KWAME NKRUMAH UNIVERSITY OF SCIENCE ANDTECHNOLOGY, KUMASI

DEPARTMENT OF CHEMISTRY

YEAR TWO (CHEM 269)

PRACTICAL CHEMISTRY III

TITLE: ELECTROCHEMISTRY

NAME: OPOKU ERNEST

EMAIL: ernopoku@gmail.com

DATE: 29TH SEPTEMBER, 2013

TITLE: ELECTROCHEMISTRY

AIMS AND OBJECTIVES:1. To compare the measured and calculated values of the EMF of the Cu2+/Cu half cell

at different concentration of Cu2+.

2. To find the end –point of the sodium chloride-silver nitrate titration by EMFmeasurements.

3. To determine the equilibrium constant of a reaction, from measurement of the EMF ofa cell when the- reaction is not at equilibrium.

INTRODUCTIONElectrochemistry is concerned with the interrelation of electrical and chemical effects.Reactions involving the reactant – the electron. Chemical changes caused by the passage ofcurrent. An electrochemical system is not homogeneous but is heterogeneous.Electroanalytical chemistry encompasses a group of quantitative analytical methods thatare based upon the electrical properties of an analyte solution when it is made part of anelectrochemical cell. These methods make possible the determination of a particularoxidation state of an element.

Ox + ne- ↔ Red

There are two general types of electrochemical methods: potentiometric (no current,equilibrium potential) and voltammetric (current measured as a function of the appliedpotential.Electrochemical cells consist of two electrodes: an anode (the electrode at which theoxidation reaction occurs) and a cathode (the electrode at which the reduction reactionoccurs).

Cu(s) + Zn+2 ↔ Cu+2

+ Zn(s) Cu(s) ↔ Cu+2

+ 2e-

(oxidation) Zn+2

+ 2e- ↔ Zn(s) (reduction)

There are two types of electrochemical cells: galvanic (ones that spontaneously produceelectrical energy) and electrolytic (ones that consume electrical energy).

E = Eo

´ +2.303 RT

[Ox]

nFlog

[Red]

The potential that develops in a cell is a measure of the tendency for a reaction to proceedtoward equilibrium.We use concentrations in the Nernst equation, but really activitiesare the proper term. The activity of a specie can be defined as the ability of aspecies to participate an equilibrium reaction involving itself.

Example: Fe+3

+ e- ↔ Fe

+2FeCl

+2, etc. depends on ionic strength

Ecell = Ecathode – Eanode

∆Grxn = - nFEcell∆Grxn = -RTlnKeq Key equations

All cell potential measurements require two electrodes!

1. AgCl(s) + e- ↔ Ag(s) + Cl-

E = Eo + (0.059/n)log1/[Cl-]

2. Hg2Cl2(s) + 2e- ↔ 2Cl- + 2Hg(l)

E = Eo + (0.059/2)log1/[Cl-]2

n = number of electrons transferred per mole, 2.303 RT/F = 0.059 V

Cu+2 + H2(g) ↔ Cu(s) + 2H+

Zn/ZnSO4 (aZn+2 = 1.00)//CuSO4 (aCu+2 = 1.00)/Cu

Anode (oxidation) Cathode (reduction)Electrode reaction kinetics are affected by the electrode surface cleanliness, surfacemicrostructure, and surface chemistry.

CHEMICALS AND EQUIPMENT1. 1M CuSO4, 0.1M CuSO4, 0.01M CuSO4 and 0.001M CuSO4 solutions2. 1M ammonia solution3. 1M KCl4. Calomel electrode5. Distilled water6. 200mL measuring cylinder7. 1mL pipette8. 4 beakers9. Banana clip10. Copper electrode11. Filter paper12. Volt meter13. 8M nitric acid

PROCEDURE

A. The effect of concentration on EMF1. Enough amounts of 1M CuSO4, 0.1M CuSO4, 0.01M CuSO4, and 0.001M CuSO4

solution were prepared and poured in clean dry different beakers.2. The calomel electrode was connected to the middle socket of the multimeter. The

rubber cap was removed from the tip of the electrode.

3. A crocodile clip lead was connected into the bottom socket of the meter. A cleanpiece of copper was placed into a crocodile clip to form a copper electrode.

4. The copper electrode was cleaned by immersion of the copper in 8M AgNO3

provided in a fumed hood followed by rinsing with water.5. A copper electrode was dipped into the beaker containing 0.001M sulphate solution

and the EMF recorded with the help of the voltmeter at 200mV range.6. The electrode and the beaker were rinsed with distilled water and the calomel

electrode dipped into a solution of 8M nitric acid.7. The procedure was repeated for the other concentrations of the sulphate solution and

the results tabulated.

