(I once held up a battery and asked a class, "What is inside this?" Most students said, "Chemicals." Then I asked, "What are the chemicals doing?" Silence. And then I explained: inside every battery, a chemical reaction is pushing electrons from one place to another. That flow of electrons is electricity. Every battery, every fuel cell, every electrochemical process in industry is built on the same principle: a reaction that transfers electrons. This chapter is about that transfer — the dance of electrons that powers our modern world. And once you understand it, you understand everything from your phone battery to the corrosion of a ship.) Think about the devices you use every day. Your phone, your laptop, your torch, your car — they all depend on batteries or fuel cells. These devices convert chemical energy into electrical energy through reactions that transfer electrons. The same principles explain why iron rusts, why silver tarnishes, and why some metals corrode faster than others. This is the world of electrochemistry — the study of the relationship between electricity and chemical reactions. At the heart of electrochemistry is a type of reaction called redox — short for reduction-oxidation. In a redox reaction, one substance loses electrons (oxidation) and another gains them (reduction). This transfer of electrons is what generates electricity in a battery and what drives the corrosion of metals. In this chapter, we are going to look at redox reactions in detail. We will learn about oxidation numbers, how to identify redox reactions, how electrochemical cells work, and how to use standard electrode potentials to predict which reactions will happen. By the end, you will understand the electron dance that powers our world. Let us begin.
Electrochemistry is the branch of chemistry that deals with the relationship between electricity and chemical reactions. It covers:
In the past, oxidation was defined as the addition of oxygen and reduction as the removal of oxygen. But a more useful definition involves electrons:
A helpful mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain.
Another way to remember: LEO GER — Loss of Electrons is Oxidation, Gain of Electrons is Reduction.
Important: Oxidation and reduction always happen together. You cannot have one without the other, because the electrons lost by one substance must be gained by another. This is why we call them redox reactions.
(Think of it like a transfer of money. If one person loses money, another person gains it. The total amount stays the same. In a redox reaction, electrons are transferred from one substance to another — one loses, one gains. This is the fundamental dance of electrochemistry.)
Oxidation number (or oxidation state) is a number assigned to an atom in a compound or ion that represents the number of electrons it has gained or lost compared to its elemental state.
Rules for assigning oxidation numbers:
| Rule | Explanation | Example |
|---|---|---|
| 1. Free elements | Oxidation number is 0. | Na = 0, O₂ = 0, Cl₂ = 0 |
| 2. Monatomic ions | Oxidation number equals the charge on the ion. | Na⁺ = +1, Cl⁻ = −1, Mg²⁺ = +2 |
| 3. Hydrogen | Usually +1 (except in metal hydrides, where it is −1). | H₂O: H = +1 |
| 4. Oxygen | Usually −2 (except in peroxides, where it is −1, and in OF₂, where it is +2). | H₂O: O = −2 |
| 5. Group 1 metals | Always +1. | NaCl: Na = +1 |
| 6. Group 2 metals | Always +2. | MgO: Mg = +2 |
| 7. Halogens | Usually −1, but can be positive when combined with oxygen or other halogens. | NaCl: Cl = −1; ClO₃⁻: Cl = +5 |
| 8. Sum of oxidation numbers | In a neutral compound, the sum is 0. In a polyatomic ion, the sum equals the charge on the ion. | H₂SO₄: (2×+1) + (+6) + (4×−2) = 0 |
Example: Assign oxidation numbers to each element in H₂SO₄.
So the oxidation numbers are H = +1, S = +6, O = −2.
A redox reaction is one in which oxidation numbers change. If the oxidation number of an element increases, it has been oxidised. If it decreases, it has been reduced.
Example: Zn + CuSO₄ → ZnSO₄ + Cu
So this is a redox reaction.
Example: NaOH + HCl → NaCl + H₂O
No oxidation numbers change, so this is not a redox reaction — it is an acid-base neutralisation.
In a redox reaction:
| Term | Definition | Example |
|---|---|---|
| Oxidising agent | Gains electrons — oxidises another substance. | KMnO₄, K₂Cr₂O₇, Cl₂, O₂ |
| Reducing agent | Loses electrons — reduces another substance. | Zn, Fe, H₂, CO, C |
(The names can be confusing at first. Here is a trick: an oxidising agent oxidises something else — so it must take electrons from it. It gains electrons itself and is reduced. A reducing agent reduces something else — so it must give electrons to it. It loses electrons itself and is oxidised. Remember: the agent does the opposite to itself.)
