chemistry

Half-reaction

Half-reaction is a chemistry topic covered in the lgStudy science library. This page brings together a partial reference excerpt, illustrations, worked examples, real-world applications and a short study plan, so you can understand Half-reaction rather than just read about it. In short: In chemistry, a half reaction (or half-cell reaction) is either the oxidation or reduction reaction component of a redox reaction. A half reaction is obtained by considering the change in oxidation states of individual substances involved in the redox reaction.

Key takeaways

Half-reaction belongs to chemistry; place it in that map before memorising details.
Learn the definition first, then one example that makes the definition concrete.
Connect Half-reaction to a quantity you can measure, compute or draw — that is where exam questions come from.
Reproduce the core statement of Half-reaction from memory before moving on to harder problems.

Reference excerpt

In chemistry, a half reaction (or half-cell reaction) is either the oxidation or reduction reaction component of a redox reaction. A half reaction is obtained by considering the change in oxidation states of individual substances involved in the redox reaction. Often, the concept of half reactions is used to describe what occurs in an electrochemical cell, such as a Galvanic cell battery at one of the two electrodes. In the case of metal plating and metal stripping the same half reaction can be written to describe both the metal undergoing oxidation (known as the anode) and the metal undergoing reduction (known as the cathode). Half reactions are often used as a method of balancing redox reactions. For oxidation-reduction reactions in acidic conditions, after balancing the atoms and oxidation numbers, H+ ions must be added to balance the hydrogen ions in the half reaction. For oxidation-reduction reactions in basic conditions, after balancing the atoms and oxidation numbers, it must be treated as an acidic solution, and then OH− ions must be added to balance the H+ ions in the half reactions (which would give H2O).

Example: Zn and Cu galvanic cell

The Galvanic cell shown in the adjacent image is constructed with a piece of zinc (Zn) submerged in a solution of zinc sulfate (ZnSO4) and a piece of copper (Cu) submerged in a solution of copper(II) sulfate (CuSO4). The overall reaction is:

Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) At the Zn anode, oxidation takes place, meaning the metal loses electrons. This is represented in the following oxidation half reaction (with the electrons on the product side):

Zn(s) → Zn2+ + 2 e− At the Cu cathode, reduction takes place, meaning electrons are gained or accepted. This is represented in the following reduction half reaction (with the electrons on the reactant side):

Cu2+ + 2 e− → Cu(s)

Example: oxidation of magnesium

Consider the example burning of magnesium ribbon (Mg). When magnesium burns, it combines with oxygen (O2) from the air to form magnesium oxide (MgO) according to the following equation:

2 Mg(s) + O2(g) → 2 MgO(s) Magnesium oxide is an ionic compound containing Mg2+ and O2− ions whereas Mg(s) and O2(g) are elements with no charges. The Mg(s) with zero charge gains a +2 charge going from the reactant side to product side, and the O2(g) with zero charge gains a −2 charge. This is because when Mg(s) becomes Mg2+, it loses 2 electrons. Since there are 2 Mg on left side, a total of 4 electrons are lost according to the following oxidation half reaction:

2 Mg(s) → 2 Mg2+ + 4 e− On the other hand, O2 was reduced: its oxidation state goes from 0 to −2. Thus, a reduction half reaction can be written for the O2 as it gains 4 electrons:

O2(g) + 4 e− → 2 O2− The overall reaction is the sum of both half reactions:

2 Mg(s) + O2(g) + 4 e− → 2 Mg2+ + 2 O2− + 4 e− When chemical reaction, especially, redox reaction takes place, the naked eye cannot see the electrons as they appear and disappear during the course of the reaction. Instead, observers see the reactants and end products. Due to this, electrons appearing on both sides of the equation are canceled. After canceling, the equation is re-written as

2 Mg(s) + O2(g) → 2 Mg2+ + 2 O2− Two ions, positive (Mg2+) and negative (O2−), exist on product side, and they combine immediately to form the magnesium oxide (MgO) compound due to their opposite charges (electrostatic attraction). In any given oxidation-reduction reaction, there are two half reactions: the oxidation half reaction and the reduction half reaction. The sum of these two half reactions is the oxidation–reduction reaction.

