Doc Brown's Chemistry KS4 science-chemistry GCSE/IGCSE/AS Revision Notes

Extra Electrochemistry Information Notes on Electrolysis, cells, batteries, fuel cells and

industrial applications of electrolysis

This section starts with an introduction to electrolysis explaining which liquids conduct

and why can some liquids conduct electricity and others cannot? - electrolytes, electrodes,

what happens on electrode surfaces, the products of electrolysis. There are tables of electrode reactions,

descriptions of experimental methods of electrolysis and a summary table of the electrolysis products from

many common melts or aqueous solutions that undergo electrolysis when a d.c. electric current is passed

through them. There is also a list of links to revision notes on the application of electrolysis to various

industrial processes. Simple cells, batteries and fuel cells are also described.

Index for this page 1. Introduction to electrolysis * 2. Summary of practical experimental arrangements

for laboratory (lab) electrolysis experiments andelectrode reactions * 3a. Summary of the products of

selected electrolysis processes * 3b. Summary of Industrial Electrolysis Process links * 4. Simple cells or

batteries * 5. Fuel cells

Associated LINKS to other pages on this site: Metal Extraction * Industrial chemistry * Metal reactivity

(redox introduction) * Types of chemical reaction *Electrolysis calculations

 

1. Introduction to electrolysis - electrolytes and non-electrolytes

Electrolysis is the process of electrically inducing chemical changes in a conducting melt or solution e.g.

splitting an ionic compound into the metal and non-metal.

SUMMARY OF COMMON ELECTRICAL CONDUCTORS

o What makes up a circuit in cells - batteries or electrolysis? What carries the current?

o These materials carry an electric current via freely moving electrically charged particles,

when a potential difference (voltage!) is applied across them, and they include:

o All metals (molten or solid) and the non-metal carbon (graphite).

This conduction involves the movement of free or delocalised electrons (e- charged

particles) and does not involve any chemical change.

o Any molten or dissolved material in which the liquid contains free moving ions is called

the electrolyte and can conduct an electrical current. (see non-electrolyte)

Ions are charged particles e.g. Na+ sodium ion, or Cl- chloride ion, and their movement or

flow constitutes an electric current, in other words the electrolyte consists of a stream of

moving charged particles.

What does the complete electrical circuit for electrolysis consist of?

There are two ion currents in the electrolyte flowing in opposite directions 

positive cations e.g. sodium Na+ are attracted to the negative cathode

electrode,

and negative anions e.g. chloride Cl- are attracted to the positive anode

electrode,

and it is possible to demonstrate this flow using a coloured ion

experiment.

remember no electrons flow in the solution, but they do flow in metal wires

or carbon (graphite) electrodes of the external circuit.

Three 'connected' sub-notes: (i) The greater the concentration of the

electrolyte ions, the lower the electrical resistance of the solution. This

is because there are more ions present to carry the current e.g. if the

voltage (V, volts) is kept constant, the current flowing (I, amps) will steadily

increase as the concentration of the electrolyte is increased. (ii) If the

electrolyte (ion) concentration is kept constant, the current will

steadily increase with increase in voltage just like any other electrical

circuit because the increase in electrical field effect from the increased p.d.

(voltage) will force the ion flow at a greater rate. (iii) So, increase in ion

concentration(salts, acids etc.) OR increase in voltage will increase the

speed of electrolysis i.e. the electrode reactions, whether it involves gas

formation or electroplating metals etc.

The circuit of 'charge flow' is completed by the electrons moving around the external

circuit e.g. copper wire, metal or graphite electrode, from the positive to the negative

electrode (this e- flow from +ve to -ve electrode perhaps doesn't make sense until

you look at the electrode reactions later on).

The molten or dissolved materials are usually acids, alkalis or salts and their electrical

conduction is usually accompanied by chemical changese.g. decomposition.

The chemical changes occur at the electrodes which connect the electrolyte liquid containing ions with the

external d.c. electrical supply.

o If the current is switched off, the electrolysis process stops.

Non-electrolytes are liquids or solutions that do not contain ions, do not conduct electricity

readily and cannot undergo the process of electrolysis e.g. ethanol (alcohol), sugar solution etc. and are

usually covalent molecules.

How do we know that ions move to the electrodes when a d.c. current is applied?

o The coloured ion experiment described below illustrates how you can show the slow movement of

ions when a voltage is applies across a conducting solution of ions - the electrolyte.

