C12 electrochemistry

LEARNING OUTCOMES

ElectrochemistryChapter 12

Describe investigations leading to the classification of substances as conductors or non-conductors

Distinguish between metallic and electrolytic conduction Define electrolysis, cathode, anode, cation and anion Define electrolytes as strong and weak based on their conductivity Predict the electrode to which an ion will drift Define oxidation and reduction reactions including reactions at electrodes Predict chemical reactions making use of electrochemical series Identify ions present in electrolytes Discuss the electrolysis of certain substances Define Faraday constant Calculate the masses and volumes of substances liberated during

electrolyses Describe industrial applications of electrolysis


Conductors and electrolytes The movement of electrically charged particles through a medium

constitutes an electric current. The medium of conduction is known as a conductor.

Electrolytes can conduct electricity in their molten and aqueous states. E.g. Metals

Non-conductors like plastic, ceramic and rubber can be used as insulation materials.

Eg. The plastic coating on copper wires


Electrical conductions in conductors Electrical conductions in conductors can be explained through how the atoms are

bonded.

In metals, the electrons become delocalised due to overlapping orbitals of the tightly packed atoms in the crystal lattice.

The valence electrons of each atom are loosely held as they are relatively distant from the nucleus. Thus they can be easily separated and move about randomly throughout, conducting electricity.

The only non-metal that conducts electricity is graphite.

In graphite, not all electrons are used in bonding. Thus, the free electrons can move along the layers and carry the electric current.

The ionic theory proposes that it is the presence of freely-moving charged particles called ions in the electrolyte that allows electrolysis to occur.

These ions are only mobile when the electrolytes are in molten or aqueous state.

These free moving ions arise only when an ionic solid melts or dissolves in water or when certain polar covalent compounds dissolve in water and their molecules ionise (dissociate into freely-moving ions).


The ionic theory

What is Electrolysis ? Electrolysis is a process by which a substance is

broken up into its components by the passage of electricity through it.

The substance must be an ionic compound and must be molten or dissolved in water in order for the ions to be mobile.

A direct current must be used for electrolysis. During this process, electrical energy is changed into

chemical energy.


One electrode is connected to the positive terminal of the battery. It is called the anode. Oxidation occurs here.

The other electrode is connected to the negative terminal of the battery. It is called the cathode. Reduction occurs here.

An electrolyte is a substance that is being electrolysed. An electrolyte is able to conduct electricity due to the presence of

mobile ions. Positive ions are known as cations. Negative ions are known as anions.


Electrolysis

At the anode,, Cl - ions give away electrons to become Cl2 gas. We say that Cl- ions are discharged.

Na+(l) + e- Na(l)

Molten sodium chloride contains Na+ and Cl- ions. The Na+ ions are attracted to the cathode, while the Cl- ions are attracted

to the anode. At the cathode, Na+ ions take in electrons to become Na atoms. We say

that the Na+ ions are discharged.

2Cl-(l) Cl2(g) + 2e-

Overall reaction:

2NaCl(l) 2Na(l) + Cl2(g)

Electrolysis of molten sodium chloride


When electricity passes through the molten ionic compound, the positive ions will migrate to the negative electrode, also known as cathode, while the negative ions will migrate to the positive electrode, known as anode.

At the cathode, the positive ions will gain electrons and become metal atoms while the negative ions will lose electrons at the anode and become non-metallic atoms.


Electrolysis of Molten Ionic Compounds

At the CathodeSodium ions gain electrons and become sodium atoms, sodium ions are discharged. Na+(l) + e- Na (l)At the AnodeChloride ions lose electrons and become chlorine gas, chloride ions are discharged. 2Cl- (l) Cl2 (g) + 2e-

Electrolysis of other molten compounds When a molten binary ionic compound is electrolysed,

the metal is always produced at the cathode and non-metal is produced at the anode.


Quick check 11. What is meant by (i) cathode, (ii) anode ?2. What is meant by an electrolyte? What type of compounds must

electrolytes be?3. During electrolysis, to which electrode do:

(a) the positive ions of the electrolyte move to; (b) the negative ions of the electrolyte move to?

4. Predict the products formed when the following substances (in the molten state) are electrolysed.

Compound Product at Anode Product at Cathode

Potassium chloride, KCl

Calcium fluoride, CaF2

Solution


1. (i) Cathode: the electrode connected to the negative terminal of the battery. (ii) Anode: the electrode connected to the positive terminal of the battery.

2. An electrolyte is a substance which conducts electricity when molten or dissolved in water. Electrolytes must be ionic compounds.

3. (a) The positive ions move to the cathode. (b) The negative ions move to the anode.

4.

Compound Product at Anode Product at CathodePotassium chloride, KCl Chlorine, Cl2 Potassium, K

Calcium fluoride, CaF2 Fluorine, F2 Calcium, Ca

Solution to Quick check 1

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When solutions are electrolysed, gases are usually produced. The gases produced can be collected in test tubes to be identified later

by simple tests.

The electrolysis of solutions is more complicated than electrolysis of molten compounds, because the products at the electrodes can come from the electrolyte as well as from water.

