Redox
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Oxidation—Reduction Reactions
(Answer questions Turn in the last two pages of this lab handwritten.)
Objectives
• Observe and record oxidation—reduction reactions between aqueous ions.
• Describe oxidation and reduction processes by writing half-reactions.
• Write and balance redox equations by balancing half-reactions.
• Identify reactants as oxidizing agents and reducing agents.
Introduction
You saw in the last two labs that oxidation—reduction reactions commonly occur when metals react with
either acids or metal ions. Redox reactions also occur between ions, one ion being oxidized and the other
being reduced. For example, permanganate ions, MnO4, undergo a redox reaction with iodide ions, I-, in the
presence of acid:
16 H+ + 2 MnO4- + 10 I-  2Mn2+ + 5 I2 + 8 H2O
This reaction is easily seen in the laboratory because permanganate has a distinctive purple color that
disappears upon formation of the nearly colorless Mn2+ ion. Similarly, the colorless iodide ion reacts to form
the iodine molecule, which appears yellow in aqueous solution. The reaction mixture turns from purple to
yellow as the reaction proceeds.
In this reaction, permanganate is reduced and iodide is oxidized. In other words, manganese gains electrons
and iodine loses electrons. Because permanganate accepts electrons from iodide and causes iodide to be
oxidized, permanganate, the reduced reactant, is also called the oxidizing agent. Similarly, because iodide
donates electrons to permanganate and causes permanganate to be reduced, iodide, the oxidized reactant, is
also called a reducing agent.
Purpose
In this lab you will carry out a number of redox reactions involving ions in aqueous solution and write
balanced equations to describe the reactions. You will also identify the oxidizing and reducing agents in each
case and become familiar with which substances usually act as oxidizing agents and which are most often
reducing agents.
Safety
• Wear your safety glasses.
• Use full small-scale pipets only for the carefully controlled delivery of liquids.
Materials
Small-scale pipets of the following solutions:
potassium iodide (KI)
sodium hypochlorite (NaOCl) (BLEACH on counter)
potassium permanganate (KMnO4)
potassium iodate (KIO3)
hydrogen peroxide (H2O2)
sodium hydrogen sulfite (NaHSO3)
sodium nitrite (NaNO2)
potassium dichromate (K2Cr2O7)
copper (II) sulfate (CuSO4)
iron (III) chloride (FeCl3)
hydrochloric acid (HCl)
starch solution
sodium hydrogen sulfate (NaHSO4)
sodium thiosulfate (Na2S2O3)
Redox
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Experimental Page – Use a file protector sheet in the red box with other sheets under the
counter in the back.
1. Mix 1 drop each of the indicated solutions in the spaces below. If no reaction is visible, add starch, an
indicator for I2. If no reaction occurs still, add HCl. Record your results in Table 36.1. Put this page
in a file protector sheet.
2. Add a few drops of NaHSO3 to each of the mixtures in Step 1, a—d. Record your results in Table 36.1.
HSO3- + I2  SO42- + I3. Add a few drops of Na2S2O3 to each of the mixtures in Step 1, e—h. Record your results in Table 36.1.
S2O32- + I2  S4O62- + I-
Redox
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Experimental Data
Record your results in Table 36.1 or a table like it in your notebook.
Cleaning up
Avoid contamination by cleaning up in a way that protects you and your environment. Carefully clean
the reaction surface by absorbing the contents onto a paper towel, wipe it with a damp paper towel, and
dry it. Dispose of the paper towels in the waste bin. Wash your hands thoroughly with soap and water.
Redox
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Questions for Analysis
Use what you learned in this lab to answer the following questions.
1. What color is aqueous I2? What color is I2 in the presence of starch?
2. Listed below are the reactants and products for each mixing. Assign oxidation numbers to each
chemical species that changes oxidation number and then write half reactions for each reaction.
a. ClO- + I-  Cl- + I2
e. Cr2O72- + I-  Cr3+ + I2
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b. MnO4- + I-  Mn2+ + I2
f. NO2- + I-  NO(g) + I2
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c. H2O2 + I-  H2O + I2
g. Cu2+ + I-  CuI(s) + I2
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d. IO3- + I-  I- + I2
h. Fe3+ + I-  Fe2+ + I2
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(Remember: Half-reactions must be balanced only for the chemical species that gains or loses electrons.)
3. What color did the mixtures turn upon addition of NaHSO3? Na2S2O3? What does each indicate?
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4. Given the reactants and products for Steps 2 and 3, write half- reactions for each: (I have started one of
the half reactions you need to finish it.)
HSO3- + I2  SO42- + I-
S2O32- + I2  S4O62- + I-
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________2S2O3-2  S4O6-2 + 2e-_________
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Redox
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5. Fill in the table below with all of the oxidizing agents and reducing agents that you observed in this
experiment.
Oxidizing agents
Reducing agents
Oxidizing agents
Reducing agents
6. Complete and balance redox equations for each reaction you observed. Multiply each half-reaction by
coefficients that balance the electrons gained and lost, and add the half-reactions. For example,
Add water to one side of the equation to balance the oxygen atoms, and add H + to the other, side to balance
the hydrogens:
16H+ + 2MnO4- + 10I-  2Mn+2 + 5I2 + 8H2O
Check to see that all atoms are balanced and that the total charge is balanced.
a. _________________________________________________________________________________
b. _________________________________________________________________________________
c. _________________________________________________________________________________
d. _________________________________________________________________________________
e. _________________________________________________________________________________
f. _________________________________________________________________________________
g. _________________________________________________________________________________
h. _________________________________________________________________________________
1. Recall that an oxidizing agent is a chemical that causes another chemical to be oxidized, and in the
process the oxidizing agent is reduced (gains electrons). Similarly, a reducing agent is a chemical that
causes another chemical to be reduced, and in the process the reducing agent is oxidized (loses
electrons.) For example, in this lab you observed that ClO- is an oxidizing agent because each Cl
atom is reduced from an oxidation number of +1 to - 1. Also, I is a reducing agent because each I
atom is oxidized from an oxidation number of -1 to 0.
ClO- + 2e-  Cl-
2I-  I2 + 2e-