S5: Chemistry

     Key unit competence

 Predict the products of given electrolytes during 
 electrolysis and work out quantitatively to determine how much is liberated at a 

 given electrode using Faraday’s law.  

Learning Objectives:

• Define electrolysis, cathode and anode.
• Explain the electrolysis of different substances.
• State Faraday’s laws and define the Faraday’s constant.
• Develop practical experimental skills related to electrolysis, interpret results, 
   and draw valid conclusions.
• Carry out a practical activity to explain the phenomenon of electrolysis.
• Compare the electrolysis of dilute solutions and concentrated solutions.
• Calculate the masses and volumes of substances liberated during electrolysis.
• Relate the nature of electrode, reactivity of metal ion in solution to the 
   products of electrolysis.

• Perform electroplating of graphite by copper metal

Introductory activity

Observe carefully the figure below and answer the following question. Record your 

answer and discuss them.

            

 1. Label the set up and give the name of this Experiment.

2. Suggest how water can be decomposed into hydrogen and oxygen.

  13.1. Definition of electrolysis and Description of electrolytic 

               cells.

  Activity 13.1:

 A.
    1. In one case, you have a source of water at the top of a hill and you want 
         to supply water to a community in the valley down the hill
    2. In another case, you have a community at the top of a hill and you 
         want to supply water to the community from a source located in the valley 
        down. Students in groups discuss how they would proceed to supply water 
        to the communities in the above two cases. 
 B. Why do we cook food by heating?
C. What is the difference between a spontaneous reaction and a non
     spontaneous reaction?

D. Have you heard about electrolysis? If yes, can you say what it is about?

     1.Definition of electrolysis.

A spontaneous reaction is a reaction that favors the formation of products without 
external energy. It is a process that will occur on its own. For example, a ball will 
roll down an incline, water will flow downhill, radioisotopes will decay, and iron will 
rust. No intervention is required because these processes are thermodynamically 
favorable.

A non spontaneous reaction (also called an unfavorable reaction) is a chemical 
reaction that necessitates external energy to occur. For example, without an 
external energy source, water will remain water forever. Under the right conditions, 
using electricity (direct current) will help to produce hydrogen gas and oxygen gas 
from water. Cooking foods is not spontaneous reaction that is why heat is used.

Electrolyte: Sodium chloride is an ionic compound in which ions arrange themselves 
in a rigid cubic lattice when in solid state. In this state, it cannot allow electric current 
to pass through it. However, when it is melted, or dissolved in water, the rigid lattice 
is broken, ions are free to move and electric current can pass. Therefore, it is classified 
as an electrolyte.

Substances which cannot allow the flow of electric current when in molten or 
in solution are referred to as non-electrolytes. When electric current (direct 
current) flows through an electrolyte, it decomposes it. This phenomenon is called 

electrolysis.

Thus electrolysis is the decomposition of an electrolyte by passage of an electric 
current through it. Therefore for electrolysis to take place, there must be a source 
of direct current. The direct current is conveyed from its source to the electrolyte 
by means of a metallic conductor and electrodes. The electrode connected to the 
positive terminal of the direct current is called the anode and the one connected to 
the negative terminal is the cathode. By convention, the electric current enters the 

electrolyte by the anode and leaves by the cathode.

When the current passes through an electrolytic solution, ions migrate and electrons 
are gained or lost by ions on the electrodes surface. Electrode that is positively 
charged has deficit of electrons is called anode and the other electrode negatively 
charged has excess of electrons and is called cathode. Chemical changes at the 

electrodes due to the passage of electric current are called electrolysis.

2. Description of electrolytic cells. 

An electrolytic cell is an electrochemical cell that drives a non-spontaneous redox 
reaction through the application of external electrical energy. They are often used to 
decompose chemical compounds, in a process called electrolysis . The Greek word 

lysis means to break up.

Important examples of electrolysis are the decomposition of water into hydrogen 
and oxygen, production of sodium metal, Na, from molten NaCl, production of 
aluminium and other chemicals. Electroplating (e.g. of copper, silver, nickel or 
chromium) is done using an electrolytic cell.

An electrolytic cell has three component parts: an electrolyte and two electrodes 
(a cathode and an anode). The electrolyte is usually a solution of water or other 
solvents in which ions are dissolved. Molten salts such as sodium chloride are also 
electrolytes. When driven by an external voltage applied to the electrodes, the ions 
in the electrolyte are attracted to an electrode with the opposite charge, where 
charge-transferring (also called faradaic or redox) reactions can take place. Only 
with an external electrical potential (i.e. voltage) of correct polarity and sufficient 
magnitude can an electrolytic cell decompose a normally stable, or inert chemical 
compound in the solution. The electrical energy provided can produce a chemical 

reaction which would not occur spontaneously otherwise.

