2B OXIDATION & REDUCTION –
REDOX
REACTIONS –
An
INTRODUCTION for GCSE/IGCSE/O
Level students
OXIDATION The oxidation
reactions you are most likely to come across at first in your chemistry
course are the reaction of metals with oxygen e.g. by heating a strip of
metal in a strongly in a bunsen flame.
The metal is considered to be oxidised because the metal
gains and combines with oxygen
atoms from the oxygen molecules in air e.g.
(i) magnesium + oxygen ===> magnesium oxide
2Mg + O2 ===> 2MgO
The magnesium ribbon burns
brightly to form the white oxide powder, oxidised – oxygen gain.
(ii) copper + oxygen ===> copper
oxide
2Cu + O2 ===> 2CuO
The copper doesn't burn like
magnesium (not as reactive), but the black oxide forms on the
copper surface.
The chemical species that gains oxygen
is oxidised. In both
cases you can consider the
oxidation in terms of electron
loss For Mg or Cu:
M ==> M2+ + 2e-
The metal atoms are oxidised
by electron loss to give the metal ion.
Oxidation is electron loss by
something (atom,
ion or molecule).
These are called half equations and show the
movement or transfer of electrons
in a chemical reaction.
Lots more on the
electron definitions of oxidation & reduction with examples
explained.
REDUCTION
The opposite of oxidation is reduction and one simple definition of it is
oxygen loss
(opposite of oxidation above!).
You most likely to come cross this first when
studying the extraction of metals from a metal oxide (maybe as an ore)
For example, if you strongly heat a mixture of copper
oxide with carbon (charcoal or graphite powder) you find bits of
brown–orange copper are formed.
copper oxide + carbon ===> copper
+ carbon dioxide
2CuO + C ==> 2Cu + CO2
The copper oxide has lost oxygen, so copper oxide has been
reduced. Any chemical species
that loses oxygen has been reduced.
Anything that removes oxygen is called
the reducing agent (carbon here).
The carbon gains
oxygen and so is oxidised.
Anything that gains oxygen is
oxidised. Anything that donates oxygen is called the oxidising
agent (copper oxide here)
Another way to look at the
reduction is from the electron transfer point of view.
Reduction is electron
gain by something (atom, ion or molecule). The carbon
donates 2 electrons to the copper ion in the copper oxide
compound.
Cu2+
+ 2e- ==> Cu
For more details see
of some of these reactions see notes on the
Metal Reactivity Series,
Metal Reactivity Series
Experiments–Observations
Extraction of
Metals
Advanced Level Chemistry Redox Reaction
Notes
(it repeats this introduction and then moves on!)
Definition reminders
and lots more EXAMPLES of OXIDATION and
REDUCTION In the examples in the first
double column section below, the equations are not meant to be complete
or balanced. What I've highlighted, is the
chemical change that
fits the definition of oxidation and reduction,
in other words the
transfer-movement of oxygen or electrons.
|
OXIDATION – definition
and examples
(based on
oxygen gain OR electron loss transfer) |
REDUCTION – definition
and examples
(based on
oxygen loss OR electron gain transfer) |
(a)
Oxidation is the gain or addition of oxygen by an atom, molecule or ion
e.g. ...
(1) S ==> SO2 [burning sulphur
– oxidised]
(2) CH4 ==> CO2
+ H2O [burning (oxidation) of methane to water and carbon dioxide, C and H
gain O]
(3) NO
==> NO2 [nitrogen monoxide oxidised to nitrogen dioxide]
(4) SO2
==> SO3 [oxidising the sulphur dioxide to sulphur trioxide in the Contact
Process for making sulphuric acid] |
(b)
Reduction is the loss or removal of oxygen from a compound etc.
e.g. ...