Half-reaction balancing method In the redox reaction between iron and chlorine:

Cl2 + 2 Fe2+ → 2 Cl− + 2 Fe3+ The two elements involved, iron and chlorine, each change oxidation state; iron from +2 to +3, chlorine from 0 to −1. There are then effectively two half reactions occurring. These changes can be represented in formulas by inserting appropriate electrons into each half reaction:

Fe2+ → Fe3+ + e− Cl2 + 2 e− → 2 Cl− Given two half reactions it is possible, with knowledge of appropriate electrode potentials, to arrive at the complete (original) reaction the same way. The decomposition of a reaction into half reactions is key to understanding a variety of chemical processes. For example, in the above reaction, it can be shown that this is a redox reaction in which Fe is oxidised, and Cl is reduced (so there is a transfer of electrons from Fe to Cl). Decomposition is also a way to simplify the balancing of a chemical equation. A chemist can atom balance and charge balance one piece of an equation at a time. For example:

Fe2+ → Fe3+ + e− becomes 2 Fe2+ → 2 Fe3+ + 2e− is added to Cl2 + 2 e− → 2 Cl− and finally becomes Cl2 + 2 Fe2+ → 2 Cl− + 2 Fe3+

Basic and acidic conditions It is also possible–sometimes necessary–to consider a half reaction in either basic or acidic conditions, as there may be an acidic or basic electrolyte involved in the redox reaction. Due to this electrolyte, it may be more difficult to satisfy the balance of both the atoms and charges; this is done by adding H2O, OH−, e−, and/or H+ to either side of the reaction until both atoms and charges are balanced. For the half reaction below:

PbO2 → PbO OH−, H2O, and e− can be used to balance the charges and atoms in basic conditions, as long as it is assumed that the reaction is in water.

2 e− + H2O + PbO2 → PbO + 2 OH− For the half reaction below:

PbO2 → PbO H+, H2O, and e− can be used to balance the charges and atoms in acidic conditions, as long as it is assumed that the reaction is in water.

2 e− + 2 H+ + PbO2 → PbO + H2O Both sides are charge balanced, meaning the overall charge is neutral, and atom balanced, meaning there are equal numbers of each element on both sides. Often there will be both H+ and OH− present in acidic and basic conditions, but the resulting reaction of the two ions will yield water.

H+ + OH− → H2O

Illustrations

Worked examples

Example 1 — a first encounter with Half-reaction

Start with the simplest possible case. Write down what Half-reaction claims or describes in one sentence, then invent the smallest concrete situation in which that sentence is true. In chemistry, the smallest case is usually a single object, a single equation or a single measurement. Check that every symbol or term in your sentence has a meaning in that case.

Example 2 — changing one variable

Take the situation from Example 1 and change exactly one quantity: double it, halve it, or set it to zero. Predict what should happen to Half-reaction before you calculate. Comparing your prediction with the result is the fastest way to find out whether you understand the idea or only the words.

Example 3 — an exam-style question

Typical questions about Half-reaction ask you to (a) state it precisely, (b) apply it to given data, and (c) explain a limitation. Practise writing all three answers in under five minutes; the third part is what separates a full-mark answer from an average one.

Applications of Half-reaction

In research: Half-reaction appears in chemistry research whenever the underlying quantities have to be modelled precisely. Papers usually cite it as a starting assumption and then explore where it breaks down.
In technology and industry: Engineering practice reuses Half-reaction in design rules, simulations and safety margins. Knowing the idea lets you read a specification sheet and understand why the numbers look the way they do.
In the classroom: Half-reaction is common in secondary-school and first-year university syllabi. It links to neighbouring topics Electrochemistry, so understanding it makes those chapters shorter.
In everyday life: Look for Half-reaction outside the textbook — in sport, cooking, traffic, electronics or the sky above you. An example you found yourself is remembered far longer than one you were given.

Ask Teacher Smith questions about this articleOpens your AI tutor with a question about “Half-reaction” →

Affiliate

Preply — study more efficiently by working with a personal tutor. 50% off.

How to study Half-reaction in 20 minutes

Read the reference excerpt below once, without taking notes.
Close the page and write down what Half-reaction means in your own words.
Compare your version with the excerpt and mark what you missed.
Work through the three examples above with pen and paper.
Explain Half-reaction out loud to somebody else — or to Teacher Smith in the lgStudy chat.

Frequently asked questions

What is Half-reaction in simple terms?

Why does Half-reaction matter?

Because it connects several chemistry ideas at once: it gives you a definition you can apply, a quantity you can calculate, and a way to check whether a result is plausible.

How should I study Half-reaction?

Read the excerpt, restate it from memory, then work through the examples and applications listed on this page. The five-step study plan above takes about twenty minutes.

What does this page cover?

It gives you a compact reference excerpt plus original lgStudy explanations, examples, applications and study material on Half-reaction.