The main components required to achieve electrolysis are:

• An electrolyte is substance containing free ions which are the carriers of 
    electric current in the electrolyte. If the ions are not mobile, as in a solid salt 
    then electrolysis cannot occur.
• A direct current (DC) supply: provides the energy necessary to create or 
   discharge the ions in the electrolyte. Electric current is carried by electronds 
   in the external circuit.
• Electrolysis depends on controlling the voltage and current.   
• Alternating current (AC) would not be appropriate for electrolysis.  Because the 
  “cathode” and “anode” are constantly switching places, AC produces explosive 
   mixtures of hydrogen and oxygen.
• Two electrodes: an electrical conductor which provides the physical interface 

   between the electrical circuit providing the energy and the electrolyte.

                               

The key process of electrolysis is the interchange of atoms and ions by the removal 
or addition of electrons from the external circuit. The products of electrolysis are in 
some different physical states from the electrolyte and can be removed by some 
physical processes. 

Electrodes of metal, graphite and semi-conductor material are widely used. Choice 
of suitable electrode depends on chemical reactivity between the electrode and 
electrolyte and the cost of manufacture.

Note: 
The suitable electrode in electrolysis should be inert (Cu, Pt, etc.) therefore it will not 
participate in the chemical reaction.
It is very easy to be confused about the names CATHODE and ANODE and what their 

properties are, both with electrochemical cells and electrolytic cells.

(To help you to remember, Cathode is the site of reduction, or, if you prefer, CCC = Cathode 

Collects Cations. Anode is the site of oxidation, or, AAA = Anode Attracts Anions.)

Checking up 13.1:

              

13.2. Electrolysis of sodium chloride

  Activity 13.2

Investigating the effect of concentration on the products formed during 
electrolysis of concentrated sodium chloride solution.

Materials: Carbon or graphite rods, connecting wire, U-tube, dry cell, glass 

syringes, concentrated sodium chloride, cork and switch.

                 

            Figure 13.2.1: Electrolysis of concentrated sodium chloride solution 

Procedure:

2. Add 10g of sodium chloride to 100cm3
    of distilled water. 
3. Warm the mixture and continue adding sodium chloride until a 
     saturated solution is formed.
4. Put the saturated solution in U-tube and fit it with carbon rods and glass 
    syringes.
5. Level the brine solution in the two arms and switch on the circuit. 
    Record any observations made after some time. Identify any gases 
    collected in the syringe.

Questions:
a. Identify the gases formed by testing them using litmus papers.
b. Using ionic equations, explain how the products are formed. 

Sodium chloride may be in different forms that can be electrolytes. It may be in its 
molten state, dilute solution or concentrated solution. In each case, the products of 
electrolysis differ because of different factors.

13.2.1. Electrolysis of molten sodium chloride 

The molten salt is introduced in a container called electrolytic cell (or electrolysis 
cell) in which there are two inert electrodes (platinum or graphite). Electrodes are 
connected to a DC generator. 

          • Cations (Na+) move toward the cathode (negative electrode), where they take 

              electrons and are reduced. On cathode metallic sodium is deposited: 

                   Figure 13.2.2: Electrolysis of molten sodium chloride solution

• Another important thing to note is that twice as much hydrogen is produced 
as oxygen. Thus the volume of hydrogen produced is twice that of oxygen. Refer 

to the equations above and note the number of electrons involved to help you 

              13.2.2. Electrolysis of Dilute Sodium Chloride Solution

An aqueous solution of sodium chloride contains four different types of ions. They 
are ions from sodium chloride: Na+ (aq) and Cl-(aq)

Ions from water: H+ (aq) and OH-(aq)

When dilute sodium chloride solution is electrolysed using inert electrodes, the 
Na+ and H+ ions are attracted to the cathode. The Cl- and OH- ions are attracted to 

the anode. 

                                         Table 13.2: Standard reduction potentials.

              

The table  shown below  is simply a table of standard reduction potentials in 
decreasing order. The species at the top have a greater likelihood of being reduced 
while the ones at the bottom have a greater likelihood of being oxidized. Therefore, 
when a species at the top is coupled with a species at the bottom, the one at the top 

will be easily reduced while the one at the bottom will be oxidized.

• At the cathode:

The H+ and Na+ ions are attracted to the platinum cathode. H+ ions gains electrons 
from the cathode to form hydrogen gas. (The hydrogen ions accept electrons more 
readily than the sodium ions. As a result, H+ ions are discharged as hydrogen gas, 
which bubbles off. Explanation why  H+ ions are preferentially discharged will be 

given later.)

                   

• At the anode:

OH- and Cl- are attracted to the platinum anode. OH- ions give up electrons to the 

anode to form water and oxygen gas. 

         

Adding the two reactions and balancing the terms:

         

If we remove 2 molecules of water on both sides, we get:

             

Note: 

• Since water is being removed (by decomposition into hydrogen and oxygen), 
the concentration of sodium chloride solution increases gradually. The overall 
reaction shows that the electrolysis of dilute sodium chloride solution is 

equivalent to the electrolysis of water.

