heat evolved by the reaction, qrxn
__________
__________
10. Moles of MgO
__________
__________
11. ΔH (kJ/mol MgO)
__________
__________
12. Average ΔH
__________
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Part D Application of Hess's Law
a.Write in the experimentally derived energy terms in the following thermochemical
equations. (These are the results from your data above.)
(1) Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
ΔH1 = _________
(2) MgO(s) + 2 HCl(aq)  MgCl2(aq) + H2O(l)
ΔH2 = _________
b.Use Hess' law to get ΔH3 for reaction (3) below from the experimental ΔH1 and ΔH2
entered immediately above. (This ΔH3 is based on your experimental data in (1) and (2)
above.) Show calculations.
(3) Mg(s) + H2O(l)  MgO(s) + H2(g)
ΔH3 = _________
c. Now do an independent (separate) calculation of ΔH3 from the standard molar
enthalpies of formation, ΔHfo, found in a reference table elsewhere (e.g. Handbook of
Chemistry and Physics). Note that you do NOT use your experimental data in this
calculation. Show calculations. Draw a box around your answer.
d.How does the theoretical ΔH3 (step c. above) compare with your experimental ΔH3
(step b. above)?
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POST-LAB
Name_________________
Experiment 13
1. Various assumptions and simplifications were made in this experiment that reduce
precision. However, considering the greater uncertainty in measuring temperature,
they are probably excusable. But now tell how these ignored errors would tend to
affect the value of ΔH obtained:
(H increases, decreases,
or no effect)
a. Not all of the "100 mL" of acid entered the
calorimeter.
b. The density of the acid solution is not exactly 1.00
g/mL, but slightly higher.
c. The specific heat of the acid solution is less than the
assumed 4.18 J/goC.
d. The "maximum temperature" observed is really lower
than if we had made a time-temperature plot to correct
for heat loss.
e. The thermometer itself has a finite heat capacity; i.e.,
it absorbs some heat.
2. In Part D, Line c, you calculated ΔH3 from the standard heats of formation (standard
molar enthalpies of formation), ΔHfo, taken from a reference table (e.g. Hand Book of
Chemistry and Physics). Now do the same for ΔH1 and ΔH2. (Again note that you do
NOT use your experimental data in this calculation.) Show both calculations below.
ΔH1
ΔH2
How can ΔH1 and ΔH2 be combined to give ΔH3? That is show the simple mathematical
calculation that gives the value of ΔH3 from the values of ΔH1 and ΔH2.
What thermo-chemical law provides the basis for what you are doing here in the simple
mathematical calculation?
Page 106
Page 107
PRE-LAB
Name_________________
Experiment 13
1. Calculate the final temperature for each of the following mixtures. Assume no heat
loss to the environment. Show your calculations.
a. 50.0 g of water at 25.0oC is mixed with 50.0 g of water at 85.0oC.
b. 75.0 g of water at 25.0oC is mixed with 25.0 g of water at 85.0oC
2. a. In Part B of this experiment, 0.20 g of Mg is added to 100 mL of 1.0 M HCl(aq).
Which is the limiting reactant? Show calculations.
b. In Part C, 0.50 g of MgO is added to 100 mL of 1.0 M HCl(aq). Which is the
limiting reactant?
3. When 50.0 mL of 1.00 M H2SO4(aq) at 26.1oC was added to 50.0 mL of 1.00 M NaOH
also at 26.1oC, the temperature rose to 32.6oC. Assume the resulting solution had a
total volume of 100.0 mL with the same density and specific heat as water. Calculate
ΔH for the reaction described by this equation.
H2SO4(aq) + 2 NaOH(aq)  Na2SO4(aq) + 2 H2O(l)
To do this calculation, work through the following questions and steps (a-f):
a. Which reactant limits?
b. How many moles of it are present?
Page 108
c. How much energy is required to heat 100.0 g solution from 26.1oC to 32.6oC?
d. How much energy is released by the reaction per one mole of the limiting
reactant?
e. How many moles of that reactant are in the given equation?
f. What is ΔH for the reaction as written above in kJ (not kJ/mol)?
Page 109
Experiment 14
PURPOSE
In this experiment you will learn to use a spectrophotometer to determine the molar
concentration of Crystal Violet in an unknown solution.
Method
The principles of light absorbance and the theory behind spectrometry are discussed
elsewhere.
chemicals
1.00 x 10-4 M Crystal Violet stock solution
equipment
Vernier Logger Pro Spectrometer and cuvettes
50 mL volumetric flasks
10 mL graduated pipet
Safety
Crystal Violet solutions can be irritating to the eyes, skin and clothing. As usual, use eye
protection while handling these solutions.
PROCEDURE
Part 1. Preparing the Standard Crystal Violet Solutions
a. Five solutions will be prepared below (as directed by instructor). The total volume for
each solution will be 50.00 mL.
Standard Solution Number
Volume of 1.00 x 10-4 M
Crystal Violet in mL
1
2
3
4
5
5.00
4.00
3.00
2.00
1.00
b. For each solution, thoroughly clean a 50 mL volumetric flask, including a final rinse
with deionized water.
c. Thoroughly clean a 10 mL graduated pipet*, again rinsing with deionized water.
Obtain about 10 mL of the 1.00 x 10-4 M Crystal Violet solution in a clean, dry 50 mL
beaker. Rinse the pipet with 2 or 3 small portions of the Crystal Violet solution.
Discard the rinsings and any solution in the beaker.
Page 110
* Graduated pipets, unlike volumetric pipets, require that you start and stop the solution
level at the desired marked line.
d. Obtain a fresh portion of about 30 mL of the 1.00 x 10-4 M Crystal Violet solution in
the 50 mL beaker. Using the pipet just rinsed with the Crystal Violet solution, pipet* the
assigned (1. a. above) volume into the 50 mL volumetric flask. Label the flask with the
appropriate solution number (see 1. a. above)
e. Add deionized water to the flask until the solution level is almost to the mark.
Stopper the flask, and mix the contents by inverting the flask 8 or 10 times. Be sure to
hold the stopper firmly! Allow solution to drain down the neck of the flask and do not
lose any of the solution as you remove the stopper at this point! Using a transfer pipet,
add deionized water to the mark. Again stopper the flask and thoroughly mix by
inverting the flask several times.
f. Prepare solutions 2, 3, 4, and 5. If there are not enough 50-mL volumetric flasks,
transfer solution 1 into a clean container and reuse the same 50-mL volumetric flask.
Do the same for the other solutions.
g. For your report, you will need to know the concentration of each of the 5 diluted
solutions that you have prepared. Calculate the concentration of each diluted solution
and enter into the table (see Data section below). Show your calculations in your
notebook.
Part 2. Determining the Absorbance Curve of the Crystal Violet solution.
a. Prepare the Vernier Spectrometer by plugging in the USB cable and opening the
Logger Pro software. If the software doesn’t immediately recognize the Spectrometer,
choose Connect Interface  Spectrometer  Scan for Spectrometers from the
Experiment menu. Allow the Spectrometer to warm up for 3 minutes before taking
readings.
b. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use a
cuvette, remember:
 All cuvettes should be wiped clean and dry on the outside with a tissue.
 Handle cuvettes only by the top edge of the ribbed sides.
 All solutions should be free of bubbles.
c. Calibrate the spectrometer by choosing Calibrate  Spectrometer from the
Experiment menu. Follow the instructions from the dialog box to complete the
calibration using your blank cuvette. You will be asked to insert the blank cuvette into
the cuvette slot. Insert it in such a way that the spectrometer light goes through the
smooth sides and not the ribbed sides of the cuvette. Click “OK.”
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d. Rinse a sample cuvette with two or three small portions of the #1 diluted Crystal
Violet solution from its volumetric flask. Fill the cuvette about ¾ full with the #1
diluted Crystal Violet solution. Place the cuvette in the cuvette slot. Click “Collect”. An
absorbance curve should appear on the screen. After viewing the absorbance curve, hit
“Stop”.
e. Click on the “Configure Spectrometer Data Collection” icon, located on the right hand
side of the toolbar to open the display. (The button looks like a rainbow graph.) Click
Abs. vs. Concentration (under Set Collection Mode). The wavelength of the maximum
absorbance will be automatically selected. Double check that 460 nm or somewhere
close (this is the max) is the only wavelength that is selected. Record this max. Click
“OK” to close the display.
Part 3. Measuring the absorbance of each standard solution and your unknown
a. With the first Beer’s law standard solution in the cuvette slot (#1), record the
absorbance value shown on the lower left-hand corner of the screen (the last digit may
fluctuate so do your best to find the average).
b. Transfer a portion of solution #2 into a clean cuvette, pre-rinsed with 2 or 3 small
portions of the solution, and insert into the cuvette slot. Record the absorbance. Repeat
this step for the remaining standard solutions.
c. Transfer a portion of your assigned unknown solution into a clean cuvette, pre-rinsed
with 2 or 3 small portions of the unknown solution. Place the cuvette in the sample
holder. Record the absorbance.
d. Pour out all Crystal Violet solutions into the waste container provided and rinse the
containers with water.
Part 4. Determination of the Concentration of an Unknown Crystal Violet Solution
a. You will use the Excel program to make your Beer’s Law Plot. Display the trend
line, the equation, and the R2 value. Make sure to set the y-intercept at zero. Use a
proper title for each axis and for the plot. Make sure to print your Beer’s Law Plot (or if
your instructor desires, send by e-mail).
b. Use the Beer’s Law Equation generated from your plot to determine the concentration
of your unknown.
Page 112
DATA
λmax __________
1) Table of Concentrations and Absorbance Values
Solution Number
Volume of 1.00 x 10-4 M
Crystal Violet in mL
Concentration (M)
1
2
3
4
5
5.00
4.00
3.00
2.00
1.00
Absorbance
Unknown Code
Absorbance of the
unknown solution
Calculated
concentration of the
unknown solution
2) Attach your Beer’s Law plot (with the equation displayed) to your lab report (or e-mail
to instructor if told to do so).
CALCULATIONS
1) Show the calculation of the concentration of each standard.
2) Using the equation determined from your Beer’s Law plot, calculate the concentration of
your unknown sample.
PRE-LAB
Experiment 14
Page 113
NAME_____________________
1. For this question, assume that you have been assigned solution number 3 to prepare.
Calculate the concentration of Crystal Violet in this solution.
2. A student has found the following %T readings for solutions of the molarities listed
below. All measurements were done at the wavelength of maximum absorbance. For each
%T reading, calculate the absorbance and enter the value in the table below.
(A = 2 – log %T )
%T
Molarity
Absorbance, A
40.9 1.6 x 10-4
_____________
50.5 1.3 x 10-4
_____________
59.0 9.6 x 10-5
_____________
70.4 6.4 x 10-5
_____________
3. Beer's law: A = εbc, where A = absorbance, ε = molar absorptivity, b = path length, and
c = concentration in M.
Plot the absorbance, A, (y-axis) versus concentration, c, (x-axis) using Excel. Print your
graph and attach to this form (Check with instructor if e-mail is OK).
4. Display the trend line for this linear relationship and the equation for the line.
5. Again following the procedure of the experiment, the same student found a %T of 56.4
for the unknown. Find the concentration of the unknown solution in two ways.
a. First, read the concentration value off the graph. Show how you read the unknown value
on the graph on the next page using dashed lines (----).
b. Second, calculate the concentration value using the slope of your line and Beer's law:
A = εbc, where A = absorbance, ε = molar absorptivity, b = path length, and
c = concentration in M. (Note that εb is the theoretical slope of the line.)
c. What is the theoretical value of the y-intercept (absorbance, A) on a Beer's Law plot?
Page 114
Page 115
Experiment AA
Measurement of Iron by Atomic Absorption (AA) Spectrometry
Prelab: Complete the prelab questions on page 121 before lab.
PURPOSE
The purpose of this experiment is to learn the technique of atomic absorption
spectrometry. The AA instrument will be used to detect and measure the iron content of
an unknown sample.
Method
Your instructor will discuss the general theory behind atomic absorption spectrometry
and how the instrument works.
Acid will be used to dissolve the iron in the unknown sample. A calibration curve will be
prepared using a series of standard Fe3+ solutions whose absorbances will be measured
using the AA instrument. The absorbance of iron in an unknown sample will be
measured using the AA instrument and the concentration calculated using the calibration
curve.
Reference: D.T. Sawyer, W. R. Heineman, and J. M. Beebe, “Chemistry Experiments for
Instrumental Methods,” John Wiley and Sons, New York. 1984.
PROCEDURE
Safety
 Be careful when boiling the sample in 8M HCl (aq). Make sure that the fume
hood is over your heating apparatus. Do not leave your flask unattended and do
not let it dry out.
 The AA instrument uses an acetylene flame. Avoid getting too close to this part of
the instrument.
A. Preparation of the Sample to be Analyzed
If your unknown sample is a solid:
1) Rinse a clean 250-mL flask 3 or more times with deionized (DI) water. Do not insert
a bottle brush into your flask.
2) Measure about 2 grams of the unknown sample and transfer it into the clean 250 mL
flask. Record the exact mass.
3) Add 25 mL of 8 M HCl (aq) to the flask and boil slowly on a hot plate for 5 minutes.
Avoid drying out!
4) Cool the mixture and then add 10 mL of deionized water.
Page 116
5) Rinse a clean glass funnel, a 125-mL erlenmeyer flask, and a 50-mL volumetric flask
3 or more times with deionized water.
6) You will carry out 3 levels of filtration. Line your glass funnel with each type of
filtering material and make sure to rinse well between each filtration. The first
filtrations should be done into an Erlenmeyer flask (ask for an extra one if you run
out). The last filtration should be done into the 50 mL volumetric flask.
a) Using a glass funnel, filter first with a double layer of cheesecloth into a clean
125-ml flask to remove the largest pieces.
b) Next, filter the mixture through a “Fast” filter paper into another clean
erlenmeyer flask. (If your filtration is taking a while, proceed to Part B and
come back to this).
c) Next filter the mixture through a Whatman No. 1 filter paper into the 50-mL
volumetric flask.
7) Once all of the liquid has filtered through (this is what you will analyze), make sure
that there are absolutely no solid particles in the solution. If there are solid particles
still left, re-filter the liquid.
8) Dilute to the mark with deionized water. You will use this solution directly for the
absorbance measurement.
B. Preparation of the Standard Iron Solutions:
1) Four solutions of known concentrations of iron will be prepared by dilution from the
1000 mg/L stock solution:
Concentration of
Standard Solution
(mg/L or ppm)
Standard 1
Standard 2
Standard 3
Standard 4
10.00
30.00
50.00
100.0
Calculate in your notebook the volume of the stock solution you will need to dilute for
each standard. Show your calculations to your instructor before proceeding.
2) Thoroughly rinse 4 100-mL volumetric flasks with deionized water. Do not use a
bottle brush! Label each one with the standard number and the concentration.
3) Using a 10-mL graduated pipet, transfer the required volume of the stock solution
into each of the 100-mL volumetric flask.
4) Dilute to the mark with DI water. Cover with parafilm or stopcock.
C. Absorbance Measurements
1)
2)
3)
4)
Your instructor will demonstrate how to use the AA instrument.
Take your standard solutions and the unknown solution to the AA instrument.
The instrument is ready to use when the flame is on. Keep away from this area of
the instrument. The aspirator tube should be sitting in and aspirating DI water
when no samples are being tested.
You will first have to do a blank calibration by aspirating the “blank” sample.
Gently transfer the aspirator to the container labeled “blank”. Avoid pinching the
5)
6)
7)
8)
9)
10)
11)
12)
13)
14)
15)
16)
17)
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fragile aspirator tube. Make sure that the tip of the aspirator is completely below
the liquid surface.
Using the pen stylus, click on the “Analyze Blank” box. As the instrument proceeds
with the multiple measurements, you can watch the progress in the window on the
upper right hand corner of the screen. (See sample screen shot on the next page).
When the instrument is done with the blank calibration, your screen should say
“Autozero performed”.
Aspirate your first standard by gently transferring the aspirator into Standard 1
flask. Make sure the tip is in the solution.
The concentrations of your standards should have been entered by your instructor.
Next to the “Analyze Standard” button, make sure that the correct Standard # and
concentration are showing.
Using the pen stylus, click on “Analyze Standard”. Watch the “progress” box on
the upper right hand corner. The measurement is done when the screen says
“Standard 1 applied”:
Record the absorbance value into the data table in your notebook.
Insert the aspirator back into the DI water beaker to clean the aspirator for a few
seconds.
The window next to “Analyze Standard” should say Standard 2 with the correct
concentration.
Transfer the aspirator from the DI water beaker into the Standard 2 flask. Click on
“Analyze Standard” as before and wait for the “Standard 2 Applied” message to
come up. Again, record the absorbance.
Follow the same steps for Standards 3 and 4, remembering to rinse the aspirator
with DI water before each standard analysis and recording the absorbance
measurement. At the end of the standards analysis, leave the aspirator in the DI
water beaker.
After the absorbance measurements are done for the 4 standards, you will now
measure the absorbance and concentration of iron in your unknown solution.
Do one last check to make sure that there are absolutely no solids in your sample.
This will clog the very tiny aperture for the nebulizer!
Transfer the aspirator into your unknown sample flask. Click on “Analyze Sample”.
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18) When the sample analysis is done, record the concentration of iron in your sample.
19) If the concentration is higher than that of the most concentrated standard, dilute the
sample ten times and repeat the analysis.
20) Remove the aspirator from your flask and gently transfer back into the DI water
beaker for the next person.
D. Determination of the Absorbance of Iron in Your Unknown.
1)
2)
3)
4)
5)
Prepare your calibration curve. Using the Excel program, make a plot of the
absorbance versus standard concentrations.
Using the “Add Trendline” function (under the Chart menu), do a linear regression
analysis.
In the same “Add Trendline” window, click on option and check the boxes to set the
intercept to zero and display the equation and the R2 value.
Use the equation to calculate the absorbance of the iron in your unknown.
Print the chart with the correct labels and staple to your report.
DATA
1)
Description of the unknown sample (brand, type of cereal, size, texture, etc.)
2)
% RDA or DV of iron per serving __________ and serving size in grams _______
3)
Mass of the unknown sample
4)
Concentration of stock solution
5)
Data table below:
Concentration of
Standard Solutions
(mg/L or ppm)
Volume of stock
needed to make 100
mL of each standard
solution
Absorbance of each
standard
Concentration of iron
in unknown from AA
result
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Standard 1
Standard 2
Standard 3
Standard 4
10.00
30.00
50.00
100.0
Calculated
absorbance
of unknown
CALCULATIONS
1)
Calculations for the volumes needed to prepare the standard solutions
2)
Calculation of the absorbance of unknown iron solution
3)
Calculate the mass of iron in 1 gram of your food sample using your concentration
result for the unknown.
4)
Calculate the mass of iron in 1 g according to the label nutritional information using
% of RDA or DV per serving. Assume an RDA of 15 mg.
5)
Do these two values agree with each other? Give some reasons as to why they may
be different.
Page 120
Page 121
PRE-LAB
Name_________________
Experiment AA
Measurement of Iron by Atomic Absorption (AA)
Spectrometry
1) Using an acceptable source, look up some information on how an atomic absorption
spectrophotometer works (online or from a text or journal article). Briefly summarize
what you learned below:
2) Explain in your own words how you would prepare Standard 1. Show your
calculation for the volume required of the stock solution with a concentration of 1.00
