What you'll learn
Oxidation and reduction is one of the most widely applied topics in CXC CSEC Chemistry, connecting metal extraction, the reactivity series, electrolysis, corrosion, the reactions of acids and the identification of oxidising and reducing agents. Its difficulty lies in the fact that three different definitions of the same process are in use — in terms of oxygen, of hydrogen, and of electrons — and a question may require any of them. Once you can move between the three, the topic becomes one of the most reliable sources of marks on the paper. By the end of this guide you should be able to apply all three definitions, identify what is oxidised and what is reduced in any reaction, write and balance half equations, name the common oxidising and reducing agents with their characteristic colour changes, and apply redox reasoning to displacement, electrolysis and corrosion.
Key terms and definitions
Oxidation — gain of oxygen, loss of hydrogen, or loss of electrons
Reduction — loss of oxygen, gain of hydrogen, or gain of electrons
Redox reaction — a reaction in which oxidation and reduction occur together
Half equation — an equation showing the loss or gain of electrons by one species
Oxidising agent — a substance that causes another to be oxidised and is itself reduced
Reducing agent — a substance that causes another to be reduced and is itself oxidised
Oxidation number — a number assigned to an atom to indicate its degree of oxidation
Spectator ion — an ion that takes no part in the reaction and appears unchanged on both sides
Ionic equation — an equation showing only the species that change, with spectator ions omitted
Corrosion — the gradual destruction of a metal by chemical reaction with its environment
Rusting — the corrosion of iron, requiring both oxygen and water
Core concepts
The three definitions
The oldest definition concerns oxygen. A substance that gains oxygen is oxidised; a substance that loses oxygen is reduced. When copper oxide is heated with carbon, the copper oxide loses oxygen and is reduced, while the carbon gains it and is oxidised.
The second definition concerns hydrogen and is the reverse in appearance. A substance that loses hydrogen is oxidised; a substance that gains hydrogen is reduced. This definition is useful in organic chemistry, where converting an alcohol to an acid involves loss of hydrogen and is therefore oxidation.
The third and most general definition concerns electrons, and it applies to every redox reaction including those with no oxygen or hydrogen present. A substance that loses electrons is oxidised; one that gains electrons is reduced. The mnemonic is OIL RIG: Oxidation Is Loss, Reduction Is Gain — of electrons.
Note carefully that OIL RIG applies only to the electron definition. Applied to oxygen it gives the wrong answer, and this is a frequent source of lost marks.
All three describe the same underlying process, because gaining oxygen or losing hydrogen both involve losing electrons.
Identifying oxidation and reduction
The most reliable method is to track charge.
When magnesium reacts with dilute hydrochloric acid, magnesium atoms begin uncharged and end as ions with a charge of plus 2, so each has lost two electrons and the magnesium is oxidised. Hydrogen ions begin with a charge of plus 1 and end as uncharged hydrogen molecules, so each has gained an electron and the hydrogen is reduced.
The general rule: an increase in positive charge means electrons lost and therefore oxidation; a decrease in positive charge, or an increase in negative charge, means electrons gained and therefore reduction.
Because electrons cannot be created or destroyed, oxidation and reduction always occur together. Identifying one guarantees the other is present.
Oxidation numbers
Oxidation numbers give a systematic way of tracking redox changes, and the rules needed are few.
An uncombined element has an oxidation number of zero. A simple ion has an oxidation number equal to its charge. In compounds, hydrogen is usually plus 1 and oxygen usually minus 2. Group 1 metals are plus 1 and Group 2 metals plus 2. The oxidation numbers in a neutral compound sum to zero, and in a polyatomic ion they sum to the charge on that ion.
An increase in oxidation number indicates oxidation; a decrease indicates reduction. This approach is particularly useful where an element appears in a compound rather than as a free ion, such as manganese in permanganate.
Half equations
A half equation shows what happens to one species, with the electrons written explicitly.
For oxidation, the electrons appear on the right, because they are lost. For reduction, they appear on the left, because they are gained.
Two conditions must be met: the atoms must balance, and the total charge must balance on both sides. Checking the charge is the step most often omitted and is where marks are lost.
Adding the two half equations, with the electrons cancelling, gives the overall ionic equation. If the numbers of electrons differ, one or both half equations must be multiplied through before adding so that the electrons cancel exactly.
Common oxidising and reducing agents
Each term describes a substance by its effect on something else, and the relationship is always inverted: the oxidising agent is itself reduced, and the reducing agent is itself oxidised. Stating this clearly is worth practising because it is so often reversed.
The common oxidising agents and their colour changes are examinable. Acidified potassium permanganate changes from purple to colourless when it acts as an oxidising agent. Acidified potassium dichromate changes from orange to green. Chlorine water and other halogens, concentrated sulfuric acid, concentrated nitric acid, and hydrogen peroxide are also oxidising agents.
The common reducing agents include carbon and carbon monoxide, used in metal extraction; hydrogen; reactive metals such as zinc and magnesium; potassium iodide solution, which turns from colourless to brown as iodine is released; sulfur dioxide; and hydrogen sulfide.
The colour changes are the practical means of identifying that a redox reaction has occurred, and questions frequently supply an observation and ask what it indicates.
Redox in displacement reactions
A more reactive metal displaces a less reactive one from a solution of its salt, and the process is redox.
When zinc is added to copper sulfate solution, zinc atoms lose electrons and are oxidised to zinc ions, while copper ions gain those electrons and are reduced to copper metal. The blue colour of the solution fades and a brown deposit of copper forms.
