Acids, bases and salts

What acids, bases and alkalis are

An acid is a proton donor - a substance that gives away hydrogen ions, H+. A base is a proton acceptor - a substance that takes hydrogen ions. When an acid dissolves in water it releases H+ ions, so every aqueous solution of an acid contains H+ ions. This is why acidic solutions conduct electricity and turn indicators to their acidic colours.

Most bases you meet are oxides or hydroxides of metals. Copper(II) oxide, CuO, and sodium hydroxide, NaOH, are both bases. A base that dissolves in water is called an alkali, so all alkalis are bases but not all bases are alkalis. The solution of an alkali contains OH- ions, hydroxide ions, which is what makes it alkaline.

The strength of an acid or alkali depends on how many H+ or OH- ions it puts into the solution. More ions means a more acidic or more alkaline solution. This is the single idea that links indicators, the pH scale and the difference between strong and weak acids.

• Acid = proton donor, so it releases H+ ions in water.
• Base = proton acceptor, and bases are usually metal oxides or metal hydroxides.
• Alkali = a base that is soluble in water; its solution contains OH- ions.
• All alkalis are bases, but only soluble bases are alkalis.

Indicators and the pH scale

An indicator is a dye that changes colour depending on whether it is in an acidic or an alkaline solution. Litmus is red in acid and blue in alkali, and it tells you only which of the two you have. Methyl orange is red in acid and yellow in alkali. Thymolphthalein is colourless in acid and blue in alkali. The two latter indicators give a sharp colour change, which is why they are used in titrations.

Universal indicator is a mixture of dyes that gives a different colour at every pH, so it does more than tell you acid or alkali - it tells you how acidic or alkaline the solution is. The colour is matched against a pH scale running from 0 to 14. A pH below 7 is acidic, pH 7 is neutral, and a pH above 7 is alkaline.

The scale is not linear in the simple sense: each step of one pH unit represents a tenfold change in hydrogen ion concentration. A solution of pH 2 has ten times the H+ concentration of a solution of pH 3, and a hundred times that of a solution of pH 4. That is why pH 1 and pH 2 are both strongly acidic even though they look like small numbers.

Universal indicator paper is used by dipping a glass rod into the solution and touching it onto the paper, then comparing the colour with the chart. Never dip the paper into the bottle, because that would contaminate the whole supply.

The characteristic reactions of acids

Acid + metal. A dilute acid reacts with a reactive metal to give a salt and hydrogen. Magnesium fizzes steadily in dilute hydrochloric acid: Mg + 2HCl -> MgCl2 + H2. Hydrogen is identified by the squeaky pop heard when a lighted splint is put into the test tube. Copper and silver do not react with dilute acids because they are below hydrogen in the reactivity series.

Acid + base. An acid reacts with a metal oxide or metal hydroxide to give a salt and water. This is a neutralisation reaction: CuO + H2SO4 -> CuSO4 + H2O. The acid and the base cancel each other out, so the solution moves towards neutral.

Acid + carbonate. An acid reacts with a metal carbonate to give a salt, water and carbon dioxide. Calcium carbonate fizzes with dilute hydrochloric acid: CaCO3 + 2HCl -> CaCl2 + H2O + CO2. The carbon dioxide turns limewater milky, which is the standard test for it.

Acid + ammonium salt. When an alkali is warmed with an ammonium salt, ammonia gas is given off: NH4Cl + NaOH -> NaCl + H2O + NH3. Ammonia is identified by its pungent smell and by turning damp red litmus paper blue. This reaction is the basis of the test for the ammonium ion.

• Acid + metal -> salt + hydrogen. Test for hydrogen: lighted splint gives a squeaky pop.
• Acid + base -> salt + water. This is neutralisation.
• Acid + carbonate -> salt + water + carbon dioxide. Test for carbon dioxide: turns limewater milky.
• Acid or alkali + ammonium salt -> salt + water + ammonia. Test for ammonia: turns damp red litmus blue.

Strong acids and weak acids

A strong acid is one that is completely dissociated in aqueous solution - every molecule splits into ions, so almost all of the acid exists as H+ ions. A weak acid is one that is only partially dissociated, so most of it stays as whole molecules and only a small fraction releases H+ ions.

Hydrochloric acid is a strong acid. Its equation shows a single arrow, because the reaction goes fully to the right: HCl(aq) -> H+(aq) + Cl-(aq). Ethanoic acid is a weak acid, and its equation uses the equilibrium symbol to show that the dissociation is incomplete: CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq).

