SK025

CHEMISTRY

LAB REPORT

EXPERIMENT 3: ELECTROCHEMICAL CELLS

DARREN LEE HUA EE

MS2225208152

K2T4A

Objective 1. To arrange Al, Zn, Mg, Fe and Cu in an electrochemical series.

2. To determine the Faraday’s constant by electrolysis of 𝐶𝑢𝑆𝑂4
solution.
Introduction Electrochemistry is a study of the relationship between electricity and
chemistry. Generally, there are two types of electrochemical cells which are
galvanic and electrolytic cells. A galvanic cell is an electrochemical cell in
which redox reaction occurs spontaneously to generate electricity. For a
galvanic cell, oxidation occurs at the anode and electrons flow to the cathode
where reduction occurs.
Half-Equation 𝑟𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛: 𝐶𝑢2+ (𝑎𝑞) + 2𝑒 − → 𝐶𝑢 (𝑠)
𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛: 𝑍𝑛(𝑠) → 𝑍𝑛2+ (𝑎𝑞) + 2𝑒 −
𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛: 𝑀𝑔(𝑠) → 𝑀𝑔2+ (𝑎𝑞) + 2𝑒 −
𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛: 𝐴𝑙(𝑠) → 𝐴𝑙 3+ (𝑎𝑞) + 3𝑒 −
𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛: 𝐹𝑒(𝑠) → 𝐹𝑒 2+ (𝑎𝑞) + 2𝑒 −
𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛: 2𝐻2 𝑂(𝑙) → 𝑂2 (𝑔) + 4𝐻 + (𝑎𝑞) + 4𝑒 −
Procedure of Part (A): 1. The metal strips were cleaned with the sandpaper.
2. A 50 ml beaker was filled with 35 ml of 0.1M 𝐶𝑢𝑆𝑂4 and the other
beaker was filled with 35 ml of 0.1M 𝑍𝑛𝑆𝑂4 .
3. The apparatus was set up as in the figure below.

4. The cell potential was recorded.

5. Steps 1 till 4 were repeated by replacing 𝑍𝑛2+ /𝑍𝑛 half cell with a
a. Magnesium strip in 0.1M 𝑀𝑔𝑆𝑂4
b. Aluminium strip in 0.1M 𝐴𝑙(𝑁𝑂3 )3
c. Iron strip in 0.1M 𝐹𝑒𝑆𝑂4
 6. The metals were arranged in ascending order of strength as reducing
agents.
7. The order was verified by calculating the standard reduction
𝑜
potential, 𝐸𝑟𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛 of each electrode.
Procedure of Part (B): 1. A copper electrode was cleaned with a piece of sandpaper.
2. The copper electrode was weighed accurately.
3. The apparatus was set up as in below figure. A 50ml beaker was filled
with 35ml 1.0M 𝐶𝑢𝑆𝑂4 .

4. The circuit was completed by connecting the wires from each

electrode to the ammeter and transformer. The transformer was set
to supply the direct current with a voltage of 3V.
5. The electrolysis was run for 15 minutes.
6. The ammeter reading and observation of each electrode were
recorded.
7. The circuit was disconnected and the exact time of electrolysis was
recorded.
8. The copper strip was dried using a hair dryer.
9. The copper strip is weighed again.
10. The mass of copper deposited was calculated. The Faraday’s
constant was calculated.
Result and Galvanic cell
Observation of Part
(A): Galvanic cell Cell Potential, 𝐸𝑐𝑒𝑙𝑙 (V)
Anode Cathode
Zn Cu 0.50
Mg Cu 0.80
Al Cu 0.60
Fe Cu 0.30
Result, Observation
and Calculation of Electrode Observation
Part (B): Cathode Brown solid is deposited on the
copper strip.
Anode Gas bubbles are formed.

Final mass of Cu electrode (g) 0.340

Initial mass of Cu electrode (g) 0.270
Mass of Cu deposited (g) 0.340 − 0.270 = 0.070
Moles of Cu (mol) 0.070
= 0.0011
63.6
Ammeter reading (A) 0.3
Time (s) 904.88
Quantity of charge, Q (C) 0.3 × 904.88 = 271.464

Calculation:
𝐶𝑢2+ (𝑎𝑞) + 2𝑒 − → 𝐶𝑢(𝑠)

2𝑒 − ≡ 2𝐹 ≡ 1 𝑚𝑜𝑙 𝑜𝑓 𝐶𝑢

0.0011 × 2 = 0.0022𝐹

271.464𝐶 ≡ 0.0022𝐹

271.464
1𝐹 =
0.0022
= 123392.7273 𝐶

Discussion 1. Reasons of errors:

a. When the circuit is closed for a period, the temperature of the
circuit increases thus the effective resistance of the circuit
increases.
b. The oxide layer on the metal strip is not fully removed.
c. The 𝐶𝑢𝑆𝑂4 solution for electrolysis is not pure.
2. Precautionary steps:
a. The switch is turned on after the apparatus is completely set up.
b. The door of weighing scale is closed while weighing the copper
strip.
c. The eye must be perpendicular to the reading scale of the
ammeter to avoid parallax error.
Conclusion From this experiment,

1. Al, Zn, Mg, Fe and Cu were able to be arranged in an electrochemical

series which is Cu, Fe, Zn, Al and Mg (in ascending order).
2. The Faraday’s constant was able to be determined by electrolysis of
𝐶𝑢𝑆𝑂4 solution which is 123392.7273C.