Experiment 1 – The Iodine “Clock” Reaction

Transcription

Experiment 1 – The Iodine “Clock” Reaction
ABSTRACT
I.
THE RATE LAW
1. The Effect of Initial Concentration of Reactants on Reaction Rate
2. Reaction Rates
3. Reaction Orders
4. The Rate Constant
II.
THE EFFECT OF TEMPERATURE ON REACTION RATE
1. Determination of Activation Energy
2. Determination of the effect of 100C temperature increase on Reaction Rate.
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Chemistry 102
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EXPERIMENT 1
REACTION RATE, RATE LAW, AND ACTIVATION ENERGY
THE IODINE “CLOCK” REACTION
PURPOSE:
I. To determine the Rate Law for the following chemical reaction:
2 I-(aq)
+
S2O82-(aq)
I2 (aq) +
iodide ion
persulfate ion
iodine
II. To determine the Effect of Temperature on the Reaction Rate
2 SO42- (aq)
sulfate ion
PRINCIPLES:
I. THE RATE LAW
1. The Effect of Initial Concentrations of Reactants on Reaction Rate
The rate of a chemical reaction is a measure of how fast a chemical reaction occurs.
The Reaction Rate can be determined experimentally by measuring the change in
concentration of the reactants of products, divided by the change in time.
During the course of the reaction, the reactants are used up to produce products.
As a result, the concentration of the reactants (I- and S2O82-) decreases and the
concentration of the products (I2 and SO42-) increases accordingly.
In this experiment the Reaction Rate will be calculated by dividing the experimentally
determined increase in concentration of one of the products (elemental iodine, I2), by the
corresponding time interval:
Δ I2
Rate =
Δt
The experimental determination of the increase in concentration of iodine (I2), during a
corresponding time interval, can be easily monitored, since the presence of even small
amounts of iodine can be detected by virtue of the intensely blue colored complex formed
between iodine and starch.
2 I-(aq)
+
S2O82-(aq)
iodide ion
persulfate ion
I2 (aq)
+
2 SO42- (aq)
iodine
sulfate ion
Reacts with starch
to form
a deep-blue complex
(Reaction 1)
One creative way of measuring the rate of formation of iodine is to couple the reaction in
which the iodine is formed (Reaction 1) with a much faster reaction that consumes all of
the iodine (Reaction 2)
I2(aq)
+
2 S2O32-(aq)
thiosulfate ion
2 I-(aq)
+
S4O62-(aq)
(Reaction 2)
Reaction 2 immediately consumes the I2 generated in the first reaction, until all of the
S2O32- (thiosulfate ion) is used up. When all of the S2O32- is consumed, I2 builds up and
reacts with starch to form the deep blue Starch-Iodine Complex, according to Reaction 1
given above:
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Chemistry 102
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EXPERIMENT 1
The appearance of the deep-blue complex tells us that at this point in time ( tcolor),
sufficient I2 has been produced by Reaction 1 to use up all of the S2O32- (thiosulfate
ion) originally added. From this observation, we can calculate the concentration of I2
(Iodine) produced by Reaction 1, by noting that according to the stoichiometry of
Reaction 2:
1 mole of I2 reacts with 2 moles of S2O32- (thiosulfate), or
½ mole of I2 reacts with 1 mole of S2O32- (thiosulfate).
It follows that at the time the deep blue color ( tcolor) appears:
[S2O32-] originally added and used up
[I2] produced =
2
2. Reaction Rates
If we know the initial concentration of the thiosulfate ion (S2O32-), that is the same for each
experiment, and remember that it is all used up when the color of the solution changes, then
we know that half the amount of I2 was also consumed in Reaction 2.
This means that the change in the I2 concentration is equal to half the starting (initial)
concentration of the thiosulfate ion S2O32-) and it remains constant throughout all the
experiments. In essence, the same amount of I2 is produced in each experiment at the time the
color changes, but it takes varying times for this to occur, since the tcolor depends on the
reaction conditions (concentration of reactants and temperature).
[S2O32-]
Δ [I2] =
2
The time of the color change (tcolor) is also the time that passed during the reaction (Δt).
It follows that the rate of any of the reactions can be calculated as:
[I2] produced at tcolor
Δ [I2]
Rate =
=
Δt
tcolor
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Chemistry 102
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EXPERIMENT 1
3. Reaction Orders
The rate of a reaction depends on the concentration of one or more of the reactants.