x 103 mg/L.
3)
Using the sample results below, calculate the concentration of iron in a sample that
has an absorbance of 0.72.
Page 122
Page 123
Experiment 15
PURPOSE
In this experiment we will
• find evidence of similar chemical behavior of elements that have a vertical
arrangement in the periodic table, and
• find progressively different behavior of elements that have a horizontal arrangement
in the periodic table.
Method
Elements that have a vertical arrangement in the periodic table are, of course, members of
the same family or group. Just as in other families we know, all members are not exactly
alike, but they often have much in common that distinguishes them from members of
other families.
Elements that have a horizontal arrangement in the periodic table are members of the
same period. An analogy to the progressive differences in these elements is the ecological
differences you might observe traveling across the state of California due east from the
Pacific coast. Whether you started the trip at Crescent City, Santa Cruz or Santa Barbara,
you would find a sequence of changes that seemed familiar.
For our brief survey of the periodic table we will focus on a dozen or so elements:
2nd
3rd
4th
5th
IA
Li
Na
K
IIA
IIIA
Mg
Ca
Al
IVA
C
VA
P
VIA
O
S
VIIA
Cl
Br
I
The chemical behavior of these elements will be investigated in the following categories:
I.
II.
III.
IV.
Reactivity with water
Reactivity with acids
Study of oxides: basic and acidic
Redox behavior of the halogens.
Because we are not trying to "reinvent" the periodic table on the basis of empirical
evidence, it will be quite appropriate to relate degree of reactivity to the knowledge of
electronic structure we take from the textbooks. If you know how—and why—ionization
energy varies in a left-to-right direction and in a top-to-bottom direction within the
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periodic table, you already have a powerful tool for predicting and explaining many
periodic trends.
PROCEDURE
Work with one, and only one, partner.
SAFETY FIRST
Wear your safety goggles throughout the experiment. Access to lithium, sodium and
potassium will be supervised by your instructor, but be aware that these metals react
explosively with water if not properly controlled. Other safety messages will be given
later in this procedure.
Note: Sprinkled throughout the procedure are numerous questions, labeled Q-1, Q-2,
etc. Answer these questions in the appropriate areas of the Report Sheet. However, be
sure that you have collected all of your data before you begin to answer these questions.
Do not repeat the question, but please do use complete sentences.
I.
A.
Reactivity with water
Three alkali metals: lithium, sodium and potassium.
Prepare three beakers. Put about 100 mL of deionized water into each. Add 4 or 5
drops of phenolphthalein solution, an acid-base indicator. Then cover each beaker
with a wire screen from your locker.
Notice that these metals are stored under oil to prevent exposure to air and moisture.
Even so, your instructor will probably have to cut off an oxide crust to get a sample
of the soft, shiny metal. Bring your instructor 3 labeled watch glasses. You will be
given a piece the size of a small bean. Return to your desk and, with both partners
watching, drop each piece into a different beaker and immediately replace the wire
screen. (A screen is used instead of a watch glass because with the screen there is
less likelihood of accumulating an explosive mixture of gases.)
Q-1. Compare the relative reactivity of these three metals. What would you
predict for the relative reactivity of rubidium? In other words, rank these
four metals.
Q-2. What is the significance of the color change?
Q-3. Note that the spherical appearance of the sodium suggests that it is now a
liquid. Look up the melting point of sodium. Suggest why the sodium has
melted.
Q-4. Why does the sphere not sink, at least part way, into the water? Give two
reasons.
B.
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Sodium and methyl alcohol
There is no water present here, but methyl alcohol can be thought of as derived
from water. Caution: methyl alcohol is flammable; keep flames away.
Obtain a small piece of sodium metal from your instructor. Drop it into a small
beaker containing about 20 mL of methyl alcohol.
Q-5. How does this reaction compare with that of Na and H2O?
Q-6. Why does the sodium not "float" this time? Two reasons.
C.
Two alkaline earth metals: magnesium and calcium
Place about 150 mL of water into each of two 250 mL beakers. Into one beaker put
a short (1 inch) strip of magnesium ribbon. Observe the behavior of the magnesium
periodically while you work with the calcium.
Prepare the second beaker for the collection of a gas as in Figure 1. Fill a large test
tube with water. Put your thumb over the end, and invert the tube into the beaker so
that little if any air is trapped at the top of the tube. (Repeat the operation until you
succeed.) Now dry your hands, take a dry watch glass to the bottle marked "Fresh
Calcium metal," and collect from 1 to 5 granules as directed. Follow any other
instructions provided by the laboratory technician on the card in the gray bin.
Figure 1
Return to the beaker. Raise the test tube only enough to trap the pieces of calcium
under the mouth of the tube. Try to collect most of the evolved gas. It may be
necessary to repeat the operation in order to obtain a full test-tube of the gas.
When the tube is full of gas, remove it, keeping it in a vertical position, mouth
down. Then, some distance from the beaker, have your partner bring a lighted
splint to the mouth of the test tube.
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When the reaction in the beaker has ceased, note the milky appearance of the
liquid. Add a few drops of phenolphthalein indicator solution and note any color
change. Now add some dilute hydrochloric acid dropwise to decolorize the
indicator. What happens to the milkiness at the same time?
Dispose of the contents of the beaker in the proper waste container. (Save the
beaker with the magnesium in it for IIA below.)
Q-7. A "barking" sound when the gas explodes is characteristic of hydrogen.
What would happen to the splint if the gas had been (a) oxygen? (b)
carbon dioxide?
Q-8. What would have happened if you had carried the tube to the flame with the
mouth up?
II.
A.
Magnesium
Place about 5 mL of 3 M HCl in a medium sized test tube. Set this test tube upright
in one of the holes in your test tube rack. The idea is to be able to invert a large test
tube over the top of the medium test tube.
Add the magnesium from the previous experiment to the smaller test tube and
immediately invert a large test tube over it. Collect the gas evolved until the
magnesium dissolves. Carefully lift up the top test tube, keeping it vertical. Bring a
match or a burning splint to the mouth of the test tube. Identify the gas.
test tube rack
M g and acid
Figure 2
Dispose of the contents of the test tube in the proper waste container.
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Q-9. Which is easier for a metal to replace: hydrogen from water or hydrogen
from an acid? What do you think would happen if you added sodium to an
acid?
B.
Aluminum
Place a small strip of aluminum in a large test tube. Add about 4 mL of 6 M HCl.
Set it aside. If there is no reaction in five minutes, set the tube in a beaker of hot
water. Is any gas evolved?
III.
Study of oxides: basic and acidic
Oxides of elements of the third period (plus carbon) will be taken to illustrate
"horizontal" trends in the periodic table. The source of the oxides will be as follows: A.
sodium: sodium peroxide, a commercial product, will be substituted, because the oxide,
difficult to make, converts to the peroxide; B. magnesium: a piece of Mg ribbon will be
burned in air; C. aluminum: the oxide presents problems because of its insolubility, so we
will study the hydroxide which is closely related; D. carbon: a piece of charcoal will be
ignited and then thrust into pure oxygen; E. phosphorus: a small sample of red
phosphorus will be ignited in an iron spoon and thrust into pure oxygen; F. sulfur: same
procedure; G. chlorine: the oxides are dangerously explosive, so the safer reaction
product of oxide with water will be substituted.
A.
Sodium
Drop about 0.2 g of sodium peroxide, Na2O2, into a large test tube containing bout
10 mL of deionized water. Set it to one side on your desk while oxygen evolves.
Then test the solution with litmus paper. (In making tests with litmus paper, use a
stirring rod to bring a drop of the solution to the paper. Do not dip the paper into the
solution.)
Dispose of the solution in the proper waste container.
B.
Magnesium
Obtain a short piece of Mg ribbon. Put a little deionized water (about 1 mL or less)
in an evaporating dish. Using your tongs, hold the Mg in a gas flame long enough to
ignite it, and then avert or shield your eyes. Let the white "ash" fall from the tongs
into the water. Stir the mixture of solid and water from time to time, and test with
red litmus until you get a color change. (The solid does not dissolve very well, and
vigorous stirring is needed to get even a slight color change.) Do your best with this
one—often there is no discernible color change. If you get no color change, place a
small sample of the oxide of magnesium in an evaporating dish with a small amount
of water and test that as above.
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Before disposing of the mixture, pour about half of it into a small beaker. To one
half add a small volume of concentrated (12 M) hydrochloric acid with stirring. To
the other half add a small volume of 50% (12.5 M) sodium hydroxide with stirring.
Report any evidence of reaction (change!) in either or both, and then dispose in the
proper waste container.
Q-10. If you were to collect the "ash" from the burning of the magnesium
(including the part that went up in smoke), would it weigh less than, more
than, or the same as the original magnesium? Explain.
C.
Aluminum
Aluminum oxide is easy to obtain from the hydroxide (by drying to remove water),
but the reverse direction requires extreme conditions we choose not to employ here.
The following exercise will convey the relevant chemistry of aluminum hydroxide.
In a large test tube place about 5 mL of 1.0 M Al(NO3)3. Start adding 6 M NaOH
dropwise. The precipitate may be almost transparent at first. Later it becomes
slushy-like. When it becomes so thick that the contents are a stiff gel, stop the
addition. (Note: Adding too much NaOH will cause the precipitate to dissolve. If
you do not get a precipitate within 20 to 30 drops, discard your mixture in the waste
bottle and try again.) Divide the residue by removing about half to another tube.
(Carry it out on a stirring rod if necessary.) Add a few drops of phenolphthalein
indicator to one tube. Shake the tube regularly as you continue to add NaOH
dropwise. Stop when the solution becomes clear pink (that is, add NaOH until the
precipitate dissolves and the indicator changes from colorless to pink). Set the tube
aside.
To the other tube containing the gel, add 6 M HCl dropwise until the solution
becomes clear and colorless.
Aluminum hydroxide is one of a number of water-insoluble hydroxides that is said
to be amphoteric, that is, it will dissolve not only in a strongly acid solution but also
in a strongly basic one.
D.
Carbon
Fill three bottles with water. Then fill these bottles with oxygen as directed by the
instructions posted near the oxygen cylinder. When the bottle is full, turn off the gas
supply (at the needle valve) and slide a glass plate over the mouth. For each of these
bottles quickly remove the glass plate, immediately add about 10 mL of deionized
water, and then recover the bottle with the plate right
Grasp one end of a wooden splint in your tongs, and hold the other end in a gas
flame until it starts to glow. Then lift the glass plate from the first bottle, and thrust
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the burning splint into the gas. Use the glass plate to confine gases as well as you
can. When burning slows down, drop the ember into the water, and seal the bottle
with the plate. Shake the bottle for a minute or so to help absorption of any gaseous
products. Test the water with blue litmus paper. (Only a slight color change is
expected.)
E.
Phosphorus
Take the two remaining bottles of oxygen—with water in the bottom—to the hood
(one for burning phosphorus, the other for sulfur). Obtain a long-handled iron
spoon, sometimes called a combustion or "deflagrating" spoon.
Use the spatula to pick up a very small amount of red phosphorus—about the size
of a very small pea—and transfer it to the spoon.
CAUTION: Although red phosphorus is not as dangerous
as the other allotrope, white phosphorus, avoid contact with
your skin, and wash your hands after handling it.
Hold the spoon in a gas flame until the phosphorus starts to burn. Then insert it into
the bottle of oxygen, using the glass plate to block some of the opening. When the
combustion dies down, remove the spoon and close the bottle with the glass plate.
Continue to heat the spoon in the flame. You must clean the spoon by burning off
excess phosphorus. If you add too much phosphorus, it will take a long time to burn
off the excess. Shake the bottle to absorb gases in the water.
Test the water solution with both colors of litmus paper. Then discard the solution
in the proper waste container.
F.
Sulfur
When the spoon is clean and cool, repeat the procedure of part E above, but
substitute sulfur for phosphorus. The same precautions for size of sample apply.
Note the choking odor of burning sulfur, due to sulfur dioxide. The solution may be
disposed of in the proper waste container.
G.
Chlorine
As was mentioned earlier, the various oxides of chlorine are dangerously explosive.
A water solution in which one of these oxides has reacted with the water has been
made available to you. It has been labeled two ways: as an acid "HOClO3" and as a
base "ClO3(OH)."
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Test the solution with both red and blue litmus paper, and decide which formula is
appropriate. The solution may be discarded in the proper waste container.
IV.
Redox behavior of the halogens
SAFETY FIRST
• All of the free halogens are toxic and irritating. Small amounts of bromine and
iodine will be generated in parts A and B respectively. Do not take more than the
specified amount of starting materials. Work under a hood. Avoid breathing the
vapors.
• Concentrated sulfuric acid is very corrosive to the skin and destructive to clothing.
Wash hands immediately after contact. Remove affected clothing, and rinse with
water.
• Hexane, a solvent something like gasoline, is extremely flammable. Keep it away
from flames and heaters.
A.
Oxidation of the bromide ion
Into a small test tube place about 0.1 g of solid KBr. (Weighing is unnecessary. An
amount the size of a pea on the tip of your spatula is about right.)
Hold the tube with your wire test tube holder. (Hold it near the bottom, rather than
at the top as you usually do.)
Drop four drops of concentrated H2SO4 straight on to the solid in the tube. Observe
evidence of reaction. Promote further reaction by rubbing the bottom of the tube on
a warm hot plate. After brown fumes have filled the entire tube, stop heating the
tube.
Dispose of the chemicals in the proper waste container.
B.
Oxidation of the iodide ion
Repeat the procedure used in part A, but substitute 0.1 g of KI for the KBr. All
other operations will be the same, except that less heating may be required. Observe
the color of the fumes.
Dispose of the chemicals in the proper waste container.
C.
Establishing the "pecking order" among the halogens
All of the halogens are potential oxidizing agents, but not all are equally powerful.
All of the halide ions can be oxidized to the free halogen, but not all with the same
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ease. From the nine test tube experiments that follow, you should be able to rank the
halogens and relate their ranking—or "pecking order"—to the periodic table.
Label three large test tubes “chlorine water," "bromine water," and "iodine water,"
and obtain about 6 mL of each of these reagents. (Note: Chlorine water is not
Cl2·H2O, that is, it does not have a formula. It is a mixture in the same way as salt
water or seawater.) Label another set of three large test tubes "0.1 M NaCl," "0.1 M
KBr," and "0.1 M KI," and obtain about 5 mL of each of these reagents. In a small
beaker obtain about 10 mL of hexane.
Now, back at your desk, place about 1 mL each of chlorine water, bromine water,
and iodine water into three separate medium test tubes. (Each tube will contain only
one of the above solutions.) Then add 1 mL of hexane to each medium test tube
(These are test tubes 1, 2, and 3 below in the table). Stopper each of the test tubes,
and shake them well. Hexane is volatile; loosen the stopper under a fume hood to
prevent vapor build up inside the tube. Record the color of the upper hexane layer
on the chart in the data section. This is the color of the ‘free’ halogen in a nonpolar
solvent.
The remaining six tests will be mixtures of a halogen (Cl2, Br2, I2) with a halide salt
(NaCl, KBr, KI). Mix about 2 mL of halogen with 2 mL of a salt solution according
to the following mixing chart (test tubes 4 through 9 below). In each case, add 1 mL
of hexane, stopper, and shake well. Loosen the stopper under a fume hood. In each
case record the color of the hexane layer.
Cl2 (aq)
Br2 (aq)
I2 (aq)
Solvent only
NaCl(aq) +
solvent
KBr(aq) +
solvent
KI(aq) +
solvent
Test tube 1:
Hexane and Cl2
only
Test tube 2:
Hexane and Br2
only
Test tube 3:
Hexane and I2
only
Test tube 4: Cl2 Test tube 5:
+ KBr(aq) +
Cl2 + KI(aq) +
hexane
hexane
Test tube 6:
Test tube 7:
Br2 + NaCl(aq)
Br2 + KI(aq) +
+ hexane
hexane
Test tube 8:
Test tube 9:
I2 + NaCl(aq) + I2 + KBr(aq) +
hexane
hexane
Dispose of the contents in the "Halogen and Solvent Waste" container.
ANHYDRIDES
The word anhydride means without water. Anhydrides are binary compounds that do not
contain hydrogen. To obtain the formula of an anhydride, subtract one or more molecules
of water from a formula. Thus, the anhydride of carbonic acid = H2CO3 – H2O = CO2.
If the original compound contains an odd number of hydrogen atoms, it will be necessary
to double its formula first. Also, enough water molecules must be removed so that no
hydrogen remains. Thus the anhydride of phosphoric acid is P2O5.
Page 132
Las Positas College
Chemistry 1A
Page 133
NAME
Dates
Report Sheet: Exp. 15: PERIODIC PROPERTIES
NOTE: All chemical equations must be balanced and must include physical states or
phases.
Name of partner:
I.
Reactivity with Water
A.
Lithium, Sodium and Potassium
Describe the reaction of each metal with water. Include specific
identification of the products formed.
Write the balanced chemical equation for each reaction:
Q-1
Q-2
Q-3
Q-4
Rank:
Page 134
B.
Sodium and Methyl Alcohol
Describe the reaction of sodium with methanol:
Q-5
Q-6
C.
Mg and Ca
Observations:
Write the balanced chemical equation for reaction of Ca and water:
Write the balanced chemical equation for reaction of HCl with the white
precipitate:
Q-7
Q-8
Judging by the reactivity of calcium and of magnesium with water, what would you
predict as to the behavior of strontium or barium with water?
________________________________________________________________________
________________________________________________________________________
II.
Page 135
A. Magnesium
Observations:
Write the balanced equation for the reaction:
Q-9
B.
Aluminum
Observations:
Write the balanced equation for the reaction:
Summarize the comparative reactivity of the metals of Group I, of Group II, and of
Group III with water and with dilute acids. What would you predict as to the relative
ease of oxidation of these metals on exposure to air?
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
Page 136
III.
Study of oxides
A.
Sodium Peroxide with Water
Observations:
Equation:
Litmus test:
B.
Magnesium
Observations:
Equation for burning of Mg:
Litmus test, mixture of ash and water:
Equation for reaction of magnesium oxide and water:
Observations on effect of adding HCl:
Equation for reaction of HCl and magnesium hydroxide:
Q-10.
C. Aluminum
Observations:
Al(NO3)3 and NaOH:
Al(OH)3 and HCl:
Al(OH)3 and NaOH:
Equation for formation of the hydroxide from the nitrate:
Equation for dissolving of the hydroxide in HCl:
Equation for dissolving of the hydroxide in NaOH:
Al(OH)3(s) + NaOH(aq)  Na[Al(OH)4](aq)
D.
Carbon
Observations:
Equation for the burning of splint (carbon):
Litmus test, water solution of combustion product:
Equation for reaction between carbon dioxide and water:
E.
Phosphorus
Observations of combustion:
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Page 138
Equation:

P4O10(s)
What is the proper systematic name for P4O10(s)?
Result of litmus test, aqueous solution:
Equation to explain this reaction:
F.
Sulfur
Observations of combustion:
Equation:
Result of litmus test, aqueous solution:
Equation to explain this reaction:
G.
Chlorine
Result of litmus test, water solution of "HOClO3" or "ClO3(OH)":
Based on the above observation, which formula is correct? _____________
What is the formula of the oxide that is the anhydride of this acid? (See
discussion at the end of the experiment for a description of anhydrides.)
General conclusion concerning the acidity or basicity of the hydroxides (oxides) of the
elements of the third period:
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
IV.
Page 139
Redox behavior of the halogens
A.
Oxidation of the bromide ion
Observations:
Complete the following set of ionic equations:
a) oxidation half-reaction:
b) reduction half-reaction:
HSO4–(aq)
 SO2(aq)
c) overall net ionic equation:
B.
Oxidation of the iodide ion
Observations:
Complete the following set of ionic equations:
a) oxidation half-reaction:
b) reduction half-reaction:
c) overall net ionic equation:
HSO4–(aq)
 SO2(aq)
Page 140
C.
"Pecking order" of the halogens: Enter the color of the hexane layer in each
box below. In addition, enter "REACTION" or "NO REACTION" in each
case. HINT: Is the color of the hexane layer that of the reactant halogen
(Cl2, Br2 or I2) or is it the color of the product halogen?
solvent only
NaCl(aq) +
solvent
KBr(aq) +
solvent
KI(aq) +
solvent
Cl2(aq)
Br2(aq)
I2(aq)
For every combination in which there was a reaction, write the net ionic equation:
____________________________________________________
____________________________________________________
____________________________________________________
On the basis of these results, circle the best
•
oxidizing agent: Cl2 Br2 I2
Cl– Br–
I–
•
reducing agent: Cl2 Br2 I2
Cl– Br–
I–
Rank each of the 6 species in this activity series:
oxidizing agents:
best
worst
worst
best
reducing agents:
Page 141
POST-LAB
Name_________________
Experiment 15
1.
Most but not all acids contain oxygen.
Name one that does not:
2.
For an acid that does contain oxygen, we can derive the formula of its anhydride.
Write the formula of the anhydride of
H2SO3
HNO2
H3PO3
HOCl
3.
Are the above oxides acidic, basic, or amphoteric?
4.
Write the formulas of the anhydrides of the following bases:
KOH ________
Ba(OH)2 ________
Al(OH)3 ________
AgOH _______
5.
Are the above oxides acidic, basic, or amphoteric?
6.
In what region of the periodic table would you expect to find elements that form
amphoteric oxides and hydroxides?
____________________________________
7.
Utilizing the general trends you have observed in this experiment, predict whether
the following reactions would occur. If a reaction is predicted to occur, write a
balanced conventional equation for the reaction. If no reaction is predicted to
occur, write the words "no reaction".
a. francium metal is
added to water.
b. strontium metal is
added to water.
c.
rubidium metal and
Page 142
chlorine gas are mixed
d. fluorine gas is added
to aqueous calcium
chloride.
e. aqueous cesium bromide
and aqueous iodine are
mixed.
Page 143
PRE-LAB
Name_________________
Experiment 15
1. Define these terms as used in chemistry:
Periodic
___________________________________________________________
Slurry
___________________________________________________________
Allotrope
___________________________________________________________
2. Give the symbol of the element that is:
a) in the second period and in Group IIIA (Group 13): ______________
b) in the third period and in Group IIA (Group 2):
3.
Give the
formula of the
oxide of this
element
of the
hydroxide
Na
4.
What two acid-base indicators are used in
this experiment?
1)
Mg
______________
Al
What is the color of each in...
Acidic solution
Basic solution
2)
5. Three members of the halogen family are studied in this experiment; give their
molecular formulas
physical state at room
temperature
color
Si
Page 144
Three halide ions are also studied in the form of salts; give their formulas:
__________
__________
__________
6. The oxide left by the removal of water from an acid or base is said to be the
anhydride of that acid or base. The process is reversible. For instance:
H2SO4 giving up H2O results in SO3 and SO3 + H2O = H2SO4
2 KOH giving up H2O results in K2O and K2O + H2O = 2 KOH
a. What is the anhydride of nitric acid, HNO3? ___________________
Remember, there can be no hydrogen in an oxide. (Hint: Start with 2 molecules of
acid.)
b. What is the anhydride of barium hydroxide, Ba(OH)2?_________________
Page 145
Experiment 16
Model Making & Geometry
PURPOSE
The purpose of this experiment is to gain experience in
a. predicting the shape, drawing stereochemical formulas, and building models of
molecules and ions;
b. predicting hybridization, numbers of sigma and pi bonds, and drawing threedimensional orbital overlap sketches (the latter for selected molecules only).
Method
Lewis structures, in conjunction with VSEPR theory and/or valence bond theory, are used
to predict molecular geometries. Molecular models are used as an aid in visualizing and
drawing the shapes/stereochemical formulas for molecules and ions. Hybridization is
determined from the Lewis structures.
PROCEDURE
For each molecule or ion listed below:
1.
2.
3.
4.
5.
6.
7.
8.
9.
Calculate AE, defined as available electrons or the total number of valence
electrons in the entire molecule or ion, adjusted for any charge in the case of ions.
Draw the complete Lewis structure, including all resonance forms. Satisfy the octet
rule if possible.
Count the number of bonded atoms (BA) and lone pairs (LP) for the central atom.
Add them to get the coordination number of the central atom. The coordination
number is used to determine the electron pair geometry (2 = linear, 3 = trigonal
planar, 4 = tetrahedral, 5 = trigonal bipyramidal, 6 = octahedral).
Build a model by selecting an atom piece that has the correct number of holes to
correspond with the coordination number. Use sticks with no atoms at the ends to
represent lone pairs of electrons. Note that it is only necessary to build one model
for any species that exhibits resonance. Use an atom piece with an extra hole if
there is a double bond.
Draw a stereochemical formula for the model. Include approximate bond angles on
the drawing. Use element symbols in the drawing and do not show lone pairs to
outside atoms.
State the electron pair geometry and the molecular shape.
Have your model checked by your instructor.
Continue in the second part to identify the hybridization of the central atom, the
number of sigma bonds and pi bonds in each species, and the formal charges on
atoms.
In the third part, you will use formal charge to select between possible structures
and you will draw three-dimensional orbital overlap sketches for a couple of
molecules.
Page 146
AE
Lewis structure
NH3
NH4+
CO2
SO2
BA
LP
Coord #
Stereochemical Drawing
Instructor approval
Electron Pair Geometry/
Molecular Geometry
AE
Lewis structure
SO3
POF3
PCl5
BA
LP
Coord #
Instructor approval
Page 147
Molecular Geometry
Page 148
AE
Lewis structure
SF6
ClF3
IF5
BA
LP
Coord #
Instructor approval
Molecular Geometry
AE
Lewis structure
BiCl52–
XeF2
XeF4
BA
LP
Coord #
Instructor approval
Page 149
Molecular Geometry
Page 150
AE
Lewis structure
TeF4
COCl2
HCN
BA
LP
Coord #
Instructor approval
Molecular Geometry
Page 151
Part II: Complete the following table. Redraw the Lewis structure for each molecule or
ion, calculate formal charges, and label any non-zero formal charge.
species
NH3
NH4+
CO2
SO2
SO3
POF3
PCl5
Lewis structure; label any non-zero
formal charges on each structure
hybridization of the central
atom
#
#
Page 152
species
SF6
ClF3
IF5
BiCl52–
XeF2
XeF4
TeF4
atom
#
#
species
COCl2
HCN
atom
Page 153
#
#
Page 154
Part III
1.
Use formal charge to choose the best arrangement for the atoms in a molecule
whose formula is usually written N2O: N—N—O or N—O—N. Include resonance
structures in your evaluation. All molecules should satisfy the octet rule for all
atoms.
2.
The following molecules do not have just one central atom. Draw Lewis diagrams,
build the models, and then draw stereochemical formulas.
C2H4Cl2 1,1-isomer: both Cl’s on the same C
C2H4Cl2 1,2-isomer: one Cl on each C
C2H2Cl2 1,1-isomer
C2H2Cl2 1,2 cis isomer: use index of textbook to find description of cis and trans
isomers
3.
Draw one or two pictures showing the overlap of orbitals to form sigma and pi
bonds in COCl2.
Page 155
POST-LAB
Name_________________
Experiment 16
1.
Model Making & Geometry
In the molecule drawn below, how many sigma bonds are there?
how many pi bonds?
what is the hybridization of atom A?
what is the hybridization of atom B?
H H H H H H
H
O C C C C C C C C C N
Cl
2.
H
B
C H
H H H
A
Draw the Lewis structure for KrCl4:
How many electrons?
What is the shape of the molecule?
What hybrid orbitals are used by Cl?
Is the molecule polar?
3.
Assign formal charges to all the atoms in the structure for sulfuric acid. What is the
oxidation number of the sulfur?
O
H
O S
O
O H
Page 156
Page 157
Experiment 17
PURPOSE
The purpose of this experiment is to construct models of the common metallic and ionic
crystal lattices and to study their characteristics.
Method
Historically, Styrofoam balls were connected with wooden toothpicks. Now we will use
the Solid State Model Kits to reproduce the following crystal lattices:
IONIC
METALLIC
A. sodium chloride (rock salt)
D. face centered cubic
(Cubic close packed)
E. body centered hexagonal
(Hexagonal close packed)
F. body-centered cubic
B. cesium chloride
C. wurtzite (Hexagonal ZnS)
By looking at these models, you will be able to identify many features of these lattices.
Figure 1 shows the unit cells for the metallic crystal lattices and Figure 2 shows the unit
cells for the ionic crystal lattices.
PROCEDURE
You will use the Solid State Model Kit and the Instructions in the kit to answer the
following questions. Refer to the diagrams both here and in the Model Kit for assistance.
1. Find the instruction manual in your Model Kit. Read “Getting Started” along with the
rest of the information on pages 1-5.
2. Do the NaCl example on pages 6-7 using the Model Kit. This example takes you
through the questions step by step, while all other examples are abbreviated. You won’t
have any problems IF you have followed instructions 1 and 2!
3. Answer all of the questions for the NaCl example in part A below, then proceed with
the other parts. Answer all of the questions while you have the model in front of you for
reference.
Page 158
Introduction
The following are the simplest crystal lattice structures found in nature.
a. face-centered cubic
(cubic close-packed)
b. body-centered hexagonal
(hexagonal close-packed)
c. body-centered cubic
Figure 1: Common metallic crystal lattice unit cells
________________________________________________________________
Black = - ion
White = + ion
a.
CsCl (simple cubic - ions)
cesium chloride
b.
NaCl (face-centered cubic – ions)
sodium chloride (rock salt)
c. ZnS (face-centered cubic – ions) d. ZnS (body-centered hexagonal – ions)
cubic zinc sulfide (zinc blende)
hexagonal zinc sulfide (wurtzite)
Figure 2. Common ionic crystal lattice unit cells for 1:1 salts
Page 159
CRYSTAL LATTICES
Construct each model using the Instructions in your Solid State Model Kit, answering
questions as you go. As you progress, have your work checked by the instructor. Be
prepared to answer questions about the information you have gathered.
IONIC LATTICES
A.
Sodium Chloride (Rock Salt), Octahedral Sites: Model A
Follow instructions on pages 6-7 of your Model Kit Instruction Manual.
Figure 3. The Na+ ion in the
octahedral hole of the NaCl lattice.
1. What is the CN of X? (That is, how many Cl- ions are around each Na+ ion?)
This is known as the coordination number (CN) or number of nearest neighbors.
_____
Figure 4. Sodium Chloride (Rock Salt).
Smaller Na+ spheres in the octahedral
holes of a face-centered cubic lattice of
Cl- spheres.
Instructor's Initials
2.
3.
Within each unit cell, how many net Cl- ions are:
a.
on the corners (net number)
_____
b,
on the edges (do not count corners again)
_____
c.
within the faces
_____
d.
within the center
_____
e.
net total
_____
Within each unit cell, how many net Na+ ions are:
a.
on the corners (net number)
_____
b.
on the edges (do not count corners again)
_____
c.
within the faces
_____
Page 160
d.
within the center
_____
e.
net total
_____
4.
The chloride ions alone describe what lattice?
_______________________
5.
The sodium ions alone describe what lattice?
_______________________
6.
Calculate the radius of the sphere that will touch all
six surrounding spheres. Hint: tilt the model until
you see the plane of 4 spheres shown to the right.
Small sphere of
radius r
7.
Calculate the radius ratio R/r (or r/R).
______
B.
Cesium Chloride. Cubic Sites: Model B
Follow the instructions beginning on page 11 of the instruction manual.
Figure 5. The Cs+ ion in the cubic hole of the CsCl
lattice
1. What is the CN of X? (That is, how many Cl- ions are around each Cs+ ion?) _____
Now complete the CsCl lattice.
Figure 6. Model B.
Cesium Chloride,
multiple unit cells
_______
2.
How many complete unit cells of CsCl are contained in this model?
3.
Within each unit cell, how many net Cl- ions are:
a.
on the corners
_____
_____
4.
Page 161
b.
on the edges
_____
c.
within the faces
_____
d.
within the center
_____
e.
net total
_____
Within each unit cell, how many net Cs+ ions are:
a.
on the corners
_____
b.
on the edges
_____
c.
within the faces
_____
d.
within the center
_____
e.
net total
_____
5.
The chloride ions alone describe what lattice?
_______________________
6.
The cesium ions alone describe what lattice?
_______________________
C.
Hexagonal Zinc Sulfide (Wurtzite), Tetrahedral Sites: Model C
Follow the instructions on page 54 of the Solid State Kit Instruction Manual.
Figure 7. The Zn2+ ion in the
tetrahedral hole of the hexagonal
ZnS lattice.
1. What is the CN of X? (That is, how many S2- ions are around each Zn2+ ion?)
2. How many complete unit cells of ZnS are contained in this model?
______
______
______
3. Within each unit cell, how net many S2- ions are:
a.
on the corners
_____
b.
on the edges
_____
Page 162
4.
b.
with in the faces
_____
c.
within the center
_____
d.
net total
_____
Within each unit cell, how many net Zn2+ ions are:
a.
on the corners
_____
b.
on the edges
_____
c.
within the faces
_____
d.
within the center
_____
e.
net total
_____
METALLIC LATTICES
D.
Face-Centered Cubic (FCC) or Cubic Close-Packed: Model D
Follow the instructions on page 27 of the Solid State Model Kit Instruction
Manual.
Figure 8. Model A.
Face-Centered
Cubic.
Instructor's Initials___________
1. How many spheres are directly touching the sphere marked with an "X" above?_____
2.
Within the unit cell, how many (net number) spheres are:
a.
on the corners (net contribution)
net no.= __________
b.
within the faces (do not count the corners again)
net no. = __________
c.
in the center
net no. = __________
d.
total all number of spheres in the unit cell
____________
3. FCC is one of the two close-packed lattices. In one of these, every other layer is the
same, called ABAB packing. This means that for every sphere in the 1st layer,-there is
Page 163
a sphere right above it in the 3rd layer. In the other type of close packing, spheres
repeat every third layer, called ABCABC packing. In this arrangement, each sphere in
the 1st layer does not have a sphere right above it until the 4th layer.
Is the packing in FCC,
ABAB
or
ABCABC
?
(Hint: the four layers you have made are NOT the close packed layers. In order to see the
packing type, use the cube made by the first three
4. What are the relationships among the sides (a, b, c) in the
FCC unit cell (equal or unequal)
_____
5.
_____
What are the angles () in the FCC unit cell?
Use the diagram of the face-centered cubic lattice at the right to
answer the following questions.
6. How many net spheres are in the unit cell?
______
7. What is the volume of a sphere in terms of R, the sphere radius?
______
8. What is the total volume occupied by the spheres in the unit cell in terms of R?
9. What is the face diagonal (corner-to-corner distance on face) in terms of R?
10. What is the value of S in terms of R? Use the Pythagorean Theorem on the right
triangle given. Show work below.
Page 164
11.
Give the volume of the cube in terms of S.
12.
Give the volume of the cube in terms of R.
13.
Calculate the percent of occupied space.
% occupied = (Vspheres/Vcube) x 100%
E. Bodv-Centered Hexagonal or Hexagonal Close-Packed (HCP): Model E
Follow the directions on page 24 of the Solid State Model Kit Instruction Manual.
Figure 9. Model E.
Body-Centered Hexagonal.
1.
2.
What is the coordination number (number of nearest
neighbors) in the hexagonal close packed lattice?
______
______
a.
net no. =
_____
b.
within the faces
net no. =
_____
c.
in the center (inside the cell)
net no. =
_____
d.
total net number of spheres in unit cell
_____
3.
Is the packing in HCP
ABAB_________
Page 165
ABABC_________?
(Hint: the four layers you have made ARE the close packed layers in the
hexagonal close-packed lattice.)
4.
What are the relationships among the sides (a, b, c)
in the HCP unit cell (answer: equal or unequal)?
_____
5.
What are the angles () in the HCP unit cell?
_____
F.
Body-Centered Cubic (BCC): Model F
Figure 10. Model F.
Body-Centered Cubic.
1.
2.
What is the coordination number (number of nearest
neighbors) in the body-centered cubic lattice?
_____
_____
a.
net no =
_____
b.
within the faces (do not count the corners again)
net no =
_____
c.
within the center
net no =
_____
d.
total net number of spheres in unit cell
_____
3.
The body-centered spheres in the neighboring unit cells are only 14% farther
away than those touching the body centered sphere in this cell. Such spheres are
called next nearest neighbors. How many next nearest neighbors are in the BCC
lattice?
_____
4.
What are the relationships among the sides (a, b, c)
in the BCC unit cell (equal or unequal)?
_____
What are the angles () in the BCC unit cell?
