Acids

Welcome to your study notes for the "Acids" topic (Section 3.1–3.21 of your Edexcel Chemistry course). Whether you are struggling to tell a strong acid from a concentrated one, or need a step-by-step walkthrough of the core practicals, you are in the right place. Let's break it down!

What you'll learn

How to define acids, bases, and alkalis using ions and pH.
Higher Tier: How concentration, strength, and pH are mathematically linked.
How to write equations for acid reactions and prepare pure, dry salts using three different techniques.

1. Acids, Alkalis, and Indicators

Before we can look at chemical reactions, we need to understand what is happening at the atomic level when these substances dissolve.

Definition

Acid

An acid is a substance that releases hydrogen ions ( $H+\text{H}^+$ ) when dissolved in water (aqueous solution).

For example, when hydrogen chloride gas ( $HCl\text{HCl}$ ) dissolves in water, it splits into its ions:

\text{HCl}(aq) \to \text{H}^+(aq) + \text{Cl}^-(aq)

Definition

Alkali

An alkali is a soluble base. In aqueous solution, alkalis release hydroxide ions ( $OH−\text{OH}^-$ ).

For example, sodium hydroxide ( $NaOH\text{NaOH}$ ) is a soluble ionic solid that releases hydroxide ions when dissolved in water:

\text{NaOH}(aq) \to \text{Na}^+(aq) + \text{OH}^-(aq)

Definition

Base

A base is any substance that reacts with an acid to form a salt and water only.

This means all alkalis are bases, but only soluble bases are alkalis!

Analogy

The Circle of Bases

Think of "bases" as a big family of chemical compounds. "Alkalis" are just the members of that family who can swim (they dissolve in water). So, all alkalis are bases, but not all bases are alkalis.

The pH Scale and Indicators

The pH scale tells us how acidic or alkaline a solution is.

Acidic solutions have a pH less than 7.
Neutral solutions have a pH equal to 7.
Alkaline solutions have a pH greater than 7.

To find the pH of a solution, we can use chemical indicators. An indicator is a dye that changes colour depending on the pH. You must memorise the colours of these three key indicators:

Indicator	Colour in Acid	Colour in Neutral	Colour in Alkali
Litmus	Red	Purple	Blue
Methyl Orange	Red	Yellow	Yellow
Phenolphthalein	Colourless	Colourless	Pink

Common Mistake

Confusing Indicator Colours

Students often misremember the colour of methyl orange. It is red in acid but yellow in both neutral and alkaline solutions. It only turns orange briefly at the transition point, so do not write "orange" as the alkaline colour!

2. pH and Hydrogen Ion Concentration (Higher Tier)

This section is Higher Tier only.

The pH of a solution is directly determined by the concentration of hydrogen ions ( $H+\text{H}^+$ ).

Key Idea

The pH-Ion Relationship

The higher the concentration of hydrogen ions ( $H+\text{H}^+$ ) in an acidic solution, the lower the pH.
The higher the concentration of hydroxide ions ( $OH−\text{OH}^-$ ) in an alkaline solution, the higher the pH.

How are they mathematically linked? The pH scale is logarithmic. This means every single step on the pH scale represents a tenfold (factor of 10) change in ion concentration.

Key Idea

The Factor of 10 Rule

As the hydrogen ion concentration ( $H+\text{H}^+$ ) in a solution increases by a factor of 10, the pH of the solution decreases by 1.

Let's look at this mathematically. If we change the hydrogen ion concentration by a factor of $10^x$ , the pH changes by $x$ units in the opposite direction:

If $[H+][\text{H}^+]$ increases by a factor of $10$ , pH decreases by 1.
If $[H+][\text{H}^+]$ increases by a factor of $100$ (which is $10^2$ ), pH decreases by 2.
If $[H+][\text{H}^+]$ decreases by a factor of $1000$ (which is $10^3$ ), pH increases by 3.

Example

Calculating pH changes after dilution

A sample of hydrochloric acid has a pH of 2.0. The acid is diluted with water until its volume is increased by a factor of 100. Calculate the new pH of the solution.