Electrolysis of solutions


At the cathode:

At the anode:

Selective discharge of ions

Positive ions from the electrolyte are discharged if they are H+ ions or ions of less reactive metals such as Cu2+, Pb2+ or Ag+.

Positive ions of reactive metals such as Na+, K+ and Ca2+ are not discharged in the presence of water. Instead, H+ ions from water are discharged and H2 gas is produced.

Negative ions from the electrolyte are discharged if they are halide ions such as I- , Br - , and Cl-.

SO42- and NO3

- ions are not discharged. Instead, OH- ions from water are discharged and O2 gas is produced.


The table shows the electrode products from solutions of ions, assuming the solutions are fairly concentrated.

Cation Product at Cathode

Ease of discharge

Anion Product at Anode

K+

Na+

Ca2+

Mg2+

Al3+

Ni2+

Pb2+

(H+) Cu2+

Ag+

Hydrogen from water

Lead

CopperSilver

Difficult

Easy

SO42-

NO32-

Cl-

Br-

I-

(OH-)

Oxygen from water

ChlorineBromine

IodineOxygen

Selective discharge of ions


Pure water will not conduct electricity, so some dilute sulphuric acid is added to make it conduct electricity. The electrolysis of dilute sulphuric acid is therefore essentially the same as the electrolysis of water.

Electrolysis of dilute H2SO4


OH- ions are discharged in preference over SO42-. Oxygen gas is formed.

2H+(aq) + 2e- H2(g)

Dilute sulphuric acid contains H+, SO42- and OH- ions.

H+ ions take in electrons to become H2 molecules; H+ ions are discharged:

4OH-(aq) O2(g) + 2H2O(l) + 4e-

Overall reaction:2H2O(l) 2H2(g) + O2(g)

At the cathode: 2 volumes of hydrogen are produced. At the anode: 1 volume of oxygen is produced.

Electrolysis of dilute H2SO4

At the cathode:

At the anode:


The electrolysis of dilute solutions is essentially the same as the electrolysis of water.

The ions present in the solution are: Na+, Cl- (from sodium chloride) and H+, OH- (from water).

Both Na+ and H+ are attracted here, but due to their relative positions in the reactivity series, H+ ions are preferentially discharged: 2H+(aq)+ 2e- H2(g)

Both Cl- and OH- are attracted here, but due to the lower position of the hydroxide ions in the reactivity series, they are preferentially discharged: 4OH-(aq) O2(g) + 2H2O(l) + 4e-

Overall reaction: 2H2O(l) 2H2(g) + O2(g) [ Electrolysis of water ]

Electrolysis of dilute sodium chloride solution

At the cathode:

At the anode:


The ions present in the solution are: Na+, Cl- and H+, OH-

Both Na+ and H+ are attracted here, but due to their relative positions in the reactivity series, H+ ions are preferentially discharged:

2H+(aq) + 2e- H2(g)

Both Cl- and OH- are attracted here, but due to the highconcentration of the chloride ions, chloride ions arepreferentially discharged: 2Cl-(aq) Cl2(g) + 2e-

Overall reaction:Hydrogen gas is produced at the cathode and chlorine gas is produced at the anode.

The Na+ and OH- ions left in the solution combine to form sodium hydroxide, thus making the solution alkaline.

Electrolysis of concentrated sodium chloride solution

At the cathode:

At the anode:


Examples of electrolysis of different solutions using inert electrodes, assuming the solutions are fairly concentrated

Electrolyte Ions in Solution Product at Cathode

Product at Anode

Aq. sodium chlorideNa+(aq), Cl-(aq) , H+(aq), OH-(aq)

hydrogen gas chlorine gas

Aq. hydrochloric acid

H+(aq), Cl-(aq), H+(aq), OH-(aq)

hydrogen gas chlorine gas

Aq. copper(II) sulphate

Cu2+(aq), SO42-(aq),

H+(aq), OH-(aq)copper metal oxygen gas

Electrolysis of Solutions


Inert and reactive electrodes Inert electrodes do not react with the product

produced or dissolved in the electrolyte. Carbon and platinum are examples of inert electrodes. Reactive electrodes can react or dissolve in the electrolyte. Copper, silver and mercury are examples of reactive

electrodes.


Neither SO42- nor OH- ions are discharged.

Instead the copper anode dissolves in the solution and produces electrons: Cu Cu2+ + 2e-

During the electrolysis the total concentration of the CuSO4 solution remains unchanged. The cathode increases in mass while the anode decreases in mass proportionately.

Electrolysis of CuSO4 solution

Copper anodeCopper

cathode

Aq. copper(II) sulphate

_ +

Cu2+

H+ SO42-

OH-

Using copper (reactive) electrodes The ions present in the solution are: Cu2+, SO4

2-, H+ and OH-.

Cu2+ ions are discharged in preference over the H+ ions: Cu2+ + 2e- Cu

At the cathode:

At the anode:


Electrolysis of CuSO4 solutionUsing copper (reactive) electrodes During the electrolysis the concentration of the H+ and

SO42- ions increases, and hence the solution becomes

more acidic. This process essentially transfers

copper metal from the anode to the cathode.