                             

                                               Figure 13.2.3: Electrolysis of dilute sodium chloride solution

13.2.3. Electrolysis of Concentrated Sodium Chloride Solution

The only difference with dilute NaCl solution is that at the anode, Cl-  ions are more 
numerous than OH- ions. Consequently, Cl- ions are discharged as chlorine gas, which 
bubbles off. 

A half-equation shows you what happens at one of the electrodes during 

electrolysis. 

                           

Sodium ions Na+ and hydroxide OH– are also present in the sodium chloride 
solution. They are not discharged at the electrodes. Instead, they make sodium 

hydroxide solution.

These products are reactive, so it is important to use inert (unreactive) materials for 

the electrodes.

One volume of hydrogen gas is given off at the cathode and one volume of chlorine 
gas is produced at the anode. The resulting solution becomes alkaline because 

there are more OH- than H+ ions left in the solution. 

                           

Checking up 13.2: 

With the help of equations of reactions which occur at each 
electrode, outline what happens during electrolysis of dilute aqueous sodium 

chloride. What happens to the pH of the solution as electrolysis continues?

13.3. Electrolysis of water

Activity 13.3: Investigate the products formed during the electrolysis of water

Materials:

• Distilled water
• Tap water
• 2 silver-colored thumb tacks
• 9V battery
• Small, clear plastic container 
• 2 test tubes
• Stopwatch
• Baking soda
• Table salt
• Lemon

• Dish washing detergent

                       

Procedure:

3. Insert the thumb tacks into the bottom of the plastic container so that 
the points push up into the container. Space them so that they’re the 
same distance apart as the two terminals of the 9V battery. Be careful not 
to harm yourself!
4. Place the plastic container with the terminals of the battery. If the cup is 
too large to balance on the battery, be sure thumb tacks are connect to 
positive and negative pushpins and do no touch each other.
5. Slowly fill the container with distilled water. If the tacks move, go ahead 
and use this opportunity to fix them before you proceed. Will distilled 
water conduct electricity on its own? Try it!
6. Add a pinch of baking soda.
7. Hold two test tubes above each push pin to collect the gas being formed. 
Record your observations. What happens? Does one tube have more 
gas than the other? What gases do you think are forming?
8. Discard the solution, and repeat the procedure with a different 
combination: 
• Distilled water and lemon juice
• Distilled water and table salt
• Distilled water and dish detergent
• Distilled water (no additive)
• Tap water (Does tap water works? If so, why?)

Question: During the electrolysis of water, which electrolyte conducts 

electricity the best?

Water can be decomposed by passing an electric current through it.  When this 
happens, the electrons from the electric current cause an oxidation-reduction 
reaction.   At one electrode, called the cathode, electrons cause a reduction.  At the 
other electrode, called the anode, electrons leave their ions completing the circuit, 
and cause an oxidation.

In order to carry out electrolysis the solution  must conduct electric current.  
Pure water is a very poor conductor.  To make the water conduct better we can 
add an electrolyte (NaCl) to the water. The electrolyte added must not be more 
electrolyzable than water.   Many electrolytes that we add electrolyze more easily 
than water.  Sulfate ions do not electrolyse as easily as water, so sulfates are often 
used to enhance the conductivity of the water

Water may be electrolyzed in the apparatus shown below. Pure water is however a 
very poor conductor of electricity and one has to add dilute sulphuric acid in order 

to have a significant current flow.

            

In pure water at the negatively charged cathode, a reduction reaction takes place, 
with electrons (e−) from the cathode being given to hydrogen cations to form 

hydrogen gas. The half reaction, balanced with acid, is:

             Reduction at cathode:

At the positively charged anode, an oxidation reaction occurs, generating 

oxygen gas and giving electrons to the anode to complete the circuit:

          Oxidation at anode:

                                           

             

Combining either half reaction pair yields the same overall decomposition of water 

into oxygen and hydrogen:

             Overall reaction:

The number of hydrogen molecules produced is thus twice the number of oxygen 
molecules. Assuming equal temperature and pressure for both gases, the produced 

hydrogen gas has therefore twice the volume of the produced oxygen gas.

  Checking up 13.3

  I understand the process of water electrolysis, that water as an electrolyte can 
be decomposed into hydrogen and oxygen via an external energy source (an 
electrical current). I know that the reduction of hydrogen takes place on the 
cathode and the oxidation reaction takes place on the anode. I also know 
that water, already partially split into H+ and OH- (though there are very few 
of these ions in pure water). 

1. Electric current (direct current) electrolyzes water. Discuss this 
     statement. 

2. Why alternative current are not used for the same process?

13.4. Electrolysis of concentrated copper (II) sulphate solution 

using inert electrode

Activity 13.4: Investigating what happens when a solution of copper (II) 

 sulphate is electrolysed using carbon and copper electrodes.

Apparatus and chemicals: Glass cell, Carbon rod, 2M copper (II) sulphate 
solution, connecting wires, dry cells, copper plates, propanone and litmus 

paper.