_____
5.
Page 166
Use the two diagrams of the body-centered cubic lattice below to-answer the following
questions.
6.
How many net spheres are in the unit cell?
_____
7.
What is the volume of all the spheres in the unit cell, in terms of R?
_____
8.
What is the body diagonal (BD) of the cube in terms of R? Note: the central
sphere touches each corner sphere.
_____
9.
What is the face diagonal (FD) of the cube in terms of S? Use Pythagorean
theorem on the right triangle which has FD as the hypotenuse. Show your
work below.
_____
10.
What is the value of S in terms of R? Use the Pythagorean theorem
on the right triangle which has BD as the hypotenuse. Use your
values of BD in terms of R and FD in terms of S. Show your work below.
_____
11.
Give the volume of the cube in terms of R.
_____
12.
Calculate the percent of occupied space.
_____
13.
Which type of packing is more efficient?
______FCC or ______ BCC
Page 167
POST-LAB
Name_________________
Experiment 17
1. Cesium chloride structure
a.
b.
2.
Why is it not proper to say that CsCl has a body-centered cubic structure?
Thallium(I) chloride crystallizes in the cesium chloride lattice, as shown in
Fig. 5. The shortest distance between the center of a Tl+ ion and the center
of a Cl- ion is 333 pm.
1)
What is the length of the edge of a unit cell of TlCl? Hint: How
many TlCl diameters equal one body diagonal?
2)
What is the density in g/cm3 of TlCl?
Superconductors. Given the idealized unit cell for the structure of a
superconductor shown below to be orthorhombic, answer the following questions.
Page 168
a.
How many copper atoms are there per unit cell? (Show calculations.)
b.
What is the number of oxygens per unit cell?and Y? and Ba?
c.
What is the formula for this superconductor?
d.
The coppers in this unit cell occupy two different kinds of sites. Based on the
number of each of these types and the formula for the compound, what are the
oxidation states for copper in this compound?
The following Questions will likely require some research and reading outside of our
textbook! Go dig out the information that you need!
3.
a.
What makes a substance a superconductor?
b.
Give the properties of superconductors.
4.
Explain why ionic solids are brittle and metals are malleable.
5.
Explain why metals decrease in conductivity with a rise in temperature while
semiconductors increase in conductivity with temperature.
6.
How does the band theory explain luster of metals and heat conductivity? Include
labeled diagrams in your answer.
Page 169
PRE-LAB
Name_________________
Experiment 17
Show your work!
1.
Write the equation for the volume, V, of a cube in terms of its side, s:
V=
2.
Write the equation for the volume, V, of a sphere in terms of its radius, R:
V=
3.
Write the Pythagorean Theorem equation for a right triangle with sides a and b
adjacent to the right angle and hypotenuse c:
4.
If a right triangle has two equal sides of length a, express the length of the
hypotenuse c in terms of a:
c=
5.
For a cube with sides of length s:
a.
Express the length of a diagonal across a face, FD, in terms of s:
FD =
b.
Express the length of a diagonal through the body, BD, in terms of s:
BD =
6.
Simplify
√8:
√8 =
Page 170
Page 171
Experiment 18
In this experiment, Temperature Probes are placed in various liquids. Evaporation occurs
when the probe is removed from the liquid’s container. This evaporation is an
endothermic process that results in a temperature decrease. The magnitude of a
temperature decrease is, like viscosity and boiling temperature, related to the strength of
intermolecular forces of attraction. In this experiment, you will study temperature
changes caused by the evaporation of several liquids and relate the temperature changes
to the strength of intermolecular forces of attraction. You will use the results to predict,
and then measure, the temperature change for several other liquids.
You will encounter two types of organic compounds in this experiment—alkanes and
alcohols. The two alkanes are n-pentane, C5H12, and n-hexane, C6H14. In addition to
carbon and hydrogen atoms, alcohols also contain the -OH functional group. Methanol,
CH3OH, and ethanol, C2H5OH, are two of the alcohols that we will use in this
experiment. You will examine the molecular structure of alkanes and alcohols for the
presence and relative strength of two intermolecular forces—hydrogen bonding and
dispersion forces.
Figure 1
Materials
Windows PC
Vernier computer interface
Logger Pro
4 Temperature Probes
4 pieces of absorbent paper (2 cm X 4 cm)
4 small rubber bands
Colored time tape
methanol (methyl alcohol)
ethanol (ethyl alcohol)
acetone
n-pentane
1-propanol
1-butanol
n-hexane
Page 172
PROCEDURE
1. Obtain and wear goggles! CAUTION: The compounds used in this experiment are
flammable and poisonous. Avoid inhaling their vapors. Avoid contacting them with
your skin or clothing. Be sure there are no open flames in the lab during this
experiment. Notify your teacher immediately if an accident occurs.
2. Prepare the computer for data collection.
Obtain 4 stainless steel temperature probes and plug them into Channels 1-4. Prepare the
computer for data collection by opening Logger Pro. You should see 4 temperature
readings corresponding to each probe. On the Graph window, the vertical (temperature)
axis should be scaled from 0 to 30C and the horizontal (time) axis from 0 to 400
seconds. In the Experiment menu, choose Data Collection, and then set the experiment
length to 400 seconds. This is different from changing the display. You must also
double click on one of the values from the axis you want to change. Choose axis options
and set the range to the correct values.
3. Wrap Probes 1-4 with identical pieces of absorbent paper secured by small rubber
bands as shown in Figure 1. Roll the paper around the probe tip in the shape of a
cylinder. Hint: First slip the rubber band up on the probe, wrap the paper around the
probe, and then finally slip the rubber band over the wrapped paper. The paper should
be even with the probe end.
4. Place 4 large test tubes in you test tube rack and label them Methanol, Ethanol,
Acetone, and Pentane. Place enough of each pure liquid in the corresponding test
tube to fill it up 2-3 cm. Put Probe 1 in the methanol test tube, Probe 2 in the ethanol
test tube, Probe 3 in the acetone test tube, and Probe 4 in the pentane test tube. Make
sure the containers do not tip over.
5. Prepare 4 pieces of colored time tape, each about 10-cm long, to be used to tape the
probes in position during Step 6.
6. After the probes have been in the liquids for at least 45 seconds, begin data collection
by clicking Collect . Monitor the temperature for at least 15 seconds to establish the
initial temperature of each liquid. Then simultaneously remove the probes from the
liquids and tape them so the probe tips extend 5 cm over the edge of the table top as
shown in Figure 1.
7. When each of the temperatures have reached minimums and have begun to increase,
click Stop to end data collection. Click the Statistics button, , then click OK
to display a box for both probes. Record the maximum (t1) and minimum (t2) values
for Temperature 1 (methanol), Temperature 2 (1-propanol), Temperature 3 (acetone),
and Temperature 4 (pentane).
8. For each liquid, subtract the minimum temperature from the maximum temperature to
determine t, the temperature change during evaporation.
9. Roll the rubber band up the probe shaft and dispose of the absorbent paper in the
“Chemically Contaminated Paper” container.
10. Based on the t values you obtained for these 4 substances, plus information in the
Pre-Lab exercise, predict the size of the t value for 1-propanol and 1-butanol.
Compare their hydrogen-bonding capabilities and molecular weights to those of
methanol and ethanol. Record your predicted t, then explain how you arrived at this
Page 173
answer in the space provided. Do the same for n-hexane. It is not important that you
predict the exact t value; simply estimate a logical value that is higher, lower, or
between the previous t values.
11. Repeat the experiment using methanol, ethanol, 1-butanol and 1-propanol. The
methanol and ethanol samples will act as a control so that you can compare this
experiment with your first experiment.
12. Now repeat the experiment using pentane and hexane.
DATA
Trial #1
Substance
methanol
ethanol
acetone
n-pentane
t1 (°C)
t2(°C)
t (t1–t2) (°C)
Trial #2
Substance
methanol
ethanol
1-propanol
1-butanol
n-pentane
n-hexane
Predicted
t (t1–t2) (°C)
Explanation
t1
(°C)
t2
(°C)
Experimental
t (t1–t2) (°C)
Page 174
DATA ANALYSIS
Answer the following questions in your notebook.
1. Two of the liquids, 1-propanol and acetone, have similar molecular weights, but
significantly different t values. Explain the difference in t values of these
substances, based on their intermolecular forces.
2. Which of the alcohols studied has the strongest intermolecular forces of attraction?
The weakest intermolecular forces? Explain using the results of this experiment.
Page 175
PRE-LAB
Name_________________
Experiment 18
Prior to doing the experiment, complete the Pre-Lab table. The name and formula are
given for each compound. Draw a structural formula for a molecule of each compound.
Then determine the molecular weight of each of the molecules. Dispersion forces exist
between any two molecules, and generally increase as the molecular weight of the
molecule increases. Next, examine each molecule for the presence of hydrogen bonding.
Before hydrogen bonding can occur, a hydrogen atom must be bonded directly to an N,
O, or F atom within the molecule. Tell whether or not each molecule has hydrogenbonding capability.
Substance
Formula
methanol
CH3OH
ethanol
C2H5OH
1-propanol
C3H7OH
1-butanol
C4H9OH
acetone
C3H6O
n-pentane
C5H12
n-hexane
C6H14
Structural Formulas
Molecular
Weight
Hydrogen Bond
(Yes or No)
Page 176
Page 177
Experiment 19
Using Freezing-Point Depression to Find Molecular Weight
When a solute is dissolved in a solvent, the freezing temperature is lowered in proportion
to the number of moles of solute added. This property, known as freezing-point
depression, is a colligative property; that is, it depends on the ratio of solute and solvent
particles, not on the nature of the substance itself. The equation that shows this
relationship is:
T = Kf • m
where T is the freezing point depression, Kf is the freezing point depression constant for
a particular solvent (3.9°C-kg/mol for lauric acid in this experiment1 ), and m is the
molality of the solution (in mol solute/kg solvent).
In this experiment, you will first find the freezing temperature of the pure solvent, lauric
acid, CH3(CH2)10COOH. You will then add a known mass of benzoic acid solute,
C6H5COOH, to a known mass of lauric acid, and determine the lowering of the freezing
temperature of the solution. In an earlier experiment, you observed the effect on the
cooling behavior at the freezing point of adding a solute to a pure substance. By
measuring the freezing point depression, T, and the mass of benzoic acid, you can use
the formula above to find the molecular weight of the benzoic acid solute, in g/mol.
Figure 1: Note-- use a second probe to measure the water bath temperature.
PURPOSE
The purpose of this experiment is to:
a. study the effect of a solute on the freezing point of a solvent,
b. to learn a method of accurate freezing point determination, and
c. to use that method to determine the molar mass of an unknown
1
“The Computer-Based Laboratory,” Journal of Chemical Education: Software, 1988, Vol. 1A, No. 2, p.
73.
Page 178
Materials
Power Macintosh or Windows PC
Logger Pro
2 Temperature Probes
400-mL beaker
ring stand, ring clamp, 2 utility clamps
Buret clamp, wire gauze
18 X 150-mm (medium sized) test tube
lauric acid
benzoic acid
Safety: Benzoic acid is moderately toxic by ingestion; irritates eyes, skin and respiratory
tract. Use the fume hoods and avoid breathing vapors. There will be open flames in the
lab. Keep the cables away from the heat of the Bunsen burner as in figure 1. This is
important becaure the cables are quickly damaged by excessive heat. Be sure to dispose
of unknowns in marked containers.
PROCEDURE
1. Obtain and wear goggles.
2. Prepare the computer for data collection by opening the Experiment 15 folder from
Chemistry with Computers. Then open the experiment file that matches the probe you
are using. The vertical axis of the graph has temperature scaled from 0°C to 100°C.
The horizontal axis has time scaled from 0 to 10 minutes.
Part I Determine the Freezing Temperature of Pure Lauric Acid
3. Clean and dry a medium test tube. Weigh accurately (to +/- 0.001g) about 7 or 8
grams of lauric acid in a preweighed weighing boat. Transfer all of the lauric acid to
the test tube. You might chose to reweigh the boat with any solid dust that remains to
correct for that which does not get into the test tube. Label your test tube and place it
into one of the hot water baths on the back counter to melt the lauric acid.
4. Add about 270 mL of tap water with a temperature between 20 and 25°C to a 300-mL
Berzelius beaker (special from the stockroom). Assemble the apparatus in Figure 1,
making sure to add a second temperature probe in the water bath.
5. Insert the Temperature Probe into the hot lauric acid. About 30 seconds are required
for the probe to warm up to the temperature of its surroundings and give correct
temperature readings. During this time, fasten the utility clamp to the ring stand so the
test tube is above the water bath. Then click Collect to begin data collection.
CAUTION: Be careful not to spill the hot lauric acid on yourself and do not touch
the bottom of the test tube.
6. Lower the test tube into the water bath. Make sure the water level outside the test tube
is higher than the lauric acid level inside the test tube. If the lauric acid is not above
50°C, reheat the lauric acid sample and begin again.
7. With a very slight up and down motion of the Temperature Probe, continuously stir
the lauric acid during the cooling. Hold the top of the probe and not its wire.
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8. Continue with the experiment until data collection has stopped after 10 minutes. Do
not attempt to pull the probe out—this might damage it. You can reuse the lauric acid
sample for Part II if you have not lost any of the lauric acid.
9. To determine the freezing temperature of pure lauric acid, you need to determine the
mean (or average) temperature in the portion of graph with nearly constant
temperature. Move the mouse pointer to the beginning of the graph’s flat part. Press
the mouse button and hold it down as you drag across the flat part of the curve,
selecting only the points in the plateau. Click on the Statistics button, . The mean
temperature value for the selected data is listed in the statistics box on the graph.
Record this value as the freezing temperature of lauric acid. Click on the upper-right
corner of the statistics box to remove it from the graph.
Part II Freezing Temperature of a Solution of Benzoic Acid and Lauric Acid
10. Prepare the computer for data collection. From the Data menu, choose Store Latest
Run. This stores the data so it can be used later. To hide the curve of your first data
run, click the Temperature vertical-axis label of the graph, and uncheck the Run 1
box. Click OK .
11. Weigh approximately 1g of benzoic acid and add it to the test tube containing lauric
acid. Put this test tube back into the hot water bath to melt the solids and mix
thoroughly with the temperature probe. Repeat Steps 3-8 to determine the freezing
point of this mixture.
12. When you have completed Step 8, click on the Examine button, . To determine the
freezing point of the benzoic acid-lauric acid solution, you need to determine the
temperature at which the mixture initially started to freeze. Unlike pure lauric acid,
cooling a mixture of benzoic acid and lauric acid results in a gradual linear decrease
in temperature during the time period when freezing takes place.
Freezing Point
As you move the mouse cursor across the graph, the temperature
(y) and time (x) data points are displayed in the examine box on
the graph. Locate the initial freezing temperature of the solution,
as shown here. Record the freezing point in the Data and
Calculations table.
13. To print a graph of temperature vs. time showing both data runs:
Time
a. Click the Temperature vertical-axis label of the graph. To display both temperature
runs, check the Run 1 and Latest boxes. Click OK .
b. Label both curves by choosing Make Annotation from the Analyze menu, and
typing “Lauric acid” (or “Benzoic acid-lauric acid mixture”) in the edit box. Then
drag each box to a position on or near its respective curve.
c. Print a copy of the Graph window. Enter your name(s) and the number of copies of
the graph you want.
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DATA
Mass of lauric acid
g
Mass of benzoic acid
g
Freezing temperature of pure lauric acid
°C
Freezing point of the benzoic acid–lauric acid
mixture
°C
Freezing temperature depression, t
Molality, m
°C
mol/kg
Moles of benzoic acid
mol
Molecular weight of benzoic acid (experimental)
g/mol
Molecular weight of benzoic acid (accepted)
g/mol
Percent error
%
CALCULATIONS
1. Determine the difference in freezing temperatures, t, between the pure lauric acid
(t1) and the mixture of lauric acid and benzoic acid (t2). Use the formula, t = t1 - t2.
2. Calculate molality (m), in mol/kg, using the formula, t = Kf • m (Kf = 3.9°C-kg/mol
for lauric acid).
3. Calculate moles of benzoic acid solute, using the answer in Step 2 (in mol/kg) and the
mass (in kg) of lauric acid solvent.
4. Calculate the experimental molecular weight of benzoic acid, in g/mol. Use the
original mass of benzoic acid from the Data and Calculations table, and the moles of
benzoic acid you found in the previous step.
5. Determine the accepted molecular weight for benzoic acid from its formula,
C6H5COOH and calculate the percent error.
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Experiment 20
Acid Rain
PURPOSE
In this experiment, you will observe the formation of four acids that occur in acid rain:

carbonic acid, H2CO3
 nitrous acid, HNO2
 nitric acid, HNO3

sulfurous acid, H2SO3
Carbonic acid occurs when carbon dioxide gas dissolves in rain droplets of unpolluted
air:
(1) CO2(g) + H2O(l) 
 H2CO3(aq)
Nitrous acid and nitric acid result from a common air pollutant,
nitrogen dioxide (NO2). Most nitrogen dioxide in our atmosphere
is produced from automobile exhaust. Nitrogen dioxide gas
dissolves in rain drops and forms nitrous and nitric acid:
CO 2
(2) 2 NO2(g) + H2O(l) 
 HNO2(aq) + HNO3(aq)
Sulfurous acid is produced from another air pollutant, sulfur
dioxide (SO2). Most sulfur dioxide gas in the atmosphere results
from burning coal containing sulfur impurities. Sulfur dioxide
dissolves in rain drops and forms sulfurous acid:
NO 2
H 2 CO 3
H NO 2
H NO 3
H2SO3
SO2
(3) SO2(g) + H2O(l) 
 H2SO3(aq)
In the procedure outlined below, you will first produce these three gases. You will then
bubble the gases through water, producing the acids found in acid rain. The acidity of the
water will be monitored with a pH Sensor.
Materials
Windows PC
Logger Pro
Vernier pH Sensor
Wash bottle with distilled water
Rubber stopper, size 00
Beakers, 600, 400, 250, 150, mL
3 20 X 150 mm test tubes
Solid NaNO2
Solid NaHCO3
Solid NaHSO3
1.0 M HCl with 2 mL plastic dropper
3 2 mL plastic droppers
3 special glass sample vials
Tap water
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PROCEDURE
Safety First! Wear goggles. Avoid ingesting dust from solids and breathing gasses
produced. The glass bulb at the tip of the pH sensor is fragile, so please do not hit it.
1. Think safety continuously!
2. Label the three clean medium (20 x 150 mm) test tubes with the formula of the solid
they will contain: “NaHCO3”, “NaNO2”, and “NaHSO3”. Label the three special (for
this experiment) clean glass sample vials with the formula of the gas that will be
dissolved in the 6 mL water that they will contain: “CO2”, “NO2” and “SO2”. You
can use a 150-mL beaker to support the test tubes.
3. Using a clean metal scoop, place some solid NaHCO3.in the tube labeled “NaHCO3”.
Add enough NaHCO3 to fill the curved bottom end of the test tube.
4. Repeat the Step 3 procedure to add solid NaNO2 and NaHSO3 to the other two
appropriately labeled test tubes. CAUTION: Avoid inhaling dust from these solids.
5. Fill the clean 250 mL beaker about half full with clean tap water.
6. Carefully unscrew the little plastic bottle on the pH sensor and carefully place it
upright on the lab bench so that the storage solution does not spill. The glass bulb at
the end of the pH sensor is very fragile, so be careful not to hit it with anything.
Carefully slip the plastic bottle cap off from the cylindrical wall of the pH sensor.
Wash the pH sensor by squirting deionized water from a wash bottle onto the shaft
and around the bulb. A 600 mL beaker serves as a good waste container to collect the
wash water as it runs off the pH sensor. After washing, place the pH sensor in the
250 mL half full beaker of clean tap water. Screw the lid on the plastic bottle and
place the 00 rubber stopper in the hole to prevent spilling the storage solution. Set
this plastic bottle, with lid on and rubber stopper in the hole, in a safe place to prevent
spilling until the end of the experiment, when you have finished with the pH sensor
and will store it for future use.
7. Using a 10 mL graduated cylinder, measure 6.0 mL of clean tap water into the clean
sample vial labeled “CO2”.
8. With the pH sensor connected to CH1 on the side of the LoggerPro interface,
calibrate the pH meter using pH 7 and pH 4 buffers. Go to the “Experiment” tab on
the menu and click on “Calibrate”. Click “Calibrate Now”. Immerse the pH meter
into the pH 7 buffer, enter 7.00, and hit “Keep”. Do the same for the pH 4
calibration.
9. Prepare the computer for data collection by opening by clicking on “File” in the upper
left corner of the menu bar, select “open”, double click on “Chemistry with
Computers”, double click on “Experiment 22 Acid Rain”, open (or double click)
“Experiment 22 pH Sensor.MBL”. A screen appears with the vertical axis of the
graph scaled with pH from 0 to 10 pH units. The horizontal axis has time scaled from
0 to 100 seconds. Change the 100 seconds time duration to 1000 seconds, by going to
the “Experiment” tab on the menu and clicking on “Data Collection”. Change the
time interval for collection to 1000 seconds. Check to see that the input display at the
lower right in the meter window shows a pH value between 6 and 9 for the water.
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10. Remove the pH sensor from the 250 mL beaker of tap water, and rinse it with
deionized water, allowing the rinse water to collect in the 600 mL beaker waste
container. Transfer the pH sensor to the sample vial (labeled “CO2”) with the 6 mL
of tap water. The pH sensor in the sample vial is very unstable vertically and tips
over, so place the sample vial, with pH sensor inserted, into a 400 mL beaker to
provide vertical stability. Very gently stir the pH sensor in the 6 mL tap water in the
sample vial. Note that in a minute or so the pH in the meter display window is
becoming somewhat constant.
11. Click on “Collect” and start collecting pH data on the tap water sample
12. Using the 2 mL plastic dropper provided with the reagent bottle 1.0 M HCl, add about
1 mL (1/2 full dropper, about 20 drops) of 1.0 M HCl to the test tube with the solid
NaHCO3. Try to wash powder stuck on the walls of the test tube to the bottom as you
add the 1.0 M HCl. Fizzing is observed as carbon dioxide, CO2, is generated in this
tube. Squeeze the bulb of a clean 2 mL plastic dropper and insert the tip and stem of
the dropper into the test tube and release the squeezing as the bulb end gets close to
the open end of the test tube. Note that the bulb of the 2 mL plastic dropper catches
at the mouth of the test tube, preventing the dropper from touching the solution at the
bottom of the test tube and providing somewhat of a seal at the open end of the test
tube. The dropper must not touch the solution at the bottom of the test tube. Very
gently swirl the test tube that now contains NaHCO3 and HCl and dropper. Place the
test tube with dropper in the 150-mL beaker, to prevent spillage.
13. With the 2 mL plastic dropper still inserted into the test tube, squeeze the bulb of the
dropper completely flat several times to completely fill the dropper with the CO2 gas
that is in the test tube. You want to collect CO2 gas only in the dropper and not any
of the solution.
14. Without depressing the bulb. gently, quickly, remove the 2 mL plastic dropper from
the test tube, wipe the stem clean with tissue and insert the stem of the 2 mL dropper
into the sample vial with the pH sensor and 6 mL tap water. Push the plastic dropper
into the sample vial along the side wall of the pH sensor, using the stem (not the bulb)
to push the tip of the dropper through the tap water to the bottom of the sample vial.
15. Very slowly completely squeeze the bulb of the dropper so that the CO2 gas slowly
bubbles through the tap water. Do this 8 to 10 times so that the tap water drawn up
into the dropper is repeatedly contacted with the residual CO2 gas in the dropper.
16. Hold the bulb on the plastic 2 mL dropper squeezed tightly as you withdraw the stem
of the dropper from the sample vial, so that no solution remains in the stopper. Very
slowly, gently stir the solution (tap water with gas bubbled through) in the glass
sample vial with the pH sensor. Look at the graph which has been monitoring the pH
during Steps #9 through #14. Observe the constant pH of the tap water only at the
earlier times, the decrease in pH as the CO2 gas was bubbled through and the falling
value of the pH with the slow gentle stirring of the solution. Continue to follow the
falling pH with very gentle stirring until the pH becomes constant (may be 6 to 8
minutes or more).
17. When the pH becomes constant, you may click on “Stop” to stop collecting data, if
the 1000 second time has not already elapsed.
18. Record on your data sheet the best constant initial value of the pH of the tap water
just before you inserted the dropper with CO2 gas (initial pH value, before CO2 was
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added) and the best constant final value for the pH after adding the CO2 gas, followed
with the gentle stirring (final pH value after CO2 was added and pH stabilized). You
will find using the “x=?” button below the menu at the top of the screen is a good
help in determining the pH values that you must record from the constant regions on
the graph. If you wish to examine these values another way, click the Statistics
button, , and examine the minimum and maximum values in the pH box displayed
on the graph. Make certain that you have recorded the initial and final pH values in
your data table.
19. After recording your data, remove the pH sensor from the glass sample vial, rinse it
with deionized water (collecting the waste water in your 600 mL waste container) and
return it to the 250 mL beaker half full of tap water. Set aside the CO2 labeled glass
sample vial and “NaHCO3” labeled test tube for washing later.
20. To gather data on sulfur dioxide, SO2, start again like at Step #7, using the 10 mL
graduated cylinder to measure 6.0 mL of clean tap water into the clean sample vial
labeled “SO2”.
21. Repeat Steps similar to #10 through #19, with this time adding about 1 mL of 1.0 M
HCl to the test tube containing solid NaHSO3, and appropriately using “NaHSO3“ in
place of “NaHCO3 and “SO2“ in place of “CO2“. Of course, SO2 is generated in this
test tube. Do not breathe the SO2 gas.
22. To gather data on nitrogen dioxide, NO2, start again like at Step #7, using the 10 mL
graduated cylinder to measure 6.0 mL of clean tap water into the clean sample vial
labeled “NO2”.
23. Repeat Steps similar to #10 through #19, with this time adding about 1 mL of 1.0 M
HCl to the test tube containing solid NaNO2, and appropriately using “NaNO2“ in
place of “NaHCO3 and “NO2“ in place of “CO2“. Of course, NO2 is generated in this
test tube. Do not breathe the NO2 gas
24. If time permits, repeat the entire experiment. Appropriately discard wastes to protect
the environment as you clean the test tubes and sample vials with tap water, followed
by deionized water rinse. Then start again with NaHCO3 as in step #3.
25. To finish this experiment: Rinse the pH sensor with deionized water, unscrew the
cap on the plastic storage bottle, take out the 00 rubber stopper, slip the cap carefully
onto the cylindrical wall of the pH sensor (be careful of the glass bulb!) and then
screw the plastic bottle with the storage solution onto the pH sensor with cap so that
the glass bulb is submersed in the storage solution. Be sure the cap is screwed on
tightly to prevent leakage of storage solution. If some loss of storage solution has
occurred so that the bulb cannot be submersed in the storage solution, please take the
sensor and bottle to the instructor for special handling. The pH sensor is permanently
damaged if it is not stored under the storage solution.
26. Appropriately discard wastes to protect the environment. Wash and rinse other glass
apparatus and return the special glass sampling vials.
Page 185
DATA
Gas
Initial pH
Final pH
Change in pH (pH)
Initial pH
Final pH
Change in pH (pH)
CO2
NO2
SO2
For repeated experiments:
Gas
CO2
NO2
SO2
CALCULATIONS AND DATA ANALYSIS
Show the following calculations and answer the questions in your notebook.
1. For each of the three gases, calculate the change in pH (pH), by subtracting the final
pH from the initial pH. Record these values in the Data and Calculations table.
2. In this experiment, which gas caused the smallest drop in pH?
3. Which gas (or gases) caused the largest drop in pH?
4. Coal from western states such as Montana and Wyoming is known to have a lower
percentage of sulfur impurities than coal found in the eastern United States. How
would burning low-sulfur coal lower the level of acidity in rainfall? Use specific
information about gases and acids to answer the question.
5. High temperatures in the automobile engine cause nitrogen and oxygen gases from
the air to combine to form nitrogen oxides. What two acids in acid rain result from
the nitrogen oxides in automobile exhaust?
6. Which gas and resulting acid in this experiment would cause rainfall in unpolluted air
to have a pH value less than 7 (sometimes as low as 5.6)?
7. Why would acidity levels usually be lower (pH higher) in actual rainfall than the
acidity levels you observed in this experiment? Rainfall in the United States generally
has a pH between 4.5 and 6.0.
Page 186
Page 187
PRE-LAB
Name_________________
Experiment 20
Acid Rain
CO2 Worksheet
1. How many gallons of gas do you use per year?
I put ______(x)_ gallons of gas every _______(y)__ days. 365x/y = ______ gallons per
year.
Convert from gallons to mL to determine mL/year. (1 gallon = 3.785 L)
2. Follow the steps to calculate how much CO2 results in one year in moles, kg, L, and
classrooms from your personal gas consumption.
a. Assume gasoline is C8H18. Write a balanced equation for the combustion reaction of
gasoline in air.
b. Determine moles of CO2 from your mL of gasoline in #1 using the balanced equation.
(dC8H18 = 0.75 g/mL)
c. Determine kg of CO2 from b.
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d. Determine Volume (in L) using PV = nRT (assume 20oC and 0.75 atm).
e. Determine the number of classrooms this volume of CO2 would completely fill
(assume 1200 sq.ft. x 10ft room).
3. The US uses about 21 million barrels of oil per day. The world used about 86 million
barrels of oil per day. Calculate how many kg and L of CO2 are produced in the world
each day from gasoline. (Neglect all natural gas, coal and biofuels burned and any
energy used to process long chains into shorter chain hydrocarbons) (1 barrel = 42
gallons).
Page 189
Experiment 21
Chemical Equilibrium: Finding a Constant, Kc
The purpose of this lab is to experimentally determine the equilibrium constant, Kc, for
the following chemical reaction:
Fe3+(aq) + SCN–(aq) 
 FeSCN2+(aq)
iron(III) thiocyanate
thiocyanoiron(III)
When Fe3+ and SCN- are combined, equilibrium is established between these two ions
and the FeSCN2+ ion. In order to calculate Kc for the reaction, it is necessary to know the
concentrations of all ions at equilibrium: [FeSCN2+]eq, [SCN–]eq, and [Fe3+]eq. You will
prepare four equilibrium systems containing different concentrations of these three ions.
The equilibrium concentrations of the three ions will then be experimentally determined.
These values will be substituted into the equilibrium constant expression to see if Kc is
indeed constant.
In order to determine [FeSCN2+]eq, you will use a Vernier spectrometer. The FeSCN2+
ion produces solutions with a red color. The computer-interfaced spectrometer measures
the amount of light absorbed by the colored solutions (absorbance, A). By comparing the
absorbance of each equilibrium system, Aeq, to the absorbance of a standard solution,
Astd, you can determine [FeSCN2+]eq. The standard solution has a known FeSCN2+
concentration.
To prepare the standard solution, a very large concentration of Fe3+ will be added to a
small initial concentration of SCN– (hereafter referred to as [SCN–]i. The [Fe3+] in the
standard solution is 100 times larger than [Fe3+] in the equilibrium mixtures. According
to Le Chatelier's principle, this high concentration forces the reaction far to the right,
using up nearly 100% of the SCN– ions. According to the balanced equation, for every
one mole of SCN– reacted, one mole of FeSCN2+ is produced. Thus [FeSCN2+]std is
assumed to be equal to [SCN–]i.
Assuming [FeSCN2+] and absorbance are related directly (Beer's Law), the concentration
of FeSCN2+ for any of the equilibrium systems can be found by:
Aeq
[FeSCN2+]eq = A
std
2
X [FeSCN +]std
Knowing the [FeSCN2+]eq allows you to determine the concentrations of the other two
ions at equilibrium. For each mole of FeSCN2+ ions produced, one less mole of Fe3+ ions
will be found in the solution (see the 1:1 ratio of coefficients in the equation on the
previous page). The [Fe3+] can be determined by:
[Fe3+]eq = [Fe3+]i – [FeSCN2+]eq
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Because one mole of SCN- is used up for each mole of FeSCN2+ ions produced, [SCN–]eq
can be determined by:
[SCN–]eq = [SCN–]i – [FeSCN2+]eq
Knowing the values of [Fe3+]eq, [SCN–]eq, and [FeSCN2+]eq, you can now calculate the
value of Kc, the equilibrium constant.
Materials
Windows PC
Logger Pro
Vernier Spectrometer
1 plastic cuvette
five 20 X 150 mm test tubes
thermometer
0.0020 M KSCN
0.0020 M Fe(NO3)3 (in 1.0 M HNO3)
0.200 M Fe(NO3)3 (in 1.0 M HNO3)
four pipets
pipet bulb or pipet pump
three 100-mL beakers
tissues (preferably lint-free)
PROCEDURE
1. Obtain and wear goggles.
2. Label four large test tubes 1-4. Pour about 30 mL of 0.0020 M Fe(NO3)3 into a clean,
dry 100-mL beaker. Pipet 5.0 mL of this solution into each of the four labeled test
tubes. Use a pipet pump or bulb to pipet all solutions. CAUTION: Fe(NO3)3
solutions in this experiment are prepared in 1.0 M HNO3 and should be handled with
care. Pour about 25 mL of the 0.0020 M KSCN into another clean, dry 100-mL
beaker. Pipet 2, 3, 4 and 5 mL of this solution into Test Tubes 1-4, respectively.
Obtain about 25 mL of distilled water in a 100-mL beaker. Then pipet 3, 2, 1 and 0
mL of distilled water into Test Tubes 1-4, respectively, to bring the total volume of
each test tube to 10 mL. Mix each solution thoroughly with a stirring rod. Be sure to
clean and dry the stirring rod after each mixing. Measure and record the temperature
of one of the above solutions to use as the temperature for the equilibrium constant,
Kc. Volumes added to each test tube are summarized below:
Test Tube
Number
Fe(NO3)3
(mL)
KSCN
(mL)
H2O
(mL)
1
2
3
4
5
5
5
5
2
3
4
5
3
2
1
0
3. Prepare a standard solution of FeSCN2+ by pipetting 18 mL of 0.200 M Fe(NO3)3
into a large test tube labeled “5”. Pipet 2 mL of 0.0020 M KSCN into the same test
tube. Stir thoroughly.
4. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use a
cuvette, remember:

All cuvettes should be wiped clean and dry on the outside with a tissue.
 Handle cuvettes only by the top edge of the ribbed sides.
 All solutions should be free of bubbles.
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6. Calibrating and Determining the Maximum Wavelength of Absorbance (max) for the
FeSCN2+ solution.
a. Prepare the Vernier Spectrometer by plugging in the USB cable and opening the
Logger Pro software. If the software doesn’t immediately recognize the Spectrometer,
choose Connect Interface  Spectrometer  Scan for Spectrometers from the
Experiment menu. Allow the Spectrometer to warm up for 3 minutes before taking
readings.
b. Calibrate the spectrometer by choosing Calibrate  Spectrometer from the
Experiment menu. Follow the instructions from the dialog box to complete the
calibration using your blank cuvette. You will be asked to insert the blank cuvette into
the cuvette slot. Insert it in such a way that the spectrometer light goes through the
smooth sides and not the ribbed sides of the cuvette. Click “Ok.”
d. Rinse the sample cuvette with two or three small portions of the standard FeSCN2+
solution in test tube #5. Fill the cuvette about ¾ full with this the #5 standard solution.
Place the cuvette in the slot. Click “Collect”. An absorbance curve should appear on the
screen. After viewing the absorbance curve, hit “Stop”.
e. Click on the “Configure Spectrometer Data Collection” icon, located on the right hand
side of the toolbar to open the display. (The button looks like a rainbow graph.) Click
Abs. vs. Concentration (under Set Collection Mode). The wavelength of the maximum
absorbance will be automatically selected. Double check that ~470 nm (max) is the
only wavelength that is selected. Click “Ok” to close the display.
f. Record the absorbance value for standard solution #5. The last digit may fluctuate so
do your best to find the average.
7.
You are now ready to collect absorbance data for the four equilibrium systems.
a. Fill a cuvette about ¾ full with the equilibrium solution in Test Tube 1. Insert the
cuvette into the slot. Record the absorbance. Repeat this step for each of the
solutions in Test Tubes 2 – 4.
b. Dispose of all solutions as directed by your instructor.
Page 192
DATA
Trial 1
Trial 2
Trial 3
Trial 4
Absorbance
_______
_______
Absorbance of standard (Trial 5)
_______
Temperature
_______
Kc expression
_______
_______ °C
Kc =
3+
[Fe ]i
–
[SCN ]i
2+
[FeSCN ]eq
3+
[Fe ]eq
–
[SCN ]eq
Kc value
Average of Kc values
Kc = ________ at ________°C
Page 193
CALCULATIONS/DATA ANALYSIS
1. Write the Kc expression for the reaction in the Data and Calculation table.
2. Calculate the initial concentration of Fe3+, based on the dilution that results from
adding KSCN solution and water to the original 0.0020 M Fe(NO3)3 solution. See
Step 2 of the procedure for the volume of each substance used in Trials 1-4. Calculate
[Fe3+]i using the equation:
Fe(NO3)3 mL
[Fe3+]i =
X (0.0020 M)
total mL
This should be the same for all four test tubes.
3. Calculate the initial concentration of SCN–, based on its dilution by Fe(NO3)3 and
water:
KSCN mL
[SCN–]i = total mL
X (0.0020
M)
In Test Tube 1, [SCN–]i = (2 mL / 10 mL)(.0020 M) = .00040 M. Calculate this for
the other three test tubes.
4. [FeSCN2+]eq is calculated using the formula:
Aeq
[FeSCN2+]eq = A X [FeSCN2+]std
std
where Aeq and Astd are the absorbance values for the equilibrium and standard test
tubes, respectively, and [FeSCN2+]std = (1/10)(0.0020) = 0.00020 M. Calculate
[FeSCN2+]eq for each of the four trials.
5. [Fe3+]eq: Calculate the concentration of Fe3+ at equilibrium for Trials 1-4 using the
equation:
[Fe3+]eq = [Fe3+]i – [FeSCN2+]eq
6. [SCN–]eq: Calculate the concentration of SCN- at equilibrium for Trials 1-4 using the
equation:
[SCN–]eq = [SCN–]i – [FeSCN2+]eq
7. Calculate Kc for Trials 1-4. Be sure to show the Kc expression and the values
substituted in for each of these calculations.
8. Using your four calculated Kc values, determine an average value for Kc. How
constant were your Kc values?
Page 194
Page 195
Experiment 24
Measuring Sulfur Dioxide in Wine
The problem of preserving food is as old as civilization. Centuries ago, people learned
that the fumes of burning sulfur inhibited the browning and rotting of fruits and
vegetables. In the 19th century, the practice of using sulfur dioxide to preserve meat and
fish became widespread, to the point that it became of concern to health scientists who
pressed for regulations on its use.
Adding sulfur dioxide to food can be done either by exposing the material directly to the
fumes of burning sulfur, by dipping the food in a solution containing sulfite ions, or by
adding sulfites directly to the food, as is done in winemaking. In the latter two cases, acid
in the food reacts with the sulfites to produce free sulfur dioxide. Sodium hydrogen
sulfite (sodium bisulfite, NaHSO3) was commonly used for a long time, because it is
more stable than sodium sulfite (Na2SO3) and produces, gram for gram, more SO2.
Another form of the sulfite ion, the metabisulfite ion (S2O52-), is even more stable and
produces slightly more SO2. However, the concern about sodium in foods has led to the
more common use of the potassium salt when additions are made directly to the food.
24.1 SULFUR DIOXIDE AND HEALTH
Food that has been freshly or heavily sulfited may cause sneezing and mild shortness of
breath. In an attempt to learn the extent of sulfur dioxide as a health problem, first rats
and dogs, and then humans, were fed increasingly larger doses of sulfites. After three
generations of rats which were fed drinking water containing 750 parts per million (ppm)
of sulfites, no abnormal effects were observed. At levels of 4-6 grams a day added to
their diet, humans would suffer abdominal pain and vomiting, but no other symptoms and
no permanent effects. On the other hand, some individuals seem to be allergic to sulfites,
and thus concern about their use remains. Regulations of the United States government
state that finished wines should not contain more than 350 ppm of sulfur dioxide.
24.2 THE COMMERCIAL IMPORTANCE OF SULFITES IN WINE
Wine has been described as one of the steps on the way to making vinegar. Until recently,
with the introduction of millipore filtering techniques, it was impossible to make wine
that would last very long unless sulfur dioxide was added, because of the difficulty in
removing every last yeast cell that might renew fermentation in wine that had been
sweetened, or every microbe that might lead to the conversion of alcohol to vinegar
(acetic acid). In addition to its low toxicity, sulfur dioxide has the advantage of
possessing more than one preservative action. It kills or inhibits microbes, such as
bacteria and molds. It deactivates the enzymes which are released from damaged fruit and
vegetables and are responsible for browning and discoloration. And, to a small extent, it
combines with oxygen that has entered from the air and prevents oxidation that might
otherwise affect flavor.
Page 196
24.3 SULFUR DIOXIDE LEVELS IN WINE
In winemaking, a quantity of sulfur dioxide is added to
the freshly crushed grapes to kill yeasts and molds that
are naturally present on the grapes. The yeasts that are
added to actually make the wine have been bred over the
years to tolerate high SO2 concentrations, typically
80-100 milligrams per liter (or, as it is often expressed,
80-100 per million (ppm)). During the rest parts of the
fermentation and finishing process, and during bulk
storage, the concentration needs to be maintained at 30-50 mg/L or higher to protect the
wine from deterioration by air and microbes. Finally, at the time of bottling, additional
sulfur dioxide is added to attain a level of 30-40 mg/L. In the case of dessert wines which
are likely to be consumed more slowly after the bottle is opened, or any wine which will
probably be used more casually, the sulfur dioxide level added to the bottle might need to
be 70-80 mg/L.
Once the sulfur dioxide has been added, it is at the mercy of environmental conditions. It
will vary with the pH of the wine, the storage temperature, ethanol (alcohol) content,
micronutrient levels, and the sanitary condition of the wine. Much of it bonds with other
chemicals in the wine, such as aldehydes, glucose and ketones, and in this form it is
“fixed”, that is, unavailable to function as a preservative. Some of the remaining “free”
sulfur dioxide evaporates directly into the atmosphere and is lost. Some is oxidized to
sulfates. As a result of these interactions, the level of sulfur dioxide in newly bottled wine
quickly declines to 20-30 mg/L. From that point, it gradually continues to dissipate, and
in older wines may persist in concentrations of 5-10 mg/L.
On the other hand, while a certain level of sulfur dioxide is necessary to preserve the
wine, the winemaker must be careful to regulate the amount present in his or her creation.
A certain amount of it will dissipate during aging, and the winemaker must be able to
trace its disappearance in order to know how much to add. Too much of it will alter the
desired aroma and flavor, or interfere with aging. Some delicate wines might be ruined by
a level of only 25 mg/L. Most individuals can recognize its odor at levels of 15-40 mg/L,
and for that reason many wine drinkers prefer to let their wines “breathe” before they
drink them by pouring them from the bottle into a decanter to allow some of the gas to
escape. Some critical judges can detect free sulfur dioxide at concentrations of 5-10
mg/L.
Furthermore, the laws of different countries limit the amount of sulfur dioxide that is
allowable in wine, and the commercial winemaker must be careful not to exceed the legal
limits.
24.4 THE CHEMISTRY OF SULFUR DIOXIDE
Sulfur dioxide dissolved in water can exist as sulfur dioxide molecules (SO2), or it can
Page 197
form the hydrogen sulfite ion (HSO3-) and the sulfite ion (SO32-). There is an equilibrium
among these various forms of sulfur dioxide, depending on the amounts present, the pH,
and the temperature.
Gaseous sulfur dioxide, which is in equilibrium with dissolved sulfur dioxide, is easily
lost by dissipation into the atmosphere, thus presenting an inherent source of error in any
attempt to measure SO2. The higher the temperature, the quicker the gas escapes.
SO2 (aq) ⇋ SO2 (g)
Sulfur dioxide reacts with water molecules to form a weak acid, the hydrogen sulfite ion:
SO2 (aq) + H2O ⇋ H+ + HSO3Hydrogen sulfite ions dissociate into sulfite ions and metabisulfite ions:
HSO3- ⇋ H+ + SO322 HSO3- ⇋ H2O + S2O52In an acid environment, the equilibrium is shifted to the left.
Thus, at a low pH, SO2 molecules predominate, as illustrated
by the graph at the right. At higher pH levels, SO32- increases.
In between, HSO3- is the prevalent species.
24.5 MEASURING SULFUR DIOXIDE IN WINE
For reasons of health, economics, aesthetics, and compliance with government
regulations, it is important for the successful winemaker to monitor carefully the sulfur
dioxide content of his or her wine. While the measurement of sulfur dioxide in a pure,
uncomplicated sulfite solution is comparatively easy, the situation is quite different in
wine. Wine is a complex mixture of hundreds of chemical compounds present in varying
proportions. In wine, sulfur dioxide reacts with acetaldehyde (CH3CHO) to form
acetaldehyde--hydroxysulfonate:
-
CH3CHO + HSO3 ⇋ CH3CHOHSO3
-
It also reacts with certain sugars, (such as glucose, C6H12O6), organic acids (such as
pyruvic acid, CH3COCOOH) and phenolic compounds (which contain the phenol
molecule, C6H5OH). Sulfur dioxide bound with these molecules is not readily measured.
Page 198
But the winemaker needs to measure the total sulfur dioxide content -- both the free,
volatile sulfur dioxide, and the fixed or bound.
The free sulfur dioxide can be measured by titrating the wine with a standard solution of
iodine. Although elemental iodine is not very soluble in water, it can be made more
soluble by adding iodide ion to the solution. The iodine enters the solution as the triiodide
ion, I3 :
-
I2 + I  I3
-
Free sulfur dioxide reacts with water, to form the sulfite ion:
SO2 + H2O  2H+ + SO32The brown triiodide ion oxidizes the sulfite, and becomes the colorless iodide ion:
-
-
H2O + SO32- + I3  SO42- + 3I + 2H+
When starch is added to the mixture, the slightest trace of triiodide ion at the endpoint
produces a conspicuous, deep blue-black color, the color of a starch-iodine complex.
However, iodine also oxidizes polyphenols, a class of compounds which are always
present in wines, but especially in red wines. Since this reaction also consumes some
iodine and produces an artificially high reading, sulfuric acid is added to the wine to
reduce the interaction between iodine and polyphenols.
After the free sulfur dioxide is measured in this way, the fixed sulfur dioxide can be
released by adding sodium hydroxide. The mixture then is reacidified, and the titration is
continued.
However, while the theory is simple, the practice is elusive. Because of its volatility,
some free sulfur dioxide is likely to escape as soon as the bottle is opened, when the
sample is measured out, and during the actual titration. During the titration some sulfur
dioxide may react with atmospheric oxygen. Some chemical other than sulfur dioxide,
such as sugars, aldehydes, ascorbic acid (Vitamin C), and the polyphenols mentioned
above, react with iodine and produce an artificially high measurement. The presence of
other chemicals in the wine causes the easily recognizable starch-iodine complex to
decompose, and the endpoint quickly fades, leading the chemist to puzzle about whether
he or she has in fact reached the endpoint. Thus, the analysis of a complex mixture is
subject to numerous sources of experimental error, which must be controlled by the
chemist’s experience and knowledge of the reactions involved. This is where the chemist
must be an artist as well as a technician.
Page 199
24. 6 ANALYZING A POTASSIUM METABISULFITE SOLUTION
Obtain about 50 mL of a standardized 0.002 M I3- (aq) solution in your clean, dry 100mL beaker. Cover the beaker with a large square of Parafilm to decrease air movement
over the solution, which will increase the loss of iodine to the atmosphere. Record the
precise molarity of the solution.
Clean a buret and rinse it with very small portions of the 0.002 M triiodide solution to
displace any water in the buret. Fill the buret with triiodide solution and drain out just
enough to fill the buret tip and bring the level of the solution just below the ‘0’ mark.
Pour about 75 mL of a potassium rnetabisulfite (K2S2O5) solution of known concentration
into your clean, 150-mL beaker. After reviewing (and practicing, if necessary) correct
pipetting techniques, use a volumetric pipet to transfer 25.00 mL of the solution into a
clean (but not necessarily dry) 250-mL conical flask. Use your wash bottle to rinse down
the inside of the flask with a minimum of demineralized water.
Because the metabisulfite ion tends to be oxidized by air, you must know what you are
doing and proceed quickly and efficiently.
S2O52- + H2O ⇋ 2H+ + 2SO322SO32- + O2  2SO42It is probably better not to use a stirring bar and mechanical stirrer during this titration.
Instead, swirl the flask by hand.
Read the initial level of triiodide solution in the buret. Add 2 droppersful (4-5 mL) of 1%
starch indicator to the flask. Rinse the inside of the flask with a small amount of water.
Control your buret with one hand. Hold the titration flask with the other. When you
begin adding the triiodide solution to the metabisulfite, you will notice a blue-black color
appearing at the moment of contact of the two solutions.
This is the starch-iodine complex which will immediately disappear as metabisulfite ions
react with the iodine. Swirl the titration mixture after each addition of triiodide just until
the blue-black color disappears. When you notice that the color of the complex begins to
linger before disappearing, add iodine solution very cautiously until the addition of one
drop, or even less than one drop, causes the entire mixture to remain a pale blue in color.
At this point, rinse down the inner walls of the flask and any triiodide solution clinging to
the buret tip. If the color persists for 10 seconds or more, record the final level of
triiodide solution.
Calculate the molar concentration of S2O52-, using the equation.
3 H2O + S2O52- + 2 I3-  2 SO42- + 6 I- + 6 H+
At this point, you may want to check with your instructor to learn if your determination is
reasonable. Then repeat the titration two more times and average your results.
Page 200
Next, obtain a sample of potassium metabisulfite solution of unknown concentration and
repeat the titration.
24.7 MEASURING THE FREE SULFUR DIOXIDE IN WINE
This time, it will not be as easy to ascertain the endpoint of the titration, because of the
tendency of the blue-black color to fade as the equilibrium shifts and iodine in the
complex is reclaimed by competing reactions in the mixture.
You will titrate two samples of the same wine. The purpose of the first sample will be to
ascertain semi-quantitatively the approximate volume of triiodide required to reach an
endpoint. You will use the second sample to titrate more quickly by adding most of the
triiodide all at once, then more meticulously approaching the final endpoint.
From your instructor, obtain your sample of wine. Using a clean, dry pipet, or one that
has been previously rinsed with the same sample of wine, transfer 50.00 mL of wine into
a clean flask. Wash down the inside walls of the flask with demineralized water.
Add 10 mL of 6 M H2SO4, which will inhibit the oxidation of polyphenols by iodine.
Then add 2 droppersful of starch indicator.
Begin titrating as before. To speed things up, you can add about 0.5mL of iodine at a
time, swirling the flask vigorously after each addition. At first, the appearance and
clearing of the blue-black color will be similar to that which you observed when titrating
the known and unknown metabisulfite solutions. However, as you approach the endpoint
this time, you will notice that the blue-black color will change to a pale pink color just
before the mixture clears. At this point, begin adding one drop of triiodide at a time, then
swirl the flask just until the colors fade back to the original straw color. Continue to do
this until the pink persists for 5 seconds. You may stop titrating when you decide that the
pink color persists for about 5 seconds, regardless of whether it subsequently fades away.
Record your final buret reading.
Now add enough 6 M NaOH to change the mixture from a pale yellow to a deep yellow
color, probably 10-30 mL, or 5-15 droppersful. The pH of the mixture should be about
13. The NaOH will hydrolyze compounds to which SO2 is bound, freeing the SO2.
Stopper the flask, and allow it to stand for about 15 minutes, or until you are ready to
return to it.
Refill your buret. At this point, you may acidify your second 50.00-mL sample of wine
and add the starch indicator. Read the initial level on your buret, and run in all but about
one milliliter of the volume of triiodide solution required by the first sample. The
sample will probably turn a deep blue-black color, but should fade to clear as soon as you
swirl it. Then titrate drop by drop until the transitory blue-black color is replaced by the
pink hue that persists for 5 seconds, no longer. Don’t forget to record your final buret
reading. This volume of titrant will allow you to calculate the free SO2 in the wine, using
the equations
:
Page 201
.
H2O + SO32- + I3-  SO42- + 3I- + 3H+
SO2 + H2O  H+ + HSO3Now add as much 6 M NaOH as you used in the first sample. Stopper the second sample
and allow it to stand for 15 minutes or until you are ready to return to it.
24.8 MEASURING THE FIXED SULFUR DIOXIDE IN WINE
Continue the analysis of your first sample. Add enough 6 M H2SO4 to return the color of
the mixture to the original pale yellow color, probably 5-10 mL, or 2-4 droppersful. If it
doesn’t, add a few milliliters more acid, or enough to return to the pale yellow. The pH of
the mixture should be about 3. You don’t need to add more starch indicator. Read your
buret, then titrate as you did before, looking for the pale pink hue that replaces the blueblack and persists for 5 seconds. Take a final reading.
Since you will need about the same volume of triiodide to titrate your second sample,
check to be sure you have enough triiodide in the buret. After the second sample has
stood for at least 15 minutes, repeat the titration more carefully with your second sample.
This volume of titrant will allow you to calculate the fixed or bound SO2 in the wine.
24.9 CALCULATIONS
1. Determine the number of moles of free SO2 that reacted with the volume of
triiodide for both samples of wine.
2. Determine the molarity of the free SO2 in both samples of wine.
3. The molarity tells you the number of moles of SO2 per liter of wine, or moles per 1000
mL. Enologists commonly express the concentration of SO2 not in moles but in grams.
Convert moles per liter to grams per liter.
4. Furthermore, enologists express the concentration of SO2 not in parts per
thousand (ppt) but in parts per million (ppm). “Grams per liter” is the same as “grams per
1,000 mL, or ppt. Express the concentration of SO2 in parts per million, either as grams
per million parts of wine or as milligrams per liter of wine.
5. Determine the concentration of fixed SO2 in the sample of wine.
6. Express the total SO2 content of the wine in ppm.
Page 202
Acknowledgements
This laboratory experiment was developed with the generous assistance of the
technicians from the wine laboratory at Wente Bros. Winery, Livermore, California:
John Jeffray
Maria Coburn
Brad Buehler
The chemistry staff at Las Positas College thanks the Wente family for helping us to
further the academic development of our students by means of this interesting and
practical application of chemistry in action.
**********
Jim Adams, 5-24-95
Las Positas College, Livermore, CA
All Rights Reserved
RELATED REACTIONS
Sulfite Equilibria
(1)
SO2 (aq) ⇋ SO2 (g)
(2)
SO2 (aq) + H2O ⇋ H2SO3 (aq) ⇋ H+ (aq) + HSO3- (aq)
(3)
HSO3- (aq) ⇋ H+ (aq) + SO32- (aq)
Metabisulfite Equilibria
(4)
S2O52- (aq) + H2O ⇋ 2 HSO3- (aq) ⇋ 2 H+ (aq) + 2 SO32- (aq)
Iodine-sulfite Reactions
(5)
3 H2O (l) + S2O52- (aq) + 2 I3-  2 SO42- (aq) + 6 I- (aq) + 6 H+ (aq)
(6)
H2O (l) + SO32- (aq) + I3- (aq)  SO42- (aq) + 3 I - (aq) + 2 H+ (aq)
Iodine Reactions
(7)
I2 (aq) + I- (aq)  I3- (aq)
(8)
I2 (aq) ⇋ I2 (g)
(9)
I2 (aq) + H2O ⇋ HIO + H+ + I-
(10)
2 HIO  2 H+ + I- (aq) + O2 (g)
(11)
3HIO + 3 OH-  2 I- + IO3- + 3 H2O
(12)
4 I- + 4H+ + O2  2 I2 + 2 H2O
Sulfite Complexes
OH
O
CH3 C H
(13)
+
HSO3-
CH3 C SO3H
Page 203
Page 204
THE CHEMICAL COMPOSITION OF GRAPES AND WINE
Ripe wine grapes are about 71 - 80 % water. Dissolved in that water are various sugars (mainly glucose, fructose, and
sucrose) to an extent of 10 - 25 % by weight, depending on the ripeness of the fruit. Another 2-10 % consists of
organic acids (such as tartaric, malic, succinic, citric, lactic and acetic), about 2-3 % minerals (calcium, phosphorus,
iron, aluminum, magnesium, manganese, iron, sodium and potassium), various anions (chloride, silicate, carbonate,
sulfate, phosphate), and 0.5-1% other compounds (proteins, ails, vitamins, esters, enzymes, starch, pectins, gums,
tannins and pigments). Furthermore, not all grapes are identical. Different varieties of grapes have their own distinctive
flavor components, present in very small amounts, some of which are more desirable than others in the production of
flavorful wine.
The addition of yeast to crushed grapes stimulates the living yeast cells to begin metabolizing the sugars in the juice,
which is called “must”. The bulk of this metabolism is the conversion of sugars to ethanol (ethyl alcohol), carbon
dioxide, and heat. However, the yeast cells also release small amounts of other metabolic products, and numerous other
chemical reactions take place during the fermentation, depending on the strain of yeast, environmental conditions, and
the presence of air, bacteria, and other chemical and biological factors present on the grapes and on the equipment.
Some of these additional reactions and factors are encouraged in order to produce a beverage that is pleasant to the
taste, while others are carefully inhibited or controlled.
Thus, to say that wine is a mixture, not a pure substance, is an understatement. The following is a list of most of the
chemicals that have been detected in wine. Some of these affect any attempt to measure accurately the sulfur dioxide
content of wine.
Hydrocarbons

