F. Electrochemistry · ElectrochemistryOxidation and Reduction: Following the Electrons
F3.1 F2.1 A1.12About 70 minutes
Learning goals- I can define oxidation and reduction as the loss and gain of electrons.
- I can split a redox reaction into an oxidation half-reaction and a reduction half-reaction.
- I can identify the oxidizing agent and the reducing agent in a reaction.
- I can tell a redox reaction from a reaction where no electrons are transferred.
Words to know
oxidation
A loss of electrons by an atom, ion or molecule. Its oxidation number goes up.
reduction
A gain of electrons by an atom, ion or molecule. Its oxidation number goes down.
redox reaction
Short for oxidation-reduction reaction: a reaction in which electrons are transferred. Oxidation and reduction always happen together.
half-reaction
An equation that shows only the oxidation or only the reduction part of a redox reaction, with the electrons written in, e.g. Zn → Zn²⁺ + 2e⁻.
oxidizing agent
The reactant that takes electrons and so causes something else to be oxidized. The oxidizing agent is itself reduced.
reducing agent
The reactant that gives up electrons and so causes something else to be reduced. The reducing agent is itself oxidized.
spectator ion
An ion that is present but does not change during the reaction, so it is left out of the net ionic equation.
Electrons on the move
A bike chain left outside all winter turns orange with rust. A phone battery lights the screen. A cut apple turns brown. Bleach removes a stain. These look like very different events, but they are the same kind of chemistry: electrons moving from one substance to another. Chemists call these oxidation-reduction, or redox, reactions.
In this unit you will learn to follow those electrons. That one skill explains batteries, fuel cells, electroplating, corrosion and the electric-vehicle batteries now being built in Ontario. It also leads straight into first-year university chemistry, materials engineering and environmental science, where electrochemistry is a core topic.
Remember from Grade 11
In Grade 11 Chemistry (SCH3U), in the unit Chemical Reactions, you studied single displacement reactions and the activity series. A more active metal displaces a less active metal from a compound:
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
You also wrote net ionic equations in the unit Solutions and Solubility. The sulfate ion does not change here, so it is a spectator ion. The net ionic equation is:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Look closely. A zinc atom (no charge) became a zinc ion with a 2+ charge. A copper(II) ion became a neutral copper atom. Two electrons moved from zinc to copper. In Grade 11 you called this single displacement. Now you will call it what it is underneath: a redox reaction.
From oxygen to electrons
The words are older than the electron. Originally, oxidation meant combining with oxygen (iron rusting, magnesium burning). Reduction meant getting a metal back out of its ore: heating iron ore with carbon "reduces" a big lump of rock to a smaller mass of metal.
Chemists later noticed that in all of these reactions, electrons are moving. So the modern definitions are about electrons, not oxygen:
- Oxidation is the loss of electrons.
- Reduction is the gain of electrons.
A common memory aid is LEO says GER: Loss of Electrons is Oxidation; Gain of Electrons is Reduction. Some people prefer OIL RIG: Oxidation Is Loss, Reduction Is Gain.
Electrons cannot just disappear. If one substance loses them, another must gain them. So oxidation and reduction always happen together, in the same reaction, in equal electron amounts.
Half-reactions
We can split a redox reaction into two half-reactions. Each one shows the electrons as e⁻.
- Oxidation: Zn(s) → Zn²⁺(aq) + 2e⁻ (electrons are a product)
- Reduction: Cu²⁺(aq) + 2e⁻ → Cu(s) (electrons are a reactant)
Add the two half-reactions. The 2e⁻ on each side cancel, and you get back the net ionic equation. A half-reaction must be balanced in atoms and in charge. In the oxidation above, the left side has charge 0, and the right side has (+2) + 2(−1) = 0. ✓
Who is the "agent"?
The names of the agents confuse many students, so go slowly here.
- The oxidizing agent causes oxidation in something else. It does this by taking electrons. So the oxidizing agent is reduced. In our example, Cu²⁺ is the oxidizing agent.
- The reducing agent causes reduction in something else. It does this by giving electrons. So the reducing agent is oxidized. In our example, Zn is the reducing agent.
Think of a travel agent: the agent does something for someone else. The oxidizing agent oxidizes the other substance.
