Redox Reactions
You are full of redox reactions. The redox reaction respiration releases the energy you need
to live, and the food you eat ultimately comes from the redox reaction photosynthesis.
Around you, mobile and laptop batteries work using redox reactions, as do fuel cells,
bleaches, old fashioned wet photography, and metal corrosion, such as rusting iron.
So what are redox reactions?
The redox comes from two words: Reduction and Oxidation. In a reaction, oxidation is the
loss of electrons from an atom or ion, and reduction is the gain of electrons by an atom
or ion. In a redox reaction there is the gain of electrons by one chemical, reduction, and
the loss of electrons by another chemical, oxidation, so both are found together. You must
remember
• oxidation is the process of electron loss
• reduction is the process of electron gain
Redox Reactions
An example of redox
When the metal sodium reacts with the green gas chlorine then the white solid sodium
chloride is made
The sodium has been oxidized and the chlorine has been reduced.
Redox Reactions
Ancient reduction
You may wonder why the word reduction is used to mean a gain of electrons. Thousands of
years ago it was noticed that the metal made by smelting had less mass than the original
ore. The ore was made smaller, reduced. The loss of oxygen made the mass less. More
recently it was realised that the metal atoms were gaining electrons.
Oxidizing agents and reducing agents
Oxidizing agents oxidize other chemicals, so oxidizing agents are themselves reduced,
so oxidizing agents gain electrons. Reducing agents reduce other chemicals, so they lose
electrons. You must remember that
• oxidizing agents are electron acceptors
• reducing agents are electron donors
Why ‘oxidation’?
Oxidation used to mean just ‘gain of oxygen’, but it was realized, more importantly, that
the other chemical was losing electrons to the oxygen. So it was decided that ‘oxidation’
should have the broader meaning ‘loss of electrons’.
Redox Reactions
White hot fire
To allow trains to travel at high speeds the rails must be welded together so there are no
gaps. This must be done in isolated places so molten iron is made using the Thermit
reaction:
Here the iron(III) oxide, Fe2O3, is reduced to iron, Fe, so
• the iron(III) oxide is reduced,
• the iron(III) oxide is the oxidizing agent
The aluminium powder, Al, is oxidized to aluminium oxide, Al2O3, so
• the aluminium is oxidized
• the aluminium is the reducing agent
Half Equations
What are half-equations?
Half equations show the gain or loss of electrons by one chemical. For example, this is the full equation for when
sodium reacts with chlorine:
Each sodium atom is losing an electron to a chlorine atom, so you could write this half equation:
Each chlorine atom in a chlorine molecule gains an electron, so you could write this half equation:
(You need to show two chloride ions because each chlorine molecule, Cl2, contains two chlorine atoms.) The way
sodium reacts does not depend on the other reactant. For example, if sodium reacted with bromine instead of
chlorine the equation would be:
Sodium is still gaining electrons in the same way;
and the half equation for bromine becomes:
Half Equations
Half-equations and redox
Half-equations involve electron gain or loss, so they always are either
oxidation or reduction. In this reaction an electron is lost, so this is oxidation:
This time electrons are gained, so this is reduction;
Why ‘half’?
In a reaction, when one chemical loses electrons then another must gain them. Half-equations only
show half the story, either the gain or the loss of electrons.
Some ions are spectators
Ions that take no part in the reaction are called spectator ions. When zinc is put into copper sulphate
solution then the full equation is:
Each zinc atom is oxidized:
Each copper ion is reduced:
Notice that the sulphate ions, SO4 2–, do not appear in the equations. They do not change during the
reaction. They are dissolved in the water, so are just floating around. As they are said to only ‘watch’ the
reaction they are called ‘spectator ions’.
Oxidation States
What is an oxidation state?
An oxidation state is the number of electrons needed to be gained or lost to make a
neutral atom. Using oxidation states is a way of working out how oxidized or reduced
something is. It is similar to the charge on ions, except that it is also used for covalent
compounds.
Different oxidation states may have different
colours. In the test tube are all the oxidation
states of vanadium from pale yellow +5 to
violet +2 at the bottom. Two oxidation states of
manganese produced the colours at the top.
How to Work out Oxidation States
Use these rules to calculate oxidation state:
• Elements always have an oxidation state of zero.
• In a compound, the sum of the oxidation numbers equals zero.
• In an ion, the sum of the oxidation numbers equals the charge.
In a compound
• Group 1 atoms always have a +1 oxidation state, e.g. Na is +1 in NaCl.