Identify the change in concentration: Increasing the volume of a solution by a factor of 100 dilutes it, meaning the concentration of hydrogen ions ( $H+\text{H}^+$ ) decreases by a factor of 100.
Express the concentration change as a power of 10: A decrease by a factor of 100 is equivalent to a decrease of $10^2$ .
Apply the pH rule: Since the concentration decreased by a power of $2$ , the pH must increase by 2 units.
Calculate the final pH:

\text{New pH} = 2.0 + 2 = 4.0

3. Strong vs. Weak, Dilute vs. Concentrated (Higher Tier)

This section is Higher Tier only.

It is extremely common to confuse "strong" with "concentrated" in everyday language, but in GCSE Chemistry they mean completely different things.

Strong vs. Weak Acids

This describes how easily the acid molecules split apart into ions when dissolved in water. This splitting process is called dissociation (or ionisation).

Definition

Strong Acid

An acid that completely dissociates (splits up) into hydrogen ions ( $H+\text{H}^+$ ) in aqueous solution.

Examples of strong acids include hydrochloric acid ( $HCl\text{HCl}$ ), sulfuric acid ( $H2SO4\text{H}_2\text{SO}_4$ ), and nitric acid ( $HNO3\text{HNO}_3$ ).

\text{HCl}(aq) \to \text{H}^+(aq) + \text{Cl}^-(aq)

(Note the single irreversible arrow $→\to$ , showing all reactant molecules turn into ions.)

Definition

Weak Acid

An acid that only partially dissociates into hydrogen ions ( $H+\text{H}^+$ ) in aqueous solution.

Examples of weak acids include ethanoic acid ( $CH3COOH\text{CH}_3\text{COOH}$ ) and citric acid.

\text{CH}_3\text{COOH}(aq) \rightleftharpoons \text{CH}_3\text{COO}^-(aq) + \text{H}^+(aq)

(Note the reversible arrow $⇌\rightleftharpoons$ , showing that the reaction is in equilibrium and most molecules remain unsplit.)

Dilute vs. Concentrated

This describes the actual amount of the acid solute dissolved in a specific volume of water, regardless of whether it has dissociated or not.

Concentrated solution: Contains a large amount of solute dissolved in a given volume of solution.
Dilute solution: Contains a small amount of solute dissolved in a given volume of solution.

Common Mistake

Strength vs Concentration

You can have a concentrated weak acid (lots of ethanoic acid molecules dissolved in a small volume of water, but most of them remain unsplit) or a dilute strong acid (very little hydrochloric acid dissolved in water, but every single molecule of it has split into ions!).

4. Reactions of Acids and Chemical Tests

Acids react with different types of bases to form ionic compounds called salts. You must learn these general equations, including their state symbols!

The General Reactions

Acid + Metal $→\to$ Salt + Hydrogen
- Example: $Mg(s)+2HCl(aq)→MgCl2(aq)+H2(g)\text{Mg}(s) + 2\text{HCl}(aq) \to \text{MgCl}_2(aq) + \text{H}_2(g)$
Acid + Metal Oxide (Base) $→\to$ Salt + Water
- Example: $CuO(s)+H2SO4(aq)→CuSO4(aq)+H2O(l)\text{CuO}(s) + \text{H}_2\text{SO}_4(aq) \to \text{CuSO}_4(aq) + \text{H}_2\text{O}(l)$
Acid + Metal Hydroxide (Base/Alkali) $→\to$ Salt + Water
- Example: $HCl(aq)+NaOH(aq)→NaCl(aq)+H2O(l)\text{HCl}(aq) + \text{NaOH}(aq) \to \text{NaCl}(aq) + \text{H}_2\text{O}(l)$
Acid + Metal Carbonate $→\to$ Salt + Water + Carbon Dioxide
- Example: $CaCO3(s)+2HCl(aq)→CaCl2(aq)+H2O(l)+CO2(g)\text{CaCO}_3(s) + 2\text{HCl}(aq) \to \text{CaCl}_2(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$

Neutralisation

A reaction between an acid and a base is called a neutralisation reaction. If the base is an alkali (a soluble base), the key ionic reaction happening in solution is the combination of hydrogen ions and hydroxide ions to form water:

\text{H}^+(aq) + \text{OH}^-(aq) \to \text{H}_2\text{O}(l)

Key Gas Tests

When acids react with metals or carbonates, gases are produced. You must know how to test for them:

Test for Hydrogen ( $H2\text{H}_2$ ): Hold a lighted splint near the mouth of the test tube. Hydrogen burns rapidly with a squeaky pop sound.
Test for Carbon Dioxide ( $CO2\text{CO}_2$ ): Bubble the gas through limewater (calcium hydroxide solution). The limewater turns cloudy (or milky) due to the formation of an insoluble white precipitate of calcium carbonate.

