Revision notes: Acids, Bases and Salts

These notes work through Acids, Bases and Salts one content standard at a time. This is one of the most calculation- and practical-heavy chapters in SPM Chemistry, so alongside the definitions we set out every formula with its units. Keep the notes beside past-paper questions and use them to check that your working, your equations and your key terms match what the marking scheme rewards.

6.1 Properties of acids and bases

An acid is a chemical that ionises in water to produce hydrogen ions, H⁺. A base is a metal oxide or metal hydroxide that reacts with an acid to form a salt and water; a base that is soluble in water is an alkali, and it ionises in water to produce hydroxide ions, OH⁻.

The role of water. Acidic and alkaline properties only appear when the substance is dissolved in water. Dry hydrogen chloride gas has no acidic properties because there are no free H⁺ ions; dissolved in water it ionises to give H⁺ and behaves as an acid. This idea is examined again and again.

Typical properties to remember:

• Acids taste sour, turn blue litmus red, react with reactive metals to give hydrogen gas, react with carbonates to give carbon dioxide, and react with bases in neutralisation to give a salt and water.
• Alkalis feel soapy, turn red litmus blue, react with acids in neutralisation, liberate ammonia gas from ammonium salts, and precipitate insoluble metal hydroxides from solutions of many metal ions.

A strong acid (such as hydrochloric, sulfuric or nitric acid) ionises completely in water; a weak acid (such as ethanoic acid) ionises only partially. The same distinction applies to strong and weak alkalis. Concentration (how much acid per unit volume) is a separate idea from strength (how fully it ionises).

6.2 The pH scale

The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is by the concentration of hydrogen ions present. A lower pH means a higher concentration of H⁺ ions.

• pH less than 7, acidic; the smaller the pH, the more acidic.
• pH equal to 7, neutral (for example pure water).
• pH greater than 7, alkaline; the larger the pH, the more alkaline.

Universal indicator gives a colour for each region, reds and oranges in acids, green at neutral, blues and purples in alkalis, while litmus only tells you acidic or alkaline. Because pH depends on H⁺ concentration, a strong acid has a lower pH than a weak acid of the same molarity.

6.3 Concentration and molarity of solutions

Concentration is the amount of solute dissolved in a given volume of solution. It can be expressed in grams per cubic decimetre (g dm⁻³) or in moles per cubic decimetre (mol dm⁻³), and the second is called the molarity.

• Molarity (mol dm⁻³) = number of moles of solute ÷ volume of solution (in dm³).
• To convert: concentration in g dm⁻³ = molarity (mol dm⁻³) × molar mass (g mol⁻¹).
• 1 dm³ = 1000 cm³, so number of moles = molarity × volume in cm³ ÷ 1000.

Carrying units through every line is the habit that protects these marks. If a volume is given in cm³, convert it to dm³ before using it in a molarity formula, or divide by 1000 at the right step.

6.4 Standard solution and dilution

A standard solution is a solution whose concentration is accurately known. It is prepared by dissolving a known mass of solute in a little water, transferring it to a volumetric flask, and adding water up to the calibration mark so the final volume is exact.

Dilution adds water to lower the concentration. The number of moles of solute does not change when you add water, which gives the dilution relationship:

M₁V₁ = M₂V₂, where M₁ and V₁ are the molarity and volume before dilution and M₂ and V₂ are the molarity and volume after. Rearrange to find whichever quantity is unknown, keeping the volume units the same on both sides.

6.5 Neutralisation and titration

Neutralisation is the reaction between an acid and a base (or alkali) to produce a salt and water only: acid + base → salt + water. In ionic terms, H⁺ + OH⁻ → H₂O.

Titration is the technique used to find an unknown concentration. A pipette measures a fixed volume of one solution into a conical flask, a few drops of indicator are added, and the other solution is run in from a burette until the indicator just changes colour at the end point. From the balanced equation:

For an acid–base reaction aH⁺ with bOH⁻, use (Mₐ × Vₐ) / (M_b × V_b) = a / b, where a and b are the mole ratios of acid to base in the balanced equation. Solve for the single unknown.

6.6 Preparation of salts

Whether a salt is soluble or insoluble decides the method. Useful solubility rules: all nitrates are soluble; all sodium, potassium and ammonium salts are soluble; most chlorides are soluble except those of silver and lead; most sulfates are soluble except those of barium, lead and calcium; most carbonates are insoluble except those of sodium, potassium and ammonium.

• Soluble salts (except sodium, potassium, ammonium): react an acid with an excess of an insoluble base, metal or carbonate, filter off the excess, then evaporate and crystallise.
• Soluble salts of sodium, potassium or ammonium: use titration, because the base is soluble and cannot be added in a filterable excess.
• Insoluble salts: use precipitation (double decomposition), mix two soluble solutions so the insoluble salt forms, then filter, wash and dry it.

6.7 Qualitative analysis of salts

Qualitative analysis identifies the ions in a salt through a set of standard tests.

• Action of heat on a solid can give clues (for example a carbonate releasing carbon dioxide, or a coloured residue).
• Gas tests: carbon dioxide turns limewater milky; oxygen relights a glowing splint; hydrogen gives a “pop” with a lighted splint; ammonia turns damp red litmus blue.
• Cation tests use sodium hydroxide solution and ammonia solution, observing which coloured or white precipitates form and whether they dissolve in excess.
• Anion tests identify carbonate, chloride, sulfate and nitrate using acid, silver nitrate, barium chloride and the brown-ring test respectively.

How to use these notes

Revise one content standard at a time and write each definition from memory. For 6.3 to 6.5, practise the calculations until the units flow automatically, that is where most marks in this chapter are won or lost. A one-to-one teacher can check that your titration working and ionic equations are complete and correctly set out. Then move on to the worked examples and the practice questions for this chapter and mark yourself against the definitions and formulae above.