C2aL14 Acids, Alkalis, and Neutralisation

Key Words

Alkali - a soluble base.

Base - a metal hydroxide or metal oxide

Excess - too much of a substance.

Eutrophication - fertilisers in waterways encouraging growth of algae.

Ionic Equation - chemical equation that describes how ionic compounds split into ions.

Neutralisation - A reaction between an acid and an alkali that results in a neutral salt.

Salt - an ionic compound formed as a result of the reaction of a metal and a non-metal.

Universal indicator - a mixture of indicators that changes colour throughout the pH range.

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Grade E

Bases are metal oxides or hydroxides.  Most bases are insoluble in water.  If they can dissolve in water, they form an alkaline solution.

If we add an acid to a alkali, there is a reaction to form a salt and water.

Acid + alkali ® salt + water

sodium hydroxide + hydrochloric acid ® sodium chloride + water

NaOH (aq) + HCl (aq) ® NaCl (aq) + H2O (l)

Since the product of this reaction is neutral, with a pH of 7, we call this a neutralisation reaction.  The alkali is neutralised by the acid.

All acids contain hydrogen (H+) ions, while all alkalis contain hydroxide ions (OH-).

Strong acids makes a lot of H+ ions.  Examples of strong acids are hydrochloric acid, sulphuric acid, and nitric acid.  Weak acids make fewer H+ ions.  Examples include citric acid and ethanoic acid (found in vinegar).

Strong acids are pH 1 and turn universal indicator red, while weak acids have a pH of about 4, and turn universal indicator orange.

Notice that the universal indicator scale is the same as the colours of the rainbow.  Aren't these chemist chappies clever?

Grade C

We can prepare a salt by neutralising a known volume of sodium hydroxide of concentration 1 mol dm-3 with the same volume of hydrochloric acid of concentration 1 mol dm-3.  We then evaporate the water, leaving sodium chloride crystals behind.

However chemists go a long way beyond this, because the neutralisation is at the heart of the chemical technique called titration.  If we know the concentration of one of the reactants, we can work out the concentration of the other.

A known volume of alkali of unknown concentration is measured into a conical flask using an accurately calibrated pipette.  Then an indicator is added, usually phenolphthalein, which turns pink when the solution is alkaline. 

Then acid of known concentration is dripped slowly into the conical flask using a graduated tube, called a burette.  This has a valve on the bottom that can be closed as soon as the reaction has finished.

As soon as the neutral point (pH 7) is reached, the phenolphthalein goes clear.  Universal indicator, which turns to green at pH 7 does not give a precise end point. 

If we know the volume of the acid, we can work out the concentration of the alkali.  If you do triple award chemistry, you will learn about the calculation.

In titrations, acid is usually added with the burette, as alkali can cause damage to the burette.

 

The diagram below gives some examples of acid and alkaline solutions:

Strong acids and strong alkalis when dissolved in water dissociate (split) completely into ions.  Weak acids (and alkalis) only partially dissociate.

HCl (aq) ®  H+ (aq) + Cl- (aq)

NaOH (aq) ® Na+ (aq) + OH - (aq)

A weak acid is NOT the same as a dilute acid.

We know that this neutralisation gives sodium chloride and water:

sodium hydroxide + hydrochloric acid ® sodium chloride + water

NaOH (aq) + HCl (aq) ® NaCl (aq) + H2O (l)

We can rewrite this as:

H+ (aq) + Cl- (aq) + Na+ (aq) + OH - (aq) ® Na+ (aq) + Cl- (aq) + H2O (l)

Notice that the sodium ions and the chloride ions have not changed.  The change that has happened is:

H+ (aq) + OH - (aq) ® H2O (l)

This is an ionic equation, and this always happens when acids and alkalis react.

All neutralisations produce a salt:

  • acid + metal ® metal salt + hydrogen;

  • acid + insoluble base ®  metal salt + water;

  • acid + soluble base ®  metal salt + water;

  • acid + ammonia ® ammonium salt.

Neutralisation reactions can be useful:

  • magnesium oxide neutralised excess hydrochloric acid in the stomach;

  • acid soils are neutralised by calcium hydroxide (slaked lime);

  • acid effluent from factories is treated with calcium hydroxide.

Grade A

Ammonia is a pungent gas that dissolves readily in water to make an alkaline solution. Dirty public lavatories often smell of ammonia.

While it could provide plants with nitrogen, the ammonia would make the soil very alkaline.  Most crops need a pH of about 6 to 7.5.

Ammonia reacts readily with nitric acid to form ammonium nitrate, a neutral salt, which is a common fertiliser:

HNO3(aq) + NH3(g) ® NH4NO3(aq)

However ammonium nitrate is a powerful oxidising agent and can decompose explosively to nitrogen and oxygen with very destructive effects.  Also it can be carried away from fields by rainwater into local watercourses.  Plants like algae thrive, and block out the light to other plants that then die off.  The lack of oxygen in the water then kills off animals.  Eutrophication processes are the commonest forms of pollution.

An alternative fertiliser is ammonium sulphate.