                     

Procedure: 

3. Determine the mass of the graphite rods and record it.
4. Put 0.5M of copper (II) sulphate solution in a glass cell with the carbon 
    (graphite) rods and set up the apparatus. Carefully observe all the 
    changes taking place at the electrodes and the solution. Test the 
     resulting solution with blue litmus paper.
5. After some time, switch off the current, remove the electrodes, wash 
    them in propanone, dry then and then weigh them.
6. Repeat the experiment using clean strips of copper metal as electrodes. 
    Weigh them and then complete the circuit using freshly prepared 

    copper (II) sulphate solution. Record your observations.

Questions: 
 1. Explain the changes observed during the electrolysis of copper (II)
     sulphate using :
              a.Carbon electrodes
              b.Copper electrodes 
2. Outline the changes that occur in the solution from the beginning 

     to the end of the experiment.

The products of electrolysing of copper sulphate solution with inert electrodes 

(carbon/graphite or platinum) are copper metal and oxygen gas.

                    

                     Figure 13.4.1: Simple apparatus to investigate the electrolysis of aqueous solutions

  Using the simple apparatus (diagram above) and inert carbon (graphite) electrodes, 
you can observe the products of the electrolysis of copper sulfate solution are a 
copper deposit on the negative cathode electrode and oxygen gas at the positive 
anode electrode. This anode reaction differs if you use copper electrodes.  You have 
to fill the little test tubes with the electrolyte (dilute copper sulphate solution), hold 
the liquid in with your finger and carefully invert them over the nearly full electrolysis 

cell. The simple apparatus (above) can be used with two inert wire electrodes.

The blue colour fades as more and more copper is deposited, depleting the 

concentration of blue copper ion in solution.

Diagram for the electrolysis of copper (II) sulphate solution with carbon 

electrodes.

              

        Figure 13.4.2: The electrolysis of copper (II) sulphate solution with carbon electrodes.

The electrode reactions and products of the electrolysis of the electrolyte copper 
sulphate solution (with inert carbon-graphite electrodes) are illustrated by the 

diagram above

(a) The electrode products from the electrolysis of copper sulphate with inert 

graphite (carbon) electrodes

The negative cathode electrode attracts ions (from copper sulphate) and 
ions (from water). Only the copper ion is discharged, being reduced into copper 
metal. The less reactive a metal, the more readily its ion is reduced on the electrode 
surface.
A copper deposit forms as the positive copper ions are attracted to the negative 

electrode (cathode)

positive ion reduction by electron gain.

The traces of hydrogen ions are not discharged, so you do not see any gas collected 
above the negative electrode. The blue colour of the copper ion will fade as the copper 
ions are converted into the copper deposit on the cathode

At the positive anode reaction with graphite electrodes

Oxygen gas is formed at the positive electrode, an oxidation reaction (electron 

loss).

The negative sulphate ions or the traces of hydroxide ions  are 
attracted to the positive electrode. But the sulphate ion is too stable and nothing 

happens. Instead hydroxide ions are discharged and oxidised to form oxygen.

                       

                                                          Figure 13.4.3: Electrolysis of copper(II) sulphate

Checking up 13.4

1. Name the product at the cathode and anode during electrolysis of:
     a. Molten lead bromide with inert electrode.
     b. Acidified copper sulphate solution with inert electrodes.
    c. Acidified water with inert electrode.
    d. Dilute hydrochloric acid with inert electrode.
    e. Concentrated hydrochloric acid with inert electrode.
2.Predict the products formed when the following molten compounds are 
    electrolysed using carbon electrodes;

     a. Lead(II) bromide

     b. Magnesium oxide

13.5. Electrolysis of concentrated copper (II) sulphate solution 

        using copper electrodes

The products of electrolysis of copper sulfate solution with copper electrodes are 
copper metal and copper ions (the copper anode dissolves).

The electrolysis of copper (II) sulphate solution using copper electrodes is shown 

below. 

                          

                          Figure 13.5.1: Electrolysis of copper (II) sulphate

Using the simple apparatus and two copper electrodes the products of the 
electrolysis of copper sulphate solution are a copper deposit on the negative 
cathode electrode and copper dissolves, at the positive anode electrode. 
This copper anode reaction differs from the one when you use an inert graphite 

electrode for the anode.

When Copper (II) sulphate is electrolysed with a copper anode electrode (the 

cathode can be carbon or copper), the copper deposit on the cathode (–) equals the 

copper dissolves at the anode (+). Therefore the blue colour of the Cu2+ ions stays 
constant because Cu deposited = Cu dissolved. Both involve a two electron transfer 
so it means mass of Cu deposited = mass of Cu dissolving for the same quantity 
of current flowing (flow of electrons). You can check this out by weighing the dry
electrodes before and after the electrolysis has taken place.

The experiment works with a carbon anode and you see the blackness of the graphite 

change to the orange-brown colour of the copper deposit.

                     

                Figure 13.5.2: The electrolysis of copper (II) sulphate solution with a copper anode and a 

                copper cathode.

