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Dr. Buckley
e-mail: [email protected]
Experiment #9 – Molar Stoichiometry in a Chemical Reaction
Laboratory Overview
CHEM 1361
August 2010
Gary S. Buckley, Ph.D.
Department of Physical Sciences
Cameron University
Table of Contents
(you may click on any of the topics below to go directly to that topic)
•The Chemical Reaction
•Experimental Scheme
•Calculations
•Safety Notes
•Notes on Oxidation-reduction Reactions
•Use of the Activity Series to Predict Redox Reactions
•Further Examples of Redox Reactions
The Chemical Reaction
The reaction you will observe here may be written as:
Cu(s) + 2 AgNO3 (aq)  Cu(NO3 ) 2 (aq) + 2Ag(s)
or, in net ionic form:
Cu(s) + 2 Ag + (aq)  Cu 2+ (aq) + 2 Ag(s)
The primary goal in this experiment is to verify the stoichiometric
coefficients in the above equation – 1 mol of Cu reacts with 2 mol of Ag+.
This reaction is an example of a displacement reaction. Notice the Cu
has lost two electrons from the reactant side to the product side and
each Ag+ has gained one electron. All displacement reactions are also
oxidation-reduction (or redox) reactions. In this case the Cu is oxidized –
loses electrons – and the Ag+ is reduced since it gains electrons.
Experimental Scheme
The reaction will be carried out by placing an accurately weighed piece of
copper wire into a solution containing a known quantity of silver nitrate.
As the reaction proceeds copper will dissolve in the solution causing a blue
color and metallic silver will form on the surface of the copper wire. This
tends to be one of the students’ favorite experiments – it is interesting to
watch.
Cu(s) + 2 Ag+ (aq)  Cu 2+ (aq) + 2 Ag(s)
Ag metal
coated on
Cu wire
Blue color due to
formation of Cu2+
ions
Calculations
Cu(s) + 2 Ag+ (aq)  Cu 2+ (aq) + 2 Ag(s)
The amount of both copper and silver reacted may be determined by
comparing the starting quantities and ending quantities of each. From the
mass reacted, the number of moles of each reacted may be determined in the
usual fashion:
# mol 
#g
molar mass
From this information, the ratio of mol of Ag reacted to mol of Cu may be
determined, as well as the mass ratio of grams of Ag reacted to grams of Cu
reacted.
Safety Notes
A few key safety notes:
1. Step 1, the making of the silver nitrate solution, must be carried out in
a fume hood.
2. You will be working with 6 M HNO3 which can burn if it comes in
contact with your skin. Wear gloves for the bulk of this experiment.
3. Another good reason to wear gloves is that the silver ion solution will
cause discoloration of your skin if it comes in contact. This is not a
safety issue, but more cosmetic in nature. The discoloration will
eventually go away.
Notes on Oxidation-Reduction Reactions
The occurrence of an oxidation-reduction (redox) reaction depends on the
relative ease with which the species involved can give up (or take on electrons).
From a terminology standpoint:
Oxidation – the loss of electrons (this species is also called the reducing agent)
Reduction – the gain of electrons (this species is also called the oxidizing agent)
In the reaction considered here:
Cu(s) + 2 Ag+ (aq)  Cu 2+ (aq) + 2 Ag(s)
Cu is oxidized (and is the reducing agent) because it loses two electrons to Ag.
Ag+ is reduced (and is the oxidizing agent) since it gains electrons from the Cu.
This reaction occurs because Cu is more easily oxidized than Ag.
Notes on Oxidation-Reduction Reactions (continued)
The reverse reaction:
Cu 2+ (aq) + 2 Ag(s)  Cu(s) + 2 Ag + (aq)
does not occur since the Cu is more readily oxidized than Ag – it will not
accept electrons from silver.
A fairly easy way to consider whether or not an oxidation-reduction
reaction occurs is to look at an activity series. You will find one in the
Appendices of your lab book or on page 136 of your textbook. The
following slide discusses the use of the activity series to determine in which
direction a redox reaction will occur.
Use of the Activity Series to Predict Redox Reactions
A small portion of an activity series is included below. The reactions are
written as oxidations - the species on the left gives up an electron. Those
higher on the list are more easily oxidized.
Notice that copper is above silver in the activity series. This means that
copper is more easily oxidized and, when the copper is in the presence of
silver ions, copper will donate electrons to silver as we have already seen.
Since copper is above silver, when writing the chemical equation for the
reaction between the two the silver reaction will be reversed.
Cu(s) + 2 Ag+ (aq)  Cu 2+ (aq) + 2 Ag(s)
Further Examples of Redox Reactions
Consider the reaction of iron and lead. Since
iron is higher in the activity series, its equation
will remain as written and the nickel chemical
equation will be reversed.
Fe + Ni2+ → Fe2+ + Ni
Balancing the redox equations requires ensuring the charges are the same on both
sides of the equation. (This always needs to be true – think about your net ionic
equations.) Consider the reaction between chromium and lead. The chromium
equation stays as written while the lead equation is reversed.
Cr  Cr3+ +3e- and Pb2+ + 2e-  Pb
Electrons that are given up (3 per chromium) must be accepted by lead (2 per
lead). Since electrons cannot be left floating around, one looks for the common
multiple of 3 and 2, which is six. Thus, the chromium equation is multiplied by 2
and the lead equation by 3 to get:
2Cr + 3 Pb2+  2Cr3+ + 3 Pb
End of Slide Show