Acids, Bases and Salts: Unit Review

Introduction

This unit has taken you from the idea of hydrogen ions in water all the way to fertiliser factories, blood buffers and acid rain. Reviewing it as a connected story, rather than a list of facts, makes it easier to remember and to apply in unfamiliar questions. This page gathers the key definitions, the essential equations and the links between topics, then tests them with mixed questions.

Core explanation

Acids and alkalis in water. Acids release H⁺(aq) ions in water; water is needed because the ions only form in solution. Alkalis are soluble bases that release OH⁻(aq). Bases are the wider group of substances that neutralise acids: metal oxides, hydroxides, carbonates and ammonia.

pH. The pH scale runs roughly from 0 to 14 at 25 °C. Below 7 is acidic, 7 is neutral, above 7 is alkaline. Each unit is a tenfold change in [H⁺]. pH is measured with universal indicator (a colour and approximate value) or a pH meter (a precise value).

Strength and concentration. Strong acids (hydrochloric, nitric, sulfuric) ionise completely; weak acids (ethanoic, citric, carbonic) ionise only partly. Strong alkalis such as sodium hydroxide ionise completely; ammonia is a weak alkali. Concentration is a separate idea: the amount dissolved per dm³. Diluting a solution moves its pH towards 7 without changing strength.

Neutralisation. The core ionic equation is H⁺(aq) + OH⁻(aq) → H₂O(l). It is exothermic. Titration, with a suitable indicator, finds exactly how much acid reacts with a given amount of alkali, and pH curves show the steep change at the end point.

The general reactions of acids.

Reactant Products Example --- --- --- Metal (above hydrogen) Salt + hydrogen Mg + 2HCl → MgCl₂ + H₂ Metal oxide Salt + water CuO + H₂SO₄ → CuSO₄ + H₂O Metal hydroxide Salt + water NaOH + HNO₃ → NaNO₃ + H₂O Carbonate Salt + water + carbon dioxide CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ Hydrogencarbonate Salt + water + carbon dioxide NaHCO₃ + HCl → NaCl + H₂O + CO₂ Ammonia Ammonium salt NH₃ + HCl → NH₄Cl

Salts. The salt's name comes from the metal (or ammonium) and the acid: chloride, sulfate, nitrate, phosphate, ethanoate. Formulae are built by balancing ionic charges. Solubility rules decide the preparation method: excess insoluble base then filtration and crystallisation; titration for two soluble reactants; precipitation for insoluble salts, summarised by ionic equations such as Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s). Some crystals contain water of crystallisation, such as CuSO₄·5H₂O. Salt solutions can be acidic, neutral or basic depending on the acid and base they came from.

Oxides. Metal oxides are generally basic, non-metal oxides are generally acidic (SO₂, CO₂, NO₂), and some, such as aluminium oxide and zinc oxide, are amphoteric, reacting with both acids and bases.

Applications. Soil is limed to raise its pH; acid rain from SO₂ and NOₓ damages lakes, forests and buildings; fertilisers are ammonium, nitrate, phosphate and potassium salts; antacids neutralise stomach acid; barium sulfate is safe for X-rays because it is insoluble; buffers hold blood near pH 7.4; sulfuric acid from the Contact process underpins industry. Throughout, corrosive substances demand eye protection and careful dilution, adding acid to water.

Step-by-step reasoning

Using the unit in an unfamiliar question:

1. Decide whether the substance is an acid, base, alkali or salt. 2. Predict the reaction type and products from the general equations. 3. Consider solubility and state symbols. 4. Link to pH: which reactant is in excess? 5. Connect to the context given (industry, environment, body).

Visual explanation

Picture a concept map with "H⁺ ions in water" at the centre. Arrows radiate out to "pH scale", "strong/weak", "neutralisation", and "reactions of acids". From "neutralisation", arrows lead to "titration", "salts" and "energy change"; from "salts", arrows lead to "naming and formulae", "solubility rules" and "preparation methods"; and outer bubbles show "soil", "acid rain", "fertilisers", "medicine", "buffers" and "safety".