(1) CuO ==> Cu [loss of oxygen from
copper(II) oxide to form copper atoms]
(2) Fe2O3
==>
Fe [iron(III) oxide reduced to iron in blast furnace]
(3) NO ==> N2 [nitrogen
monoxide reduced to nitrogen, catalytic converter in car exhaust]
(4) SO3 ==> SO2
[sulphur trioxide reduced to sulphur dioxide] |
(c)
Oxidation is the loss or removal of electrons from an atom, ion or molecule
e.g.
the balanced half equations (1) Fe ==> Fe2+ +
2e– [iron
atom loses 2 electrons to form the iron(II) ion, start of rusting
chemistry]
(2) Fe2+ ==> Fe3+
+ e– [the iron(II) ion loses
1 electron to form the iron(III) ion, also part of rusting chemistry]
(3) 2Cl–
==> Cl2
+ 2e– [the loss of electrons by chloride ions to form chlorine molecules e.g.
in electrolysis] |
(d)
Reduction is the gain or addition of electrons by an atom, ion or molecule
e.g. the balanced half equations (1)
Cu2+ + 2e–
==> Cu
[the copper(II) ion gains 2 electrons to form neutral copper
atoms, electroplating or displacement reaction)
(2)
Fe3+ + e–
==> Fe2+ [the iron(III) ion gains an electron
and is reduced to the iron(II) ion]
(3) 2H+ + 2e– ==>
H2 [hydrogen ions gain electrons to form neutral hydrogen molecules,
electrolysis of acids or metal–acid reactions] |
(e)
An
oxidising agent is the species that gives the oxygen
to an atom, ion or molecule OR
an
oxidising agent accepts
electrons i.e. an oxidising agent removes
electrons from some atom, ion or molecule.
In either process
the oxidising agent gets reduced. |
(f)
A
reducing agent is the species that removes the oxygen
from an atom, ion or molecule OR
a
reducing agent acts as
the electron donor i.e. it gives electrons to some atom, ion or
molecule.
In either case the
reducing agent gets oxidised. |
REDOX REACTIONS – in a
reaction overall, BOTH oxidation and reduction must go together |
(g)
Redox
reaction analysis based on the oxygen definitions of oxidation and
reduction |
-
(1) copper(II) oxide +
hydrogen
==> copper + water
-
CuO(s) + H2(g)
==> Cu(s) + H2O(g)
-
- Copper oxide reduced (oxygen loss) to copper by
hydrogen, hydrogen is
oxidised to water,
- hydrogen is the reducing agent (removes O from CuO)
and gets oxidised to water in the process,
- copper oxide is the oxidising agent (donates O to hydrogen).
-
(2) iron(III) oxide +
carbon monoxide
==> iron + carbon dioxide
-
Fe2O3(s) +
3CO(g)
==> 2Fe(l) + 3CO2(g)
-
- The iron(III) oxide is reduced to iron
(oxygen loss)
by the CO and
the carbon monoxide is oxidised to carbon dioxide,
- CO is the reducing
agent (removes oxygen from Fe2O3) and gets
oxidised in the process to CO2,
- the Fe2O3
is the oxidising agent (O donator to CO).
-
(3) nitrogen
monoxide + carbon monoxide ==> nitrogen + carbon dioxide
-
2NO(g) + 2CO(g) ==> N2(g)
+ 2CO2(g)
-
- The nitrogen monoxide is reduced to
nitrogen by oxygen loss,
- carbon
monoxide is oxidised to carbon dioxide by oxygen gain,
- CO is the reducing agent (accepts
oxygen) and gets oxidised in the process,
- and the NO is
the oxidising agent (donates oxygen to the CO) and is reduced to
nitrogen in the process.
-
(4) iron(III) oxide +
aluminium ==> aluminium oxide + iron (the Thermite reaction)
-
Fe2O3(s)
+ 2Al(s) ==> 2Fe(s) + Al2O3(s)
-
- Iron(III) oxide is reduced, oxygen
loss, and is the oxidising
agent, it oxidises aluminium,
- aluminium is
oxidised, oxygen gain, and Al is the reducing agent, it removes oxygen from
the iron oxide.
|
(h)
Redox
reaction analysis based on the electron definitions of oxidation and
reduction |
- (1)
Examples metal – metal salt solution
displacement reactions are described and explained in detail.