Consider a reaction with a pattern similar to the reaction studied in this experiment.
aA
+
bB
cC
+
dD
REACTANTS
PRODUCTS
As long as the reverse reaction is negligibly slow, the relationship between the Rate of
Reaction and the Concentration of Reactants can be expressed by a mathematical
expression called the Rate Law:
Rate = k [A]m [B]n
where:
k:
is a proportionality constant called the Rate Constant,
m
is the Reaction Order with respect to Reactant A, and
n:
is the Reaction Order with respect to Reactant B.
The values of the reaction orders (“m” and “n”) determine the dependence of the reaction
rate on concentration of the respective reactants. Reaction orders commonly have one of
the following values: 0, 1, -1, or 0.5. Reaction orders (such as “m” and “n”) can only be
determined experimentally and they are NOT related to the coefficients of the balanced
chemical equation (such as “a”, “b”, “c” and “d”). The examples below illustrate how the
Reaction Orders (“m” and “n”) can be determined experimentally for a reaction involving
two reactants (A and B), such as the reaction studied in this experiment.
The reaction orders (“m” and “n”) with respect to the two reactants (A and B) are
determined by measuring the initial rate for several reaction runs with varying
concentrations of one reactant (for example A) independently of the concentration of the
other reactant (B). This allows us to determine the dependence of the rate on the
concentration of [A] and the numerical value of the Reaction Order with respect to
reactant A (“m”).
Temperature
Room
Temperature
Room
Temperature
Reaction
Run
1
[A}
M
0.0300
[B]
M
0.0450
Initial Rate
M/s
4.83 x 10-6
2
0.0600
0.0450
9.66 x 10-6
x2
Constant
x2
Between the first two experiments (1 & 2), the concentration of [A] doubles, the
concentration of [B] stays constant and Reaction Rate doubles. It follows that the initial
rate is directly proportional to the initial concentration of [A]. The reaction is therefore of
the first order with respect to [A] and m = 1.
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Chemistry 102
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EXPERIMENT 1
Experimental data seldom provide such obvious and easy to interpret numbers. Reaction
orders calculated from experimental data can be calculated by substituting two initial
concentrations and the corresponding initial rates into a ratio of the rate laws.
To calculate “m”
Rate 2
k [A]m[B]n
9.66 x 10-6 M/s
k [0.0600 M]m [0.450 M]n
=
=
Rate 1
k [A]m[B]n
4.83 x 10-6 M/s
k [0.0300 M]m [0.450 M]n
Canceling out similar terms and doing the calculations yields: 2.00 = 2.00m
To find “m”, take the log of both sides of the equation and solve for “m”
log 2.00
0.301
m
log 2.00 = log (2.00 ) log 2.00 = m log 2.00
m=
=
=1
log 2.00
0.301
The Reaction Order with respect to [B], (“n”) can be calculated in a similar manner by
using the data obtained for Reaction Runs 3 & 4.
Temperature
Room
Temperature
Room
Temperature
Reaction
Run
3
[A}
M
0.0500
[B]
M
0.0150
Initial Rate
M/s
1.33 x 10-6
4
0.0500
0.0300
2.66 x 10-6
Constant
x2
x2
Between the last two experiments (3 & 4), the concentration of [B} doubles, the
concentration of [A] stays constant and Reaction Rate doubles. It follows that the initial
rate is directly proportional to the initial concentration of [B]. The reaction is therefore of
the first order with respect to [B] and n = 1.
It follows that for this reaction the Rate Law is:
Rate Law = k [A]1[B]1
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Chemistry 102
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EXPERIMENT 1
4. The Rate Constant
The Rate Constant “k” is characteristic for every reaction and it is independent of the
concentrations of the reactants. However, “k” is temperature dependent.
The rate constant, “k” can be calculated, after the reaction orders have been determined
experimentally, from any of the reactions that have been run at the same temperature (for
example at room temperature, commonly abbreviated R.T.).
In this experiment data for four reactions run with different initial concentrations of
reactants but at the same temperature (room temperature) provide the experimental data
needed for the calculation of the Rate Constant, k, at R.T. In order to obtain an accurate
value for the rate constant, “k” will be calculated for every one of the four experiments
and the average value of the rate constant will be reported in the final expression of the
Rate Law.