The electrode reactions and products of the electrolysis of copper sulphate solution 
(with a copper anode) are illustrated by the diagram above.

(a) The electrode products from the electrolysis of copper sulphate with copper 

electrodes

The negative cathode electrode attracts ions (from copper sulphate) and 
ions (from water). Only the copper ion is discharged, being reduced to copper metal.
A reduction electrode reaction at the negative cathode: 

(copper deposit, reduction 2 electrons gained) reduction by electron gain.

The positive anode reaction with copper electrodes
It’s the copper anode that is the crucial difference than electrolysing copper sulphate 

solution with an inert carbon (graphite) or platinum electrode.

The negative sulphate ions  (from copper sulphate) or the traces of hydroxide 
ions   (from water) are attracted to the positive electrode. But both the sulphate 
ion and hydroxide ion are too stable and nothing happens to them because the 
copper anode is preferentially oxidised to discharge  copper ions.

An oxidation electrode reaction at the positive anode: copper dissolves, two 

electrons are lost following the half reaction:

                    

Copper atoms oxidised to copper (II) ions: dissolving of copper in its electrolytic 

purification or electroplating (must have positive copper anode).

Copper (II) ions reduced to copper atoms: deposition of copper in its electrolytic 
purification or electroplating using copper (II) sulphate solution, so the electrode 

can be copper or other metal to be plated or any other conducting material.

This means for every copper atom that gets oxidised, one copper ion is reduced, 

therefore 

when copper electrodes are used in the electrolysis of copper sulphate solution, the 
mass loss of copper from the positive anode electrode should equal to the mass of 

copper gained and deposited on the negative cathode electrode.

You can show this by weighing both electrodes at the start of the experiment. After 
the current has passed for some time, carefully extract the electrodes from the 
solution, wash them, dry them and reweigh them. The gain in mass of the cathode 

should be the same as the loss of mass from the anode.

Checking up 13. 5

1.
     a. Predict the products formed (i) at the cathode, (ii) at the anode, when 
         copper (II) sulphate solution is electrolysed using carbon electrodes.
    b. Describe the colour changes in the electrolyte.
 2. What will you observe
    a. At cathode
    b. At anode
    c. In electrolytic, during the electrolysis of copper (II) sulphate solution 
        with copper electrode.
3.Write equations for the reaction taking place at cathode and at anode 
      during the electrolysis of:
    a. Acidified copper sulphate solution with copper electrode.
    b. Acidified water with inert electrode.
    c. Molten lead bromide with inert electrode.
4.Using a table, highlight the similarities and differences between using 
   graphite electrodes and copper electrode for the electrolysis of copper 

   sulphate.

13.6. Faraday’s Laws

Activity 13.6

 Comparison of the amounts of different substances liberated by the same 

quantity of electricity.

                    

                   Figure 13.6: Comparison of the amounts of different substances
                    liberated by the same quantity of electricity

Set up the circuit containing a copper voltmeter and a silver voltmeter (a 
voltmeter is a vessel containing two electrodes immersed in a solution of ions 

through which a current is to be passed.)

Identify the copper and silver cathodes, clean and dry them, and after weighing 
them return them in their respective voltmeters. Pass a current of about 0.5A 
for 20 or 30 minutes, after which the cathodes should be removed, cleaned and 

dried, and reweighed.

Compare the masses of copper and silver deposited. Note that care must be 
taken in removing the silver cathode from the solution as the metal does not 

always adhere well to the cathode.

The relationship between the mass of product formed at an electrode and the 
quantity of electricity passed through an electrolyte is given by Faraday’s laws of 

electrolysis.

Michael Faraday (1791-1867) did the first work on electrolysis and formulated what 
is known today as Faraday’s laws of electrolysis.

These laws express the quantitative results of electrolysis. They assert that the 
amount (expressed in moles) of an element liberated during electrolysis depends 

upon:

      1. The time of passing the steady current, t
      2. The magnitude of the steady current passed, I

     3. The charge of the ions

13.6.1. Faraday’s first law

According to this law, “The amount of substance liberated at an electrode is 

directly proportional to the quantity of electricity passed”

                       

1F = 96500 coulomb
 So, 1 Faraday [96500 coulomb] of electricity will produce 1 gm equivalent of Ag, 
Cu and Al at cathode.

13.6.2. Faraday’s second law 
According to this law, “if same quantity of electricity is passed through different 
electrolytes, then the amount of substances liberated at the respective electrodes 

are in the ratio of their equivalent masses”.

Or 

When the same quantity of electricity passes through solutions of different electrolytes, 
the amounts of the substances liberated at the electrodes are directly proportional to 

their chemical equivalents.

         

Equivalent mass is the mass of a substance especially in grams that combines with 
or is chemically equivalent to eight grams of oxygen or one gram of hydrogen; the 

atomic or molecular Mass divided by the valence.

Example: 
Calculate the amount of electric charge in coulombs which can 
deposit 5.2g of aluminium when a current was passed through a solution of 

aluminium sulphate for some time.

