explaining electrolysis and experimental method
Post on 30-Nov-2015
33 Views
Preview:
DESCRIPTION
TRANSCRIPT
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'.
*
*
*
*
*
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.
*
*
*
*
*
*
top related