An example of calculation of “k” at room temperature is given below for a reaction with
a pattern similar to the reaction studied in this experiment.
aA
+
bB
cC
+
dD
To calculate “k”:

First solve the Rate Law for k:
Rate
Rate = k [A]m [B]n

Solve for k
k=
[A]m [B]n
Next, substitute the Initial Concentrations of the Reactants and the experimentally
determined Reaction Orders and Initial Rate
Assume:
 [A]
= Initial Concentration of Reactant A = 0.300 M
 [B]
= Initial Concentration of Reactant B = 0.450 M
 m
= Reaction Order with respect to Reactant A = 1
 n
= Reaction Order with respect to Reactant B = 1
 Initial Rate = 4.83 x 10-6 M/s
4.83 x 10-6 M/s
= 3.58 x 10-3 M-1 . s-1
k=
1
[0.0300 M] [0.0450 M]
1
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Chemistry 102
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EXPERIMENT 1
II. THE EFFECT OF TEMPERATURE ON REACTION RATE
Changing the temperature changes the Rate of Reaction, by the effect the temperature change
has on the Rate Constant, k. In fact, the Rate Constant, k is only constant when the
temperature remains constant.
The Arrhenius equation quantifies the temperature dependence of the rate constant in the
following form:
Ea
RT
k=Ae
where:
k = The Rate Constant
Ea = The Activation Energy
 is an energy barrier that must be surmounted for the reactants to be
transformed into products.
 at a given temperature, the higher the Activation Energy, the slower the
reaction rate.
A = The Frequency Factor
 represents the number of times that the molecules of reactants approach the
activation barrier per unit time.
R = The universal gas constant = 8.314 J/K . mol
T = The absolute temperature
e = 2.718
The factors included in the Arrhenius equation are important quantities in understanding the
kinetics of reactions. This experiment aims to:
1. To determine the numerical value of the Activation Energy, Ea for the reaction
being studied and,
2. To determine the ratio by which reaction rate increases when the temperature
increases by 100C (arbitrarily chosen from 200C to 300C)
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Chemistry 102
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EXPERIMENT 1
1. Determination of Activation Energy, Ea
Since the temperature dependence of the reaction rate is contained in the rate constant,
we need to focus on the effect of temperature on the Rate Constant, “k”.
This is achieved by obtaining, analyzing and interpreting kinetic data obtained in the
laboratory. The experimental data is obtained by:
 Running several experiments with the same initial concentrations of reactants but at
different temperatures and determining the corresponding Reaction Rates
 Calculating from experimental data the values of the Rate Constants, k, for these
experiments.
The experimental data obtained is analyzed and interpreted by using the Arrhenius
equation.
Recall that according to the Arrhenius Equation:
Ea
Taking the natural log of both sides of this equation yields:
Ea
Ea
1
ln k = ln A OR
ln k = + lnA
RT
R
T
RT
k=Ae
The equation we have obtained has the form of the equation of a straight line: y = b + mx
Ea
1
ln k = -
+ lnA
R
y
= (m
T
x)
+
b
Dependent Slope Independent
Variable
Variable
Vertical
Horizontal
Coordinate
Coordinate
A plot of the natural log of the Rate Constant, k (ln k) versus the inverse of the Absolute
Temperature (1/T) yields a straight line with a slope = – Ea/R.
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EXPERIMENT 1
Such a plot is called an Arrhenius plot and is commonly used in the calculation of kinetic data,
such as the Activation Energy, Ea.
ln k
Ea
Slope = R
1
(K -1)
T
Figure 1
The Activation Energy, Ea can be calculated from the slope (slope = - Ea/R) of the straight line.
2. Determination of the effect of a 100C temperature increase on Reaction Rate.
The steps involved in this calculations are are listed below:
 Convert t1 = 200C and t2 = 300C respectively to 0K (T1 and T2)
 Calculate 1/T1 and 1/T2 (K-1)
 Read from the Arrhenius plot, the corresponding values for ln k (ln k1 and ln k2):
a
ln k2
ln k1



1
1
T2
T1
(K-1)
Figure 2
Take the anti ln of the above values to obtain:
 “k1” at 200C, and
 “k2” at 300C
Calculate the two reaction rates by substituting the respective “k” values and the
initial reaction concentrations (recall that the concentrations of reactants are the
same as in Reaction Run # 1)
Divide the reaction rate at 300C by the reaction rate at 200C
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Chemistry 102
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EXPERIMENT 1
PROCEDURE:
PART I: DETERMINATION OF THE RATE LAW
The table below summarizes the volumes of reactants and the reaction conditions under which
the experimental data will be obtained for all parts of the experiment.
