Who gives and who takes: redox reactions, one electron at a time

How to use this simulation

1. Start with “Zinc in copper sulfate” selected. Read the top row first: Zn sits at 0 and Cu²⁺ at +2. Before pressing anything, guess which one wants electrons.
2. Let it run. Amber dots stream along the dashed curve from zinc to copper: 2 electrons in total. Then the bottom row fades in with the new numbers, +2 for zinc and 0 for copper.
3. Tick “Hide oxidation numbers” and switch reaction. Every atom shows a question mark, so you must work each number out from the rules before you let it run.
4. Compare the red label (the species that is oxidised, which is the reducing agent) with the blue label (the species that is reduced, which is the oxidising agent) on every reaction.
5. Read the two half-equations under the bottom row. Add them together, cancel the electrons, and you get the full equation at the top.

An apple turns brown, an iron gate turns red

Cut an apple and leave it on the table. Within minutes the white flesh browns. Leave an iron gate in the rain and a month later it is orange and flaking. Light a gas stove and a blue flame appears. Three very different scenes, and chemistry files all of them under one heading: oxidation. Something is giving electrons away, and something else is taking them.

The trouble is that you cannot see electrons. So chemists invented a bookkeeping system, the oxidation number, that lets you point at an atom and say whether it has lost or gained electrons in a reaction. This page lets you watch the bookkeeping and the real movement side by side.

What redox means, in three steps

The word has changed meaning twice. First, oxidation meant combining with oxygen: iron plus oxygen gives rust, so iron is oxidised. Reduction meant the opposite, losing oxygen, as when iron ore is heated with carbon to give iron metal.

Second, chemists noticed the same pattern with hydrogen. Adding hydrogen to something is reduction, removing hydrogen is oxidation. Chlorine taking hydrogen from other molecules fitted neatly into “oxidising”.

Third, and best, is the electron definition, which covers every case at once. Oxidation is loss of electrons. Reduction is gain of electrons. Remember it as OIL RIG: Oxidation Is Loss, Reduction Is Gain. Since electrons cannot vanish, the two always happen together. That is why we say redox, short for reduction and oxidation, and why one half cannot exist alone.

Oxidation numbers: bookkeeping for electrons

An oxidation number is the charge an atom would have if every shared pair of electrons went completely to the more electronegative atom. In an ion like Cu²⁺ it is the real charge. In a covalent molecule like HCl it is a bookkeeping device, but it still tells you where electrons moved. Use these rules in order:

Reading the animation

The canvas has two rows. The top row, Before, shows the reactants. Each atom is a circle with its symbol inside and its oxidation number above it. A coloured ring means this atom is about to change; a grey ring means its number stays the same. The number in front of a formula tells you how many units there are, and one of each is drawn.

Electrons are the amber dots. They leave the species that is oxidised (red label, “Reducing agent”) and travel along the dashed curve to the species that is reduced (blue label, “Oxidising agent”). The count over the curve is the total number of electrons that move for the equation as written.

Then the bottom row, After, fades in with the products. Compare the numbers above each coloured atom. When a number rises, that atom was oxidised. When it falls, that atom was reduced. The two half-equations at the foot of the canvas say the same thing in electron language.

Agents: the species that makes the other one change

This is where almost everybody trips. The oxidising agent is the species that causes oxidation, and it does so by being reduced itself. The reducing agent causes reduction by being oxidised itself. They are named for what they do to the other party, not for what happens to them.

Check it in the first reaction. Cu²⁺ takes electrons from zinc. Zinc is oxidised because of Cu²⁺, so Cu²⁺ is the oxidising agent, and Cu²⁺ is itself reduced. Zinc hands electrons over, so it is the reducing agent. Here are all five reactions in the simulation:

Half-equations and how to balance with them

A half-equation shows only one half of the transfer, with the electrons written in explicitly. Oxidation puts the electrons on the right (products), reduction on the left (reactants). The magic is that when you add the two half-equations the electrons cancel, so you get the full equation.

Here is the method, using permanganate and iron(II). Step 1: write each half. Fe²⁺ → Fe³⁺ + e⁻ (iron rises from +2 to +3) and MnO₄⁻ → Mn²⁺ (manganese falls from +7 to +2). Step 2: balance oxygen with H₂O and hydrogen with H⁺, then add electrons until the charges match: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. Step 3: make the electrons equal. Manganese takes 5 electrons, iron gives 1, so multiply the manganese half by 1 and the iron half by 5. Step 4: add and cancel.

Result: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O. Pick the “KMnO₄ with Fe²⁺” reaction and watch 5 electrons move on the curve; that is exactly the 5 iron ions each giving one.

