CHEM 125 Yale University Oxidation and Reduction Reactions Chemistry Lab Report

ChemCollective: Exploring Oxidation-Reduction ReactionsDr. Jursich
UIC, Department of Chemistry
Learning Objectives:
To identify reducing and oxidizing agents in a redox reaction and their relation to electrons transferred.
To observe series of potential redox reactions and determine relative reactivity of metals as reducing
Experimental Objectives:
Mix various combinations of reducing and oxidizing agents to see which ones lead to reaction. From these
observations, rank the reducing agents in order of increasing strength.
For the reaction of Ag+ and Cu(s), determine the equilibrium constant of this redox reaction and compare
with theoretical value obtained from standard potentials of half-reactions.
Background Information:
Additional background information on redox reactions and relative reactivity may be obtained from
following resources.
Chem 124 textbook, Chemistry: the molecular Nature of Matter and Change, Advanced Topics,
Silberberg, Amateis, 8ed , Chapter 4(Section 5), Chapter 21(Sections 3-4), Appendix D for Table of
Standard Reduction Potentials
Go to . You should see workbench page as below. Read over the
information in Redox for background and general experimental approach. However, we will be doing
Parts I and II slightly differently. Follow experimental guidelines below. For your report you’ll need to
write out the procedure used. You can give just one procedure to describe what was done for each
Part I: Qualitatively Determining Reactivity of Metals
In this part, you will want to look for reactions between various metals and metal cations. All cations are
provided as aqueous solutions of nitrate salts. The NO3- will of course be a spectator ion and not part of
the redox chemistry we will examine. Let’s look for reactivity between the metals Cu, Mg, Zn, and Ag and
the different cations Cu2+, Mg2+, Zn2+, Ag+. Of course, there is no reaction between any metal and its own
cation because there is no change of products and reactants after electron transfer. But what about each
metal with the other cations in aqueous solution? That is what will determine by the simulation
Write out a procedure how you would mix the two reactants together (metal + cation). Define what
glassware and amounts are used so another person can repeat what you did. You may write one general
procedure for the mixing of each metal with each cation and then list metals and cations used. Here you
have the advantage that in the virtual world, the stockroom has unlimited amounts of chemicals and we
don’t need to be concern about disposing chemical waste. In the real world, chemical safety, costs and
disposal are key factors in designing chemistry experiments. But you need not concern yourself with that
Summarize your results by preparing a 5×5 table with metals along 1st vertical column and cations along
the 1st row for the header information the table. Then do the experiments mixing the two. In each case,
see if a reaction occurs by examining species in your glassware when the two are mixed in the left-side
panel of display. An example of you wont be doing here – adding Pb(s) to 0.10 M Cu(NO3)2 is shown below.
Note the Cu2+ ion goes to essentially zero (~10-7 M) whereas initial concentration was 0.100 M. So reaction
takes place.
For each combination where no reaction takes place indicate with a “NR” and when reaction takes place
indicate with “Rx”. Do use a fairly large amount of metal ~1 g or more as lesser amounts may not show
reaction when there should be a reaction. Try using different amounts of metal.
From your table list the reducing agents in order of strongest to weakest (that is decreasing reactivity).
Let’s check our experimental results with theoretical prediction. Using a table of standard reduction
potentials, determine E0 for reactions of each of the metals tested with one cation of Cu2+(aq). Show one
example calculation and enter all your calculated E0’s in your table with your observations of all the metals
with Cu2+(aq).
Part II: Quantitative Analysis of the Ag+(aq) and Cu metal reaction
Rather than looking at Ag + Cu2+ let’s examine a more interesting reaction Ag+ and Cu(s). Add Cu(s) to
AgNO3 solution. Make sure you added enough Cu to reach equilibrium. You know you reached equilibrium
when there is unreacted Cu(s) left after reaction and adding more Cu(s) does not change ion
concentrations. Be sure to look for that! Once the Ag+ concentration settles down (it is drifting due to
temperature changes from the reaction), record the concentrations of Cu2+ and Ag+ you observe.
Write out the two balanced half-reactions (without spectator ions, please) involved in this redox reaction
and show how they can be added to give the overall balanced net ionic equation for this reaction.
Determine the theoretical E0 and from E0 determine equilibrium constant, K, for the redox reaction.
In principle, the equilibrium constant can also be calculated from the ion concentrations you observe after
the reaction if one considers the just one simple redox equilibrium itself. From your observed ion
concentrations calculate equilibrium constant, K, for this reaction. Note: Don’t worry if the two
equilibrium constants don’t agree well. Just make sure the calculations are correct. There are other
chemical equilibria taking place which become important at such extremely low concentrations of Cu2+
ions which is beyond the scope of this class.
Lab Report Expectations:
Report to be typed or neatly hand written/printed or combination of both. In the end, the entire
report needs to be composed into a single pdf document.
For your lab report, make sure to organize into the following components as a single file document such
as Word. Then convert it into a single pdf file and upload into Labflow.
Header Information: Your FULL name, Lab section (TA/day/time),
FULL name of expt, date of expt
Purpose: One or two sentences in your own words on the purpose of the experiment.
Procedure: Give general description of procedure used to test for reaction. Indicate the amounts
(concentration or mass) and glassware used. Don’t need to describe for each reaction.
Results: Part I: Have a table summarizing the different combinations tested. Include the E0 values for all
reactions with Cu2+(aq). Have a descriptive title above table just as you would for a graph. Show one
example calculation of E0 value.
Part II: For the reaction with Ag+ and Cu, calculate K based on theoretical E0 of the reaction and
on the Ag+ and Cu2+ ion concentrations given in the program. Show your calculations. The two values will
likely differ quite a bit due to approximations on the calculations which are beyond this course.
Conclusion: List the four metals listed in order of decreasing reducing agent strength (decreasing
reactivity). How does this compare to E0 of the reduction half reaction potentials?
Post Lab Questions:
1) In the Ag+ and Cu(s) reaction:
a) which reactant is oxidized and which reactant is reduced?
b) which reactant is the reducing agent and which is the oxidizing agent?
c) which reactant loses electrons and which gains electrons?
2) Based on a standard reduction potential, determine which metals can be oxidized be 1.0 M HCl?
Mg, Zn, Cu, Sn, Ag, Fe, Pb
3) Based on a standard reduction potential, which species can act as an oxidizing agent for Fe(s)?
Co2+, Sn4+, Al3+, Zn2+, Pb2+
4) Based on a standard reduction potential, which species reduce Sn2+(aq)?
Pb, Cu, Zn, Ag, Al, Cr

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