Dr. White
Chem 1B
Saddleback College
1 Experiment 19 – Electrolytic Cells (Demonstration)
Objectives
 To learn about electrolytic cells and how they
differ from galvanic cells.
 To learn about Faraday’s Law.
 To learn about electrochemical plating.
 To explore the relationship between current
and mass for a chemical reaction.
Introduction
Electrolytic cells use batteries or electricity to power
reactions that are not normally spontaneous. One
use of electrolytic cells is for electroplating, the
process of coating the surface of a metal object with
a layer of a different metal through electrochemical
means. In this experiment, an electrolytic cell will
be used to plate nickel onto copper.
The
electrolytic cell will be set up as shown below.
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metal plated onto the copper, Faraday’s law must
be used. Faraday’s Law establishes a relationship
between the quantity of electricity (in Coulombs, C)
and the quantity of electrons (in moles). Faraday’s
Law states that 96485 C = 1 mole e . The number
of Coulombs provided by an electrolytic cell is
determined by its amperage and the number of
seconds it operates:
amperes x seconds =
Coulombs.
For example, if an electrolytic cell that plates
chromium was operated for 2.0 minutes at 0.25
amperes, the Coulombs = 120 s x 0.25 amperes =
30. Coulombs. Then, Faraday’s Law can be used
to determine the mass of chromium plated as
follows:
30C ×
1 mol e- 1 mol Cr 52.0 g Cr
×
×
= 0.0054g Cr
96485C 3 mol emol Cr
The amount of actual metal plated compared to the
theoretical (stoichiometric) mass plated that is
calculated by calculating the current efficiency, as
shown by the following equation:
Current Efficiency =
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In our experiment, the anode is a strip of nickel
metal and the cathode is a strip of copper metal.
The cathode is hooked up to the negative pole of
the power supply and the anode is hooked up to the
positive poll of the power supply. As the cell
operates the electrons are forced through the
external circuit from the anode to cathode by the
power supply. The following equations describe the
reactions in the cell:
2+
Anode: Ni(s) → Ni (aq) + 2e
2+
Cathode: Ni + 2e → Ni (s)
-
Nickel ions form at the anode, migrate through the
solution, and deposit at the cathode as nickel solid.
As the cell runs, the nickel anode dissolves into
nickel ions and electrons and nickel plates onto the
copper cathode.
In order to calculate the expected amount of nickel
Actual Mass
x 100%
Theoretical Mass
Procedure
The class will perform the experiment together.
1. The instructor will immerse a copper strip in 6 M
HCl for about 3 minutes and rinse it with deionized
water. Then, the instructor will dry it and weigh it.
Write down the mass.
Write down your
observations of the unplated copper strip.
2. Your instructor will set up the electrolytic cell.
3. Your instructor will turn on the electrolytic cell
and report the current. Write down this value.
4. After the cell has run for about 20 minutes, the
instructor will turn off the cell and report the exact
amount of time the cell operated. Write down this
time.
5. The instructor will weigh the copper strip and
report the mass to the class. Write down this mass.
Also, write down your observations of the plated
strip.
6. Calculate the theoretical mass of nickel plated
and the current efficiency of the cell.
Dr. White
Chem 1B
Saddleback College
2 EXP. 19 Electrolytic Cells:
Pre-Lab Questions
1. In an electrolytic cell that plated zinc, the reactions described by the following equations
occurred:
Anode: Zn (s) → Zn2+ (aq) + 2eCathode: Zn2+ (aq) + 2e- → Zn (s)
The cell operated for 3 minutes and 45 seconds at 315 milliamperes.
a. Calculate the theoretical mass, in milligrams, of zinc deposited.
_________________
b. The actual mass of zinc deposited was 20.5 mg. Calculate the current efficiency.
_________________
Dr. White
Chem 1B
Saddleback College
Name
3 Exp. 19 Electrolytic Cells: Data and Results
Mass of the unplated copper strip
Appearance of the unplated copper strip
Mass of the nickel-plated copper strip
Mass of the plated nickel, in grams
Electroplating current, in milliamperes
Time of electroplating, in minutes and seconds
Appearance of the plated copper strip
Calculate the mass, in milligrams, of nickel that should have been deposited by the
electrochemical cell. Show your calculation below.
Mass of Nickel: _____________
Calculate the current efficiency of the electrolytic cell. Show your calculation below.
Current Efficiency:_____________
Dr. White
Chem 1B
Saddleback College
Follow Up Questions
1. What current would be needed to electroplate 45.0 g of zinc in 3.00 hours from a solution
of zinc nitrate?
2. For problem 1 above, what current would be needed if the current efficiency is 92.0%?
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