Agents are whole reactants
Name the agent as the whole species as it appears on the reactant side: "Cu²⁺(aq)" or "Zn(s)", never "copper" alone and never a product. Products cannot cause a reaction that has already made them.
Worked exampleExample 1: Identify everything
Magnesium ribbon is placed in silver nitrate solution: Mg(s) + 2Ag⁺(aq) → Mg²⁺(aq) + 2Ag(s). Write the half-reactions and name the oxidizing and reducing agents.
- Find what lost electrons: Mg (charge 0) became Mg²⁺. It lost 2 electrons, so magnesium was oxidized: Mg(s) → Mg²⁺(aq) + 2e⁻.
- Find what gained electrons: each Ag⁺ became Ag. Each gained 1 electron, so silver ions were reduced: Ag⁺(aq) + e⁻ → Ag(s).
- Make the electrons equal: magnesium gives 2 e⁻, so 2 silver ions must take them: 2Ag⁺(aq) + 2e⁻ → 2Ag(s). Adding the half-reactions gives the original equation.
- The reactant that was reduced (Ag⁺) is the oxidizing agent. The reactant that was oxidized (Mg) is the reducing agent.
Oxidation: Mg → Mg²⁺ + 2e⁻; reduction: 2Ag⁺ + 2e⁻ → 2Ag. Oxidizing agent: Ag⁺(aq). Reducing agent: Mg(s).
Worked exampleExample 2: Balancing the electrons
Aluminium foil reacts with copper(II) chloride solution. Write the half-reactions and the balanced net ionic equation.
- Aluminium forms Al³⁺: Al(s) → Al³⁺(aq) + 3e⁻.
- Copper(II) ions become copper: Cu²⁺(aq) + 2e⁻ → Cu(s).
- The electrons lost must equal the electrons gained. The lowest common multiple of 3 and 2 is 6. Multiply the oxidation by 2 and the reduction by 3: 2Al → 2Al³⁺ + 6e⁻ and 3Cu²⁺ + 6e⁻ → 3Cu.
- Add and cancel the 6e⁻: 2Al(s) + 3Cu²⁺(aq) → 2Al³⁺(aq) + 3Cu(s).
- Check the charge: left 3(+2) = +6; right 2(+3) = +6. ✓
2Al(s) + 3Cu²⁺(aq) → 2Al³⁺(aq) + 3Cu(s). Aluminium is the reducing agent; Cu²⁺ is the oxidizing agent. The chloride ions are spectators.
Worked exampleExample 3: Redox without a solution
Sodium metal burns in chlorine gas: 2Na(s) + Cl₂(g) → 2NaCl(s). Is this a redox reaction?
- In NaCl, sodium is present as Na⁺ and chlorine as Cl⁻ (an ionic compound).
- Each sodium atom went from 0 to +1: it lost one electron. Oxidation: 2Na → 2Na⁺ + 2e⁻.
- Each chlorine atom went from 0 to −1: it gained one electron. Reduction: Cl₂ + 2e⁻ → 2Cl⁻.
- Electrons were transferred, so yes, it is redox, even though no solution is involved.
Yes. Na is the reducing agent and Cl₂ is the oxidizing agent. Any reaction in which an element becomes part of an ionic compound is a redox reaction.
Not every reaction is redox
In a double displacement reaction, ions just change partners. Nothing gains or loses electrons. For example, in AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq), silver is Ag⁺ before and after, and chloride is Cl⁻ before and after. The same is true of acid-base neutralization: H⁺(aq) + OH⁻(aq) → H₂O(l).
A quick test: is any element changing from uncombined to combined, or the reverse? If yes (like Zn → Zn²⁺), it is redox. If all the ions keep their charges, it is not.
What about reactions between molecules, like the combustion of methane? No ions form, so it is harder to see the electrons. In the next lesson, Oxidation Numbers: Electron Bookkeeping, you will learn a bookkeeping system that reveals the electron shifts in any reaction.
| Term | What happens to electrons | What happens to the oxidation number | In Zn + Cu²⁺ → Zn²⁺ + Cu |
|---|
| Oxidation | lost | goes up | Zn → Zn²⁺ |
| Reduction | gained | goes down | Cu²⁺ → Cu |
| Reducing agent | gives electrons (is oxidized) | goes up | Zn(s) |
| Oxidizing agent | takes electrons (is reduced) | goes down | Cu²⁺(aq) |
Just for funWhy did the sodium atom finally feel calm after meeting chlorine?