TABLE I
ALL KINETIC RUNS
250 mL Reaction Flask
Run
Nr.
1
2
3
4
5
6
7
Temp
KI
0.200 M
Reactant 1
Room 20.00mL
Temp
Room 10.00 mL
Temp
Room 20.00 mL
Temp
Room 20.00 mL
Temp
About 20.00 mL
00C
About 20.00 mL
100C
About 20.00 mL
40oC
50 mL Flask
Na2S2O3
0.00500 M
KCl
0.200 M
Soluble
Starch
10.00 mL
0.00 mL
10.00 mL
10.00 mL
10.00 mL
0.00 mL
10.00 mL
0.00 mL
10.00 mL
0.00 mL
10.00 mL
0.00 mL
10.00 mL
0.00 mL
3-4
drops
3-4
drops
3-4
drops
3-4
drops
3-4
drops
3-4
drops
3-4
drops
(NH4)2S2O8 (NH4)2SO4
0.100 M
0.100 M
Reactant 2
20.00 mL
0.00 mL
CuSO4
0.100 M
0.00 mL
20.00 mL
0.00 mL
0.00 mL
10.00 mL
10.00 mL
0.00 mL
5.00 mL
15.00 mL
0.00 mL
20.00 mL
0.00 mL
0.00 mL
20.00 mL
0.00 mL
0.00 mL
20.00 mL
0.00 mL
0.00 mL
The table below summarizes the volumes of reactants to be used in making up the four
reaction mixture for the reactions run at room temperature
TABLE II
Kinetics Runs at Room Temperature and Varying Concentrations
250 mL Reaction Flask
50 mL Flask
Run
KI
Na2S2O3
KCl
Soluble (NH4)2S2O8 (NH4)2SO4 CuSO4
Nr.
0.200 M
0.00500
0.200 M
Starch
0.100 M
0.100 M
0.100 M
M
Reactant 1
Reactant 2
1
20.00 mL 10.00 mL
0.00 mL 3-4 drops
20.00 mL
0.00 mL 0.00 mL
2
10.00 mL 10.00 mL 10.00 mL 3-4 drops
20.00 mL
0.00 mL 0.00 mL
3
20.00 mL 10.00 mL
0.00 mL 3-4 drops
10.00 mL 10.00 mL 0.00 mL
4
20.00 mL 10.00 mL
0.00 mL 3-4 drops
5.00 mL 15.00 mL 0.00 mL
The actual procedure for carrying out each reaction run the same for all runs and it is described
in detail for Reaction Run Nr. 1
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Chemistry 102
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EXPERIMENT 1
a. Filling the 250 mL Erlenmeyer Flask (Reaction Flask)
 Use a buret to add exactly 20.00 mL of 0.200 M KI into the 250 – mL Erlenmeyer Flask,
which we shall refer to as the Reaction Flask.
 Use a buret to add exactly 10.00 mL of 0.00500 M Na2S2O3 into this flask.
 Add 3 or 4 drops of starch solution to the flask.
 Mix the contents of the flask by swirling the flask on the bench top.
b. Filling the 50 – mL Erlenmeyer Flask
 Use a buret to add exactly 20.00 mL of 0.100 M (NH4)2S2O8 into this flask.
c. Temperature readings
 Read the temperature of the solution in the 250 – mL flask with a rinsed, dried
thermometer. Record the temperature to the nearest degree.
 Remove the thermometer, rinse and dry it, and measure the temperature of the solution in
the 50 – mL flask. The solutions in both flasks should be at the same temperature ± 20C,
since both solutions have been kept at room temperature. Record the temperature as
Room Temperature (RT) to the nearest degree. If there is a slight difference between the
temperature readings in the two flasks (due to the limited accuracy of the thermometers),
record the average temperature as the initial temperature (Room Temperature).
d. Mixing and Timing
Figure 1
Have a timer available.