An electrochemical cell is a device that converts chemical energy into electrical energy through a redox reaction. There are two main types:
In this chapter, we will focus on galvanic cells.
A simple galvanic cell consists of:
Example: The Daniell cell — zinc and copper electrodes in their respective sulphate solutions.
The standard electrode potential (E°) of a half-cell is a measure of the tendency of the half-cell to gain electrons (be reduced). It is measured relative to the standard hydrogen electrode, which is assigned a potential of 0.00 V.
Standard conditions: 298 K (25°C), 1 atm pressure, and 1 mol/dm³ concentration of all solutions.
Standard electrode potentials are listed in tables. The more positive the E° value, the greater the tendency to be reduced (gain electrons). The more negative the E° value, the greater the tendency to be oxidised (lose electrons).
| Half-Reaction | E° (V) |
|---|---|
| F₂ + 2e⁻ → 2F⁻ | +2.87 |
| Cl₂ + 2e⁻ → 2Cl⁻ | +1.36 |
| O₂ + 4H⁺ + 4e⁻ → 2H₂O | +1.23 |
| Ag⁺ + e⁻ → Ag | +0.80 |
| Cu²⁺ + 2e⁻ → Cu | +0.34 |
| 2H⁺ + 2e⁻ → H₂ | 0.00 |
| Pb²⁺ + 2e⁻ → Pb | −0.13 |
| Fe²⁺ + 2e⁻ → Fe | −0.44 |
| Zn²⁺ + 2e⁻ → Zn | −0.76 |
| Mg²⁺ + 2e⁻ → Mg | −2.37 |
| Na⁺ + e⁻ → Na | −2.71 |
| K⁺ + e⁻ → K | −2.93 |
To predict whether a redox reaction will occur, compare the standard electrode potentials of the two half-cells:
Example: Will zinc displace copper from copper(II) sulphate solution?
| Application | Description |
|---|---|
| Batteries | Portable sources of electrical energy based on galvanic cells. Examples: dry cells, lead-acid batteries, lithium-ion batteries. |
| Fuel cells | Convert chemical energy from fuel (usually hydrogen) directly into electrical energy, with water as the only by-product. Used in spacecraft and some vehicles. |
| Electroplating | Using electrolysis to coat one metal with another (e.g., chrome-plating, silver-plating). |
| Corrosion prevention | Understanding electrochemistry helps prevent corrosion through methods like cathodic protection (using a sacrificial anode). |
| Extraction of metals | Electrolysis is used to extract reactive metals like aluminium and sodium from their ores. |
| Electrolysis | Used to produce chlorine, hydrogen, and sodium hydroxide from brine. |
(Electrochemistry is the bridge between chemistry and electricity. It powers your phone, protects your car from rust, and helps industry produce the chemicals we need. Once you understand redox, you understand the electron dance that makes all of this possible.)
Corrosion is an electrochemical process. For example, iron rusts because of a redox reaction involving oxygen and water.
Cathodic protection: A more reactive metal (like zinc or magnesium) is connected to the iron. The more reactive metal becomes the anode and corrodes in preference to the iron. This is called sacrificial protection.
Examples:
Redox equations can be balanced using the half-reaction method:
Example: Balance the equation for the reaction between Fe²⁺ and MnO₄⁻ in acidic solution.
Step 1: Write the half-reactions.
Step 2: Balance electrons. Multiply the oxidation half-reaction by 5.
Step 3: Add and cancel electrons.
5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O
Question: Assign oxidation numbers to each element in K₂Cr₂O₇.
Solution:
K = +1 (Group 1 metal).
O = −2 (usually).
Let Cr = x.
2(+1) + 2x + 7(−2) = 0 → 2 + 2x − 14 = 0 → 2x = 12 → x = +6.
So K = +1, Cr = +6, O = −2.
Question: Is the reaction 2KI + Cl₂ → 2KCl + I₂ a redox reaction? Explain.
Solution: Yes, it is a redox reaction. Chlorine goes from 0 to −1 (reduction — gains electrons). Iodine goes from −1 to 0 (oxidation — loses electrons). Since oxidation numbers change, it is a redox reaction.