-
(a) magnesium + iron(II)
sulphate
==> magnesium sulphate + iron
-
Mg(s) + FeSO4(aq)
==> MgSO4(aq) + Fe(s)
- this is the 'ordinary molecular' equation for a typical metal displacement reaction, but this does not really show what happens in terms of atoms,
ions and electrons, so we use ionic equations like the one shown
below.
- The sulphate ion SO42–(aq)
is called a spectator ion, because it doesn't change in the
reaction and can be omitted from the ionic equation.
- No electrons show
up in the full equations because electrons lost by x = electrons gained
by y!!
- magnesium + iron(II)
ion ==> magnesium ion + iron
- The fully balanced symbol ionic equation
is ...
-
Mg(s) + Fe2+(aq)
==> Mg2+(aq) + Fe(s)
- the magnesium atom, Mg, loses 2 electrons
(oxidation) to form the magnesium ion.
- The iron(II) ion, Fe2+, gains 2
electrons (reduced) to form iron atoms.
- You can think of it as two 'half
reactions' involving the movement and
transfer of electrons
- Mg is the reducing agent
(electron donor) and the
Fe2+
is the oxidising agent (electron remover or acceptor).
- The more reactive metal (magnesium)
always displaces the less reactive metal (iron).
- Displacement reactions
involving metals and metal ions are electron transfer
reactions.
- In metal – soluble metal salt
displacement reactions, the metal atom always loses electrons
(oxidation) and the metal ion always gains electrons
(reduction).

-
(b) Exactly the same reaction occurs if you
add iron filings to copper sulfate solution.
- A brown precipitate of copper forms on
the surface of the iron filings and the blue colour fades as the
less reactive copper is displaced by the more reactive iron.
- The
iron sulfate formed in solution is a very pale green so the
final solution is almost colourless if excess iron is added.
- iron + copper(II) sulfate ==> iron(II)
sulfate + copper
-
Fe(s) +
CuSO4(aq)
==> FeSO4(aq) + Cu(s)
- iron + copper(II)
ion ==> iron(II) ion + copper
- The fully balanced symbol ionic equation
is ...
-
Fe(s)
+ Cu2+(aq) ==> Fe2+(aq) + Cu(s)
- The sulfate ion is a spectator ion
and is NOT shown in the ionic equation.
- Again, to fully understand what is
happening, you can think of it as two 'half–reactions''
involving the movement and transfer of electrons
- Iron (Fe) is the reducing agent
(electron donor) and the
copper(II) ion (Cu2+)
is the oxidising agent (electron remover or acceptor).
-
- -

-
(c) Adding magnesium to blue copper(II)
sulfate solution
- The blue colour fades as
colourless magnesium sulfate is formed and brown bits of copper metal form a
precipitate around the magnesium ribbon:
- magnesium + copper sulfate ==> magnesium
sulfate + copper
-
- -
(d)
copper + silver nitrate
==> silver + copper(II) nitrate
- Even some of the less reactive metals still
facilitate displacement reactions.
-
Cu(s) + 2AgNO3(aq) ==> 2Ag + Cu(NO3)2(aq)
- the nitrate ion NO3– is the spectator
ion
- copper atom + two silver ions
==> two silver atoms + copper(II) ion
- The fully balanced symbol ionic equation
(redox equation)
is ...
-
Cu(s) + 2Ag+(aq) ==> 2Ag(s) + Cu2+(aq)
- The more reactive copper atoms are oxidised by the silver
ion by electron loss
- Electrons are transferred from
the copper atoms to the silver ions, which are reduced
- the silver ions are the
oxidising agent and the copper atoms are the reducing agent
- Again, you can think of it as
two 'half–reactions'
-
Cu
==> Cu2+ + 2e–
(the oxidation, electron loss, copper atom oxidised)
-
2Ag+
+ 2e– ==> 2Ag
(the reduction, electron gain, silver ion reduced)
- again, the electron loss and
gain cancel out, so you don't see them in the full equation.
- The full redox ionic equation is
derived from combining the half equations in the right
proportions.
- Note that in this case, you need two
lots of the silver half equation to one of the copper half
equation to give the correctly balanced full ionic equation.
Note
2 electrons lost
balances 2 x 1 electron gained
- The more reactive metal (copper)
always displaces the less reactive metal (silver).