A simple experiment to show the movement of coloured ions

A rectangle of filter paper is soaked in an ammonia-ammonium chloride solution and mounted on a microscope

slide. The paper is connected to a d.c. supply with clips. A 'line' of copper chromate solution is placed in the middle

of the paper and the current switched on. The copper chromate is green-brown in solution but gradually it

disappears and separates, in different directions, into a yellow and blue bands. The yellow band is due to negative

chromate ions, CrO42--, moving towards the positive electrode. The blue band is due to positive copper ions,

Cu2+, moving towards the negative electrode. All due to opposite charges attracting in the electric field

produced by the potential difference (the voltage!).

Liquids that conduct must contain freely moving ions to carry the current and complete the circuit.

o You can't do electrolysis with an ionic solid!, the ions are too tightly held by chemical bonds and

can't flow from their ordered situation!

o When an ionically bonded substances are melted or dissolved in water the ions are free to

move about.

However some covalent substances dissolve in water and form ions.

e.g. hydrogen chloride HCl, dissolves in water to form 'ionic' hydrochloric acid H+Cl-(aq) 

The solution of ions (e.g. salts, acids etc.) or melt of ions (e.g. chlorides, oxides etc.) is called

the electrolyte which forms part of the circuit. The circuit is completed by e.g. the external copper wiring and

the (usually) inert electrodes like graphite (form of carbon) or platinum AND electrolysis can only happen

when the current is switched on and the circuit complete.

ELECTROLYSIS SPLITS a COMPOUND:

o When substances which are made of ions are

dissolved in water, or melted material, they can

be broken down (decomposed) into simpler

substances by passing an electric current

through them.

o This process is called electrolysis.

o Since it requires an 'input' of energy, it is an

endothermic process.

During electrolysis in the electrolyte (solution or melt of free moving ions) ...

o ... positive metal or hydrogen ions move to the negative electrode (cations attracted to cathode), e.g.

in the diagram, sodium ions Na+ , move to the negative electrode (-ve),

o and negatively charged ions move to the positive electrode (anions attracted to anode), e.g. in the

diagram, chloride ions Cl-, move to the positive electrode (+ve). 

The diagram shows the industrial electrolysis process (in a Down's Process Cell) to extract sodium metal

from sodium chloride (common salt).

During electrolysis, gases may be given off, or metals dissolve or are deposited at the electrodes.

o Metals and hydrogen are formed at the negative electrode from positive ions by electron gain

(reduction), e.g. in molten sodium chloride

sodium ions change to silvery grey liquid sodium

Na+ + e- ==> Na  (a reduction electrode reaction)

o and non-metals e.g. oxygen, chlorine, bromine etc. are formed from negative ions changing on the

positive electrode by electron loss (oxidation), e.g. in molten sodium chloride

chloride ions change to green chlorine gas

2Cl- - 2e- ==> Cl2   or  2Cl- ==> Cl2 + 2e-  (an oxidation electrode reaction)

o The electrons released by the oxidation at the positive anode, flow round through the anode and

wire to the positive cathode and so bring about the reduction i.e. of the sodium ion.

In a chemical reaction, if an oxidation occurs, a reduction must also occur too (and vice versa) so these

reactions 'overall' are called redox changes.

o You need to be able to complete and balance electrode equations or recognise them and derive an

overall equation for the electrolysis.

o Below is list of some of the most common electrode equations you will encounter and

the experimental arrangements are shown below them in section 2.

 

2a. Summary of arrangements for laboratory electrolysis experiments

Examples of types of apparatus that can be used in schools and colleges

I'm happy to add other examples if they seem appropriate

Gas tests and this section followed by examples of electrode reactions to which the equation numbers refer.

Method 1a

Aqueous solutions with

inert electrodes (carbon or

platinum)

Graphite electrodes 'upwardly'

dipped into an solution of the electrolyte. The

cell can be made from plastic pipe and a big

rubber bung with two holes in it. A more

elaborate format is to use aHoffman

Voltammeter (see below) using platinum

electrodes and accurately calibrated collecting

tubes like burettes.

Method 2

Molten salts with

carbon electrodes

Carbon (graphite)

electrodes dipped

into molten salt

which has been

strongly heated in a

crucible. It is difficult to collect the gases

at the electrodes! The salts may be very

high melting, so sometimes a small

amount of another salt impurity is added

to lower the melting point.

