2gAcids, bases and salt preparations

You need to know which compounds dissolve, because it decides how a salt can be made and whether a precipitate will form.

The rules

How to use them

Learn the exceptions, not the rules. Two rows have none at all — sodium/potassium/ammonium compounds, and nitrates — which makes them reliable starting points in any question.

Note that the first row overrides the others. Sodium carbonate is soluble even though most carbonates are not, because sodium compounds are always soluble.

Predicting a precipitate

When two solutions are mixed, swap the ions to find the two possible products, then check each against the rules.

Silver nitrate + sodium chloride gives silver chloride and sodium nitrate. Silver chloride is insoluble, so it appears as a precipitate; sodium nitrate is soluble and stays in solution.

Mix sodium nitrate with potassium chloride and every possible product is soluble, so nothing happens.

Write down both possible products before deciding. Guessing which one precipitates is where the marks go.

Why this matters later

The rules decide the method for making a salt:

• Soluble salt → react an acid with an insoluble base, or titrate acid with alkali.
• Insoluble salt → mix two soluble solutions and filter off the precipitate.

So getting the solubility right is the first step of every salt preparation question.

Acids and bases as proton transfer

An acid is a proton donor. A base is a proton acceptor.

This works because a hydrogen ion, H⁺, is a proton. A hydrogen atom has one proton and one electron; remove the electron and only the proton remains.

That is true of hydrogen alone — no other element's ion is a bare proton.

Base or alkali?

A base neutralises an acid. An alkali is a base that dissolves in water.

Solubility is the whole difference.

• Copper(II) oxide neutralises acids but does not dissolve → a base, not an alkali.
• Sodium hydroxide neutralises acids and dissolves → a base and an alkali.

Every alkali is a base; not every base is an alkali.

Ammonia is the base worth remembering separately, because it contains no metal at all. It accepts a proton to become the ammonium ion, NH₄⁺ — which is exactly why the proton definition is taught, since a hydroxide-based definition would miss it.

The three reactions of acids

Carbonates give three products. Forgetting the water is the standard omission.

Naming the salt

The acid decides the second half of the name:

Sulfuric acid gives sulfates, SO₄²⁻, never sulfides. Those are different ions.

One exclusion worth knowing

The specification excludes reactions between nitric acid and metals. Nitric acid is an oxidising agent and behaves differently — it reacts with copper, for instance, without producing hydrogen.

So when applying the "metal must be above hydrogen" rule, use hydrochloric or dilute sulfuric acid.

How you make a salt depends on whether your starting materials dissolve.

Method 1: acid + insoluble base

Use this when the base is insoluble — a metal oxide, hydroxide or carbonate that does not dissolve.

To make copper(II) sulfate from copper(II) oxide and dilute sulfuric acid:

1. Warm the acid, to speed the reaction up.
2. Add copper(II) oxide a little at a time, stirring, until some remains undissolved.
3. Filter to remove the excess solid.
4. Evaporate the blue filtrate to the point of crystallisation.
5. Leave to cool slowly so crystals form.
6. Dry the crystals between filter papers.

Why excess, and why filter

The excess is what guarantees all the acid has reacted. Undissolved solid remaining is the signal to stop — the blue colour appears long before the acid is used up, so stopping at the colour change leaves acid in the final crystals.

Filtering then removes the surplus, which works only because it is insoluble. An excess of a soluble base could not be removed and would contaminate the product.

Never evaporate a hydrated salt to dryness

For a salt that crystallises with water of crystallisation — copper(II) sulfate is the standard example — strong heating drives that water off, leaving a white anhydrous powder instead of blue crystals.

The warning is about hydrated salts. Not every salt has water of crystallisation to lose.

Evaporate only until crystals begin to form, then let it cool. Slow cooling gives larger, better-formed crystals.

Method 2: acid + alkali (Separate Chemistry)

When both reactants are soluble, neither can be filtered off. The trick is to do the reaction twice:

1. Titrate with an indicator to find the exact volume of acid that neutralises the alkali.
2. Repeat with the same volumes and no indicator.
3. Evaporate and crystallise as before.

The indicator is essential for finding the end point and unacceptable in the final crystals, so the measurement and the preparation are separated.

Why warming helps

Warming increases the rate, so the solid dissolves faster. It does not change the amount of product — the yield is fixed by how much acid there was.

Separate Chemistry only.

An insoluble salt cannot be made by the excess-solid method, because the product would appear as a solid immediately and could not be separated from the leftover reagent — filtering would collect both together.

Instead it is made by precipitation.

The method

To make lead(II) sulfate from lead(II) nitrate and sodium sulfate solutions:

1. Mix the two solutions.
2. A white precipitate of lead(II) sulfate forms.
3. Filter to collect it.
4. Wash the residue with distilled water.
5. Dry it in a warm oven or between filter papers.

Pb(NO₃)₂ + Na₂SO₄ → PbSO₄ + 2NaNO₃

Two sodium nitrates are needed to balance the two nitrate groups from Pb(NO₃)₂.

Why washing matters

Filtering collects the solid, but a film of the soluble by-product — sodium nitrate here — remains on its surface. Washing removes it.

Use distilled water. Tap water contains dissolved ions that would themselves contaminate the product.

Missing the washing step is the most common omission in this preparation.

Choosing the starting solutions

Three conditions:

1. Both starting materials must be soluble, so their ions are free in solution.
2. Between them they must supply the two ions of the salt you want.
3. The other product must be soluble, so it stays in solution and washes away.

That third condition is easy to forget. If both products were insoluble, the solid you collected would be a mixture.

So to make barium sulfate you need a soluble barium compound — barium chloride or nitrate — and a soluble sulfate such as sodium sulfate.

Comparing the two methods

The question to ask at the start of any salt preparation is simply what is soluble here? That determines what can be filtered, what stays dissolved, and therefore which method can work at all.

This sub-topic carries more marks than almost any other in Section 2, and the errors are consistent year to year.

1. Forgetting water from the carbonate reaction

Acid + carbonate gives three products: salt, water and carbon dioxide. The water is the one that gets left out.

2. Sulfate and sulfide

Sulfuric acid gives sulfates, SO₄²⁻. A sulfide, S²⁻, is a different ion entirely and comes from different chemistry.

3. Bracket subscripts

A number after a bracket multiplies everything inside. Ca(NO₃)₂ contains one calcium ion and two nitrate ions — so two nitrogens and six oxygens.

Examiners report this specifically, and it costs marks in formula writing, Mr calculations and equation balancing alike.

4. Stopping the salt preparation too early

The blue colour appears as soon as any copper(II) sulfate forms, long before all the acid has reacted. The signal to stop adding solid is undissolved solid remaining, not the colour change.

Stop early and unreacted acid ends up in your crystals.

5. Evaporating to dryness

This destroys the crystals by driving off the water of crystallisation. Evaporate to the point of crystallisation and let it cool.

6. Using water instead of sodium carbonate to remove acid

Water dilutes leftover acid; it does not remove or neutralise it. Where a method calls for neutralising residual acid, an alkali or carbonate is needed.

7. Charges on ions in equations

Examiners report incorrect charges as a recurring problem — writing S²⁻ where SO₄²⁻ was needed, or getting the sign wrong.

Check that the total positive charge equals the total negative charge before moving on. It takes seconds and catches most of these.