Pour the solution from the
50 – mL flask into the 250 mL
reaction flask.

Figure 2
While you start swirling the
solutions, your partner should
start the timer.
You should leave the 50 – mL
flask over the mouth of the
reaction flask, as shown above.
Figure 3
Continue swirling the
solution until the blue
color appears.
Note and record the time
at which the blue color
appears, to the nearest
0.1 of a second
Photos by Andrew Huertas
Figures 3a, 3b & 3c
Check the temperature of the mixture. Note and record the temperature as the final
temperature. The temperature of the reaction mixture should be recorded as the average
of the initial and final temperatures
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Chemistry 102
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EXPERIMENT 1
 If the timing was faulty, repeat the entire run.
e. Repeat the experiment with the other three mixtures in the table
 Both flasks should be rinsed with tap water and deionized water and drained between
experiments.
 Burets should be used in measuring the volumes of the following five solutions:
KI, Na2S2O3, KCl, (NH4)2S2O8 and (NH4)2SO4
NOTE:
Although KCl and (NH4)2SO4 are not reactants in the reaction we are studying, these
reagents serve to maintain the effective concentrations of all ions (“ionic strength”) at a
constant level. Holding the ionic strength constant removes the dependence of the reaction
rate on variations in the solvent.
PART II. THE EFFECT OF TEMPERATURE ON REACTION RATE
To obtain experimental data reflecting the dependence of the Reaction Rate on Temperature, the
reaction will be carried out at several different temperatures, but with the same concentration of
all reactants as in ReactionRun # 1.. Table III below indicates the reaction conditions:
TABLE III
Reaction Mixtures at different temperatures
Reaction Temperature
Reaction Mixture
Run
Range
5
About 00C
Same as in Rxn Run # 1
0
6
About 10 C
7
About 400C
Photo by Andrew Huertas
Reaction Run # 5 (about 00C)
Figure 4
 The reaction is carried out by adding the same volumes of the same solutions as in
Reaction Run # 1.
 Prepare mixtures of ice-cold water mixtures in a 250 mL beaker (for the 50 mL flask)
and a 600 mL beaker (for the 250 mL flask), which will be used to cool the contents
of the two flasks, when immersed in the ice/cold water slurry.
 Place one thermometer in each of the flasks and swirl the contents of the flasks.
 Immerse the flasks in the beakers containing the ice/cold water slurry. While swirling
the two flasks gently, follow the temperature readings on the two thermometers and
allow the temperature of the solutions to stabilize. Make sure that the bulbs of both
thermometers are immersed in the respective solutions. This is particularly important
for the temperature reading in the 50 mL flask and it may require to tilt this flask,
while immersed in the ice/cold water slurry. Your aim is to reach a temperature close
to 0 0C (± 20 C) for both solutions and that this temperature be the same (or almost the
same) for the two solutions contained in the two flasks. It is not imperative that the
temperature be exactly 00C (difficult to achieve) but it is important that the
temperature of the two solutions be as close as possible to each other.
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Chemistry 102
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EXPERIMENT 1
HINTS




The ice/cold water slurry should be mixed with a glass rod or a spatula.
The flasks should be fully immersed in the ice/cold water slurry.
The contents of the flasks should be gently mixed by swirling.
If the temperature reached is not low enough, the temperature of the solutions
may be further lowered by adding and mixing a spoonful of rock salt into the
ice/cold water slurry.
 When you are confident that the temperatures of the two solutions are about the same
and within ± 20 C from 00 C
 Note and record this temperature (Initial temperature), and
 Proceed to pour the solution from the 50 – mL flask into the 250 mL reaction
flask, while keeping the 250 mL flask immersed in the ice/water slurry and the
thermometer immersed in the flask.
 At this point, it is best to divide the tasks of the two team members as follows:
 One team member is responsible for:
o Starting the stopwatch,
o Observing and recording the time at which the color change occurs
to the nearest 0.1 of a second.
 The other team member is responsible for:
o Keeping the 250 mL flask immersed in the ice/cold water slurry for the
entire reaction interval.
o Mixing the contents of the 250 mL reaction flask by swirling the flask,
while keeping the thermometer immersed in the reaction mixture.
o Keeping track of the temperature of the reaction mixture, and
o Recording the temperature at the exact time at which the color change
occurred (final temperature) of the reaction mixture)
o The temperature of the reaction mixture should be recorded as the
average of the initial and final temperatures for this reaction run.