Meth

od 3

Metal

electrodes dipped in aqueous salt

solutions

 M = metal, usually dipped in the metal

sulphate solution.

For electroplating in general: The (-)

is made the metal/conducting surface to

be coated, and the (+) is made of the

plating metal which dissolves and

replaces any deposit formed on the (-)

electrode (for more details).

Brine - concentrated sodium chloride Molten lead(II) bromide gives silver Copper(II)

solution (brine) with carbon (graphite) gives

equal volumes of hydrogen gas (-) and green

chlorine gas (+) with sodium hydroxide left in

solution, via equations 2 and 3. However in

very dilute solution, reaction 8 occurs too giving

oxygen gas as well as chlorine gas.

beads of lead (-) and brown fumes of

bromine (+), illustrated above,

seeequations 2 and 10 .

sulphate with copper electrodes, the

copper deposits (-) and the copper

dissolves (+), equations 4 and 5. The

blue colour of the Cu2+ ions stays

constant because Cu deposited = Cu

dissolved. Both involve a 2 electron

transfer so it means mass of Cu

deposited (eq 4 ) = mass of Cu

dissolving (eq 5 ). Compare with carbon

(graphite) electrodes.

Copper(II) sulphate with carbon (graphite)

electrodes, the copper deposits (-) and oxygen

gas (+), equations 4 and 8 .

Molten sodium chloride gives silvery

sodium (+) and pale green chlorine gas

(-), see equations 2 and 11 .  See

alsoextraction of sodium in

introduction. Compare with brine

solution.

Copper(II) chloride with carbon (graphite)

electrodes, the copper deposits (-) and chlorine

gas (+),  equations 2  and 4 .  Compare

withcopper electrodes.

Molten anhydrous zinc chloride gives

zinc (+) and chlorine (-), equations 1

and 2 .  

 

Method 1b

Dilute sulphuric acid

(= acidified

water) gives hydrogen

and oxygen gases (2 :

1 volume ratio)

via equations

3 and8 .    This is one

way of showing water

is a compound i.e. by

splitting into two

gaseous elements.

Molten anhydrous calcium

chloride gives calcium (+) and chlorine

(-), equations 13 and 2 .

Chemical Tests for some of the

Gases formed by these electrolysis

experiments.

hydrogen - colourless gas

that gives a squeaky pop

when ignited with a lit splint.

oxygen - colourless gas that

relights a glowing splint.

chlorine - green gas that

bleaches damp litmus

paper. Best done with blue

litmus which changes to red

before being bleached.

For more details and tests see the

qualitative chemical analysis page

Concentrated hydrochloric acid gives equal

volumes of hydrogen (-) and chlorine (+)

via equations 2 and 3. However in very dilute

solution,equation 8  occurs too, giving oxygen as

well as chlorine.

   

For most sulphate salts of reactive metals e.g.

sodium/magnesium sulphate the electrolysis

products of the aqueous salt solution

   

arehydrogen at the negative (-) cathode

electrode (Eq 3 ) and oxygen at the positive (+)

anode (Eq 8 ) with inert electrodes such as

carbon or platinum.

 

2b. Summary of ELECTRODE REACTIONS - half-cell electrode equations

 

 

 

Equation

number

(-) negative cathode electrode where reduction of

the attracted positive cations is by electron gain to form

metal atoms or hydrogen [from Mn+ or H+, n = numerical

positive charge]. The electrons come from the positive

anode (see below).

(+) positive anode electrode where the oxidation of the

atom or anion is by electron loss. Non-metallic negative

anions are attracted and may be oxidised to the free

element. Metal atoms of a metal electrode can also be

oxidised to form positive metal ions which pass into the

liquid electrolyte. The released electrons move round in the

external part of the circuit to produce the negative charge on

the cathode electrode.

So, before each electrode equation is a (-) for a negative

The electrode equations are shown on the

left with examples of industrial processes

where this electrode reaction happens

below on the right. Unless otherwise stated,

the electrodes are inert i.e. they do not

chemically change e.g. platinum or carbon-

graphite.