The sulfate ions take no part and appear unchanged on both sides, so they are spectator ions. Removing them gives the ionic equation, which displays the electron transfer clearly.
Halogen displacement follows the same pattern with the direction of electron transfer reversed, since halogens gain electrons. Chlorine displaces bromide and iodide; bromine displaces iodide only. The colour changes in the solution identify which halogen has been released.
Redox in electrolysis
Electrolysis is redox driven by an electric current, and the electrode names map directly onto the processes.
At the cathode, the negative electrode, positive ions gain electrons, so reduction occurs at the cathode.
At the anode, the positive electrode, negative ions lose electrons, so oxidation occurs at the anode.
A memory aid worth adopting is that reduction happens at the cathode — both relate to the negative electrode — while oxidation occurs at the anode.
Corrosion and rusting
Rusting is the corrosion of iron and is a redox process. Iron atoms lose electrons and are oxidised, while oxygen is reduced. Both oxygen and water must be present; the classic experiment uses three test tubes, one with iron in air and water, one with boiled water and a layer of oil to exclude air, and one with dry air over a drying agent, and only the first shows rust.
Rusting is accelerated by dissolved salt, which is why corrosion is a particular problem in coastal Caribbean environments.
Prevention works either by excluding air and water — painting, greasing, coating with plastic, or electroplating — or by sacrificial protection, in which a more reactive metal such as zinc or magnesium is attached. The more reactive metal loses electrons in preference to the iron, so it is oxidised instead and the iron is protected. Galvanising combines both, since the zinc coating excludes air and water and also protects sacrificially if scratched.
Worked examples
Example 1: Applying two definitions (4 marks)
Iron oxide is heated with carbon monoxide, producing iron and carbon dioxide. Identify what is oxidised and what is reduced, using both the oxygen and the electron definitions.
In terms of oxygen, the iron oxide loses oxygen to become iron, so it is reduced. The carbon monoxide gains oxygen to become carbon dioxide, so it is oxidised and is acting as the reducing agent.
In terms of electrons, the iron ions in iron oxide carry a positive charge and end as uncharged iron atoms, so they have gained electrons and are reduced. The carbon in carbon monoxide has lost electrons in becoming more highly combined with oxygen, so it is oxidised.
Example 2: Writing half equations (4 marks)
Write half equations for the reaction between magnesium and copper sulfate solution, and identify the oxidising agent.
For the oxidation, each magnesium atom loses two electrons to form a magnesium ion with a charge of plus 2, so the two electrons appear on the right-hand side of the half equation.
For the reduction, each copper ion with a charge of plus 2 gains two electrons to form a copper atom, so the two electrons appear on the left-hand side.
Both half equations balance for atoms and for charge, and both involve two electrons, so they combine directly with the electrons cancelling. The copper ions are the oxidising agent, since they take electrons from the magnesium and are themselves reduced.
Example 3: Interpreting a colour change (3 marks)
Acidified potassium permanganate solution is added to a solution of iron(II) sulfate. The purple colour disappears. Explain what has happened.
Acidified potassium permanganate is a strong oxidising agent, and its purple colour disappears when it is reduced. The loss of the purple colour therefore shows that the permanganate has been reduced by the iron(II) sulfate.
The iron(II) ions have been oxidised to iron(III) ions, losing one electron each. Since the iron(II) ions have caused the permanganate to be reduced, they are acting as the reducing agent and are themselves oxidised, which is the standard inverted relationship.
Common mistakes and how to avoid them
The most frequent error is applying OIL RIG to the oxygen definition, which reverses the answer. OIL RIG concerns electrons only; gaining oxygen is oxidation.
Students routinely confuse the oxidising agent with the substance that is oxidised. The oxidising agent is the one reduced, and writing the relationship out before answering prevents the error.
In half equations, many candidates balance atoms but not charge. Both must balance, and the charge check is where the marks are.
Another common slip is placing electrons on the wrong side. Oxidation produces electrons, so they go on the right; reduction consumes them, so they go on the left.
Finally, in rusting questions, candidates often state that air alone causes rust. Both oxygen and water are required, and salt accelerates the process without being necessary.
Exam technique for "Oxidation and Reduction"
Decide which definition applies before you begin. Oxygen present and no charges shown means the oxygen definition; ions or charges shown means the electron definition.
Write the charge on every species before and after the reaction. The direction of electron transfer then becomes obvious without further reasoning.
Learn the oxidising and reducing agents together with their colour changes, since observation questions are common and the colour is usually the whole of the evidence.
For electrolysis, state the electrode, the process and the direction of electron movement in a single sentence: reduction occurs at the cathode as positive ions gain electrons.
When an ionic equation is requested, identify and remove the spectator ions explicitly, since recognising them often carries a mark of its own.
Quick revision summary
Oxidation is gain of oxygen, loss of hydrogen or loss of electrons; reduction is the reverse in each case, with OIL RIG applying to electrons only. Track charge to identify each: increased positive charge means oxidation, decreased means reduction, and the two always occur together. Oxidation numbers give a systematic alternative, with an increase indicating oxidation. Half equations place electrons on the right for oxidation and the left for reduction and must balance for both atoms and charge, combining to give the ionic equation once electrons cancel. The oxidising agent is itself reduced and the reducing agent is itself oxidised; acidified permanganate goes purple to colourless and acidified dichromate orange to green when reduced, while potassium iodide turns colourless to brown when oxidised. Displacement is redox with the anion as spectator ion, reduction occurs at the cathode and oxidation at the anode in electrolysis, and rusting requires both oxygen and water, is accelerated by salt, and is prevented by barriers or by sacrificial protection with a more reactive metal.