Difficulty: Intermediate | Prerequisites: General chemistry fundamentals, familiarity with ions and electron configuration.
Oxidation-reduction (redox) reactions sit at the heart of chemistry, from batteries to metabolism. This topic covers the transfer of electrons between species, how to track those transfers using oxidation numbers, and how to balance the resulting equations. If you are comfortable with ionic compounds and basic equation balancing, you are ready for this. If not, revisit those first.
Redox reactions involve the transfer of electrons from one species to another. Oxidation is the loss of electrons; reduction is the gain. You assign oxidation numbers using a set of rules, then balance redox equations by splitting them into half-reactions and equalising both atoms and charge.
Oxidation
The loss of one or more electrons by a species. In simple terms, when a substance is oxidised it gives electrons away, and its oxidation number goes up.
Reduction
The gain of one or more electrons by a species. Think of it as the opposite of oxidation: the species "reduces" its charge by accepting electrons, and its oxidation number goes down.
Oxidising agent (oxidant)
The species that causes another to be oxidised. It accepts electrons from the other species and is itself reduced. Typically contains oxygen or another highly electronegative atom.
Reducing agent (reductant)
The species that causes another to be reduced. It donates electrons and is itself oxidised. Typically contains hydrogen or a metal.
Oxidation number (oxidation state)
A bookkeeping number assigned to an atom that tracks electron ownership in a compound. It represents the hypothetical charge an atom would have if all bonds were completely ionic.
Half-reaction
One of the two parts of a redox equation, showing either the oxidation or the reduction step in isolation, with electrons written explicitly.
Formal charge
A different bookkeeping system that assumes electrons in bonds are shared equally. Oxidation number, by contrast, assigns both bonding electrons to the more electronegative atom. The two are often confused but measure different things.
Spectator ion
An ion present in solution that appears on both sides of a reaction and takes no part in the redox process. Removed when writing net ionic equations.
Net ionic equation
The equation left after all spectator ions are cancelled out, showing only the species that change during the reaction.
Disproportionation (dismutation)
A redox reaction in which one element is simultaneously oxidised and reduced, producing at least two products with different oxidation states for that element.
Oxidation numbers let you track which atoms gain or lose electrons in a reaction. Apply the rules in the order below; earlier rules take priority.
Free (uncombined) elements have an oxidation number of 0. Examples: N₂, P₄, S₈, O₂.
Monatomic ions take the charge of the ion as their oxidation number. Na⁺ is +1; Cu²⁺ is +2.
Group IA elements in a compound are always +1.
Group IIA elements in a compound are always +2.
Hydrogen is usually +1, but is -1 when bonded to a less electronegative element (Groups IA, IIA). Example: H is +1 in HCl but -1 in NaH.
Oxygen is usually -2, with two exceptions:
In peroxides (O₂²⁻), oxygen is -1.
In OF₂, oxygen is +2 (fluorine is more electronegative).
Group VIIA (halogens) are usually -1, except when combined with a more electronegative element. Example: Cl is -1 in HCl but +1 in HOCl.
The sum of all oxidation numbers in a neutral compound equals 0.
The sum of all oxidation numbers in a polyatomic ion equals the ion's charge. Example: in SO₄²⁻ the oxidation numbers sum to -2.
A useful shortcut: compounds are typically written cation-first, anion-second. If hydrogen appears at the end of a formula (as in NaH), it is likely the anion and carries a negative oxidation number.
Metals tend to lose electrons (oxidised), making them good reducing agents. Non-metals tend to gain electrons (reduced), making them good oxidising agents.
Oxidation number vs. formal charge
Formal charge assumes electrons in a bond are split equally between the two atoms. Oxidation number gives both bonding electrons to the more electronegative atom. This is why carbon in CO₂ has an oxidation state of +4 (both bonding pairs assigned to oxygen) but a formal charge of 0.
To balance a redox equation, both the number of atoms and the net charge must be equal on each side. The half-reaction (ion-electron) method breaks the problem into manageable pieces.
Separate the half-reactions. Identify which species is oxidised and which is reduced. Write each process as its own equation.
Balance atoms other than O and H first, then balance oxygen and hydrogen:
In acidic solution: add H₂O to balance O, then add H⁺ to balance H.
In basic solution: add OH⁻ and H₂O to balance O and H.