Solution:

3 moles of electrons are needed to deposit 1 mole of aluminium (24g of 

aluminium).

Checking up 13.6

1. A current of 0.65A was passed through sulphate solution of metal X 
     for 35 minutes between platinum electrodes. If 0.449 g of the metal is 
    deposited at the cathode, calculate the charge on the element X (atomic 
     mass is 63.5)

2. When a constant current was passed through an aqueous solution 
    of copper (II) nitrate for one hour the mass of the copper cathode 
   increased by 15.24 g. Calculate the current in amperes which was used( 

   F= 96500 Cu = 63.5)

13.7. Factors affecting Electrolysis

Activity13.7: Investigating how the nature of electrodes affects the discharge                                                                                              of ions during electrolysis.
Apparatus and chemicals: 
• Beaker (100 cm3)
• 6V Battery
• 2 connecting wires with crocodile clips
• 2 carbon electrodes
• 2 copper electrodes
•   solution
Caution: Do not allow the electrodes to touch each other while the power 

supply is switched on, otherwise this may damage the equipment.

A. Electrolysing of copper (II) sulphate solution using Carbon electrodes.

                 

          Figure 13.7.1: Electrolysis of copper (II) sulphate solution using carbon electrodes.

 Procedure: let the current flow for 5 minutes

 Observe what happens at each electrode

Questions:

1. What product is formed on the cathode?
2. What product is formed at the anode?
3. Has the colour of the solution changed?

4. Explain the observations in 3.

B.Electrolysing of copper(II) sulphate solution using Copper electrodes 
                                  

Procedure: let the current flow for 5 minutes

 Observe what happens at each electrode

Questions:

1. What product is formed on the cathode?
2. What product is formed at the anode?
3. Has the colour of the solution changed?

4. Explain the observations in 3.

In an electrolysis where there are more than one species which can be discharged 
at the same electrode, only one of them is discharged at a time; for example, in an 
aqueous sodium chloride solution, we have four species that is, and 
 ions from sodium chloride and and ions from water.
During electrolysis ions and ions migrate to the cathode while   ions and 
   migrate to the anode.
Now the question is, which species of ions will be discharged at the cathode and 
which ones will be discharged at the anode first?

The factors which decide the selective discharge of ions are:

• Nature of electrodes
• Position of the ion in electrochemical series
• Concentration 

• The state of the electrolyte

13.7.1. Nature of Electrodes

In the electrolysis of sodium chloride solution using a platinum cathode, ions 

are discharged first in aqueous solution.

       

However, if the cathode is mercury, sodium is discharged. The sodium atom 

discharged combines with mercury cathode to form sodium amalgam.

             

  Electrolysis of copper sulphate using copper anode  

In this electrolysis, the anions,   migrate to the anode but none of them 

is discharged. Instead the copper atoms of the anode ionise and enter the solution.

                 

The copper (II) ions are attracted to the cathode and copper is deposited as a brown 

layer. The use of platinum anode gives oxygen as the product due to the reaction;

             

13.7.2. Position of ion in Electrochemical Series

When solving for the standard cell potential, the species oxidized and the species 
reduced must be identified. This can be done using the activity series. The table 13.2 
is simply a table of standard reduction potentials in decreasing order. The species 
at the top have a greater likelihood of being reduced while the ones at the bottom 
have a greater likelihood of being oxidized. Therefore, when a species at the top is 
coupled with a species at the bottom, the one at the top will become reduced while 

the one at the bottom will become oxidized. 

During electrolysis of solution containing a mixture of ions, the ion lower in 

electrochemical series is discharged first in preference to the one high in the series.

Let us look at the role of water in electrolysis products.

Water molecules to a small extent (degree) ionize as

            

Due to the above ionization, aqueous solutions of electrolytes contain two sets 
of ions that is those from the salt dissolved and the   ions from water 

molecules.

 Example:

Electrolysis of aqueous copper (II) sulphate solution using platinum electrodes.

Ions present in solution:

                

       Cathode

      

          The cathode gets coated with a brown layer of copper as the solution loses its blue 

           colour.

           Anode 

           

13.7.3. Concentration of electrolyte solution

Increase of concentration of an ion tends to promote its discharge, for example in 
concentrated hydrochloric acid, containing 
 as negative ions, the highly concentrated  is discharged in preference.

However, if the acid is very dilute, some discharge of will also occur. It is important 
to know that as the acid is diluted, there will not be a point at which chlorine ceases 
to be produced and oxygen replaces it. Instead a mixture of the two gases will come 

off, with the proportion of oxygen gradually increasing.

Another case in which the order of discharge according to the electrochemical series 
is reversed by a concentration effect is that of sodium chloride solution.

In concentrated solution of sodium chloride called brine, the following reactions 
occur.
              