Myrcene
Alcohols

methanol

ethanol

1-propanol

2-propanol

1 –butanol

2-butanol

2- methylpropanol

2-methyl-2-propanol

3-methylthiopropanol

(-)2,3-butanediol

furfuryl alcohol

isopentyl alcohol

2-methyl-1-butanol

1 -pentanol

2-pentanol

cis-3-hexen-1-ol

1-hexanol

2-hexanol

3-methyl-1-pentanol

4-methyl-1-pentanol

benzyl alcohol

1 -heptanol

2-heptanol

2-phenethanol

p-hydroxyphenylethanol

1 -octanol

1 -nonanol

2-nonanol

glycerol

2-indolylethanol

2-bornanol

Linalool

Citranellol

1-decanol

2-decanol

1 –undecanol

1-dodecanol

2-dodecanol

-terpineol












3,7-dimethyl-1,5,7-octatrien-3-ol
2-ethylhexanol
trans-2-hexen-1-ol
trans-3-hexen-1-ol
3-octanol
(±)2,3-butanediol
1-octen-3-ol
1-phenylethanol
nerol
geraniol
4-terpineol
farnesol
Aldehydes

acetaldehyde

propionaldehyde

2-methylpropanal

2-furaldehyde

Isovaleraldehyde

Hexenal

2-hexenal

Benzaldehyde

Vanillin

Cinnamaldehyde

phenylacetaldehyde
Ketones

-ionone

-ionone

2-nonanone

3-hydroxy-2-butanone

2,3-butanedione

acetone

2-butanone
Acids

formic

oxalic

acetic

pyruvic

3-hydroxypyruvic















































propionic
lactic
succinic
malic
butyric
isobutyric
tartaric
citric
2-oxoisovaleric
4-(methylthio)-2-oxobutyric
glutaric
2-methylbutyric
3-methylbutyric
2-hydroxy-3-methylbutyric
pentanoic
tricarballylic
3-methyl-2-oxopentanoic
4-methyl-2-oxopentanoic
hexanoic
4-methylpentanoic
2-hydroxyhexanoic
2-hydroxy-4-methylpentanoic
benzoic
4-hydroxybenzoic
salicylic
protocatechuic
gallic
heptanoic
phenylacetic
vanillic
octanoic
3-hydroxyoctanoic
p-hydroxyphenylpyruvic
1-phenyllactic
syringic
azelaic
nonanoic
9-decenoic
decanoic
undecanoic
dodecanoic
tetradecanoic
pentadecanoic
hexadecanoic
heptadecanoic
2-furoic
2-oxoglutaric
Esters

ethyl formate

ethyl acetate

ethyl propionate

propyl acetate

isopropyl acetate

ethyl 3-hydroxypropionate

ethyl lactate

1,3-propanediol monoacetate

ethyl acetoacetate

ethyl 4-oxobutyrate

ethyl acid succinate

diethyl oxalate





























































ethyl acid malate
ethyl acid tartrate
butyl acetate
ethyl butyrate
ethyl isobutyrate
isobutyl acetate
ethyl pyruvate
ethyl 3-hydroxybutyrate
ethyl 4-hydroxybutyrate
ethyl isovalerate
ethyl valerate
pentyl acetate
isopentyl acetate
2-methylbutyl acetate
ethyl 2-hydroxy-3-methylbutyrate
methyl salicylate
diethyl succinate
diethyl malate
diethyl tartrate
ethyl hexanoate
hexyl acetate
isobutyl isobutyrate
ethyl-2-hydroxy-4-methylpentanoate
isopentyl lactate
2-phenethyl formate
ethyl heptanoate
isobutyl valerate
dimethyl phthalate
2-phenethyl acetate
ethyl octanoate
hexyl butyrate
hexyl butyrate
isobutyl hexanoate
isopentyl isovalerate
ethyl isopentyl succinate
ethyl 3-phenyllactate
ethyl nonanoate
hexyl valerate
hexyl isovalerate
isopentyt hexanoate
2-methylbutyl hexanoate
propyl octanoate
diethyl phthalate
2-phenethyl butyrate
ethyl 9-decenoate
ethyl decanoate
hexyl hexanoate
isobutyl octanoate
ethyl undecanoate
isopentyl octanoate
(-)2,3-butanediol monoacetate
(±)2,3-butanediol monoacetate
2-phenethyl hexanoate
diisopentyl succinate
ethyldodecanoate
hexyl octanoate
isobutyl decanoate
isopentyl decanoate
2-methylbutyl pentanoate
2-phenethyl octanoate
ethyl tetradecanoate
Page 205
Page 206








ethyl pentadecanoate
isopentyl dodecanoate
2-methylbutyl dodecanoate
ethyl hexadecanoate
isopentyl tetradecanoate
2-methylbutyl tetradecanoate
ethyl methylmalate
2,6,6-trimethyl-2-vinyl-4-acetoxytetrahydropyran
Carbohydrates