-
- -
-
(e)
Displacement of iron from its
salts by more reactive metals.
- Iron forms more than one series of
compounds - a characteristic of
transition metals.
- e.g. if you add more reactive zinc or
magnesium to iron(II) sulfate OR iron(III) sulfate
solution, iron is deposited on the zinc or magnesium surface.
- You should also see that the magnesium
reacts faster than the zinc.
- I am assuming you can write out the word
equations.
- (i) Iron(II) sulfate displacement
reactions
-
Mg(s) + FeSO4(aq)
===> MgSO4(aq) + Fe(s)
(already done in (a) above)
-
Zn(s) + FeSO4(aq)
===> ZnSO4(aq) + Fe(s)
- The half
equations are
- oxidations
- electron loss:
Zn ==> Zn2+
+ 2e-
- reduction
- electron gain:
Fe2+
+ 2e- ==> Fe
- These can now be combined to give the
complete balanced ionic equation
(redox equation)
-
Zn(s) + Fe2+(aq)
===> Zn2+ + Fe(s)
-
- -
- (ii) iron(III) sulfate displacement
reactions: (these are a bit more tricky to
balance!)
-
3Mg(s) + Fe2(SO4)3(aq)
===> 3MgSO4(aq) + 2Fe(s)
-
3Zn(s) + Fe2(SO4)3(aq)
===> 3ZnSO4(aq) + 2Fe(s)
- The half
equations are
- oxidations
- electron loss:
Mg ==>
Mg2+ + 2e-
and
Zn ==> Zn2+
+ 2e-
- reduction
- electron gain:
Fe3+
+ 3e- ==> Fe
- Again, these can now be combined to give
the complete balanced ionic equations
(redox equations)
-
3Mg(s) +
2Fe3+(aq) ===> 3Mg2+ +
2Fe(s)
-
3Zn(s) +
2Fe3+(aq) ===> 3Zn2+ +
2Fe(s)
- -
-
(2) Metal - acid reaction:
- The reaction between a metal and an
acid is also, technically, what is called a REDOX reaction.
-
This means the reaction takes place
in two parts, an oxidation involving electron loss and a reduction involving electron gain.
- If a metal is above hydrogen in reactivity,
it can displace hydrogen from the acid - theoretically anything at
least reactive as lead - which reacts VERY slowly with acids AND you
never add sodium or potassium to acid - check out the chart !!!
- Here the metal atoms like Zn lose electrons to
form positive ions (oxidation).
- The hydrogen ions, H+, gain electrons to
form hydrogen gas molecules (reduction).
- The two simultaneous changes
occur on the surface of the metal where the positive hydrogen ions
hit the metal surface and pinch
electrons from the metal and so metal ions pass into solution e.g.
- zinc + hydrochloric
acid
==> zinc chloride + hydrogen
-
Zn(s) + 2HCl(aq)
==> ZnCl2(aq)
+ H2(g)
- the chloride ion Cl– is the spectator
ion which doesn't take part in the chemical changes.
- zinc + hydrogen ion
==> zinc ion + hydrogen
- The fully balanced symbol ionic equation
(redox equation)
is ...
-
Zn(s) +
2H+(aq) ==> Zn2+(aq)
+ H2(g)
- The zinc transfers electrons to the hydrogen
ions.
- Zinc atoms are oxidised to zinc ions by electron
loss, so zinc is the reducing agent (electron donor)
- Hydrogen ions are the oxidising agent (gaining the electrons) and
are reduced to form hydrogen molecules
- Again, you can think of it as
two 'half–reactions'
-
Zn ==> Zn2+
+ 2e–
(the oxidation half equation, electron loss, zinc atom is
oxidised)
-
2H+
+ 2e– ==> H2
(the reduction half equation, electron gain, hydrogen ion
reduced)
- again, the electron loss and
gain cancel out, so you don't see them in the full equation.
-
-
More on
acid reactions including with metals
-
(3)
Halogen displacement
reactions
- halogen (more reactive)
+ halide salt (of less reactive halogen) ==> halide salt
(of
more reactive halogen) + halogen (less reactive)
-
X2(aq) + 2KY(aq)
==> 2KX(aq) + Y2(aq)
- The potassium ion is a spectator ion, so
the ionic equation is ...