 Repeat the experiment at about 100C.
 Note and record the initial temperature before mixing, the time required for the
reaction and the final temperature of the reaction mixture at the time the solution
turned blue.
 Record the temperature of the reaction mixture for this run as the average of the
initial and the final temperature.
Reaction Run # 1 (Room Temperature)
 Data for this Reaction Rate is already available from Part I.
 Repeat the experiment at about 400C.
 Note and record the initial temperature before mixing, the time required for the reaction
and the final temperature of the reaction mixture at the time the solution turned blue.
 Record the temperature of the reaction mixture for this run as the average of the initial
and the final temperature.
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Chemistry 102
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EXPERIMENT 1
REPORT FORM
NAME: ______________________ Date: _____________ Partner: ______________________
PART I: DETERMINATION OF THE RATE LAW
1. Initial Concentrations of Reactants
For each kinetic run, calculate the initial concentration of the reactants:
[I-] and [S2O82-].
These calculations should be done prior to performing the
experiment!
Since the reaction takes place in a total volume of 50.00 mL, this volume must be taken
into account in calculating the initial concentration of the two reactants.
For example, in Run 1, since the 20.00 mL of 0.200 M KI added reacts in a total volume
of 50.00 mL, the initial concentration of [I-]0 can be calculated as follows:
20.00 mL
[I-]0 = 0.200 M KI
= 0.0800 M KI
50.00 mL
Similarly, in Run 1, the initial concentration of [S2O82-]0 is calculated as follows:
20.00 mL
[S2O82-]0 = 0.100 M (NH4)2S2O8
= 0.0400 M (NH4)2S2O8
50.00 mL
On the next page, carry out similar calculations for all other initial values of the two
reactants and complete the appropriate columns in DATA TABLE III.
You are required to:
 Show all your calculations
 Include units in your calculations, and
 Express all measured quantities (including your answer) in the appropriate number of
significant figures.
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Chemistry 102
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EXPERIMENT 1
DATA TABLE I
[I-]0
[S2O82-]0
Run 1
Run 2
Same as in Run 1
Run 3
Same as in Run 1
Run 4
Same as in Run 1
2. Reaction Rates
In order to determine the Reaction Rates, the following quantities must be known:
 Initial concentration of [S2O32-] added and completely used up (see below)
 Concentration of [I2] produced, at the time the deep blue color (tcolor) appears.
(see below); Enter this value in the Data Table I on the next page.
 The time in seconds when the deep blue color (tcolor) appears for each Reaction Run
(determined experimentally). Enter these data in DATA TABLE III.
DATA TABLE II
[S2O32-]
[S2O32-]added
[I2]produced
=
2
All Runs
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Chemistry 102
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EXPERIMENT 1
DATA TABLE III
Initial Concentrations and corresponding Reaction Rates
RATE*: Express Rate as (A x 10-6) throughout the entire experiment
Run
Nr.
Temp.
(0C)
(Room
Temp)
[I-]0
[S2O82-]0
TIME
(tcolor)
RATE
Expressed as A x 10-6
[I2]produced
tcolor
(M)
(M)
(s)
(M . s-1)
1
2
3
4
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Chemistry 102
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EXPERIMENT 1
3. Reaction Orders with respect to Reactants
The general formula for the Rate Law for the reaction studied is:
Rate1 = k [I-]m [S2O82-]n
Rate2 = k [I-]m [S2O82-]n
When performing your calculations you are required to:
 Show all calculations neatly, in a well-organized manner and in detail
(follow the format of the example shown on page 4 or the sample calculation presented in
your textbook (page 605)
 Round off your answer to an integer.
a. Reaction Order with respect to [I-}
From Run 1 and Run 2
Value of “m” rounded off to an integer =
b. Reaction Order with respect to [S2O82-]
From Run 3 and Run 4
Value of “n” rounded off to an integer =
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Chemistry 102
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EXPERIMENT 1
4. Rate Constant “k” at Room Temperature
The Rate Constant, “k” can be calculated by substituting the known values of reactant
concentrations, the reaction orders and the corresponding reaction rates in the formula
of the Rate Law. Refer to the example on page 5.