 PLEASE NOTE - all electrode equations are

a summary-simplification of what happens on

an electrode surface in electrolysis. There

may be e.g. two equations which are totally

equivalent to each other to describe WHAT

IS ACTUALLY FORMED e.g. the formation

of hydrogen or oxygen and in some cases

other products may be formed too.

cathode electrode = a reduction reaction equation or

a (+) for a positive anode electrode = an oxidation

reaction equation

1 a reduction electrode reaction

(-) Na+(l) + e- ==> Na(l) (sodium metal)

sodium ion reduced to sodium metal

atoms: typical of electrolysis of molten

chloride salts to make chlorine and the metal

2 an oxidation electrode reaction

(+) 2Cl-(l/aq) - 2e- ==> Cl2(g)

 or  2Cl- ==> Cl2 + 2e-

Note that you can write these anode oxidation reactions

either way round

chloride ion oxidised to chlorine gas

molecules: electrolysis of molten chloride

salts(l) or their concentrated aqueous

solution(aq) or conc. hydrochloric acid(aq) to

make chlorine

3 a reduction electrode reaction

(-) 2H+(aq) + 2e- ==> H2(g) (hydrogen gas)

or 2H3O+(aq) + 2e- ==> H2(g) + 2H2O(l)

or 2H2O(l) + 2e- ==> H2(g) + 2OH-(aq)

All three equations amount to the same overall change

i.e. the formation of hydrogen gas molecules and as far

hydrogen ion or water

reduced to hydrogen

gas molecules: electrolysis of many salt or

acid solutions to make hydrogen

as I know any is acceptable in an exam?

4

a reduction electrode reaction

(-) Cu2+(aq) + 2e- ==> Cu(s) (copper deposit)

copper(II) ion reduced to copper

atoms: deposition of copper in its electrolytic

purification or electroplating using copper(II)

sulphate solution, electrode can be copper or

other metal to be plated

5 an oxidation electrode reaction

(+) Cu(s) - 2e- ==> Cu2+(aq) (copper dissolves)

or  Cu(s) ==> Cu(s) + 2e-

copper atoms oxidised to copper(II)

ions: dissolving of copper in its electrolytic

purification or electroplating (must have

positive copper anode)

6 a reduction electrode reaction

(-) Al3+(l) + 3e- ==> Al(l) (aluminium)

aluminium ions reduced to aluminium

atoms: extraction of aluminium in the

electrolysis of its molten oxide ore(l) 

7 an oxidation electrode reaction

(+) 2O2-(l) - 4e- ==> O2(g) (oxygen gas)

or  2O2-(l)  ==> O2(g) + 4e-

oxide ion oxidised to oxygen gas

molecules: electrolysis of molten oxides e.g.

anode reaction in the extraction of aluminium

from molten bauxite.

8 an oxidation electrode reaction

(i) (+) 4OH-(aq) - 4e- ==> 2H2O(l) + O2(g) (oxygen gas)

There are two equations that describe the

formation of oxygen in the electrolysis of

water.

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

(ii) (+) 2H2O(l) - 4e- ==> 4H+(aq) + O2(g) (oxygen gas)

or 2H2O(l) ==> 4H+(aq) + O2(g) + 4e-

Both equations amount to the same overall change i.e.

the formation of oxygen gas molecules and as far as I

know either is acceptable in an exam?

hydroxide ions or water molecules are

oxidised to oxygen gas

molecules: electrolysis of many salt

solutions such as sulphates, sulphuric acid

etc. gives oxygen (chlorides ==> chlorine in

concentrated solution, but can also give

oxygen in diluted solution)

9  a reduction electrode reaction

(-) Pb2+(l) + 2e- ==> Pb(l) (lead deposit)

lead(II) ions reduced to lead

atoms: electrolysis of molten lead(II)

bromide(l) 

10 an oxidation electrode reaction

(+) 2Br-(l/aq) - 2e- ==> Br2(g/l) (bromine)

or  2Br- ==> Br2 + 2e-

bromide ions oxidised to gas/liquid

bromine molecules:electrolysis of molten

bromide salts(l) or their concentrated aqueous

solution(aq) or conc. hydrobromic acid(aq) to

make bromine

11a reduction electrode reaction

(-) Zn2+(aq) + 2e- ==> Zn(s) (zinc deposit)

zinc ions reduced to zinc

atoms: galvanising steel (the electrode) by

electroplating from aqueous zinc sulphate

solution, (or from molten zinc chloride?)

12 a reduction electrode reaction silver ions reduced to silver atoms: silver

electroplating from silver salt solution(aq),

electrode can be other metal

(-) Ag+(aq) + e- ==> Ag(s) (silver deposit)

13 a reduction electrode reaction

(-) Ca2+(l) + 2e- ==> Ca(s) (calcium metal)

calcium ions reduced to calcium

atoms e.g. in molten calcium chloride or

bromide etc.