-
X2(aq) + 2Y–(aq)
==> 2X–(aq) + Y2(aq)
- where halogen X is more
reactive than halogen Y, F > Cl > Br > I
- Halogen X is reduced by electron
gain and halide ion Y– is oxidised by electron
loss.
- in principle: X + e–
==> X– (electron gain, reduction of
halogen atom X)
- but more correctly:
X2 +
2e– ==> 2X–
(electron gain, reduction of halogen molecule X2)
- in principle: Y–
==> Y + e– (electron loss,
oxidation of halide ion Y–)
- but more correctly:
2Y–
==> Y2 + 2e–
(electron loss, oxidation of halide ion Y– to give
molecule Y2)
- X2
is the oxidising agent
(electron acceptor)
- KY (via
Y–)
is the reducing agent
(electron donor)
-
See
GCSE Group 7 The Halogens
– displacement reaction notes for more details
-
(4) If
chlorine is bubbled into
iron(II) chloride solution, iron(III) chloride is formed
- Or you can just mix chlorine water
with iron(II) chloride solution, a bit safer!
-
iron(II) chloride + chlorine ==>
iron(III) chloride
-
2FeCl2 + Cl2
==> 2FeCl3
-
2FeCl2(aq) + Cl2(aq)
==> 2FeCl3(aq)
- The iron(II) is oxidised to the
iron(III) ion via the chlorine by electron loss:
-
Fe2+
– e– ==> Fe3+ (an oxidation
half equation)
- The chlorine molecule is reduced to
the chloride ion by electron gain:
-
Cl2
+ 2e– ==> 2Cl– (a reduction half
equation)
- Chlorine is the oxidising agent
because it gains electrons, the electron acceptor or remover,
- and the iron(II) ion is the reducing
agent because it donates or give electrons to the chlorine molecule.
- If you put two of the two
'half–equations' together, you get the following full redox–ionic
equation to match the 'molecular equation above. The original
chloride ions are effectively spectator ions.
- The full balanced symbol ionic equation
is ...
-
2Fe2+ + Cl2
==> Fe3+ + 2Cl–
-
2Fe2+(aq)
+ Cl2(aq)
==> 2Fe3+(aq)
+ 2Cl–(aq)
-
(5) Electrode reactions in
electrolysis are electron transfer redox changes
- at the negative cathode
positive
ions are attracted:
- metal ions are reduced to the metal by electron gain.
-
Mn+
+ ne– ==> M
- n = the numerical
charge of the ion and the number of electrons transferred
and gained by the cation (positive ion).
- or
2H+(aq)
+ 2e– ==> H2(g) (for the
discharge of hydrogen)
- at the positive anode
negative ions are attracted:
- negative non–metal ions are
oxidised by electron loss e.g. the half equations are
- for oxide ions:
2O2–
– 4e– ==> O2 or
2O2–
==> O2 + 4e–
- for hydroxide ion:
4OH–
– 4e–
==> O2 + 2H2O
or
4OH– ==> O2 + 2H2O
+ 4e–
- for halide ions (X = F, Cl,
Br, I):
2X– – 2e–
==> X2
or
2X– ==> X2 + 2e–
|
2C Miscellaneous
Extra Redox
Notes
Advanced Level Chemistry Redox Reaction
Notes (it repeats this introduction and then moves on!)
TOP OF PAGE
and sub-index
OTHER ASSOCIATED PAGE
LINKS
SEE ALSO
2. RUSTING &
Introducing REDOX reactions
and 3.
Metal Reactivity Series
Experiments–Observations
and GCSE/IGCSE m/c QUIZZES on metal
reactivity
Foundation–tier Level
(easier) multiple choice quiz on the Reactivity Series of Metals
or
Higher–tier Level (harder) multiple choice quiz on
the Reactivity Series of Metals
and
GCSE/IGCSE
reactivity gap–fill worksheet or
Rusting
word–fill worksheet
Rates of Reactions Experiments (e.g.
metal–acid)
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