Reactions of Acids and Making Salts

Acids do far more than turn an indicator red. They react with metals, with bases and with carbonates, and in every one of these reactions they produce a salt. Knowing the three reactions of acids — and being able to name the salt and make a pure sample of it in the lab — is a core part of Topic C3 of OCR Gateway Combined Science. This lesson sets out the three reactions, shows how to build the name of a salt from its acid, and describes the required practical for making a pure, dry salt.

By the end of this lesson you should be able to write general equations for the three reactions of acids, name the salt produced, describe the gas tests involved, and describe the method for making a pure, dry sample of a soluble salt.

This lesson builds AO1 recall of the three reactions of acids and their gas tests, strong AO2 application through the required practical for making a pure, dry salt and naming the salt from its acid, and AO3 evaluation when you judge how to obtain a pure sample and why each step is needed.

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The Three Reactions of Acids

Acids react in three characteristic ways, and you should learn each general equation:

1. Acid + metal → salt + hydrogen

A reactive metal (such as magnesium or zinc) reacts with an acid to give a salt and hydrogen gas:

$\text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}_2 + \text{H}_2$Mg+2HCl→MgCl2​+H2​

You see bubbles of hydrogen, and the metal slowly disappears.

2. Acid + base → salt + water

A base is a metal oxide or metal hydroxide. It neutralises the acid to give a salt and water (no gas):

$\text{CuO} + \text{H}_2\text{SO}_4 \rightarrow \text{CuSO}_4 + \text{H}_2\text{O}$CuO+H2​SO4​→CuSO4​+H2​O $\text{NaOH} + \text{HCl} \rightarrow \text{NaCl} + \text{H}_2\text{O}$NaOH+HCl→NaCl+H2​O

This is the neutralisation reaction from the last lesson, now with the salt written in.

3. Acid + metal carbonate → salt + water + carbon dioxide

A metal carbonate reacts with an acid to give a salt, water and carbon dioxide:

$\text{CaCO}_3 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{O} + \text{CO}_2$CaCO3​+2HCl→CaCl2​+H2​O+CO2​

You see fizzing as the carbon dioxide is given off.

Reaction	Products	Tell-tale sign
Acid + metal	salt + hydrogen	bubbles of gas; metal dissolves
Acid + base (oxide/hydroxide)	salt + water	no gas; solid dissolves
Acid + metal carbonate	salt + water + carbon dioxide	fizzing (CO₂)

Exam Tip: Memorise the three "endings": metal → hydrogen, base → water (only), carbonate → water + carbon dioxide. Whether a gas is given off — and which gas — is the quickest way to tell which reaction has happened.

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Testing for the Gases

Two gas tests turn up with these reactions:

• Hydrogen (from acid + metal): a lighted splint held at the mouth of the tube gives a squeaky pop.
• Carbon dioxide (from acid + carbonate): bubble the gas through limewater (calcium hydroxide solution) and it turns milky (cloudy white).

Exam Tip: "Limewater turns milky/cloudy" is the test for carbon dioxide; the "squeaky pop with a lighted splint" is the test for hydrogen. Quote the observation, not just the name of the test. A common misconception is to mix these up, so keep them firmly apart.

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Naming the Salt

A salt's name has two parts: the metal part comes from the metal, base or carbonate, and the second part comes from the acid used. The acid decides the ending:

Acid	Salt ending	Example salt
Hydrochloric acid (HCl)	chloride	sodium chloride, $\text{NaCl}$NaCl
Sulfuric acid ($\text{H}_2\text{SO}_4$H2​SO4​)	sulfate	copper sulfate, $\text{CuSO}_4$CuSO4​
Nitric acid ($\text{HNO}_3$HNO3​)	nitrate	potassium nitrate, $\text{KNO}_3$KNO3​

So magnesium reacting with hydrochloric acid gives magnesium chloride; copper oxide reacting with sulfuric acid gives copper sulfate; and a carbonate reacting with nitric acid gives a nitrate.

Exam Tip: Build the salt name in two steps: metal first (from the metal/base/carbonate), then the ending from the acid (chloride / sulfate / nitrate). For example, zinc + sulfuric acid → zinc sulfate.

