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Quantitative Determination of the Purity and Dissociation Constant of Potassium Hydrogen Phthalate by Potentiometric Titration

By Nicole-Picart Apr 22, 2015 798 Words
DETERMINATION OF ELECTRODE POTENTIALS

N.M. PICART1 and A.J. EDUARDO2
1INSTITUTE OF BIOLOGY, COLLEGE OF SCIENCE
2INSTITUTE OF BIOLOGY, COLLEGE OF SCIENCE
UNIVERSITY OF THE PHILIPPINES, DILIMAN, QUEZON CITY 1101, PHILIPPINES DATE SUBMITTED: 22 APRIL 2015
DATE PERFORMED: 15 APRIL 2015

ANSWERS TO QUESTIONS

1. Discuss the differences between galvanic and electrolytic cells. A galvanic cell uses a spontaneous reaction to generate electrical energy. In the cell reaction, some of the difference in free energy between higher energy reactants and lower energy products is converted into electrical energy, which operates the load (surroundings) – flashlight, MP3 player, car starter motor, and so forth. Thus, the system does work on the surroundings. All batteries contain galvanic cell. Moreover, the direction of electron flow is from anode to cathode, and two containers are needed. An electrolytic cell uses electrical energy to drive a nonspontaneous reaction. In the cell reaction, an external source supplies free energy to convert lower energy reactants into higher energy products. Thus, the surroundings do work on the system. Electroplating and the recovery of metals from ores utilize electrolytic cells. Moreover, direction of electron flow is from anode to cathode so a direct current source is needed to reverse the flow, and only one container is needed.

2. What is the measured value of Ecell for each cell and its relevance to the EOred for the iron half-cell and halogen half-cells? The measured value of Ecell for the zinc, iron, bromine, and iodine half-cells were -0.71, 0.61, 0.40, and 0.51 respectively. The standard cell potential is the difference between the standard electrode potential of the cathode (reduction) half-cell and the standard electrode potential of the anode (oxidation) half-cell. Thus, to get the EOred, EOcell (which can be computed using the Nernst Equation) should be added to EOoxid (or EOreference).

3. Compare the EOred for each half-cell with their literature value.
The % error of the standard reduction potential of zinc, iron, bromine, and iodine half-cells as compared to their book values were 6.95%, 20.88%, 63.20%, and 17.07% respectively. The differences between the zinc, iron and iodine half-cells, and their book values were relatively small unlike the bromine half-cell, which has a large % error.

4. Based on your observations, what are the half reactions which occur at the anode and cathode during the electrolysis.
Anode: 2Ha- ⟶ Ha2 +2e- (Ha=Halide)
Cathode: 2H2O + 2e- ⟶ 2OH- + H2 (Effervescence)

5. Based on the calculated values, what are the most effective reducing and oxidizing agent? The higher the EOcell value, the greater is its ability to undergo reduction; thus, it is a more effective oxidizing agent. Based on the calculated values from the results of the experiment, the most effective reducing agent is Zn while the most effect oxidizing agent is Fe3+. However, there are discrepancies between the experimental and theoretical results. Theoretically, the most effective reducing agent is Zn while the most effective oxidizing agent is Br2.

6. Are values of standard reduction potentials helpful in determining the spontaneity of a reaction? Electrons flow spontaneously from the negative to the positive electrode, which is toward the electrode with the more positive electrical potential (anode to cathode). Thus, when the cell operates spontaneously, there is a positive cell potential. A positive Ecell arises from a spontaneous reaction. The more positive it is, the more work the cell can do, and the farther the reaction proceeds to the right. A negative Ecell is associated with a nonspontaneous cell reaction. If Ecell = 0, the reaction has reached equilibrium and the cell can do no more work.

7. What are the possible sources of errors and their effect on the calculated parameters? Rationalize. Uncertainties in the measurements of the volumes and weights of the components of the experiment can cause errors. The uncertainties of the mass of the samples can also contribute to the errors in the solution. The uncertainties in measurements of the solutions can cause increase/decrease in the molarities of the solutions and can change the calculated EOcell. Another source of error is a faulty multimeter which could cause various changes on the voltage obtained. More sources of errors are the orientation of the electrode such that negative readings were obtained and upon switching these electrodes an unexpected mixing of the chemicals occurs in a way that the reading obtained deviates from the supposed reading.

REFERENCES

[1] Silberberg, M. (2007). Principles of General Chemistry. Boston: McGraw-Hill Higher Education.

[2] Petrucci, R. (2011). General Chemistry Principles and Modern Applications (10th ed.). Toronto, Ontario: Pearson Canada

[3] Skoog, D., West, D. Holler, F.J., Crouch, S. (2014). Fundamentals of Analytical Chemistry (9th ed.). Brooks/Cole. USA d Calculations

Determination of Standard Reduction Potentials

I. Zn2+|Zn

II. Fe3+|Fe2+|C

III. Br -|Br2|C

IV. I-|I2|C

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