The consequence is the concentration of H+ ions. At the same concentration in mol/dm3, a strong acid gives a much higher concentration of H+ ions than a weak acid, so it has a lower pH, reacts faster and conducts electricity better. A weak acid still reacts in all the same ways - it simply does so more slowly and to a lesser extent.

Do not confuse strong with concentrated. Strong describes how completely the acid dissociates; concentrated describes how much acid is dissolved in the water. A dilute solution of hydrochloric acid is still a strong acid, and a concentrated solution of ethanoic acid is still a weak acid.

Neutralisation and titration

Neutralisation is the reaction between an acid and an alkali to produce a salt and water. The important ions are the hydrogen ions from the acid and the hydroxide ions from the alkali, and they combine in the ionic equation: H+(aq) + OH-(aq) -> H2O(l). The other ions are spectator ions - they are present but unchanged - which is why the same ionic equation describes every acid-alkali neutralisation.

A titration is used when the base is soluble, because there is no solid to filter off. A measured volume of alkali is pipetted into a conical flask with a few drops of indicator. Acid is added from a burette, which delivers it accurately, until the indicator changes colour. That is the end-point.

A single titration only shows you when the neutralisation is complete, not the true end-point, because the indicator is itself slightly acidic or alkaline. The method is therefore to do a rough titration, then repeat without the indicator, adding acid drop by drop until the exact volume is reached. Repeating and taking a mean of the concordant results makes the value reliable.

The volume of acid used tells you the amount of acid that reacted, and from that you can calculate the concentration of the acid or the alkali. Titration is the quantitative side of neutralisation.

The four types of oxide

Oxides are classified by how they behave towards acids and bases, and the classification follows the metallic or non-metallic character of the element. Acidic oxides are formed from non-metals and react with bases to form a salt and water. Carbon dioxide and sulfur dioxide are acidic oxides, which is why they dissolve in water to give acidic solutions and why CO2 turns limewater milky.

Basic oxides are formed from metals and react with acids to form a salt and water. Copper(II) oxide and calcium oxide are basic oxides. This is the reverse of the acid-plus-base reaction you already know: CuO + H2SO4 -> CuSO4 + H2O.

Amphoteric oxides react with both acids and bases to form a salt and water, so they behave as an acid in one reaction and as a base in the other. Aluminium oxide, Al2O3, and zinc oxide, ZnO, are the two examples to learn. With an acid, Al2O3 + 6HCl -> 2AlCl3 + 3H2O; with an alkali, Al2O3 + 2NaOH -> 2NaAlO2 + H2O.

Neutral oxides are oxides that react with neither acids nor bases. Carbon monoxide, CO, nitrogen monoxide, NO, and water itself are the common examples. Classifying an oxide is therefore a matter of asking what it reacts with, not what it looks like.

Preparing soluble salts

There are three routes to a soluble salt, and which one you use depends on whether the starting material is soluble. Acid plus excess metal: add the metal to warm dilute acid until no more will react, then filter off the unreacted metal. Acid plus excess insoluble base or insoluble carbonate: add the solid a little at a time until it stops dissolving, then filter off the excess. In all three cases the excess solid is used so that all the acid is used up, leaving a neutral solution rather than one containing leftover acid.

Once you have a clear solution of the salt, you evaporate some of the water to concentrate it, then leave it to crystallise slowly. Slow cooling gives larger, purer crystals because the ions have time to arrange themselves in a regular lattice, while any soluble impurity stays behind in the solution. The crystals are filtered off and dried between filter papers or in a warm oven. Never heat the crystals strongly, because that would drive off the water of crystallisation and leave a powder.

If the base is soluble - for example sodium hydroxide or potassium hydroxide - you cannot use the excess method, because there is no solid to filter off and the salt would be contaminated with the leftover base. Instead you must use a titration to find the exact volumes of acid and alkali that just neutralise each other, repeat without the indicator, and then evaporate the neutral solution to crystallise the salt.

For an insoluble salt such as barium sulfate or silver chloride, neither method works, because the salt would dissolve as you tried to purify it. Instead it is made by precipitation: mix solutions of two soluble salts chosen so that the required ions meet and form the insoluble product, then filter, wash and dry the precipitate. For example, BaCl2(aq) + Na2SO4(aq) -> BaSO4(s) + 2NaCl(aq).