TABLE OF RESULTS

Concentration of CuSO4

( mol/L)Measured Ecell

(mV)0.001 73.930.01 100.900.1 104.971.0 119.73

CALCULATIONelectrode potential of the saturated calomel electrode = 0.274V at 18oCThe reaction that occurred is as shown belowCu2+ + 2e- Cu LHECu Cu2+ + 2e- RHECu2+ + Cu Cu + Cu2+ cell reactionFrom the equationE(Cu2+/Cu)= 0.247-E(cell)

E(Cu2+/Cu) (0.001M) = 0.247- 73.93×10-3

=0.17307V

E(Cu2+/Cu) (0.01M) = 0.247- 100.90×10-3

=0.14610V

E(Cu2+/Cu) (0.1M) = 0.247- 104.97×10-3

=0.14203VE(Cu2+/Cu) (1.0M) = 0.247- 119.73×10-3

=0.12727V

Concentration of CuSO4

( mol/L)E(cell)/mV E(Cu2+/Cu)/V E(Cu/Cu2+)/V

0.001 73.93 0.17307 -0.173070.01 100.90 0.14610 -0.14610

0.1 104.97 0.14203 -0.142031.0 119.73 0.12727 -0.12727

FromE= Eo -0.059 log [reduced form]

n [oxidized form]where n= number of electrons in one half-cell= 2Eo= the standard electrode potential of the copper electrode = 0.337V[reduced form]= 1M

Concentration of CuSO4

(mol/L)[reduction form][oxidation form]

0.059Log [reduction form][oxidation form]

E/V

1.000 1 0.0000 0.33700.100 10 0.059 0.27800.010 100 0.118 0.21900.001 1000 0.177 0.1600

B. Determination of an Equilibrium constant1. 50mL of 1M ammonia solution and 2mL of 0.01M copper sulphate solution was

measured into a beaker and swirled. A reasonable amount was afterwards poured intoa beaker.

2. A reasonable amount of 0.1M of copper sulphate was put into another cleanedbeaker.

3. Two copper electrodes were dipped into the two solutions.4. A filter paper was soaked with 1M KCl and used as salt bridge.5. The EMF of the two half-cells was taken at 2.5V range and the result recorded.

The recorded EMF was 0.394V at 250mV range, hence the voltage was 250mV

Cu/Cu2+/ (salt bridge) Cu (NH3)42+/Cu.

FromE(cell)= Eo

(cell) – 0.059 log [Cu2+][NH3]4

2 [Cu(NH3)42+]

Eo(cell)=E(cell) + 0.059 log [Cu2+][NH3]4

2 [Cu(NH3)42+]

Eo(cell)= 0.025+ 0.0295 log (0.1)(0.96)4 = 0.0625V=62.5 mV

3.85×10-4

log K = 2Eo(cell)/0.059 = 2×0.0625/0.059 = 2.1186

K = log-1 2.1186 = 131.401

DISCUSSIONBy carefully examining the EMF values obtained, it was observed that as the concentration ofthe copper sulphate solution increases the EMF values of the cells also increases. Thisobservation is due to the fact that there are more ions in solutions of higher concentrations

than solutions of lower concentration which consequently led to the attainment of suchresults. The results calculated for E (Cu/Cu2+) are negative because every reduction half-cellreaction has a corresponding standard reduction potential Eored. If the reaction is reversed, itbecomes an oxidation half-cell with a corresponding standard oxidation potential, Eoox. Forany given half-cell reaction, Eoox = -Eored. When the equilibrium constant K>>>1, theequilibrium position of the reaction lies to the right and favors the products and vice versa.Thus the equilibrium constant of this experiment showed that the reaction favors theproduction of the Cu2+ ions in solution since the calculated value for K>>>1.

PRECAUTIONS1. A crucial care was taken to avoid the electrode getting into contact with the solution in

order to obtain accurate values of the EMF.2. The salt bridge was well immersed into both of the Cu (NH3)4

2+ solution and the CuSO4

solutions.3. The copper electrode was immersed in 8M nitric acid followed by rinsing after each

use.4. Reading of values from the pipette were taken from the eye level to avoid parallax

errors.

CONCLUSIONFrom close observation of the values obtained from the experiment and making referencefrom the understated references, EMF of a cell increases as the concentration of the solutionincreases. Also the equilibrium constant of the reaction can be determined from its EMFvalues.

REFERENCES1. General Chemistry, Raymond Chang, 5th edition, pages 630-640.2. Chemical Technician Ready Reference Handbook 2nd edition by Shugar and

Bauman pages 592- 593.3. Health Chemistry Laboratory Experiment 2nd edition by Micheal .A. Dispezio

page 310-312.4. Textbook of Practical Chemistry 4th edition (2001) by Wesley. D. Smith pages 620-633.5. Practical Chemistry 4th edition (1990) by Francis Kinkier pages 745-748.