Types of Chemical Reactions: Study Guide (Chemistry 4.7)

Three questions that sort most reactions

Thousands of reactions happen in water every day in a lab. You do not need to learn them one by one. Instead, ask what is being passed between the particles:

1. Are electrons transferred? Then it is an oxidation-reduction (redox) reaction.
2. Is a proton (an H⁺ ion) transferred? Then it is an acid-base reaction.
3. Do ions in solution come together to form an insoluble solid? Then it is a precipitation reaction.

Each type has its own evidence. Electron transfer shows up as a change in oxidation numbers. Proton transfer shows up as one species losing an H and another gaining it. Precipitation shows up as a solid, marked (s), forming from ions marked (aq).

The best equation to classify is the net ionic equation from Topic 4.2. Spectator ions do nothing, so leaving them out makes the real change easier to see.

Oxidation numbers: bookkeeping for electrons

You cannot see electrons move, so chemists use oxidation numbers. An oxidation number is the charge an atom would have if every bond in the substance were fully ionic. It is a bookkeeping tool, not a measured charge. Assign oxidation numbers with these rules, in this order of priority:

1. An atom in a pure element is 0 (Mg, O₂, Cl₂, S₈).
2. A monatomic ion has an oxidation number equal to its charge (Na⁺ is +1, S²⁻ is −2).
3. The oxidation numbers in a formula add up to the overall charge: 0 for a neutral compound, the ion charge for a polyatomic ion.
4. Fluorine is always −1 in compounds. Group 1 metals are +1 and group 2 metals are +2.
5. Hydrogen is usually +1. The exception is a metal hydride such as NaH, where it is −1.
6. Oxygen is usually −2. The main exception is a peroxide such as H₂O₂, where it is −1.

Use rule 3 to find the one atom left over. Some examples:

Species	Working	Oxidation number
NH₄⁺	N + 4(+1) = +1	N is −3
SO₃²⁻	S + 3(−2) = −2	S is +4
H₂O₂	2(+1) + 2(O) = 0	O is −1
NaH	+1 + H = 0	H is −1
ClO₄⁻	Cl + 4(−2) = −1	Cl is +7
CO₂	C + 2(−2) = 0	C is +4

Write oxidation numbers with the sign first (+4, −2). This keeps them separate from ion charges, which are written with the number first (2+, 2−).

Redox reactions: electrons move

In a redox reaction, one species loses electrons and another gains them.

• Oxidation is loss of electrons. The oxidation number increases.
• Reduction is gain of electrons. The oxidation number decreases.
• Electrons travel from the species oxidised to the species reduced. Oxidation and reduction always happen together, and the electrons lost equal the electrons gained.

So the quickest test for redox is: assign oxidation numbers on both sides. If any atom’s oxidation number changes, the reaction is redox.

You may meet the terms “oxidising agent” and “reducing agent” in textbooks. The course does not assess them, so this guide simply says which species is oxidised and which is reduced. That is also the clearer way to write an exam answer.

Combustion is a kind of redox

Combustion is a redox reaction in which a substance reacts with oxygen gas, O₂. Oxygen starts at 0 and ends at −2 in the products, so it is reduced. Something else must be oxidised.

For a hydrocarbon (a compound of only carbon and hydrogen), complete combustion gives only carbon dioxide and water. For example, butane: 2C₄H₁₀ + 13O₂ → 8CO₂ + 10H₂O. When the fuel contains another element, other oxides form too. You will see an example with sulfur in the practice set.

Acid-base reactions: protons move

An acid-base reaction transfers one or more protons (H⁺ ions) from one species to another. A hydrogen nucleus moves; its electron stays behind. That is why the oxidation numbers do not change. Hydrogen is +1 before and after.

For example, when ethanoic acid (acetic acid) meets hydroxide ions:

CH₃COOH(aq) + OH⁻(aq) → CH₃COO⁻(aq) + H₂O(l)

The acid loses an H⁺ and becomes CH₃COO⁻. The hydroxide ion gains that H⁺ and becomes water. Topic 4.8 builds the full Brønsted-Lowry model of acids and bases from this idea.

Precipitation reactions: ions come together

A precipitation reaction usually happens when you mix two solutions of ionic compounds. If one possible combination of cation and anion is insoluble (or only very slightly soluble), those ions leave the solution as a solid, the precipitate.

To decide which combination forms the solid, you need only one rule from this course:

All sodium (Na⁺), potassium (K⁺), ammonium (NH₄⁺) and nitrate (NO₃⁻) salts are soluble in water.

So if a mixture forms a precipitate, the solid cannot contain any of these ions. You do not need to memorise other solubility rules. If a question needs other solubility data, it will give it to you.

In a precipitation reaction nothing changes oxidation number, and no proton moves. The ions keep their charges; they simply pack together into an ionic solid.

A decision path you can use every time

The order matters. Question 1 comes first because redox can hide behind other clues. A reaction can contain H⁺ or produce a solid and still be redox, as Worked example 3 shows.

Worked example 1: classify three reactions

Question. Classify each reaction as acid-base, redox or precipitation. Give the evidence.

(a) Mg(s) + 2Ag⁺(aq) → Mg²⁺(aq) + 2Ag(s)

(b) CH₃COOH(aq) + OH⁻(aq) → CH₃COO⁻(aq) + H₂O(l)

(c) A student mixes nickel(II) chloride solution, NiCl₂(aq), with sodium hydroxide solution, NaOH(aq). A green solid forms.