Write the equation for the Rate Law:
Rate =
DATA TABLE IV
Calculate “k” for Reactions 1 through 6. Show your calculations and include units!
Run 1
Run 2
kk21 ==
Run 4
k3 =
k4 =
DATA TABLE V
Summary of Rate Constants “k” for Reaction Runs at Room Temperature
1
2
3
4
Units
Run Nr.
k
(
k1 =
Run 3
)
k
(Average)
(
)
Write the complete form of the Rate Law:
 Include the formulas of both Reactants. Do not include numerical values
 Include the respective Reaction Orders for both Reactants
 Include the experimentally determined numerical value of “k”, expressed in the correct units.
RATE =
18
Chemistry 102
______________________________________________________________________________
EXPERIMENT 1
PART II: THE EFFECT OF TEMPERATURE ON REACTION RATE
To determine the quantitative relationship between Reaction Rate and Temperature you will
be plotting an Arrhenius Plot (ln k versus 1/T) to determine the:
1. Activation Energy (Ea) and
2. The ratio by which the reaction rate increases when the temperature is increased by 100C
(arbitrarily chosen from 200C to 300C)
1. Determination of the Activation Energy, Ea
 Summarize your experimental data for the Reactions Runs with the same
concentration of all reactants (as used in Reaction Run # 1), but run at different
temperatures. In this manner, the corresponding reaction rates for these runs will be
affected by temperature only.
DATA TABLE VI
Run Temp.
Initial
Final
Average
Time
RATE
Nr.
Range Temperature Temperature
Temperature
(tcolor)
[I2]produced
of
of
of
tcolor
Reactants
Reaction
Reaction
Mixture
Mixture
0
0
( C)
( C)
(0C)
(s)
(M . s-1)
5
About
00C
6
About
100C
1
Room
Temp.
7
About
400C
 Calculate the Rate Constants for Runs 5, 6 & 7 from the initial concentrations of
reactants and the corresponding reaction rates.
DATA TABLE VII
(Calculation do not need to be shown)
Run
Recorded
[I-]0
[S2O82-]0
RATE
Rate Constant
Nr.
Average
k
Temp.
(M-1 . s-1)
(K)
(M)
(M)
(M . s-1)
(Expressed as A x 10-3)
5
About
00C
6
About
100C
1
Room
Calculated Average
Temp
7
About
400C
19
Chemistry 102
______________________________________________________________________________
EXPERIMENT 1
 Calculate and collect the data that will be used for plotting ln k versus 1/T (The
Arrhenius Plot)
DATA TABLE VIII
(Calculations do not need to be shown)
Run
1
ln k
Nr.
T
(K-1)
5
About
00C
6
About
100C
1
Room
Temp
7
About
400C
 Plot a graph of ln k vs. 1/T.
A sample graph is attached to your Report Form. Please follow the format, scale, data
reporting style and all the other details included in the sample graph.
Calculate the slope of the graph and show all the calculations on the graph.
 Calculate the Activation Energy (Ea) from the slope of the graph
Please show calculations and include units.
J
Slope: __________
Recall: R = 8.314
mol . K
Ea =
J/mol
Ea =
20
KJ/mol
Chemistry 102
______________________________________________________________________________
EXPERIMENT 1
2. Determination of the effect of 100C temperature increase on Reaction Rate
 Collect your data:
t1 = 200C
T1 = _____ K
1/T1 = _______ (K-1)
t2 = 300C
T2 = _____ K
1/T2 = _______ (K-1)
 Read the corresponding values for ln k1 and ln k2 from your Arrhenius plot.
Indicate these values and the source of these readings on your graph
(as shown on Figure 2, page 8)
 Follow the guidelines given below Figure 2 (page 8) to calculate the ratio between the
Reaction Rate at 300C and the Reaction Rate at 200C
 Show all calculations neatly, in a well-organized manner and in detail
 Include units in your calculations
 Round off your answer to the nearest integer.
State your conclusion:
______________________________________________________________________________
______________________________________________________________________________
Bibliography:
1. Nivaldo J. Tro, “Chemistry: A Molecular Approach”, Third Edition
2. R.A.D. Wentworth “Experiments in General Chemistry”, Sixth Edition
3. James M. Postma & all, “Chemistry in the Laboratory”, Seventh Edition
21

Experiment 1 – The Iodine “Clock” Reaction

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