Real-world analogy

The unit is like a tree. The roots are H⁺ and OH⁻ ions in water; the trunk is neutralisation; the branches are the reactions of acids and the methods of making salts; the leaves are the applications, from farming to medicine. Every leaf is fed by the same roots.

Real-world example

A single field shows the whole unit at work. Acid rain and ammonium fertilisers lower the soil pH; the farmer measures it and spreads limestone, which neutralises the acid and releases carbon dioxide. The fertiliser itself is a salt made by neutralising acids with ammonia in a factory that depends on sulfuric and nitric acids.

Why?

Why does a single ionic equation, H⁺ + OH⁻ → H₂O, underlie so many different reactions? In every acid–alkali neutralisation, the metal ions and acid anions are spectators. Only H⁺ and OH⁻ change, which is why all such reactions release a similar amount of energy per mole of water formed.

Common misconception

"Base and alkali mean the same thing." All alkalis are bases, but not all bases are alkalis. Copper(II) oxide is a base that does not dissolve in water, so it is not an alkali.

Worked example

Question: Zinc carbonate is added to dilute nitric acid until no more reacts. Name the salt, write the equation, state how the pH changes and describe how to obtain the solid salt.

Reasoning: Carbonate + acid → salt + water + CO₂. Zinc is Zn²⁺, nitrate is NO₃⁻. The acid is used up, so pH rises to about 7. Zinc carbonate is insoluble, so the excess can be filtered off.

Answer: Zinc nitrate; ZnCO₃ + 2HNO₃ → Zn(NO₃)₂ + H₂O + CO₂; pH rises towards 7; filter off excess zinc carbonate, then evaporate and crystallise the filtrate.

Quick check

1. What is the ionic equation for every acid–alkali neutralisation? Answer: H⁺(aq) + OH⁻(aq) → H₂O(l)

Exam focus

Make sure you can: write all six general reactions of acids with balanced equations; name and write formulae of salts; choose and justify a salt preparation method; distinguish strength from concentration; interpret pH values as tenfold steps; and explain at least one application in agriculture, the environment, medicine and industry.

Advanced insight

At a higher level, acids are defined as proton donors and bases as proton acceptors (the Brønsted–Lowry theory). This explains why ammonia is a base even though it contains no hydroxide: it accepts a proton to become NH₄⁺. It also explains buffers and acidic or basic salt solutions, which will be developed further in later units.

Summary

Acids release H⁺ in water, alkalis release OH⁻, and pH measures [H⁺] on a tenfold scale. Strength (ionisation) differs from concentration. Neutralisation, H⁺ + OH⁻ → H₂O, is exothermic and is followed by titration. Acids react with metals, oxides, hydroxides, carbonates, hydrogencarbonates and ammonia to give salts, prepared by methods chosen from solubility. These ideas explain liming, acid rain, fertilisers, medicines, buffers, industry and safety.

Practice questions

1. Explain the difference between a base and an alkali, giving one example of each that is not the other. Answer: A base neutralises an acid; an alkali is a base that dissolves in water to give OH⁻. Copper(II) oxide is a base but not an alkali; sodium hydroxide is an alkali. 2. Solution A has pH 3 and solution B has pH 5. How many times greater is [H⁺] in A than in B? Answer: 100 times, since the pH differs by 2 units and each unit is a factor of 10. 3. Write the ionic equation for the precipitation of silver chloride. Answer: Ag⁺(aq) + Cl⁻(aq) → AgCl(s) 4. Explain why limestone is used both to treat acidic lakes and in power-station chimneys. Answer: Calcium carbonate is a base; it neutralises acids in lake water and reacts with the acidic gas sulfur dioxide in flue gases, reducing acid rain. 5. Name the salt formed when potassium hydroxide neutralises sulfuric acid and write its formula. Answer: Potassium sulfate, K₂SO₄.