glucose

fructose

mannitol
Lactones

-butyrolactone

4-acetyl-4-hyroxybutyric acid- -lactone

pantolactone

4,5-dihydroxyhexanoic acid- -lactone

4-carboethoxy-4-hydroxybutyric acid- -lactone

trans-4-hydroxy-3-methyloctanoic acid- -lactone

2-methyl-4-hydroxybutyric acid- -lactone

4-hydroxypentanoic acid- -lactone

4-hydroxyhexanoic acid- -lactone

4-hydroxynonanoic- -lactone

2-vinyl-2-methyltetrahydrofuran-5-one
Acetals

2,4,5-trimethyl-1,3-dioxolane

1,1-diethoxyethane

2.4-dimethyl-5-ethyl-1,3-dioxolane

1-ethoxy-1-propoxyethane

1,1-diethoxypropane

1,1-dipropoxyethane

1,1-diethoxybutane

1,1-diethoxy-2-methylpropane

1-ethoxy-1 (3-methylbutoxy)ethane

1-ethoxy-1 (2-methylbutoxy)ethane

1,1-diethoxy-3-methylbutane

1,1-diethoxy-2-methylbutane

1-ethoxy-1-(2-phenethoxy)ethane

1,1-diisopentoxyethane

1,1-di(2-methylbutoxy)ethane

1 -isopentoxy-1 -(2-methylbutoxy)ethane

1-isopentoxy-1-(2-phenothoxy)ethane

1 -(2-methylbutoxy)-1 -(2-phenthoxy)ethane
Phenols

Phenol

3-methylphenol

p-ethylphenol

2-methylphenol

p-vinylphenol

1-naphthol

acetovanillon

tyrosol

p-vinylguaiacol
Nitrogen-containing compounds

methylamine

dimethylamine

ethylamine

ethanolamine

propylamine

isopropylamine

pyrrolidine

N-ethylacetamide

butylamine

isobutylamine

putrescine

histamine

pentylamine

isopentylamine

hexylamine

methyl anthranilate

4-carboethoxy-4-aminobutyric acid- -lactam

N-acetylalanine ethyl ester

N-isopentylacetamide

2-phenethylamine

alanine

arginine

glutamic acid

proline

tyramine

N-(2-phenethyl)acetamide

N-isobutylacetamide

N-(2-methylbutyl)acetamide

N-3-(methylthio)propylacetamide
Miscellaneous

Methanethiol

ethanethiol

dimethylsulfide

dimethyldisulfide

diethylsulfide

diethyldisulfide

diallylsulfide

butylethylsulfide

diisopropylsulfide

diisopropyldisulfide

di(3-methylbutyl)sulfide

cis-linalool oxide

trans-linalool oxide

2,6,6-trimethyl-2-vinyl-4-hydroxytetrahydropyran

nerol oxide

cis-roseoxide

trans-roseoxide

2-methylthiophane-3-one phthalide
Gases

carbon dioxide

sulfur dioxide

hydrogen sulfide
Page 207
Lab Report 24
DATA
24.6 Analyzing a Potassium Metabisulfite Solution
a. Accepted concentration of known K2S2O5 (aq) (as M)
_______________
b. Volume of K2S2O5 (aq) taken (in mL)
_______________
c. Concentration of standardized I3- (aq) solution (as M)
_______________
d. Titration data for a sulfite solution of known concentration:
Trial
1
2
3
4
Final volume (mL)
Starting volume (mL)
Delivered volume (mL)
e. Observations:
f. Sources and sizes of errors:
g. Calculations: Molarity of known K2S2O5 (aq) based on volume of I3- (aq):
Trial
1
2
3
4
Moles of I3- (aq)
Moles of S2O52- (aq)
M of S2O52- (aq)
h. Average experimental value of M of S2O52- (aq) ____________
Page 208
i. Spread:
high - low
x 100% = __________
average
j. Compare your experimental value to the accepted concentration of S2O52- (aq):
average experimental M - accepted M
x 100% = __________
accepted M
l. Code number of unknown K2S2O5 (aq)
___________________
m. Volume of K2S2O5 (aq) taken (in mL)
___________________
n. Concentration of standardized I3- (aq) solution (as M)
___________________
o. Titration data for a sulfite solution of unknown concentration:
Trial
1
2
3
4
Final volume (mL)
p. Observations:
q. Sources and sizes of errors:
r. Ca1cuations: Molarity of unknown K2S2O5 (aq) based on volume of I3- (aq):
Trial
1
2
3
4
Moles of I3- (aq)
Moles of S2O52- (aq)
M of S2O52- (aq)
s. Average experimental value of M of S2O52- (aq):
t. Spread:
high - low
x 100% = __________
average
____________
Page 209
24.7 Measuring the Free Sulfur Dioxide in Wine
a. Name of the wine _______________________________________________
b. Volume of wine taken each time (in mL)
__________________
c. Concentration of standardized I3- (aq) solution (as M)
__________________
Acidify the sample of wine with 10 mL of 6 M H2SO4
d. Titration data for free SO2:
Sample
Trial
1
2
Final volume (mL)
e. Observations:
f. Sources and sizes of errors:
g. Calculations: Molarity of free SO2 based on volume of I3- (aq):
Trial
Moles of I3- (aq)
Moles of HSO3- (aq)
M HSO3- (aq)
M of SO2 (aq)
mg SO2/L
1
2
Page 210
24.8 Measuring the Fixed Sulfur Dioxide in Wine
Hydrolyze with 10-30 mL of 6 M NaOH
Stopper and let stand for 15 minutes.
Re-acidify with 5-l0 mL 6M H2SO4
a. Titration data for fixed SO2:
Sample
Trial
1
2
Final volume (mL)
b. Observations:
c. Sources and sizes of errors:
d. Calculations: Molarity of fixed SO2 based on volume of I3- (aq):
Trial
1
2
Moles of I3- (aq)
Moles of HSO3- (aq)
M HSO3- (aq)
M of SO2 (aq)
mg SO2/L
e. Total sulfites (add the free and fixed SO2 determined for Sample 2):
mg SO2/L
Parts SO2 per million (ppm)
Reference Material
Lab Report Format
Periodic Table
Errors, Precision and Accuracy
Treatment of Experimental Data
Statistics and Uncertainty in the Laboratory
Names, Formulas and Oxidation Numbers of Some Common Ions
Net Ionic Equations
Solubility Tables (4 versions)
Colors of Ions in Aqueous Solutions
Common Oxidation States of Six Elements Important in Redox Chemistry
Activity Series of Metals and Nonmetals
Acids and Bases
Properties of Water: Density and Vapor Pressure
Las Positas College
R1
Las Positas College
General Chemistry
Ansell/Deleray
Lab Reports
(Information which might be included)
-Experiment Title:
(brief, descriptive)
-Experimenter’s Name (include lab partners)
-Date:
Place:
-Purpose: (Objectives, Anticipated results, 1-2 complete sentences)
-Experimental procedures and Equipment Set Up: (brief outline, sketches and diagrams, may include by
reference, another experimenter should be able to repeat your experiment from your report.)
-Equipment: (name, type, number, accuracy, precision, limitations, etc.)
-Outside Data and References: (e.g. density or boiling points found in the Merck Index)
- Notes, Potential Problem areas, Reminders, Safety Cautions
- Ambient Conditions: IF RELEVANT (time start, time end, temperature, humidity, barometric pressure,
changes in conditions during experiment.) Record on data sheet during experiment.
- Experimental data: (Prepare data tables in advance, if possible)
(Record in INK on data sheet during experiment)
-Observations, Notes: (Record in INK during the experiment, in detail!)
----------------------------------------------------------------------------------------------------------------------------- --Calculations: (may be done during or after experiment, pencil okay)
-Discussion – (thoughtful explanation, analysis, discussion of results)
-Compare with theoretical/other experimental, expected, unanticipated)
-Sources of error, magnitudes of error, effects of error: (record during experiment)
-Conclusion, Recommendations: (quality of results, brief status of results and technique,
recommendations for improvements/changes in experiment, future experiments suggested,
applications.)
NOTES:
There are many ways to set up an experimental record. A record style should be appropriate for the
specific experiment.
The reports should be brief, concise and complete—one/several word notes or descriptions are desirable.
Set up and do as much as you can before starting experimental work (i.e. before lab period). Especially in a
crowded lab, completed pre-lab work is essential for safety, time.
Use ink for recording all data, sources of error, observations (pencil is okay for calculations). Immediately
record all data directly on data sheet, NOT on scratch paper, paper towels or the palm of your hand!
Record data, observations, sources of error in duplicate (using recorded lab notebooks.) Always keep a
copy for your records.
If a piece of data is recorded incorrectly, a single thin line (no erasure or obliterations) and a brief note
describing the error may be used to delete such data. Do not throw away or destroy any
data.
The experimental reports and-post experiment assignments are due at the beginning of your next lab
meeting following the scheduled completion of the experiment.
2
IIA
4
Be
9.012
12
Mg
24.31
20
Ca
40.08
38
Sr
87.62
56
Ba
137.3
88
Ra
(226)
CWA 12/12/93
Revised 5//12/97
7
6
5
4
3
2
1
1
IA
1
H
1.008
3
Li
6.941
11
Na
22.99
19
K
39.10
37
Rb
85.47
55
Cs
132.9
87
Fr
(223)
3
IIIB
21
Sc
44.96
39
Y
88.91
57
La
138.9
89
Ac
(227)
Las Positas College
5
VB
23
V
50.94
41
Nb
92.91
73
Ta
180.9
105
Db
(262)
6
7
8
9
VIB VIIB VIIIB VIIIB
24
25
26
27
Cr
Mn
Fe
Co
52.00 54.94 55.85 58.93
42
43
44
45
Mo
Tc
Ru
Rh
95.94 (98) 101.1 102.9
74
75
76
77
W
Re
Os
Ir
183.9 186.2 190.2 192.2
106
107
108
109
Sg
Bh
Hs
Mt
(263) (262) (265) (266)
14
IVA
6
C
12.01
14
Si
28.09
32
Ge
72.61
50
Sn
118.7
82
Pb
207.2
15
VA
7
N
14.01
15
P
30.97
33
As
74.92
51
Sb
121.8
83
Bi
209.0
16
VIA
8
O
16.00
16
S
32.07
34
Se
78.96
52
Te
127.6
84
Po
(209)
17
VIIA
9
F
19.00
17
Cl
35.45
35
Br
79.90
53
I
126.9
85
At
(210)
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.2 173.0 175.0
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
232.0 (231) 238.0 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (260)
4
IVB
22
Ti
47.88
40
Zr
91.22
72
Hf
178.5
104
Rf
(261)
13
IIIA
5
B
10.81
13
10
11
12
Al
VIIIB IB
IIB 26.98
28
29
30
31
Ni
Cu
Zn
Ga
58.69 63.55 65.39 69.72
46
47
48
49
Pd
Ag
Cd
In
106.4 107.9 112.4 114.8
78
79
80
81
Pt
Au
Hg
Tl
195.1 197.0 200.6 204.4
110
111
112
Ds
Rg
(269) (272)
Periodic Table of the Elements
R2
18
VIIIA
2
He
4.003
10
Ne
20.18
18
Ar
39.95
36
Kr
83.80
54
Xe
131.3
86
Rn
(222)
R12
Las Positas College
Chemistry 30A, 31
JB—7/30/01
DECISION TREE: Should a formula be written as ions?
Is the substance an acid (formula begins with H), a base (contains hydroxide ion), a salt
(cation anion) or other?
If it is an acid, go to A.
If it is a base, go to B.
If it is a salt, go to S.
If it is an other, do not write it as ions.
Example: Zn stays Zn
(Note: There are three substances that decompose yielding gases when
formed in chemical reactions: H2CO3, H2SO3 and NH4OH.)
A.
Is it a strong acid: H2SO4(aq), HClO4(aq), HNO3(aq), HCl(aq), HBr(aq), or
HI(aq)?
If yes, write it as ions. Example: H2SO4(aq) becomes 2 H+ + SO42If no, do not write it as ions. Example: HNO2(aq) stays HNO2
B.
Is it a strong base: NaOH, LiOH, KOH, Ba(OH)2, or Ca(OH)2?
If yes, write it as ions. Example: NaOH becomes Na+ + OHIf no, do not write it as ions. Example: Fe(OH)3 stays Fe(OH)3
(Note: Most hydroxides not listed above are weak or nonelectrolytes because they
are insoluble in water—always check this out when writing net ionic equations.)
S.
Is it a soluble salt? Consult a table of solubilities.
If soluble, write it as ions.
Example: MgCl2 becomes Mg2+ + 2 Cl-
If not soluble, do not write it as ions.
Example: AgCl stays AgCl(s)
R18
Las Positas College
Chemistry 1A
10/15/94 Juliet Bryson
COMMON OXIDATION STATES OF SIX ELEMENTS
IMPORTANT IN REDOX CHEMISTRY
Much of the redox chemistry in this course can be derived from about six elements. The six
are Cl, Cr, Mn, N, O, and S. The common oxidation states of each element, in turn, will be
displayed on a vertical number scale, with compounds and ions listed that exhibit each
particular oxidation number. In general, compounds at the very highest oxidation level are
likely to be oxidizing agents. (Certainly not reducing agents. Why?). Compounds at the very
bottom are likely to be reducing agents. (Certainly not oxidizing agents. Why?) Compounds
exhibiting intermediate oxidation states can go either way—it depends on the other reagent
involved. If that reagent is a stronger reducing agent (RA), the first will become the
oxidizing agent (OA), and vice versa.
Chlorine (Bromine and iodine behave similarly)
Oxidation No Chlorine Compound
Characteristics
+7
Cl2O7 is unstable.
(Cl2O7), HClO4, ClO4HClO4 is a strong OA, reduced to Cl+5
strong OA, reduced to ClHClO3, ClO3+3
HClO2, ClO2good OA, reduced to Cl+1
(Cl2O), HClO, ClO-, OCl- good OA, reduced to Cl0
Cl2
good OA, reduced to Cl-1
ClChromium
Oxidation No
+6
+3
0
Oxygen
Oxidation No
0
-1
-2
Chromium Compound
acidic
basic
2CrO42Cr2O7
(orange)
(yellow)
3+
Cr(OH)4Cr
(green or violet) (green)
Cr
Oxygen Compound
O2
H2O2, HO2H2O, OH-
Characteristics
both are strong OA’s
Cr(OH)3 is amphoteric (soluble in
either acid or base)
the metal
Characteristics
the ‘original’ and most abundant OA
OA or RA
R19
Manganese
Oxidation No
+7
+6
+4
+2
0
Nitrogen
Oxidation No
+5
Manganese Compound
MnO4- (purple)
(permanganate ion)
MnO42- (green)
(manganate ion)
MnO2
Mn2+
Mn
Nitrogen Compound
N2O5, HNO3, NO3-
+4
+3
NO2 (N2O4)
(N2O3), HNO2, NO2-
+2
NO
+1
0
-3
N2O
N2
NH3, NH4+
Sulfur
Oxidation No
+6
+4
+2
(ave!)
0
-2
Sulfur Compound
SO3, H2SO4, HSO4-, SO42SO2, H2SO3, HSO3-, SO32S2O32(thiosulfate ion)
S
H2S, HS-, S2-
Characteristics
strong OA, reduced to
a) Mn2+ in acid
b) MnO2 in neutral or slightly basic
c) MnO42- in strongly basic
easily reduced to MnO2
brown solid, not soluble
pale pink to colorless
the metal
Characteristics
strong OA, reduced to:
a) NO2 in conc acid (> 8 M)
b) NO in dil acid (<6 M), but can go all
the way to NH3 with strong OAs
brown gas
only NO2- is stable
active as OA or RA
colorless but oxidized by O2 in air to NO2
(brown)
supports combustion about like oxygen
RA, but not commonly used in this
capacity in Chem 1A
Characteristics
conc acid is strong OA
OA or RA
used as RA in analytical chemistry
yellowish powder
strong RA; usually oxidized to S
SOME COMMON REDUCING AGENTS
1) The metals: oxidized to their positive ions; e.g. Sn oxidized to Sn2+ or Zn oxidized to Zn2+
2) Ions in which the metal has another higher oxidation state; e.g. Sn2+ oxidized to Sn4+ or Fe2+
oxidized to Fe3+ or Hg22+ oxidized to Hg2+
3) Carbon and organic compounds may be oxidized to other organic compounds and to CO2 and
H2O; e.g. C (coke, much used in industry) oxidized to CO or CO2 or CH3CH2OH (an alcohol,
ethanol) oxidized to CH3CHO (an aldehyde, ethanal) and then further oxidized to CH3COOH ( an
organic acid, ethanoic acid or acetic acid)
R20
ACTIVITY SERIES OF METALS
The metals are listed in order of decreasing strength as reducing agents.
Li
K
Ba
Very active with H2O
Ca
or acids. H2 formed.
Na
Mg
Al
Active with acids or
Mn
with steam. H2 formed.
Zn
Cr
Fe
Cd
Ni
Less active with acids.
Sn
Pb
H2 formed.
H2
Active only with stronger
Cu
oxidizing acids, as HNO3. Sb
Bi
No H2 formed.
Hg
Ag
React only with aqua regia: Au
Pt
HCl+HNO3, 3:1
best RA
Oxides reduced by electrolysis, but
not by H2 or CO.
Oxides reduced by C
or Al; not by H2 or CO.
Oxides reduced by heating
with H2 or CO.
Oxides reduced to metal
(decomposed) by heat alone.
poorest RA
ACTIVITY SERIES OF NONMETALS
The nonmetals are listed in order of decreasing strength as oxidizing agents.
F
Cl
O
Br
I
S
strongest OA
less strong

Chemistry 1A: General Chemistry

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