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Required Practical: Making a Pure, Dry Salt

A key required practical is to make a pure, dry sample of a soluble salt from an acid and an insoluble base (a metal oxide or carbonate that does not dissolve). Using an insoluble base is the trick that lets you remove the excess easily by filtering. A typical example is making copper sulfate from sulfuric acid and copper oxide:

$\text{CuO} + \text{H}_2\text{SO}_4 \rightarrow \text{CuSO}_4 + \text{H}_2\text{O}$CuO+H2​SO4​→CuSO4​+H2​O

Method (numbered):

1. Measure a fixed volume of dilute sulfuric acid into a beaker and warm it gently (this speeds up the reaction). Do not boil it.
2. Add the insoluble base, copper oxide, a little at a time, stirring, until no more dissolves and some solid copper oxide remains. The copper oxide is now in excess, which makes sure all the acid has reacted.
3. Filter the mixture to remove the excess copper oxide, leaving a clear blue solution of copper sulfate (the filtrate).
4. Pour the copper sulfate solution into an evaporating basin and heat it gently to evaporate about half the water, leaving a saturated solution (the "crystallisation point").
5. Leave the solution to cool and crystallise slowly. Crystals of copper sulfate form as the water evaporates.
6. Remove the crystals and dry them by patting between filter paper (or leaving in a warm place).

Adding the base in excess and then filtering it off is the central idea: it guarantees all the acid is used up (so the salt is not contaminated with acid), and the leftover solid is easy to remove because it never dissolved.

Exam Tip: The two marks examiners most often award here are: (1) add the base in excess so all the acid reacts, and (2) filter to remove the unreacted (excess) base. Then crystallise — evaporate some water and let it cool — to get pure, dry crystals. Do not boil the solution dry, which spoils the crystals.

Why this method needs an insoluble base

The whole method depends on the base being insoluble. If you used a soluble base (an alkali such as sodium hydroxide), you could not see when you had added enough, because it would simply dissolve — there would be no leftover solid to signal that the acid was all used up, and you could not filter any excess out. With an insoluble base, the moment some solid stops dissolving you know the acid has all reacted, and that leftover solid is trivial to remove by filtering. This is why soluble salts of metals such as copper, zinc and magnesium are made from their insoluble oxides or carbonates, not from an alkali.

Making a salt from a soluble base (titration method)

For a salt of a very reactive metal — for example sodium chloride from sodium hydroxide — the base is soluble, so the excess cannot be filtered off. Instead a titration is used to find the exact volume of acid that just neutralises a measured volume of the alkali:

1. Use a pipette to measure a fixed volume of the alkali into a conical flask and add a few drops of indicator (phenolphthalein is ideal — pink in alkali, colourless in acid).
2. Add acid from a burette, swirling, until the indicator just changes colour, showing exact neutralisation. Record this volume.
3. Repeat without indicator, adding exactly that recorded volume of acid to the same volume of fresh alkali, so the salt is not contaminated by the indicator.
4. Crystallise the resulting salt solution as before (evaporate some water, cool, dry).

Type of base	Soluble?	Method to make the salt
Metal oxide / carbonate (e.g. CuO, $\text{CaCO}_3$CaCO3​)	Insoluble	Add in excess, filter, crystallise
Alkali (e.g. NaOH, KOH)	Soluble	Titration to find the exact volume, then crystallise

Exam Tip: Choose the method from the base: an insoluble base → "add to excess then filter"; a soluble base (alkali) → titration with an indicator. Both finish with crystallisation to get a pure, dry salt.

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Soluble and Insoluble Salts

Whether a salt is soluble also affects how it is made and whether a precipitate forms in a reaction. The general solubility rules for GCSE are worth knowing:

Salts	Solubility
All sodium, potassium and ammonium salts	Soluble
All nitrates	Soluble
Most chlorides	Soluble (except silver chloride and lead chloride)
Most sulfates	Soluble (except barium, calcium and lead sulfate)
Most carbonates and hydroxides	Insoluble (except those of sodium, potassium and ammonium)