It had lost its one negative thing. (Chlorine, meanwhile, was thrilled to gain it. Redox is a very fair trade.)
Hands-on activityThe browning apple: a kitchen redox test
You need: one apple, a knife and cutting board (a grown-up should do or supervise the cutting), 3 small plates, lemon juice (or a crushed vitamin C tablet dissolved in 2 tablespoons of water), water, a spoon, and a clock.
- Cut the apple into 6 slices of similar size. Put 2 slices on each plate.
- Plate A: leave the slices bare. Plate B: brush the cut faces with water. Plate C: brush the cut faces with lemon juice or the vitamin C solution.
- Predict which plate will brown first and last. Write your prediction as a hypothesis ("If..., then..., because...").
- Check and record the colour every 15 minutes for one hour. Use a simple scale: 0 = no browning, 3 = very brown.
- Explain: Browning happens when compounds in the apple (polyphenols) are oxidized by oxygen from the air, helped by an enzyme. Vitamin C (ascorbic acid) is a good reducing agent: it is oxidized more easily and it can reduce the first oxidation products back again, so browning is delayed until the vitamin C is used up.
- Which is the oxidizing agent in browning? (Oxygen, O₂.) Which substance acts as a sacrificial reducing agent? (Vitamin C.)
What to look for: Plate C should stay pale longest. Plate B usually browns at about the same rate as plate A, which shows that it is the vitamin C, not just the liquid, that protects the apple. Wash the knife carefully and eat the apple slices afterwards if you like.
Your turn
1Define oxidation and reduction in terms of electrons.
Oxidation is the loss of electrons; reduction is the gain of electrons.
2Write the half-reaction for iron(III) ions becoming iron(II) ions. Is it oxidation or reduction?
Fe³⁺(aq) + e⁻ → Fe²⁺(aq). Electrons are gained, so it is reduction.
3Write the half-reaction for bromide ions forming bromine, Br₂. Is it oxidation or reduction?
2Br⁻(aq) → Br₂(l) + 2e⁻. Electrons are lost, so it is oxidation.
4In Zn(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂(g), name the substance oxidized, the substance reduced, the oxidizing agent and the reducing agent.
Zn is oxidized (it is the reducing agent); H⁺ is reduced (it is the oxidizing agent).
5Write the two half-reactions for Zn(s) + 2H⁺(aq) → Zn²⁺(aq) + H₂(g).
Oxidation: Zn → Zn²⁺ + 2e⁻. Reduction: 2H⁺ + 2e⁻ → H₂.
6Is Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) a redox reaction? Explain.
No. Barium stays 2+ and sulfate stays 2−; no electrons are transferred. It is a precipitation (double displacement) reaction.
7Why can an oxidizing agent never appear on its own, without a reducing agent?
The electrons it takes have to come from somewhere. Electrons are not created or destroyed, so another substance must lose them: oxidation and reduction always happen together.
More practice: warm-up
1LEO says GER. What does LEO stand for?
Loss of Electrons is Oxidation.
2Is Cl₂ + 2e⁻ → 2Cl⁻ oxidation or reduction?
Reduction (electrons gained).
3Is Na → Na⁺ + e⁻ oxidation or reduction?
Oxidation (electron lost).
4The oxidizing agent is ___ (oxidized/reduced).
5In Fe(s) + Cu²⁺(aq) → Fe²⁺(aq) + Cu(s), which species is the reducing agent?
Fe(s), because it gives electrons to Cu²⁺.
More practice: on your own
1Write the half-reactions and the balanced net ionic equation for magnesium reacting with iron(III) ions to form magnesium ions and iron metal.
Mg → Mg²⁺ + 2e⁻ (×3) and Fe³⁺ + 3e⁻ → Fe (×2). The common multiple is 6 electrons. 3Mg(s) + 2Fe³⁺(aq) → 3Mg²⁺(aq) + 2Fe(s). Charge check: +6 = +6.
2Chlorine water is added to potassium iodide solution and the colourless solution turns brown: Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq). Identify the oxidizing and reducing agents.
Cl₂ gains electrons (0 → −1): it is the oxidizing agent. I⁻ loses electrons (−1 → 0): it is the reducing agent. (This is the halogen series you met in Grade 11: chlorine displaces iodine.)
