Electrolysis of molten compounds

2 types of electrochemical cell:

a) Electrolytic Cell

b) Voltaic Cell /Galvanic Cell

Electrolysis cell : Electricity energy chemical

energy

Voltaic cell : Chemical energy electricity

energy

The Electrolytic Cell

Electrode connected to the positive terminal of the cell is

positive electrode and is given a name, anode.

The electrode connected to the negative terminal of the

cell is negative electrode and is called the cathode.

When electricity is passed through an electrolyte,

chemical reaction happens.

In this reaction, chemical is splitting up into 2 new

substances.

All electrolytes are ionic, which means they are composed

of positively and negatively charged ions.

On passing an electric current through the electrolyte,

these ions move towards the oppositely charged

electrode.

Most negatively charged ions are non-metal ions, such as

oxide (O2-, chloride (Cl-), Iodide (I-), etc.

During electrolysis, negatively charged ions move towards

the positive electrode(anode). The negative ions lose

their electron(s) to the anode, which is positively charged.

The electron(s) is then move to the cathode through the

external circuit (the wire).

The positively charged ions move towards the negative

electrode(cathode').

These positive ions are metal ions, such as copper (Cu2+),

silver (Ag+), lead (Pb2+), etc, or hydrogen (H+).

At cathode, positive ions gain electron(s) from the

cathode, which has an excess of electrons and therefore

an overall negative charge.

This process results in the chemical decomposition of the

electrolyte. It also allows electrons to travel from the

cathode to the anode and hence allows conduction of

electricity.

During the electrolysis, electrical energy is supplied to the

system to produce a chemical reaction.

Therefore, during electrolysis, electrical energy convert

into chemical energy.

Example 1:Electrolysis of MOLTEN Lead (II) Bromide

This is composed of lead(II) ions, Pb2 + , and bromide ions,

Br-. Its chemical formula is therefore PbBr2.

A suitable apparatus which could be used to carry out this

electrolysis is shown in Figure above.

The bulb helps to show when electricity is flowing in the

circuit, and until the lead(II) bromide is completely molten,

the bulb does not light up . This confirms that electrolytes

have to be molten for the ions to start to move to the

electrodes and thereby conduct electricity.

At the Cathode At the Anode

Observation

When electricity is

Observation

When electricity

flowing, a silvery

deposit of lead metal

forms on the

cathode. In fact, as it

is molten, it is more

likely to drip off in a

molten blob.

is flowing, brown

fumes of

bromine gas are

seen at the

anode.

Half equation

Pb2+ + 2e ---> Pb

Half equation

2Br- ---> Br2 + e

Explanation

The lead(II) ions,

as they are

positive, move to

the negative

cathode, where

each ion gains

two electrons to

form a lead atom.

Any reaction at a

cathode involved

is again in

Explanation

The bromide

ions, as they are

negative, move

to the positive

anode, where

each loses an

electron to form

a bromine atom.

Then two of

these newly

formed atoms

electrons. This is

called reduction

or more exactly,

cathodic

reduction .

combine to form

bromine gas.

Any reaction at

an anode

involves a loss

of electrons.

In summary, the lead(II) bromide is split into its

component elements :

PbBr2 ---> Pb + Br2

Electrolysis Of Molten Lead(II) Oxide

(electrolysis of aqueous solutions)

Introduction

We have learnt that electrolyte can be molten ionic

compound or aqueous solution of ionic compound, acid or

alkali.

An aqueous solution is solution of water of a substance.

For example, if you heat sodium chloride until it melts, it is

called molten sodium chloride, but if you dissolve sodium

chloride in water, it is called aqueous sodium chloride.

Electrolysis of aqueous solution is different from

electrolysis of molten electrolyte.

This is mainly because an aqueous solution contain more

types of ions.

Let us take the example of molten sodium chloride and

sodium chloride aqueous.

In molten sodium chloride, the ion present are sodium ion

(Na+) and chloride ion (Cl-), due to the decomposition of

the solid sodium chloride.

NaCl ---> Na+ + Cl-

In sodium chloride aqueous, other than the

decomposition of sodium chloride solid to form

sodium and chloride ions, some of the water

molecule will also disassociates to form hydrogen

(H+) and hydroxide (OH-) ions.

NaCl ---> Na+ + Cl-

H2O ---> H+ + OH-

Which means in an aqueous solution, it can be

more than 1 positive and negative ions.

When the ions move to the anode and cathode,

only 1 negative ion and 1 positive ion will be

selected to be discharged, and this is called

selective discharge.

There are a few factors that determine which ion

will be selected to be discharge, and this will be

discussed in next section.

Factors Affecting Electrolysis

There are three main factors that can affect the

electrolysis products, there are:

1. position in the electrochemical series

2. the concentration and

3. the type of electrode

Electrochemical series

The chart above lists the ions in order of difficulty of

discharge.