• Group 2 atoms always have a +2 oxidation state, e.g. Mg is +2 in MgCl2.
• Group 3 atoms always have a +3 oxidation state, e.g. Al is +3 in AlCl3.
• Fluorine always has a –1 oxidation state, e.g. F is –1 in KF.
• Oxygen has a –2 oxidation state, unless it is in a peroxide compound, such as H2O2, when
O is –1, or with fluorine (as F is more electronegative than O); e.g. O is –2 in MgO, but is –1
in Na2O2, and +2 in OF2
• Chlorine has a –1 oxidation state, unless it is with F or O (as they are more
electronegative than Cl), e.g. Cl is –1 in NaCl, but +1 in Cl2O, and +3 in ClF3.
• Hydrogen is +1 except in metal hydrides where it has an oxidation state of –1, e.g. H is +1
in HCl, +1 in H2O, but –1 in NaH.
Oxidation State Examples
Here are some examples of common compounds with all the oxidation numbers:
Sodium Chloride (common salt), NaCl, Na = +1, Cl = –1
Sodium Carbonate (washing soda), Na2CO3, Na = +1, C = +4, O = –2
Calcium Fluoride (fluorspar), CaF2, Ca = +2, F = –1
Calcium Hydroxide (lime water), Ca(OH)2, Ca = +2, O = –2, H = +1
Potassium Nitrate (saltpetre), KNO3, K = +1, N = +5, O = –2
Iron(III) Oxide (haematite), Fe2O3, Fe = +3, O = –2
Copper(II) Sulfate, CuSO4, Cu = +2, S = +6, O = –2
Compound Names
Old and new names
Compounds used to be named differently. At one time each writer would have their own
names for compounds. It was very confusing, so internationally chemists agreed standard
names. Later it was thought that the words used were difficult or confusing, so
internationally it was agreed to use numbers. Here are some examples;
• KNO3 was called common nitre or saltpetre, then potassium nitrate, but now is called
potassium nitrate(V), because the N has an oxidation state of +5.
• KNO2 was called potassium nitrite, but now is called potassium nitrate(III), because the N
has an oxidation state of +3. As further examples here are some chlorine compounds;
Some of the old names are still used. For example, KMnO4 should be called
potassium manganate(VII), but is was known as potassium permanganate.
How To Work Out Oxidation States
You need to be able to calculate oxidation states in various situations. These worked
examples will help you when you meet more difficult questions.
Step 1 Write down the formula.
Step 2 For the oxidation states known, write the oxidation states above the symbol.
Remember an oxidation state is for one atom.
Step 3 For the oxidation states known, write the sum of the oxidation states below the
symbol.
Step 4 Work out the oxidation state of the unknown. For a compound, the sum of the
oxidation states must equal zero. For an ion, the sum of the oxidation states must equal
the charge. If there is more than one atom of the element, then its number is the sum
of the oxidation states.
These are the colours of the oxidation states of
the radioactive element plutonium which vary
from Pu(III) to Pu(VII). Plutonium could be used in
nuclear power stations or to make nuclear
bombs. Understanding the oxidation states of
plutonium will help to clear up the waste from
the Cold War.
How To Work Out Oxidation States
Combining half-equations
When you are given two half-equations, sometimes you will need to join them to make
one full equation. The aim here is make sure that the same number of electrons that are
donated by one half-equation are accepted by the other.
How to combine two half-equations
The easiest way to combine two half-equations is to work in steps. These worked
examples will help you when you meet more difficult questions. Refer to the examples
while reading these steps;
Step 1 Write out the two half-equations.
Step 2 Note the number of electrons each half-equation gains or loses.
So that both equations involve the same number of electrons, you
may have to multiply up one or both equations.
Step 3 Multiply up the reactants and products.
Step 4 Write all the reactants together, and the products together.
Step 5 There should be the same number of electrons on both sides of this
equation. Cancel them. What is left is the full balanced equation.
How To Work Out Oxidation States
How To Work Out Oxidation States
How To Work Out Oxidation States
The cell in the photograph uses zinc and silver oxide to store the energy.
The half-equations are:
Zn(s) + 2OH–(aq)  Zn(OH)2(s) + 2e–
Ag2O(s) + H2O(l) + 2e–  2Ag(s) + 2OH–(aq)
The overall discharge equation is:
Zn(s) + Ag2O(s) + H2O(l)  Zn(OH)2(s) + 2Ag(s)
The cell has a high energy density, but is very expensive.
How To Work Out Oxidation States