3Write the net ionic equation for copper metal in silver nitrate solution, and say which ion is a spectator.
Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s). Copper loses 2 e⁻ and two silver ions each gain 1 e⁻. The nitrate ion is the spectator.
4A student says, "In Zn + Cu²⁺ → Zn²⁺ + Cu, the oxidizing agent is Zn²⁺ because it has been oxidized." Correct the student.
Zn²⁺ is a product, and agents are reactants. The substance oxidized is Zn, which makes Zn the reducing agent. The oxidizing agent is Cu²⁺, the reactant that took the electrons.
5Is the reaction H₂(g) + Cl₂(g) → 2HCl(g) redox? Give a reason you can see now, without oxidation numbers.
Yes. Two uncombined elements (H₂ and Cl₂) end up combined in a compound. Any reaction that turns an element into part of a compound (or the reverse) changes how electrons are shared, so it is redox. You will measure the change with oxidation numbers in the next lesson.
Challenge
1Silver tarnish is silver sulfide, Ag₂S. One way to remove it is to put tarnished silver in hot baking-soda water touching aluminium foil: 3Ag₂S(s) + 2Al(s) → 6Ag(s) + Al₂S₃ (the aluminium sulfide then reacts with water). Write the electron-transfer half-reactions and state how many electrons are transferred in the equation as written.
Silver goes from Ag⁺ (in Ag₂S) to Ag: 6Ag⁺ + 6e⁻ → 6Ag (reduction). Aluminium goes from Al to Al³⁺: 2Al → 2Al³⁺ + 6e⁻ (oxidation). Sulfide stays S²⁻. 6 electrons are transferred. The silver metal is recovered, unlike polishing, which rubs silver away.
2When hydrogen peroxide decomposes, 2H₂O₂(l) → 2H₂O(l) + O₂(g), some oxygen atoms end up in water and some in O₂. Explain how the same substance can be both oxidizing agent and reducing agent. (Hint: in H₂O₂ each oxygen is in a −1 state; in water it is −2 and in O₂ it is 0.)
Some oxygen atoms go from −1 to −2 (they gain electrons: reduced) and others go from −1 to 0 (they lose electrons: oxidized). So H₂O₂ is both the oxidizing agent and the reducing agent. A reaction like this is called disproportionation.
3A strip of metal X is placed in a solution of Y²⁺ ions. A coating of Y forms on X and the solution gains X³⁺ ions. Write the balanced net ionic equation and identify the strongest reducing agent among X, Y, X³⁺ and Y²⁺ in this reaction.
X → X³⁺ + 3e⁻ (×2); Y²⁺ + 2e⁻ → Y (×3): 2X(s) + 3Y²⁺(aq) → 2X³⁺(aq) + 3Y(s). X gave its electrons to Y²⁺, so X is the reducing agent, and it is a stronger reducing agent than Y (otherwise Y would give electrons back). Among those listed, X is the strongest reducing agent.
Quiz
2. In Mg(s) + Cu²⁺(aq) → Mg²⁺(aq) + Cu(s), what is the oxidizing agent?
3. Which half-reaction is correctly balanced?
4. Which reaction is NOT a redox reaction?
5. In 2Al(s) + 3Cu²⁺(aq) → 2Al³⁺(aq) + 3Cu(s), how many electrons are transferred?
6. A reducing agent is a substance that:
Big idea (F3.1): redox reactions are electron transfers. Oxidation is loss of electrons, reduction is gain, and they always happen together in equal electron amounts.
How to help: Ask your teen to explain the apple activity to you in their own words, using the words oxidized, reducing agent and oxidizing agent. If they can teach it, they understand it. Point out redox around the house: rust on a shovel, a tarnished silver spoon, batteries, bleach.
Common difficulties:
- Mixing up the agent with the substance it acts on. The oxidizing agent is reduced. Practise with the travel-agent idea: the agent does something to someone else.
- Putting electrons on the wrong side of a half-reaction. In oxidation the electrons are products; in reduction they are reactants. Checking the charge on both sides catches this every time.
- Forgetting to multiply half-reactions so the electrons lost equal the electrons gained.
- Naming a product as an agent.
Safety: the activity uses a knife. A grown-up should cut or supervise.