Electrolysis calculations  section 13. of the  Chemical Calculations pages index

 

3a. Summary of the products of electrolysis of various electrolytes

What are the products of the electrolysis of molten aluminium oxide, aqueous copper sulphate solution, aqueous sodium chloride

solution (brine), hydrochloric acid, sulphuric acid, molten lead(II) bromide, molten calcium chloride?

Electrolyt

e

cathode

productcathode equation

anode

produc

t

anode equation comments

molten

aluminiu

m oxide

Al2O3(l)

molten

alumini

um

Al3+(l) + 3e- ==> Al(l)

oxygen

gas

2O2-(l) - 4e- ==> O2(g)

or  2O2-(l)  ==> O2(g) + 4e-

Method 2 in principle:

The industrial method

for the extraction of

aluminium its ore

aqueous copper any conducting electrode e.g. oxygen inert electrode like carbon Method

copper(II)

sulfate

CuSO4(aq)

deposit

carbon rod, any metal including

copper itself

Cu2+(aq) + 2e- ==> Cu(s)

gas

(graphite rod) or platinum

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

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

(ii) 2H2O(l) - 4e- ==> 4H+(aq) + O2(g)

or 2H2O(l) ==> 4H+(aq) + O2(g) + 4e-

1a or Adapt   Method 3 -

e.g. use carbon rods:

The blue colour of the

copper ion will fade as

the copper ions are

converted to the copper

deposit on the cathode

aqueous

copper (II)

sulphate

CuSO4(aq)

copper

deposit

any conducting electrode e.g.

carbon rod, any metal including

copper itself

Cu2+(aq) + 2e- ==> Cu(s)

copper(

II) ions

- the

copper

anode

dissolv

es

copper anode

Cu(s) - 2e- ==> Cu2+(aq)

or  Cu(s) ==> Cu(s) + 2e-

See Method 3: This is

the basis of the method

of electroplating any

conducting solid with a

layer of copper. When

using both copper

cathode and anode, the

blue colour of the

copper ion does not

decrease because

copper deposited at the

(-) cathode = the copper

dissolving at the (+)

anode.

molten

sodium

molten

sodiumNa+

(l) + e- ==> Na(l)chlorin

e gas

2Cl-(l) - 2e- ==> Cl2(g) Method 2 in principle:

This a method used to

chloride

NaCl(l) or  2Cl-

(l) ==> Cl2(g) + 2e-

manufacture sodium

and chlorine. See

Down's Cell

aqueous

sodium

chloride

solution

(brine)

NaCl(aq)

hydroge

n

2H+(aq) + 2e- ==> H2(g)

or 2H3O+(aq) + 2e- ==> H2(g) + 2H2O(l)

or 2H2O(l) + 2e- ==> H2(g) + 2OH-(aq)

chlorin

e gas

2Cl-(aq) - 2e- ==> Cl2(g)

 or  2Cl-(aq) ==> Cl2(g) + 2e-

Method 1a: This is the

process by which

hydrogen, chlorine and

sodium hydroxide are

manufactured

hydrochlo

ric acid

HCl(aq)

hydroge

n gas

2H+(aq) + 2e- ==> H2(g)

or 2H3O+(aq) + 2e- ==> H2(g) + 2H2O(l)

For theoretical calculations of

gas volumes formed see Section

14 of Chemical Calculations

chlorin

e gas

2Cl-(aq) - 2e- ==> Cl2(g)

 or  2Cl-(aq) ==> Cl2 + 2e- 

Method 1a and method

1b: All acids give

hydrogen at the

cathode.

Theoretically the gas

volume ratio is H2:Cl2 is

1:1, BUT, chlorine is

slightly so there seems

less chlorine formed

than actually was.

sulphuric

acid

hydroge

n gas

2H+(aq) + 2e- ==> H2(g)

For theoretical calculations of

oxygen

gas

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

Method 1a and method

1b: All acids give

hydrogen at the

sulfuric

acid

H2SO4(aq)

gas volumes formed see Section

14 of Chemical Calculations

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

(ii) (+) 2H2O(l) - 4e- ==> 4H+(aq) + O2(g)

or 2H2O(l) ==> 4H+(aq) + O2(g) + 4e- 

cathode. Whereas

hydrochloric acid gives

chlorine at the anode,

the sulfate ion does

nothing and instead

oxygen is formed. This

is the classic

'electrolysis of water'.