Question 3

A university student set up three different electrolytic cells. The substances that 
were electrolysed were NaCl (l), 0.05 M NaCl (aq) and 5.0 M NaCl (aq). Which of 
the following statements correctly describes the results of the experiment? 

a. The reactions occurring for the aqueous solutions will produce the same 
     products at the anode and cathode. 
b. Chlorine gas is the major product when molten NaCl (aq) and 0.05 M 
     NaCl (aq) are electrolysed. 
c. The pH at the cathode increases when solutions of NaCl are electrolysed. 
d. The only means by which different products can be produced for 

     varying concentrations of NaCl is to alter the voltage.

13.7.4. The state of the electrolyte

The half reactions taking place at the electrode depends on whether the electrolyte 
is in a molten or an aqueous state, and if in aqueous state its concentration. For 
example, the electrode reactions that take place during the electrolysis of molten 

potassium iodide are:

           

However, if aqueous potassium iodide is used, the following electrode reactions 
take place:
            
                                              Electrode signs and reactions

             

 Checking up 13.7

Question 1 

A student set up the following experiment. How would the rate of electrolysis 

in beaker 2 compare to that in beaker 1?

                           

13.8. Application of electrolysis

Activity 13.8 

Copper-Plated Key

Materials:

• 1.5-volt D battery with battery holder
• Two alligator clip leads or insulated wire
• Beaker or glass
• Copper sulphate
• Copper electrode (or coil of copper wire)
• Brass key

• Safety equipment

Procedure:

5. Prepare the key for copper-plating by cleaning it with toothpaste or 
    soap and water. Dry it off on a paper towel.
6. Stir copper sulphate into some hot water in a beaker until no more will 

    dissolve. Your solution should be dark blue. Let it cool.

3. Use one alligator clip to attach the copper electrode to the positive terminal 
     of the battery (this is now the anode) and the other to attach the key to the 
    negative terminal (now called the cathode).

4. Partially suspend the key in the solution by wrapping the wire lead loosely 
    around a pencil and placing the pencil across the mouth of the beaker. The 
     alligator clip should not touch the solution.

5. Place the copper strip into the solution, making sure it doesn’t touch the key 
     and the solution level is below the alligator clip. An electrical circuit has now 
    formed and current is flowing.

6. Leave the circuit running for 20-30 minutes, or until you are happy with the 
     amount of copper on the key.

Question: Observe carefully electrolysis process and records what happened 

during the electrolysis process.

Electrolysis has a number of important industrial applications. These include 
the extraction and purification of metals, electroplating and anodizing and the 

manufacture of other chemicals for example sodium hydroxide (NaOH).

Extraction of metals

Metals in group I and II of the periodic table cannot be reduced by chemical 
reducing agents; they are extracted from their fused halides by electrolysis. Sodium 
is obtained by electrolysis of molten sodium chloride in the Dawncell. Magnesium is 
obtained by electrolysis of  genera ted from dolomite and sea water.
Extraction of aluminum

The chief ore of aluminum is bauxite, it contains silica and iron (III) 

oxide as impurities. Bauxite is dissolved in a strong solution of sodium hydroxide:

             

The impurities are not affected by the presence of sodium hydroxide because they 
are not amphoteric and they are thus filtered off. The solution is diluted, cooled and 

seed by adding a few crystals of pure Al 

On seeding, the aqueous tetrahydroxoaluminate is precipitated as pure Al

from the solution:

                          

The electrolytic cell is an iron tank lined with carbon, which acts as the cathode. The 
anodes are blocks of carbon dipped into the electrolyte. The electrolyte is a solution 
of molten aluminum oxide in molten cryolite. Cryolite acts as a solvent to dissolve 
aluminum oxide and as an impurity to lower the melting point of aluminum oxide. 

The electrolytic cell is maintained at around 900°C.

                  

Aluminum ions are discharged at the cathode, forming a pool of molten aluminum 
at the bottom of the tank.

At high temperature, oxygen reacts with the carbon anode to form carbon dioxide 
gas. Hence, the anodes are slowly burnt away as carbon dioxide gas and needs to be 

replaced frequently.

Manufacture of NaOH and extraction of  gas in Downcell

Construction of Down’s cell 

• Down’s cell consists of a rectangular container of steel.
• Inside of the tank is lined with firebricks.
• Anode is a graphite rod which projects centrally up through the base of the 
   cell.
• Cathode is a ring of iron, which surrounds the anode
     The anode and cathode are separated from each other by a cylindrical steel 
     gauze diaphragm; so that and are kept apart.

• A bell like hood is submerged over the anode

             

            
These are manufactured by electrolysis of concentrated sodium chloride, called 

brine. Hydrogen is also obtained as by products. In solution, sodium chloride ionizes:

       

Using carbon electrodes, the products of electrolysis are chlorine at the anode and 

hydrogen at the cathode. Hydrogen is discharged in preference to sodium.

           

The sodium discharged at the mercury cathode forms a solution of sodium amalgam 
in mercury. The sodium amalgam is collected in a reservoir in which it reacts 
with water to form sodium hydroxide solution and hydrogen gas. Mercury is also 

recovered and returned to the electrolytic cell to pass through the process again.