This process is used in the industry for the purification of impure copper to obtain pure copper.


Quick check 21. Place in order the ease of discharge of the following cations (starting from the easiest

first): Ca2+, Na+, H+, Al3+, Mg2+, K+, Cu2+, Pb2+, Ag+.2. Place in order the ease of discharge of the following anions (starting from the easiest

first): Br-, Cl-, NO3-, OH-, I-, SO4

2-,.

3. (a) State the products obtained when a solution of dilute sulphuric acid is electrolysed using platinum electrodes.(b) Write the ionic equations for the reactions taking place at the cathode and anode.

4. (a) State the products obtained when a concentrated solution of sodium chloride is electrolysed using inert electrodes.(b) Write the ionic equations for the reactions taking place at the cathode and anode.

5. (a) State the products obtained when a solution of sodium hydroxide is electrolysed using platinum electrodes.(b) Write the ionic equations for the reactions taking place at the cathode and anode.

Solution


1. Ag+, Cu2+, H+, Pb2+, Al3+, Mg2+, Ca2+, Na+, K+

2. OH-, I-, Br-, Cl-, NO3-, SO4

2-

3. (a) At cathode: hydrogen; At anode: oxygen (b) At cathode: 2H+(aq) + 2e- H2(g) At anode: 4OH-(aq) 2H2O(l) + O2(g)

4. (a) At cathode: hydrogen; At anode: chlorine(b) At cathode: 2H+(aq) + 2e- H2(g) At anode: 2Cl-(aq) Cl2

5. (a) At cathode: hydrogen; At anode: oxygen(b) At cathode: 2H+(aq) + 2e- H2(g) At anode: 4OH-(aq) 2H2O(l) + O2(g) Return



The Faraday constant, F, is the quantity of electricity carried by one mole of electrons and is equivalent to 96500 Cmol-1.

The amount of a substance deposited on each electrode of an electrolytic cell is directly proportional to the quantity of electricity passing through the cell.

The quantity of electricity contained in a current running for a specified time can be calculated by Q= l x t

The quantity of electricity required to deposit an amount of metal can be calculated by Q = n(e) x F


Faraday’s Law of Electrolysis

Electroplating

Electroplating with copper

Electroplating is the process in which a metallic object is coated with another metal by electrolysis.


Cu2+ ions are discharged as copper metal is deposited on the object:

Cu(s) Cu2+(aq) + 2e-

The electrolyte is copper(II) sulphate (CuSO4)solution.

The object to be plated is made the cathode; copper is made the anode.

Cu2+(aq) + 2e Cu(s)

There is a net transfer of copper from the anode to the cathode. The concentration of the CuSO4 solution remains unchanged.

Electroplating with copper

At the cathode:

At the anode:

Copper dissolves:


Anodising is the process of making the oxide layer on the surface of the aluminium thicker. This will protect the aluminium even better.

The aluminium object is made the anode while the cathode could be copper or lead or aluminium.

When current is applied, the water in the electrolyte breaks down and oxygen is deposited at the anode. This oxygen then combines with the aluminium to form aluminium oxide and thus provides a protective layer for the aluminium. This prevents corrosion.

Anodising


Electroplating with other metals Many metallic objects can be electroplated in

the same way: The object is made the cathode and the metal used for

electroplating is made the anode. The electrolyte is a solution of ions of the metal used for the plating.

Electroplating can be used to protect iron objects from corrosion by covering it with a less reactive metal like chromium or copper.

Electroplating also makes the object more attractive and increases its value e.g. plating it with gold, silver and platinum.


Electroplating Metal Uses

Chromium Water taps, motorcar bumpers and bicycle parts

Tin Tin cans

Silver Silver-plated sports trophies, plaques, ornaments, knives and forks

Gold Gold-plated watches, plaques, ornaments

Examples of electroplating


1. State two uses of electrolysis in the industry.2. (a) What is electroplating?

(b) State two advantages of electroplating an iron object with chromium.

3. A metal spoon is to be coated with silver. Sketch a diagram to show how you would set up the electrolytic cell for this to be carried out. Label the material that can be used for the cathode and anode. Also state a suitable solution for the electrolyte.

Quick check 3

Solution


1. (i) To extract reactive metals like sodium, magnesium and aluminium from their ores;(ii) To electroplate metallic objects with less reactive metals for attractiveness and protection from corrosion.

2. (a) Electroplating is the coating of a more reactive metal with a less reactive metal by electrolysis.

(b) Electroplating with chromium protects the iron from corrosion and the silver colour of the chromium improves its appearance.


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3. Cathode: the metal object; Anode: silver metalElectrolyte: silver nitrate solution.

Solution to Quick check 3 (cont’d)

Electroplating with silver

1. http://www.matter.org.uk/schools/Content/Electrolysis/ElectrolysisExplainApplet.html

2. http://www.nmsea.org/Curriculum/7_12/electrolysis/electrolysis.htm3. http://inventors.about.com/library/inventors/blelectroplating.htm

To learn more about Electrolysis, click on the links below!


Education

C12 electrochemistry