Types of Reaction: Redox, Acids and Neutralisation
Oxidation and Reduction
Scope note: redox, acids and neutralisation are topic C3.3, examined on Chemistry Paper 1 (Foundation) or Paper 3 (Higher), and assumed knowledge for the Paper 2 topics.
Oxidation is the gain of oxygen or the loss of electrons. Reduction is the loss of oxygen or the gain of electrons. These two processes always happen simultaneously in redox reactions. A useful mnemonic is OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).
For example, when copper oxide is reduced by carbon: \(2CuO + C \rightarrow 2Cu + CO_2\). The copper oxide loses oxygen (it is reduced); the carbon gains oxygen (it is oxidised). The carbon is the reducing agent and copper oxide is the oxidising agent.
In a redox reaction, the oxidising agent causes the oxidation of another substance but is itself reduced. The reducing agent causes the reduction of another substance but is itself oxidised. This is a common source of confusion — always track both substances separately.
At Higher Tier you must also explain redox in terms of electron transfer. In the reaction between magnesium and oxygen: \(Mg \rightarrow Mg^{2+} + 2e^-\) (magnesium is oxidised — it loses electrons); \(O_2 + 4e^- \rightarrow 2O^{2-}\) (oxygen is reduced — it gains electrons). The substance that loses electrons is the reducing agent; the substance that gains electrons is the oxidising agent.
The Reactivity Series
The reactivity series lists metals in order of their tendency to react — specifically, how readily each metal is oxidised (loses electrons). Metals above hydrogen in the series can displace hydrogen from dilute acids. Metals higher in the series can displace metals lower in the series from their salt solutions — these are called displacement reactions.
The order of metals in the reactivity series can be remembered using a mnemonic. From most reactive to least reactive, the series runs: potassium, sodium, calcium, magnesium, aluminium, zinc, iron, tin, lead, hydrogen, copper, silver, gold, platinum. Metals above hydrogen react with dilute acids to produce hydrogen gas. Metals below hydrogen (copper, silver, gold) do not react with dilute acids. This knowledge allows chemists to predict whether a given displacement reaction is possible: if a more reactive metal is added to a salt solution of a less reactive metal, a displacement will occur. If the added metal is less reactive than the metal in solution, no reaction takes place.
When zinc is added to copper sulfate solution: \(Zn + CuSO_4 \rightarrow ZnSO_4 + Cu\). Zinc is more reactive than copper, so it displaces copper from the solution. Zinc is oxidised (loses electrons, forms \(Zn^{2+}\)) and copper ions are reduced (gain electrons, form Cu atoms). The blue colour of the copper sulfate solution fades as colourless zinc sulfate forms, and a reddish copper deposit appears. This is both a displacement reaction and a redox reaction.
Acids, Alkalis and the pH Scale
The pH scale runs from 0 to 14. A pH below 7 indicates an acidic solution (high concentration of \(H^+\) ions); pH 7 is neutral; above 7 is alkaline (high concentration of \(OH^-\) ions). At Higher Tier you also need the terms strong, weak, concentrated and dilute. Strong acids (such as hydrochloric acid, HCl) fully dissociate in water into ions, whereas weak acids (such as ethanoic acid) only partially dissociate. It is important to distinguish between a strong acid (degree of ionisation) and a concentrated acid (amount of acid dissolved per litre) — a concentrated weak acid is not the same as a strong acid.
Strong acid: fully dissociates into ions in water (e.g. HCl \(\rightarrow H^+ + Cl^-\)). Weak acid: partially dissociates (e.g. ethanoic acid). Concentrated: many moles of acid per litre. Dilute: few moles per litre. These are independent properties — a strong acid can be dilute.
Neutralisation and Making Salts
Neutralisation occurs when an acid reacts with a base or alkali. The general equation is: acid + base → salt + water. For example, hydrochloric acid and sodium hydroxide: \(HCl + NaOH \rightarrow NaCl + H_2O\). At the ionic level, the \(H^+\) ion from the acid reacts with the \(OH^-\) ion from the alkali: \(H^+ + OH^- \rightarrow H_2O\).
Acids also react with metals, carbonates and metal oxides. Key reactions to know:
- Acid + metal → salt + hydrogen (e.g. \(Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2\))
- Acid + metal carbonate → salt + water + carbon dioxide (e.g. \(CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2\))
- Acid + metal oxide → salt + water (e.g. \(CuO + H_2SO_4 \rightarrow CuSO_4 + H_2O\))
In your practical work (PAG C6), you will carry out acid-alkali titrations to find the volume of acid needed to exactly neutralise a given volume of alkali. A suitable indicator (e.g. methyl orange or phenolphthalein) is added to the alkali in a conical flask. Acid is added from a burette drop by drop until the indicator changes colour permanently at the end point. The titration is repeated without indicator to produce a pure salt solution, which can be evaporated carefully to obtain crystals. Concordant results (within 0.10 cm3 of each other) are used to calculate the mean titre.