Question: Will chlorine displace bromine from potassium bromide solution? Use standard electrode potentials.
Solution: Cl₂ + 2e⁻ → 2Cl⁻ E° = +1.36 V. Br₂ + 2e⁻ → 2Br⁻ E° = +1.07 V. Chlorine has the more positive E°, so it is reduced. Bromide has the more negative (less positive) E°, so it is oxidised. The reaction is: Cl₂ + 2KBr → 2KCl + Br₂. E°cell = 1.36 − 1.07 = +0.29 V. Positive, so the reaction is spontaneous. Yes, chlorine will displace bromine.
Question: Balance the equation for the reaction between Zn and Ag⁺.
Solution:
Oxidation: Zn → Zn²⁺ + 2e⁻
Reduction: Ag⁺ + e⁻ → Ag
Multiply the reduction half-reaction by 2: 2Ag⁺ + 2e⁻ → 2Ag
Add: Zn + 2Ag⁺ → Zn²⁺ + 2Ag
The equation is balanced.
1. Define oxidation and reduction in terms of electrons.
Oxidation is the loss of electrons. Reduction is the gain of electrons.
2. What is an oxidising agent? Give one example.
An oxidising agent is a substance that gains electrons and oxidises another substance. Example: chlorine (Cl₂) or potassium manganate(VII) (KMnO₄).
3. Assign oxidation numbers to each element in H₂O₂.
H = +1. O = −1 (peroxide, so oxygen is −1 instead of the usual −2).
4. Is the reaction Zn + CuSO₄ → ZnSO₄ + Cu a redox reaction? Explain.
Yes. Zinc goes from 0 to +2 (oxidation — loses electrons). Copper goes from +2 to 0 (reduction — gains electrons). Since oxidation numbers change, it is a redox reaction.
5. What is a galvanic cell?
A galvanic cell is an electrochemical cell that converts chemical energy into electrical energy through a spontaneous redox reaction. It consists of two half-cells, a salt bridge, and an external circuit.
6. Use standard electrode potentials to predict whether zinc will displace copper from copper(II) sulphate solution.
Zn²⁺ + 2e⁻ → Zn E° = −0.76 V. Cu²⁺ + 2e⁻ → Cu E° = +0.34 V. Zinc is more negative (more easily oxidised), copper is more positive (more easily reduced). E°cell = 0.34 − (−0.76) = +1.10 V. Positive, so the reaction is spontaneous. Yes, zinc will displace copper.
7. What is cathodic protection? Give one example.
Cathodic protection is a method of preventing corrosion by connecting a more reactive metal (a sacrificial anode) to the metal being protected. The more reactive metal corrodes in preference to the protected metal. Example: galvanising iron with zinc, or attaching zinc blocks to a ship's hull.
Question 1 [10 marks]
(a) Define oxidation and reduction in terms of electron transfer. [2]
(b) Assign oxidation numbers to each element in MnO₄⁻ and Cr₂O₇²⁻. [4]
(c) Identify which of the following is a redox reaction: (i) NaOH + HCl → NaCl + H₂O, (ii) Zn + CuSO₄ → ZnSO₄ + Cu. [4]
(a) Oxidation is the loss of electrons; reduction is the gain of electrons. (2 marks: 1 for each.)
(b) MnO₄⁻: Mn = +7, O = −2. Cr₂O₇²⁻: Cr = +6, O = −2. (4 marks: 2 for each ion.)
(c) (i) NaOH + HCl → NaCl + H₂O is not a redox reaction — oxidation numbers do not change. (ii) Zn + CuSO₄ → ZnSO₄ + Cu is a redox reaction — Zn is oxidised (0 to +2) and Cu is reduced (+2 to 0). (4 marks: 2 for each identification + reason.)
Question 2 [10 marks]
(a) Describe the structure and working of a simple galvanic cell. [4]
(b) Explain the role of the salt bridge. [2]
(c) Use standard electrode potentials to predict whether the reaction between Zn and Fe²⁺ will occur spontaneously. [4]
(a) A galvanic cell consists of two half-cells — each with an electrode dipped in an electrolyte solution. The two half-cells are connected by a salt bridge, and the electrodes are connected by an external circuit. At the anode, oxidation occurs (electrons are released). At the cathode, reduction occurs (electrons are gained). Electrons flow from anode to cathode through the external circuit. (4 marks: 2 for structure, 2 for working.)