Theoretically the gas

volume ratio is H2:O2 is

2:1 which you see with

the Hofmann

Voltameter

molten

lead(II)

bromide

PBr2(l)

molten

leadPb2+

(l) + 2e- ==> Pb(l)

bromin

e

vapour

2Br-(l) - 2e- ==> Br2(g)

 or  2Br-(l) ==> Br2(g) + 2e-

Method 2: A good

demonstration in the

school laboratory -

brown vapour and

silvery lump provide

good evidence of

what's happened

molten

calcium

chloride

solid

calcium

Ca2+(l) + 2e- ==> Ca(s) chlorin

e gas

2Cl-(aq) - 2e- ==> Cl2(g)

 or  2Cl-(aq) ==> Cl2(g) + 2e-

Method 2 in principle:

The basis of the

industrial method for

CaCl2(l)

the manufacture of

calcium metal

           

************

**

**********

**

****************************************

************

*********

*

*******************************************

************

*****************************

********

3b. Summary of Industrial Processes using Electrolysis

Below is a list of links to other web pages on this site with sections on electrolysis.

  extraction of aluminium  from purified molten bauxite ore.

  Anodising aluminium  to strengthen the protective oxide layer.

  extraction of sodium  from molten sodium chloride using the 'Down's Cell'.

  purification of copper  using copper electrodes.

  purification of zinc

Making chlorine, hydrogen and sodium hydroxide from sodium chloride salt solution by

o   brine electrolysis  and also an    electrolysis Q on the Halogens task sheet

  Electroplating  conducting surfaces to coat it with a metal layer.

 

4. Simple Cells or batteries

In electrolysis, electrical energy is taken in (endothermic) to enforce

the oxidation and reduction to produce the products.

The chemistry of simple voltaic cells or batteries is in principle

the opposite of electrolysis.

A redox reaction occurs to produce products and energy is given out

because it is an exothermic reaction, BUT the energy is released

as electrical energy NOT heat energy so the system shouldn't heat up.

A simple cell can be made by dipping two different pieces of metal (of different reactivity) into a solution of

ions e.g. a salt or dilute acid. If you use the same metal for both strips, they 'cancel' each other out, so no

potential difference (voltage) so no current of electrical energy.

All you need is a solution of charged positive and negative particles called ions e.g. sodium Na+,

chloride Cl-, hydrogen H+, sulphate SO42- etc.

The greater the difference in reactivity, the bigger the voltage produced. However this is not a

satisfactory 'battery' for producing even a small continuous current.

BUT a simple demonstration cell can be made by dipping strips of magnesium and copper into an a salt

solution and connecting them via a voltmeter (e.g. as in diagram) and a voltage is readily recorded.

o The electrode reactions are:

at the (+) electrode 2H+(aq) + 2e- ==> H2(g) (hydrogen ions reduced)

here the copper is inert and the hydrogen ions come from water.

at the (-) electrode Mg(s) - 2e- ==> Mg2+(aq) (magnesium atoms oxidised)

the magnesium dissolves into solution by a chemical reaction

overall the redox reaction is 2H+(aq) + Mg(s) ==> Mg2+

(aq) + H2(g) 

and the electrons from the oxidation of the magnesium move round through the magnesium

strip, along the external wire to the copper electrode.

Note the (+) and (-) polarity of the electrodes in a cell, is the opposite of electrolysis because

the process is operating in the opposite direction i.e.

in electrolysis electrical energy induces chemical changes,

but in a cell, chemical changes produce electricity.

One of the first practical batteries is called the 'Daniel cell' which is illustrated below.

o

o This 'voltaic 'or galvanic' electrochemical cell uses a half-cell of copper dipped in copper(II) sulphate,

o and in electrical contact with a 2nd half-cell of zinc dipped in zinc sulphate solution.

o The zinc is the more reactive, and is the negative electrode, releasing electrons because

on it zinc atoms lose electrons to form zinc ions, Zn(s) ==> Zn2+(aq) + 2e-

o The less reactive metal copper, is the positive electrode, and gains electrons from the negative

electrode through the external wire connection and here ..

the copper(II) ions are reduced to copper atoms, Cu2+(aq) + 2e- ==> Cu(s)

o Overall the reactions is: Zn(s) + CuSO4(aq) ==> ZnSO4(aq) + Cu(s)

or ionically: Zn(s) + Cu2+(aq) ==> Zn2+

(aq) + Cu(s)

o The overall reaction is therefore the same as displacement reaction, and it is a redox reaction

involving electron transfer and the movement of the electrons through the external wire to the bulb

or voltmeter etc. forms the working electric current.