(a) Assign oxidation numbers. Mg goes from 0 (element) to +2 (ion). Ag goes from +1 (ion) to 0 (element). Oxidation numbers change, so this is redox.

• Mg’s oxidation number rises, so Mg is oxidised. Each Mg atom loses 2 electrons.
• Ag’s oxidation number falls, so Ag⁺ is reduced. Each Ag⁺ gains 1 electron, and there are 2 Ag⁺, so 2 electrons are gained.
• The 2 electrons move from Mg to the Ag⁺ ions. Lost = gained = 2.

(b) Oxidation numbers: H stays +1, O stays −2, and each C keeps its value. Nothing changes, so it is not redox. CH₃COOH loses one H⁺ and OH⁻ gains it, so this is acid-base.

(c) The mixture contains Ni²⁺, Cl⁻, Na⁺ and OH⁻ ions. Each cation could meet the other anion. The new pairs are NaCl and Ni(OH)₂.

• NaCl is a sodium salt, so it is soluble. It cannot be the solid.
• So the green solid must be Ni(OH)₂. The net ionic equation is Ni²⁺(aq) + 2OH⁻(aq) → Ni(OH)₂(s).
• Ni is +2 before and after; O and H keep −2 and +1. OH⁻ does not gain or lose a proton. This is precipitation. Na⁺ and Cl⁻ are spectator ions.

Worked example 2: combustion as electron transfer

Question. Ethane, C₂H₆, burns completely in oxygen.

(a) Write the balanced equation. (b) Use oxidation numbers to show that this is a redox reaction, and say which element is oxidised and which is reduced. (c) Show that the electrons lost equal the electrons gained.

(a) Complete combustion of a hydrocarbon gives CO₂ and H₂O. Balance C, then H, then O:

C₂H₆ + O₂ → 2CO₂ + 3H₂O needs 7 O atoms on the right, so 3.5 O₂. Doubling to clear the half gives

2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O

Check: C 4 = 4; H 12 = 12; O 14 = 4(2) + 6 = 14.

(b) Oxidation numbers:

Atom	Before	After	Change
C	−3 in C₂H₆ (2C + 6(+1) = 0)	+4 in CO₂	rises by 7: oxidised
H	+1 in C₂H₆	+1 in H₂O	none
O	0 in O₂	−2 in CO₂ and H₂O	falls by 2: reduced

Carbon is oxidised and oxygen is reduced, so the reaction is redox.

(c) In the balanced equation there are 4 C atoms, each losing 7 electrons: 4 × 7 = 28 electrons lost. There are 14 O atoms, each gaining 2 electrons: 14 × 2 = 28 electrons gained. They match, as they must.

Check. If you had used 3.5 O₂ per ethane and found unequal electron totals, that would signal an arithmetic slip. The equality is a useful test of any balanced redox equation.

Worked example 3: two reactions that look like something else

Question. Classify each reaction and explain the trap.

(a) Mg(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂(g)

(b) Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s)

(a) It is tempting to call this acid-base because H⁺ is a reactant. Check oxidation numbers first. Mg goes from 0 to +2, so Mg is oxidised. H goes from +1 to 0 (in H₂), so H⁺ is reduced. This is redox. No proton is passed to a base; instead, the H⁺ ions gain electrons.

(b) It is tempting to call this precipitation because a solid forms. But the solid is silver metal, not an ionic compound made from two ions in solution. Cu goes from 0 to +2 (oxidised) and Ag from +1 to 0 (reduced). This is redox.

Compare CaCO₃(s) + 2H⁺(aq) → Ca²⁺(aq) + CO₂(g) + H₂O(l). Here C stays +4, O stays −2, H stays +1 and Ca stays +2. No oxidation number changes. The carbonate ion accepts protons, so this one really is acid-base, even though it gives off a gas.

Common misconceptions

• “Any reaction with H⁺ is acid-base.” Metals react with H⁺ by redox: the H⁺ gains electrons and becomes H₂.
• “Any reaction that makes a solid is precipitation.” A metal deposited from solution is a redox product. Precipitation forms an ionic solid from aqueous ions.
• Confusing oxidation numbers with ion charges. Carbon in CO₂ has an oxidation number of +4, but CO₂ contains no C⁴⁺ ion.
• Getting the direction of electron flow backwards. Electrons leave the species oxidised (oxidation number goes up) and arrive at the species reduced (oxidation number goes down).
• Thinking only one of oxidation and reduction can happen. They always happen together, with equal numbers of electrons lost and gained.
• Forgetting that combustion is redox. Burning is electron transfer to oxygen, even though no ions appear in the equation.
• Picking a sodium, potassium, ammonium or nitrate salt as the precipitate. These salts are always soluble, so they stay in solution as spectator ions.

Where this leads

Next, Topic 4.8 develops the Brønsted-Lowry model of acid-base reactions, including conjugate acid-base pairs. Topic 4.9 then balances redox equations with half-reactions, using the oxidation numbers you assigned here. Acid-base equilibria return in Unit 8, solubility equilibria in Unit 7, and electron transfer in electrochemistry in Unit 9. Try the practice questions now, then use the revision notes and the checklist to consolidate.