The ions at the top of the list is more difficult to be

discharged, but as we go down the table, they

become easier to be discharged. For example, Cu2+

easier to be discharged compare with H+ and OH- is

easier to be discharged compare with I-.

This series of ions is called the Electrochemical

Series. The lower the ion in the electrochemical

series, the easier the ion to be discharged during

electrolysis.

Electrolysis of Aqueous Sulphuric Acid

As sulphuric acid is aqueous, it is composed not

only of hydrogen ions (H+ ) and sulphate ions

(SO42-), but also of hydroxide ions (OH-) from the

water.

H2SO4 + H2O --> 2H+ + SO42- + H+ + OH-

The apparatus used to carry out this electrolysis

and collect the gases given off is shown in Figure

9 .8 .

When we have more than one type of ion moving

to an electrode, selective discharge (or

preferential discharge) takes place.

This means that the ion which can lose or gain

electrons with the greatest ease is discharged,

and the other ions, which are harder to

discharge, remain in solution .

With the electrolyte aqueous sulphuric acid,

migration of ions to the electrodes also occurs.

At the Cathode At the Anode

Here we have only

one ion, the

hydrogen, H+ (aq),

and each ion gains

an electron to

become a

hydrogen atom.

Two of these newly

formed atoms then

combine to form a

hydrogen gas

molecule .

Here we have a

choice of either

sulphate, SO42-

(aq), or hydroxide

OH- (aq) ions.

Hydroxide is easier

to discharge, so

oxygen gas is given

off at the anode.

Equation:

2H+ + 2e ---> H2

Equation:

OH- + 4e ---> O2 + H2O

Notes

With electrolysis of aqueous solutions of

dilute acids or alkalis, the volume of

hydrogen given off at the cathode is roughly

twice that of the oxygen gas at the anode.

Accordingly, the elements of water are lost

and as the electrolysis continues, the

concentration of the acid or alkali

increases .

Essentially, the electrolysis of aqueous

sulphuric acid is the electrolysis of water,

with hydrogen and oxygen gas being given

off in a ratio of 2 : 1 .

[edit] Concentration

If the concentration of a particular ion is high, it

will be selected to be discharged even though it

is higher in the electrochemical series compares

with another ion present in the solution.

For example, if dilute hydrochloric acid is

electrolysed, hydrogen gas is given off at the

cathode and oxygen gas at the anode.

However, when concentrated hydrochloric acid is

electrolysed, hydrogen gas is still given off at the

cathode, but chlorine rather than oxygen gas will

be released at the anode, even though chloride

is in a higher position in electrochemical series.

Electrolysis Of Diluted Or Concentrated

Hydrochloric Acid

Electrolysis of Concentrated Sodium Chloride

Solution (Brine)

The electrolytic cell used for electrolysis of

concentrated sodium chloride solution is designed to

collect gaseous products at both electrodes as

shown in Figure above.

At Cathode At Anode

The sodium and

hydrogen ions

move to the

cathode .

As the hydrogen

ion (H+), is lower in

the reactivity series

than the sodium

ions (Na+ ), it

Both the chloride

ions (Cl-) and the

hydroxide ions

(OH-) migrate to the

anode .

The chloride ions

(Cl-) are

preferentially

discharged

accepts electrons

more easily.

The hydrogen ions

(H+) are

discharged.

because of their

higher

concentration

Equation:

H+ + e ---> H

Hydrogen atoms join in

pairs to give

molecules :

H + H ---> H2

Equation:

Cl- ---> Cl + e

Chlorine atoms join in

pairs to give

molecules:

Cl + Cl ---> Cl2

Changes in Solution

As the hydrogen ions and chloride ions are

discharged, sodium ions and hydroxide ions

remain in the solution . The solution

becomes sodium hydroxide .

Type of Electrode

This is best shown if we consider the electrolysis

of aqueous copper(II) sulphate solution.

Electrolysis of Copper(II) Sulphate by Using

Carbon Electrode

Anode

1. If we use carbon electrodes, they are inert

electrodes and do not affect the electrolysis .

2. Therefore, at the anode, we have a choice of

sulphate or hydroxide ions .

3. The hydroxide ions are easier to discharge,

so oxygen gas is given at the anode :

Partial equation 40H- (aq) O2(g) + 2H2O (l) +

4e- (oxygen gas given off)

Cathode

At the cathode, we have a choice of copper or

hydrogen ions .

The copper ions are easier to discharge, so we

see a pink deposit of copper metal on the carbon

electrode.

Partial equation Cu2+ (aq) + 2e- Cu (s) (copper

metal deposited)

Electrolysis of Copper(II) Sulphate by Using

Copper Electrode

However, if we use copper electrodes, these are

active electrodes and do affect the electrolysis.