The cell voltage can be predicted by subtracting the less positive

voltage from the more positive voltage:

o e.g. referring to the list of electrode potentials on the right

and choosing two different metals coupled together in the

electrolyte solution ...

o a magnesium and copper cell will produce a voltage of

(+0.34) - (-2.35) = 2.69 Volts

o or an iron and tin cell will only produce a voltage of (-0.15) -

(-0.45) = 0.30 Volts.

o Note (i) the bigger the difference in reactivity, the bigger the cell voltage produced

o and (ii) the 'half-cell' voltages quoted in the diagram are measured against the H+(aq)/H2(g) system

which is given the standard potential of zero volts.

Cells or batteries are useful and convenient portable sources of energy for torches, radios, shavers

and other gadgets BUT they are expensive compared to what you pay for 'mains' electricity. On the other

hand you have no choice for a car battery!

  Electrolysis and cell theory for Advanced Level Chemistry Students

 

5. Fuel Cells - another sort of battery

Hydrogen gas can be used as fuel.

o It burns with a pale blue flame in air reacting with oxygen to be oxidised to form water.

hydrogen + oxygen ==> water

2H2(g) + O2(g) ==> 2H2O(l) 

o It is a non-polluting clean fuel since the only combustion product is water. 

o It would be ideal if it could be manufactured by electrolysis of water e.g. using solar cells.

o Hydrogen can be used to power fuel cells. see    GCSE/IGCSE-AS notes on fuels

o  It all sounds wonderful BUT, still technological problems to solve for large scale manufacture and

distribution of 'clean' hydrogen gas or use in generating electricity AND its rather an inflammable

explosive gas!

Fuel cells are 'battery systems' in which two reactants can be continuously fed in. The consequent

redox chemistry produces a working current.

Hydrogen's potential use in fuel and energy applications includes powering vehicles, running turbines or fuel

cells to produce electricity, and generating heat and electricity for buildings and very convenient for remote

and compact situations like the space shuttle.

When hydrogen is the fuel, the product of its oxidation is water, so this is potentially a clean non-polluting

and non-greenhouse gas? fuel.

Most fuel cells use hydrogen, but alcohols and hydrocarbons can be used.

A fuel cell works like a battery but does not run down or need recharging as long as the 'fuel' supply is there.

It will produce electricity and heat as long as fuel (hydrogen) is supplied.

DIAGRAM and CHEMISTRY below: A fuel cell consists of two electrodes consisting of a negative electrode

(or anode) and a positive electrode (or cathode) which are sandwiched around an electrolyte (conducting

salt/acid/alkali solution of free ions).

Hydrogen is fed to the (-) anode, and oxygen is fed to the (+) cathode.

The platinum catalyst activates the hydrogen atoms/molecules to separate into protons (H+) and electrons

(e-), which take different paths to the (+) cathode.

o The electrons go through an external circuit, creating a flow of electricity e.g. to light a bulb.

o The protons migrate through the electrolyte and pass through the semi-permeable membrane to the

cathode, where they reunite with oxygen and the electrons to produce water.

Each cell only produces a small voltage (typically 0.4 to 1.0V) so many cells can be put together in series

to give a bigger working voltage.

Note on reverse action:

o If there is spare electricity from another source available, you can run the fuel cell in reverse and

electrolyse the water to make hydrogen and oxygen (acting as an electrolyser).

o The two gases are stored, and when electricity or heat needed, the fuel cell can then be re-run using

the stored gaseous fuel.

o This is called a regenerative fuel cell system.

o You can use solar energy from external panels on the space shuttle to do this, and use the fuel

when in the 'darkness of night'.

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equation

Description of a hydrogen-oxygen fuel cell

It uses costly platinum electrodes and an acid electrolyte such as

phosphoric acid, H3PO4

1. oxidation 2H2(g) ==> 4H+(aq) + 4e-  (at negative anode electrode*)

2. reduction O2(g) + 4H+(aq) + 4e- ==> 2H2O(l) (at positive cathode electrode*)

3 = 1 + 2 redox 2H2(g) + O2(g) ==> 2H2O(l)

* Note the +ve and -ve electrode charges are reversed compared to

electrolysis, because the system is operating in the opposite direction.

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