Balance charge by adding electrons. Add e⁻ to whichever side is more positive in each half-reaction.
Equalise electrons between the two half-reactions. Multiply each half-reaction by the factor needed so that both transfer the same number of electrons.
Add the half-reactions and cancel species that appear on both sides. Verify that atoms and charges balance in the final equation.
Sum of oxidation numbers in a neutral compound = 0
Sum of oxidation numbers in a polyatomic ion = charge of that ion
Electrons lost (oxidation half-reaction) = Electrons gained (reduction half-reaction), after balancing
Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain) or LEO the lion says GER (Lose Electrons = Oxidation, Gain Electrons = Reduction).
Redox reactions power batteries and fuel cells: one electrode is oxidised and the other is reduced, generating electrical current. Corrosion (rusting of iron) is a redox process where iron loses electrons to oxygen and water. In biology, cellular respiration is a long chain of redox reactions that transfer electrons from glucose to oxygen, releasing energy your cells use as ATP.
Students often think "oxidation" requires oxygen. It does not. Oxidation simply means loss of electrons, regardless of whether oxygen is involved.
Confusing oxidation number with formal charge. Formal charge splits bonding electrons equally; oxidation number assigns them to the more electronegative atom. They give different numbers for the same atom.
Forgetting that the oxidising agent is the species that gets reduced (and vice versa). The agent is named for what it does to the other species, not what happens to itself.
Assuming hydrogen is always +1. In metal hydrides such as NaH or CaH₂, hydrogen is -1 because the metal is less electronegative.
⚠️ Assigning oxidation numbers is a near-certain exam topic. Know the rules cold, especially the exceptions for hydrogen (-1 in metal hydrides) and oxygen (-1 in peroxides, +2 in OF₂).
⚠️ Expect at least one question asking you to identify the oxidising and reducing agents in a given reaction. Remember: the oxidising agent is reduced; the reducing agent is oxidised.
⚠️ Balancing half-reactions in acidic vs. basic solution uses different steps for O and H. Do not mix them up. Acidic = H₂O + H⁺. Basic = OH⁻ + H₂O.
⚠️ The difference between oxidation number and formal charge is a common short-answer or multiple-choice trap.
True or false: Oxidation always involves oxygen.
False. Oxidation is the loss of electrons, regardless of whether oxygen is present.
Fill in the blank: The oxidation number of oxygen in H₂O₂ is ___.
-1.
True or false: The reducing agent in a reaction is the species that gets reduced.
False. The reducing agent is oxidised (it donates electrons).
Fill in the blank: In a neutral compound, the sum of all oxidation numbers equals ___.
True or false: In NaH, hydrogen has an oxidation number of +1.
False. Hydrogen is -1 in NaH because sodium is less electronegative.
Q: What is the oxidation number of Cl in HOCl?
A: +1. Hydrogen is +1, oxygen is -2, and the compound is neutral, so Cl must be +1 to make the sum zero.
Q: In the reaction 2Mg + O₂ → 2MgO, identify the oxidising agent and the reducing agent.
A: O₂ is the oxidising agent (it is reduced, gaining electrons). Mg is the reducing agent (it is oxidised, losing electrons).
Q: What is the first step when balancing a redox equation using the half-reaction method?
A: Separate the overall equation into two half-reactions, one for oxidation and one for reduction.
Q: What is the oxidation state of carbon in CO₂?
A: +4. Each oxygen is -2 (total -4), and the molecule is neutral, so carbon must be +4.
Q: How do you balance oxygen atoms in a half-reaction in acidic solution?
A: Add H₂O molecules to the side that needs more oxygen atoms, then add H⁺ ions to balance the hydrogen atoms introduced by the water.
Redox reactions connect directly to electrochemistry (galvanic and electrolytic cells), where half-reactions occur at separate electrodes and generate measurable voltage. They also link to thermodynamics, since the free energy of a redox reaction determines whether it is spontaneous. In biology, electron transport chains in cellular respiration and photosynthesis are extended series of redox reactions.
Redox, oxidation-reduction, electron transfer, oxidation number, oxidation state, half-reaction, ion-electron method, oxidising agent, oxidant, reducing agent, reductant, OIL RIG, LEO GER, formal charge, free element rule, peroxide, metal hydride, balancing redox equations, acidic solution balancing, basic solution balancing, Chemistry Ch. 11, Michigan State University