4.2 Reactions of acids

4.2.1 Reactions of acids with metals

- Acids react with some metals to produce salts and hydrogen.
- acid + metal → salt + hydrogen
- e.g. hydrochloric acid + magnesium → magnesium chloride + hydrogen
- These are redox reactions, as the metal is oxidised and acid is reduced.
- Magnesium, Iron, and Zinc all react with hydrochloric and sulphuric acids.

4.2.2 Neutralisation of acids and salt production

- Acids react with bases (e.g. insoluble metal hydroxides and metal oxides) to produce salts and water.
- acid + base → salt + water
- e.g. hydrochloric acid + sodium hydroxide → sodium chloride + water

- Acids also react with alkalis (e.g. soluble metal hydroxides) to produce salts and water.
- acid + alkali → salt + water
- e.g. sulphuric acid + potassium hydroxide → potassium sulphate + water

- Acids also react with metal carbonates to produce salts, water, and carbon dioxide.
- acid + metal carbonate → salt + water + carbon dioxide
- e.g. sulphuric acid + calcium carbonate → calcium sulphate + water + carbon dioxide

- All of these reactions are examples of neutralisation, as the acid is neutralised by the base, alkali, metal oxide, or metal carbonate.

The particular salt produced in any reaction between an acid and a base or alkali depends on:

• the acid used (hydrochloric acid produces chlorides, nitric acid produces nitrates, sulfuric acid produces sulfates)
• the positive ions in the base, alkali or carbonate.

More examples

- hydrochloric acid + copper oxide → copper chloride + water
- nitric acid + potassium hydroxide → potassium nitrate + water
- sulfuric acid + sodium carbonate → sodium sulfate + water + carbon dioxide

4.2.3 Soluble salts

_{(stolen from spec)}
- Soluble salts can be made from acids by reacting them with solid insoluble substances, such as metals, metal oxides, hydroxides or carbonates.
- The solid is added to the acid until no more reacts and the excess solid is filtered off to produce a solution of the salt.
- Salt solutions can be crystallised to produce solid salts.

How to make pure, dry samples

- Add the chosen acid to a beaker.
- Add the solid, insoluble reactant (metal, metal oxide, hydroxide, or carbonate) gradually to the acid, stirring until no more reacts (excess solid remains).
- Filter the mixture to remove the excess unreacted solid.
- Heat the filtrate (salt solution) gently to evaporate some water.
- Leave the solution to cool and crystallise.
- Filter off the crystals and dry them to obtain pure, dry samples of the soluble salt.

4.2.4 The pH scale and neutralisation

- Acids produce hydrogen ions (H⁺) in aqueous solution.
- Aqueous solutions of alkalis contain hydroxide ions (OH^-).
- The pH scale measures how acidic or alkaline a solution is, ranging from 0 to 14.
- A pH of 7 is neutral, below 7 is acidic, and above 7 is alkaline.
- The lower the pH, the more acidic the solution (the higher the concentration of H⁺ ions).
- The higher the pH, the more alkaline the solution (the higher the concentration of OH^- ions).
- Each change of one pH unit represents a tenfold change in H⁺ ion concentration.

Neutralisation

- In neutralisation reactions between an acid and an alkali, hydrogen ions react with hydroxide ions to produce water.
- H⁺ + OH^- → H₂O
- As the acid is neutralised, the pH of the solution increases towards 7.

Indicators

- Indicators are substances that change colour depending on the pH of the solution they are in.
- Litmus paper turns red in acidic solutions and blue in alkaline solutions.
- Methyl orange turns red in acidic solutions and yellow in alkaline solutions.
- Phenolphthalein is colourless in acidic solutions and pink in alkaline solutions.
- Universal indicator is a mixture of indicators that shows a range of colours depending on the pH of the solution.

4.2.5 Titrations

- Titrations are a quantitative method used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.
- A few drops of indicator are added to the solution being titrated.
- The titrant (known solution) is added from a burette until the endpoint is reached, indicated by a colour change.
- The volume of titrant used is measured and can be used to calculate the concentration of the unknown solution.

Steps for a titration (using strong acids and strong alkalis only)

1. Use a pipette to measure a fixed volume of the solution of unknown concentration (either a strong acid: sulfuric, hydrochloric, or nitric acid, or a strong alkali) and transfer it to a conical flask.
2. Add a few drops of a suitable indicator (e.g. methyl orange or phenolphthalein) to the conical flask.
3. Fill a burette with the solution of known concentration (the titrant: a strong acid or strong alkali) and record the initial volume.
4. Slowly add the titrant from the burette to the conical flask, swirling the flask to mix, until the indicator changes colour (the endpoint).
5. Record the final volume of the titrant in the burette and calculate the volume used.
6. Repeat the titration to obtain concordant results (two or more results within 0.1 cm³ of each other).
7. Use the average volume of titrant used to accurately calculate the concentration of the unknown strong acid or alkali.

4.2.6 Strong and weak acids

- A strong acid is completely ionised in aqueous solution.
Examples of strong acids are hydrochloric, nitric and sulfuric acids.
- A weak acid is only partially ionised in aqueous solution.
Examples of weak acids are ethanoic, citric and carbonic acids.
For a given concentration of aqueous solutions, the stronger an acid, the lower the pH.
- Strong acids have a higher concentration of H⁺ ions than weak acids.
- The pH of a strong acid is lower than the pH of a weak acid of the same concentration.
- The pH scale is logarithmic.
- As the pH decreases by one unit, the hydrogen ion concentration of the solution increases by a factor of 10.
- For example, HCl (a strong acid) has a pH of 1, and acetic acid (a weak acid) has a pH of about 3.
- This means that HCl has a hydrogen ion concentration 100 times greater than that of acetic acid.