The five reactions, one by one

Each of these shows a slightly different side of redox. Switch between them while you read.

• Zinc in copper sulfate: the classic. A grey zinc strip in blue copper sulfate slowly grows a brown coating of copper while the blue fades. 2 electrons per Zn, from 0 to +2 for zinc and +2 to 0 for copper.
• Iron rusting: 4 iron atoms and three O₂ molecules make two Fe₂O₃. Iron goes from 0 to +3, oxygen from 0 to −2, and 12 electrons cross. Rust is slow redox, helped by water and salt.
• Magnesium and HCl: magnesium goes to +2, hydrogen falls from +1 to 0, and the bubbles you see are H₂. Chlorine stays at −1 the whole time, so it is a spectator and carries no colour in the animation.
• KMnO₄ with Fe²⁺: purple permanganate (manganese +7) is reduced to almost colourless Mn²⁺ (+2), which is why it works as its own indicator in a titration.
• Hydrogen and chlorine: both elements start at 0. Hydrogen ends at +1, chlorine at −1. There is no ion in HCl gas, only a shared pair pulled towards chlorine, but oxidation numbers still say which atom effectively lost the electron.

Try it yourself in the simulation

Do these in order. Write your guess down first, then look.

• Experiment 1: with numbers hidden, predict the oxidation number of every atom in Mg + 2HCl. Run it, then show the numbers and score yourself.
• Experiment 2: pick Iron rusting. Count the electrons on the curve (12) and divide by the number of iron atoms (4). You should get the charge on each iron ion.
• Experiment 3: pick KMnO₄ with Fe²⁺ and find the species that has no coloured ring (H⁺ and the oxygens). Its oxidation number never changes, so it is not part of the electron transfer.
• Experiment 4: for each of the five reactions, name the reducing agent before reading the label. If you are wrong, find which rule you slipped on.
• Experiment 5: pause at the moment the dots arrive and look at both half-equations. Add them by hand and check the electrons really cancel.

Solved problems

The numbers below are computed with the same arithmetic the animation uses.

Mistakes everybody makes

These cost marks every year.

• Mixing up the agent with the species that changes: the oxidising agent is the one that is reduced.
• Forgetting that oxidation numbers are written with the sign first: +2, not 2+. A charge is written the other way round: Cu²⁺.
• Treating the oxidation number of oxygen as −2 everywhere. In peroxides such as H₂O₂ it is −1.
• Counting only one atom when a formula has several: in Cr₂O₇²⁻ you must solve for two chromiums together.
• Thinking all reactions are redox. In NaOH + HCl → NaCl + H₂O, no oxidation number changes.
• Writing electrons on the wrong side of a half-equation: oxidation loses them, so they appear in the products.

Redox around you

You live inside redox reactions all day.

• Batteries: every cell stores a spontaneous redox reaction and forces the electrons to travel through your phone before they reach the other side.
• Respiration: glucose is oxidised to carbon dioxide while oxygen is reduced to water. Cells capture the energy step by step.
• Bleach and disinfectant: chlorine compounds are oxidisers that damage the chemistry of germs and stains.
• Metal extraction: iron ore is reduced with carbon in a blast furnace; aluminium is reduced by electrolysis.
• Rust prevention: galvanising puts zinc on steel so zinc, the stronger reducing agent, is oxidised first and the steel survives.

Exam corner

Try these without looking, then compare.

• Find the oxidation number of Mn in KMnO₄ (answer: +7) and Cr in K₂Cr₂O₇ (answer: +6).
• In Mg + 2HCl → MgCl₂ + H₂, name the oxidising agent, the reducing agent and the species that is oxidised. Answer: HCl is the oxidising agent, Mg the reducing agent, and Mg is oxidised.
• Write the two half-equations for Zn + Cu²⁺ and state how many electrons are transferred (2).
• Explain why H₂O₂ can act as both oxidiser and reducer (oxygen is −1, between 0 and −2).
• How many moles of Fe²⁺ does one mole of MnO₄⁻ oxidise in acid? Answer: 5.

Revision: the whole page in a minute

Read this before the exam.

• Oxidation = loss of electrons = oxidation number goes up. Reduction = gain = number goes down.
• The oxidising agent is reduced; the reducing agent is oxidised.
• Pure elements are 0; O is −2, H is +1, and the numbers add up to the species charge.
• In a half-equation, oxidation has electrons in the products and reduction has them in the reactants.
• To balance: split into halves, balance O with H₂O and H with H⁺, add electrons, equalise them, add.
• Electrons lost always equal electrons gained.

Frequently asked questions