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Exercise 8 Page 1

Illinois Central College 132 Name:______Laboratory Section: ______Redox Titration

Equipment

1-25.00 mL burette 0.100 N KMnO4

2-50 mL beakers KHSO3 solution of unknown Normality

Objectives The objectives of this experiment are to develop an understanding of oxidation-reduction titration and the use of the "eqivalent" concept as applied to oxidizing and reducing agents.

Background Oxidation Reduction reactions are chemical reactions in which substances undergo changes in oxidation state.

Oxidation is defined as the loss of electrons (or an increase in oxidation state) and reduction as the gain of electrons (or a decrease in oxidation state). In acid base titrations, equivalent amounts of acid and base must be used for exact neutralization at the titration endpoint. In oxidation-reduction reactions, there is a similar equivalence between oxidizing and reducing agents.

In acid base reactions, one "equivalent" corresponded to 1 of H+1 (for an acid) or 1 mole of OH-1 (for a base). In oxidation-reduction reactions, one "equivalent" refers to 1 mole of electrons either provided (by a reducing agent) or taken (by an oxidizing agent). So, just as with acids and bases, one equivalent of a reducing agent will reduce one equivalent of an oxidizing agent. At the endpoint of a titration;

number of equivalents of oxidizing agent = number of equivalents of reducing agent. Oxidizing Agent: the substance which takes up electrons Reducing Agent: the substance which gives up electrons Alternatively, we can write:

"milliequivalents" of oxidizing agent = "milliequivalents" of reducing agent

where a "milliequivalent is 1/1000 of an equivalent. Exercise 8 Page 2

Consider the reaction of potassium permanganate with oxalic acid in the presence of sulfuric acid. The balanced chemical equation and net ionic equations are;

2 KMnO4 + 5 H2C2O4 + 3 H2SO4 10 CO2 + 2 MnSO4 + 8 H2O

(oxidation: loss of 2 e-1) x 5 = 10 e-1

-1 +1 +2 2 MnO4 + 5 H2C2O4 + 6 H 10 CO2 + 2 Mn + 8 H2O

(reduction: gain of 5 e-1) x 2 = 10 e-1

The molecular mass of KMnO4 is 158.0 g/mol

Since one mole of KMnO4 actually removes 5 moles of electrons in this redox reaction, we can say that 1 mol KMnO4 = 5 equivalents of KMnO4 (at least, in this particular reaction).

This allows us to define the "" of a compound, that is, the mass which would provide 1 equivalent of oxidizing power. For KMnO4, that would be;

158.0 grams KMnO ( 4 )x(1 mol KMnO4 ) = 31.60 grams/equivalent 1 mol KMnO4 5 equivalents

The molecular mass of oxalic acid is 90.0 grams. Since one mole of oxalic acid actually provides

2 moles of electrons in this redox reaction, we can say that 1 mol H2C2O4 = 2 equivalents of

H2C2O4 (again, in this particular reaction).

So, the equivalent weight of this reducing agent can also be calculated, that is, the mass of ocxalic acid that would provide 1 equivalent of reducing power. For oxalic acid, that would be;

90.0 grams H C O ( 2 2 4 )x(1 mol H2C2O4 ) = 45.0 grams/equivalent 1 mol H2C2O4 2 equivalents

This leads us to the logical end of defing the "Normality" of oxidizing and reducing agents just as we did with acids and bases. Since different oxidizing and reducing agents can provide various numbers of equivalents, we simply redefine the strength of an oxidizing or reducing agent as "the number of equivalents per liter" (or milliequivalents per milliliter) ,or Normality. N = milliequivalents of oxidizing agent milliequivalents of reducing agent oxidizing agent milliliter Nreducing agent = milliliter

Since the endpoint of a redox titration demands that the equivalents of oxidizing agent equal the equivalents of reducing agent, then;

Noxidizing agent x mLoxidizing agent = Nreducing agent x mLreducing agent Exercise 8 Page 3

Safety Precautions

Safety goggles must be worn in the lab at all times. Any skin contacted by chemicals should be washed immediately.

Procedure In this experiment, you will determine the normality of the reducing agent, potassium

bisulfite (KHSO3) by titrating with a standard 0.100 N KMnO4 solution.

1. Set up a 25.00 mL burrette and fill it with your unknown of reducing agent

(KHSO3). Be sure to record the letter of your unknown on your Report Sheet.

2. Draw 15-25 mL of the 0.100 N KMnO4 solution (from the dispensing table) into a clean, dry 50 mL beaker. Carefully read the burette to the nearest 0.01 mL before and after the withdrawal and record these values on your Report Sheet.

3. Add approximately 5 mL of concentrated (12 M) H2SO4 to your permanganate solution. If the solution turns brown during the titration with your reducing agent, add additional acid.. 4. Dispense the reducing agent into your permanganate solution while stirring drop by drop as the endpoint is approached, until one drop of the reducing agent completely decolorizes the permanganate. (White paper or some other white background makes the color change more discernable.) Record your final burette reading on your Report Sheet.

5. Starting with a different volume of permanganate, repeat the titration for Trial 2.

6. Calculate the Normality of the KHSO3 and record it on your Report Sheet. 7. Report the average Normality on your Report Sheet.

8. Write a balanced oxidation-reduction equation for the reaction of KMnO4 with KHSO3. The unbalanced equation is:

KMnO4 + KHSO3 + H2SO4 MnSO4 + KHSO4 + K2SO4 + H2O

9. Based on their Normalities, calculate the weights of KMnO4 and KHSO3 present in one liter of each solution and record these values on your Report Sheet. Exercise 8 Page 4 Exercise 8 Page 5

Illinois Central College CHEMISTRY 132 Name:______Laboratory Section: ______REPORT SHEET Redox Titration

Unknown Number ______

Oxidizing Agent: KMnO4 Final burette reading

Initial burette reading

Volume used

Normality of KMnO4 0.100 N 0.100 N

Reducing Agent: KHSO3 Final burette reading

Initial burette reading

Volume used

Normality of KHSO3

Average Normality

1. Write the balanced redox equation for the reaction of KMnO4 with KHSO3.

(over) Exercise 8 Page 6

2. Mass of KMnO4 per liter in a 0.100 N solution. ______g/L

3. Mass of KHSO3 per liter ______g/L (based on your experimental normality)

Show All Calculations Below: Exercise 8 Page 7

Illinois Central College CHEMISTRY 132 Name:______Laboratory Section ______

PRELAB: Exp. 8 Redox Titration

1. Balance the following Redox reactions.

a) _____Zn + _____AgNO3 _____Zn(NO3)2 + _____Ag

b) _____Sn + _____HNO3 _____SnO2 + _____NO2 + _____H2O

c) _____MnO + _____PbO2 + ____HNO3 ____HMnO4 + ____Pb(NO3)2 + ____H2O

d) _____Zn + _____NaNO3 + _____NaOH _____Na2ZnO2 + _____NH3 + ____H2O

e) ___K2Cr2O7 + ___H2S + ___H2SO4 __K2SO4 + __Cr2(SO4)3 + ___S + ____H2O

f) ___MnCl2 + ___PbO2 + ___HNO3 ___HMnO4 + ___Cl2 + ___Pb(NO3)2 +___H2O Exercise 8 Page 8