      

The sodium hydroxide produced is crystallized. It is used in:
• Manufacture of soap
• The paper industry-wood contains lignin, which is a nitrogenous compound, in 
addition to cellulose. Wood chips are converted into pulp by boiling the chips 
with sodium hydroxide solution to remove the lignin. The digested material is 

bleached with chlorine.

Purification of metals

Metals such as copper, zinc and aluminium can be purified by electrolysis. The 
purification of metals is known as refining. The copper obtained after the reduction 
process is not very pure. It contains small amounts of impurities such as iron. This 
copper is called blister copper and is refined by an electrolytic method. It is cast 

into bars which are used as anodes in acidified copper (II) sulphate solution.

The cathode is made of thin pure copper. During the electrolysis, ions are 
transferred from the anode to the cathode where they are discharged and copper is 

deposited.

            

The net effect is to dissolve the anode made of impure copper and thicken the 

cathode (pure copper) with more pure copper.

Electroplating

This is a process of coating a metal with another of interest mainly to prevent it from 
rusting, or/ and to improve its appearance, for example, in silver plating articles as 
cake dishes, made of base alloy, for example cupronickel, are made the cathode in 
plating bath of potassium (or sodium) dicyanoargentate(I), solution. This 
contains some silver ions, The anode is pure silver. When direct current passes, 

the following reactions occur.

       

In general;

• The metal being coated is made the cathode and the metal coating is the anode 
• The solution used is made of the ions of a metal that is coating, so that the 

    anode can dissolve. Anode is the pure plating metal.

                 

 A good electroplating requires steady electric current, appropriate concentration of 

electrolyte and temperature. The metal to electroplate must be clean.

Checking up 13.8

1. a. What is the difference between electrolytic extraction of a metal and 
          electroplating?
    b. Draw a set up used to electroplate a spoon by silver.
2.What is the material for cathode and anode during electro refining of impure 
        copper? 

3. Electrometallurgy is the process of reduction of metals from metallic 
     compounds to obtain the pure form of metal. Given elements: aluminum, 
     lithium, sodium, potassium, magnesium, calcium, Zinc, Iron and copper 
     which ones can be reduced by chemical agents such as carbon and which 

     ones are produced by electrolysis only?

13.9. End Unit Assessment

1. Choose from a list of words and fill in the missing words in the text below: 
      electrolysis, cathode and anode. When the current passes through an 
       electrolytic solution, ions migrate and electrons are gained or lost by ions 
      on the electrodes surface. Electrode that is positively charged has deficit 
      of electrons is called…………and other negatively charged has excess of 
      electrons is called……………. chemical changes at the electrodes due to the 
       passage of electric current are called ………………...

2. Answer by true or false
      a. The Avogadro (1791-1867) did the first work on electrolysis and 
         formulated what is known today as faraday’s laws of electrolysis.
     b. Electroplating is a process of coating a metal with another of interest 
         mainly to prevent it from rusting, and to improve its appearance.
     c. A non electrolyte is a solution or molten compound which cannot be 
        decomposed by an electric current.
    d. An ion is an atom or a group of atoms (radical) which has an electric 
       charge.
  e. An anode is the negative electrode through which electrons leave 
       the electrolyte and a cathode is the positive electrode through which 
        electrons enter the electrolyte or leave the external circuit.

3. Which of the following involves electrolysis?
      a. Photosynthesis and Respiration
      b. Purification of Copper and Sea Water
      c. Purification of copper and extraction of reactive metals 
     d. Extraction of reactive metals and Respiration

4. Which of the following is not an inert electrode? 
        a. Carbon 
        b. Copper
        c. Platinum 
        d. Mercury 

5. What ions are present in the electrolysis of aqueous Copper (II) Sulfate with 
         Copper electrodes?
     a. Copper (II) ions, Sulfate ions
     b. Hydrogen ions, Oxygen ions, Copper (II) ions
     c. Copper (II)ions
     d. Hydrogen ions, Hydroxide ions, Copper (II) ions, Sulfate ions

6. What happens during the electrolysis of a diluted sodium chloride solution?
a. Hydrogen ions and chlorine ions are discharged.
b. Hydrogen ions and hydroxide ions are discharged.
c. Sodium ion and chlorine ions are discharged.
d. Sodium ions and hydroxide ions are discharged.

7. State three applications of electrolysis on a large scale and describe one of 
them briefly.

8. Describe how the factor of concentration affects electrolysis.

9. Calculate the amount of electricity required (in Faradays) to deposit one mole 
of lead ions if a current of 2.5A was passed for 20 minutes through molten lead 
(II) bromide and 3.20 g of lead metal was deposited.
 
10. An element X has a relative atomic mass of 88. When a current of 0.5A is 
     passed through the molten chloride of X for 32 minutes and 10 seconds, 0.44 
     g of X deposited at the electrode.

    • Calculate the number of Faradays needed to liberate 1 mol of X.

    • Write the formula of the ion of X (1F= 96500C).