Topic 7: Acids, bases and salts

Acids, bases and salts sit at the heart of everyday chemistry, from the citric acid in fruit to the limestone that neutralises acidic soil. In this topic you will learn how these substances behave, how to measure and explain acidity, and how to make pure salts in the laboratory. You will also build a reliable toolkit of chemical tests for identifying unknown ions and gases.

Properties of acids and alkalis

Acids are substances that release hydrogen ions (H+) when dissolved in water. Their characteristic reactions follow predictable patterns. Acids react with reactive metals to produce a salt plus hydrogen gas, for example zinc with dilute sulfuric acid gives zinc sulfate and hydrogen. Acids react with bases and with metal carbonates: with carbonates they produce a salt, water and carbon dioxide, which is why fizzing is seen. Acids turn blue litmus red and have a pH below 7. Bases are the chemical opposites of acids; a base that dissolves in water is called an alkali and releases hydroxide ions (OH-). Alkalis feel soapy, turn red litmus blue and have a pH above 7. Alkalis react with acids in neutralisation and react with ammonium salts on warming to release ammonia gas, which can be used to test for ammonium ions.

The pH scale and indicators

The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is. A pH of 7 is neutral, values below 7 are acidic and values above 7 are alkaline. The lower the pH, the higher the concentration of hydrogen ions; the higher the pH, the higher the concentration of hydroxide ions. Indicators are dyes that change colour depending on the pH. Litmus is a simple indicator that is red in acid and blue in alkali. Thymolphthalein is colourless in acid and blue in alkali, while methyl orange is red in acid and yellow in alkali. Universal indicator is a mixture of dyes that shows a range of colours from red in strong acid, through green at neutral, to purple in strong alkali, allowing an approximate pH value to be estimated by matching the colour to a chart.

Strong and weak acids

The strength of an acid describes how completely it ionises, or splits into ions, when dissolved in water. A strong acid is fully ionised in aqueous solution, so almost every acid molecule releases its hydrogen ion. Hydrochloric, sulfuric and nitric acids are strong acids. A weak acid is only partially ionised, so at any moment only a small fraction of the molecules have released hydrogen ions. Ethanoic acid (found in vinegar) and carbonic acid are common weak acids. It is important to separate strength from concentration: strength is about the degree of ionisation, while concentration is about how much acid is dissolved in a given volume. A dilute strong acid can still be a strong acid, and a concentrated weak acid is still a weak acid. At the same concentration, a strong acid has a lower pH than a weak acid because it produces more hydrogen ions.

Neutralisation and the ionic equation

Neutralisation is the reaction between an acid and a base that produces a salt and water. The acidic and alkaline properties cancel out, and if exactly the right amounts react the product solution is neutral. A typical word equation is hydrochloric acid plus sodium hydroxide gives sodium chloride plus water. Because the sodium and chloride ions are spectator ions that do not change during the reaction, the essential chemistry of every acid-alkali neutralisation can be written as a single ionic equation: H+ (aq) + OH- (aq) gives H2O (l). This shows that hydrogen ions from the acid combine with hydroxide ions from the alkali to form water. Neutralisation is used in everyday life, for example to treat acidic soil with lime, to relieve indigestion with antacids, and to treat acidic industrial waste before it is released.

Classifying oxides

Oxides are compounds of an element with oxygen, and they can be sorted by how they behave towards acids and alkalis. Basic oxides are metal oxides that react with acids to form a salt and water; copper(II) oxide and calcium oxide are examples, and those that dissolve in water form alkalis. Acidic oxides are usually non-metal oxides that react with alkalis to form a salt and water; carbon dioxide and sulfur dioxide are examples, and they form acidic solutions in water. Amphoteric oxides are special because they react with both acids and alkalis to form a salt and water; aluminium oxide and zinc oxide are the key examples to remember. Neutral oxides do not react with either acids or alkalis; water, carbon monoxide and nitrogen monoxide are examples. Knowing these categories helps predict reactions and choose suitable reactants when making salts.

Preparing soluble salts

Soluble salts can be made by reacting an acid with a suitable base, metal or carbonate. When the base is insoluble, the excess-solid method is used: warm acid is added to excess solid (such as copper(II) oxide or zinc), so all the acid is used up; the leftover solid is then removed by filtration, and the salt solution is evaporated partly and left to crystallise. When the base is a soluble alkali such as sodium hydroxide, the excess-solid trick cannot be used because the reactants are both solutions and nothing is left over to filter. Instead a titration is carried out: an indicator shows the exact volume of acid needed to neutralise a fixed volume of alkali. The titration is then repeated using the same volumes but without indicator, and the pure salt solution is crystallised to obtain the salt. After crystals form they are removed by filtration and dried between filter papers or in a warm oven.

Preparing insoluble salts by precipitation

Salts that do not dissolve in water cannot be crystallised from solution, so they are made by precipitation. Two soluble solutions are chosen so that mixing them brings together the ions of the insoluble salt. For example, to make insoluble lead(II) sulfate, a solution containing lead ions (such as lead(II) nitrate) is mixed with a solution containing sulfate ions (such as sodium sulfate or dilute sulfuric acid). The insoluble salt forms instantly as a fine solid called a precipitate. The mixture is filtered to collect the precipitate as the residue, the solid is washed with distilled water to remove any soluble impurities clinging to it, and it is then dried. Choosing the correct pair of soluble reactants relies on knowing solubility rules, such as that most nitrates are soluble while many sulfates, carbonates and halides of certain metals are not.

Tests for cations

Positive ions (cations) can be identified using sodium hydroxide solution and, for some, aqueous ammonia. Adding sodium hydroxide drop by drop gives coloured precipitates: copper(II) ions give a light blue precipitate, iron(II) ions give a green precipitate that slowly turns brown, and iron(III) ions give a red-brown precipitate. Aluminium, calcium and zinc ions all give white precipitates with sodium hydroxide; aluminium and zinc hydroxides dissolve in excess sodium hydroxide (because they are amphoteric) while calcium hydroxide does not, helping to tell them apart. Ammonium ions do not form a precipitate; instead, warming the substance with sodium hydroxide releases ammonia gas, which turns damp red litmus paper blue. Flame tests identify some metals by colour: lithium gives red, sodium gives yellow, potassium gives lilac, calcium gives orange-red, barium gives light green and copper(II) gives blue-green.

Tests for anions and gases

Negative ions (anions) each have a distinctive test. Carbonates fizz when dilute acid is added, releasing carbon dioxide. Chloride, bromide and iodide ions are tested by adding dilute nitric acid then silver nitrate solution: chloride gives a white precipitate, bromide a cream precipitate and iodide a yellow precipitate. Sulfate ions are tested by adding dilute nitric acid then barium nitrate solution, which gives a white precipitate. Nitrate ions are identified by adding sodium hydroxide and aluminium foil and warming, which releases ammonia gas. Common gases also have simple tests. Carbon dioxide turns limewater milky. Hydrogen gives a squeaky pop with a lighted splint. Oxygen relights a glowing splint. Ammonia turns damp red litmus paper blue. Chlorine bleaches damp litmus paper, turning it white. Sulfur dioxide turns acidified potassium manganate(VII) from purple to colourless.