(b) The salt bridge allows ions to flow between the two half-cells, completing the circuit and maintaining electrical neutrality. Without it, charge would build up and the reaction would stop. (2 marks: 1 for ion flow, 1 for "completes the circuit".)
(c) Zn²⁺ + 2e⁻ → Zn E° = −0.76 V. Fe²⁺ + 2e⁻ → Fe E° = −0.44 V. Zinc is more negative (more easily oxidised), iron is more positive (more easily reduced). E°cell = −0.44 − (−0.76) = +0.32 V. Positive, so the reaction is spontaneous. Zn + Fe²⁺ → Zn²⁺ + Fe. (4 marks: 2 for identifying which is oxidised and reduced, 2 for E° calculation and conclusion.)
Question 3 [10 marks]
(a) Describe how electrochemistry is used in batteries and fuel cells. [4]
(b) Describe the process of corrosion of iron, and explain how cathodic protection prevents it. [4]
(c) State one application of electrolysis in industry. [2]
(a) Batteries are galvanic cells that convert chemical energy into electrical energy through redox reactions. Fuel cells use a continuous supply of fuel (e.g., hydrogen) and oxygen to produce electricity, with water as the only by-product. (4 marks: 2 for batteries, 2 for fuel cells.)
(b) Iron corrodes through an electrochemical process. At anodic areas, iron is oxidised: Fe → Fe²⁺ + 2e⁻. At cathodic areas, oxygen is reduced: O₂ + 2H₂O + 4e⁻ → 4OH⁻. The Fe²⁺ ions are further oxidised to rust. Cathodic protection prevents this by connecting a more reactive metal (e.g., zinc or magnesium) to the iron. The more reactive metal becomes the anode and corrodes in preference to the iron (sacrificial protection). (4 marks: 2 for corrosion process, 2 for cathodic protection.)
(c) Any one of: extraction of aluminium, manufacture of chlorine and sodium hydroxide from brine, electroplating, refining of copper. (2 marks: 1 for use, 1 for clarity.)
Challenge 1: Explain why a reaction with a positive E°cell is spontaneous, but may not happen quickly.
A positive E°cell indicates that the reaction is thermodynamically favourable — the products have lower free energy than the reactants, so the reaction can occur spontaneously. However, thermodynamic favourability does not guarantee that the reaction will happen quickly. The rate of a reaction depends on kinetics — specifically, the activation energy. If the activation energy is high, the reaction may be very slow, even though it is thermodynamically favourable. For example, the reaction between hydrogen and oxygen to form water has a very positive E°cell and is highly spontaneous, but it does not happen at room temperature without a spark or catalyst because the activation energy is high. So a positive E°cell tells you the reaction can happen, but not how fast.
Challenge 2: Explain why the salt bridge is essential in a galvanic cell.
The salt bridge is essential in a galvanic cell because it completes the electrical circuit and maintains electrical neutrality in the two half-cells. Without the salt bridge, the flow of electrons from the anode to the cathode would cause a buildup of positive charge at the anode (as Zn²⁺ ions are produced) and a buildup of negative charge at the cathode (as Cu²⁺ ions are removed). This charge buildup would oppose the further flow of electrons and stop the reaction. The salt bridge allows ions to migrate between the two half-cells — positive ions move to the cathode compartment and negative ions move to the anode compartment — neutralising the charge buildup and allowing the reaction to continue. Without it, the cell would quickly stop working.
Challenge 3: Explain the difference between a galvanic cell and an electrolytic cell, giving one example of each.
A galvanic cell (also called a voltaic cell) converts chemical energy into electrical energy through a spontaneous redox reaction. The reaction occurs naturally and produces a voltage. An example is a dry cell battery or the Daniell cell (zinc and copper electrodes). An electrolytic cell uses electrical energy to drive a non-spontaneous redox reaction. The reaction does not occur naturally — it requires an external power supply. An example is the electrolysis of molten aluminium oxide to extract aluminium, or the electrolysis of brine to produce chlorine and sodium hydroxide. The key difference is the direction of energy conversion: galvanic cells produce electricity from chemistry, while electrolytic cells use electricity to drive chemistry.
"Electrochemistry is the dance of electrons — the transfer that powers batteries, prevents rust, and builds the modern world. Now you know the steps."