Anode At the anode, the copper electrode dissolves

into solution :

Partial equation Cu(s) Cu 2+ (aq) + 2e (copper

electrode dissolves )

Cathode At the cathode, the copper ions are

deposited as pink copper metal:

Partial equation Cu2+ (aq) + 2e- Cu (s) (copper

metal deposited)

Evaluating electrolysis in industry

[edit] Industrial Applications of Electrolysis

Electrolysis has many varied industrial applications.

The major applications of electrolysis in industry are

1. Extraction of Metals

2. Purification of Metal

3. Electroplating

[edit] Extraction of Metal

The extraction of metals from their ores, in particular

aluminium and sodium, is important industrial uses of

electrolysis.

The diagram below shows the methods of extraction for

different metals.

We can see that those metals which are less reactive than

carbon in reactivity series are extracted from their ore by

displacement reaction using carbon. This will be

discussed in detail in chapter 3, form 5, Oxidation and

Reduction.

Copper and mercury can be extracted from their ore by

burning directly in air.

Silver (Ag) and gold (Au) need no extraction because they

exist as element in nature.

Those metals which are more reactive than carbon are

extracted by electrolysis.

[edit] Extraction of Aluminium

Aluminium is the most abundant metal found in the earth's

crust. It makes up about 8% by weight of the Earth’s solid

surface.

It is also a very useful metal due to its low density and

ability to resist corrosion.

The main source of aluminium is bauxite ore (Aluminium

Oxide).

In industry, aluminium is extracted by electrolysis from

bauxite ore.

Adding Cryolite

In electrolysis, molten aluminium oxide must be used to

extract aluminium. Aluminium oxide decompose to form

aluminium and oxide ions when melted.

Al2O3 ---> 2Al3+ + 3O2-

However, the melting point of aluminium oxide is very

high (over 2 000°C), so another aluminium compound

called cryolite (Na3AIF6) is added to lower down the

melting point (about 980oC).

The diagram above shows how aluminium is extracted

from molten aluminium oxide by electrolysis.

Graphite is used as the anode and cathode.

During electrolysis, the aluminium ions are attracted

towards the graphite cathode.

The ions is discharged and become molten aluminium

metal.

The partial equation of this reaction is as follow:

Al3+ + 3e ---> Al

At the anode, oxygen gas which also has

commercial value is collected. The partial

equation of this reaction is as follow:

2O2- ---> O2 + 4e

At the temperature of 980 °C, the oxygen

burns the carbon anode. Therefore the

anode has to be replaced periodically.

Also, this cell uses large quantities of

electricity, and therefore needs cheap

sources of power.

[edit] Extraction of sodium chloride

In industry, sodium is extracted from molten

sodium chloride. Molten sodium chloride is

put into the apparatus as showing in the

diagram above.

When sodium chloride is melted, the sodium

and chloride ions disassociate to become

freely move ions, as shown in the chemical

equation below.

NaCl ---> Na+ + Cl-

In thhis electrolytic cell, graphite was

used as anode while iron is used as

cathode.

The negative chloride ions are attracted

to the anode and then discharged to

form chlorine gas.

2Cl- ---> Cl2 + 2e

Since chlorine gas is also

significant in industry, it is collected

and stored.

In cathode, the sodium ions are

discharged to form sodium atom.

Na+ + e ---> Na

Due to high temperature, the

sodium metal formed is in

molten form.

Metal sodium have lower

density. Therefore it moves

upward and been collected.

[edit] Purification Of Copper

In the refining or purification of

copper, the impure copper is made the anode and a thin,

pure copper plate is used as a cathode.

The electrolyte is usually acidified copper(II) sulphate

solution.

When electricity flows, the

copper dissolves from the

impure anode and goes into

solution as copper ions.

Impurities in the copper do not

dissolve, and instead fall off

the anode as anode sludge. At

the cathode, the copper ions

are deposited as pure copper

metal.

Reaction in anode (impure

copper)

In anode, the copper atoms from

the electrode are ionised to form

copper(II) ions.

Cu ---> Cu2+ + 2e

Reaction in cathode (pure

copper)

Cu2+Cu ---> Cu + 2e

[edit] Electroplating

Electroplating: Coating

with a Thin Protective

Layer of Metal

A very common use

of electrolysis is to

form a thin protective

coating of a metal on

the surface of another

which is likely to

corrode.

The diagram above

illustrate the

electroplating of a key

with copper.

In this process, we

need to make the

cathode the object for

plating (the key.

The anode is then

made of the metal we

wish to plate with

(copper), and the

electrolyte needs to

be a solution of a salt

of this metal

(copper(II) sulphate).

Anode

In anode, the copper

atoms from the

electrode are ionised

to form copper(II)

ions.

Cu ---> Cu2+ + 2e

Cathode

In cathode, the

copper ions are

discharged to

form copper

atom and then

deposit on the

surface of the

key

Cu2+ ---> Cu + 2e