5. Core Practical: Investigating pH Changes (3.6)

In this practical, you investigate how pH changes when you gradually neutralise an acid. Specifically, you add an insoluble base (powdered calcium hydroxide, $Ca(OH)2\text{Ca(OH)}_2$ ) to a fixed volume of dilute hydrochloric acid ( $HCl\text{HCl}$ ).

The Method

Measure $cm3150\text{ cm}^3$ of dilute hydrochloric acid into a beaker.
Measure and record the initial pH of the acid using a pH probe or universal indicator paper.
Weigh out $g0.3\text{ g}$ of calcium hydroxide powder.
Add the powder to the beaker, stir thoroughly, and record the new pH once the reaction is complete.
Repeat steps 3 and 4 multiple times until the pH no longer changes (showing the base is in excess and the solution is fully neutralised).
Plot a graph of pH (y-axis) against the total mass of calcium hydroxide added (x-axis).

6. Solubility Rules and Preparing Salts

To prepare a pure, dry sample of a salt, you must first check whether the salt you want to make is soluble or insoluble in water.

The Solubility Rules

You must memorise these rules:

Compound Type	Solubility Rules
Sodium, Potassium, Ammonium	All common salts are soluble
Nitrates	All are soluble
Chlorides	Common chlorides are soluble (except silver chloride and lead chloride)
Sulfates	Common sulfates are soluble (except lead sulfate, barium sulfate, and calcium sulfate)
Carbonates & Hydroxides	Common ones are insoluble (except sodium, potassium, and ammonium ones)

Method A: Preparing a Soluble Salt from an Insoluble Reactant (Core Practical 3.17)

Use this method when making a soluble salt (like copper sulfate) from an acid and an insoluble base (like copper oxide).

Preparation of copper sulfate crystals

Why do we do these steps?

Add the insoluble reactant in excess: This ensures that all of the acid reacts completely and is used up, so your final salt isn't contaminated with unreacted acid.
Filter the mixture: To remove the leftover, unreacted excess solid base.
Heat gently to evaporate some water: Heat the filtrate in an evaporating basin over a water bath until crystals just begin to form at the edges (the crystallisation point).
Leave to crystallise: Let it cool slowly so large, pure hydrated crystals form. Filter the crystals, rinse with cold distilled water, and pat dry with filter paper.

Method B: Preparing a Soluble Salt from a Soluble Reactant (Titration)

If both your acid and your base are soluble (e.g., reacting hydrochloric acid with sodium hydroxide), you cannot simply add "excess base" because it would dissolve in the water, contaminating your product! Instead, you must use a titration to find the exact proportions needed.

Titration setup

The Method

Use a volumetric pipette to measure an exact volume (usually $cm325.0\text{ cm}^3$ ) of the alkali into a conical flask. Add a few drops of indicator (e.g., phenolphthalein).
Fill a burette with the acid and note the starting volume.
Place the conical flask on a white tile so you can see the colour change clearly.
Slowly run the acid from the burette into the flask, swirling constantly, until the indicator changes colour permanently (the end point). Record the final volume of acid added.
Repeat the experiment without any indicator, adding the exact volume of acid determined in step 4 to the same volume of alkali. This leaves you with only salt and water in the flask.
Heat the solution gently in an evaporating basin to crystallise the salt.

Method C: Preparing an Insoluble Salt (Precipitation)

When two soluble solutions react to form an insoluble solid, that solid is called a precipitate. You can prepare a pure, dry sample of an insoluble salt by mixing two soluble salt solutions.

Definition

Precipitation Reaction

A reaction in which soluble substances react to form an insoluble product (a precipitate).

The Method

Mix the two soluble salt solutions together in a beaker to form the precipitate.
Filter the mixture; the insoluble precipitate will be trapped as the residue on the filter paper.
Wash the residue thoroughly with distilled water while still in the filter paper to wash away any traces of the soluble reaction products.
Leave the pure residue to dry in a warm oven or pat dry between dry filter papers.

Example

Predicting a precipitate using solubility rules

Predict whether a precipitate will form when aqueous solutions of silver nitrate ( $AgNO3\text{AgNO}_3$ ) and sodium chloride ( $NaCl\text{NaCl}$ ) are mixed. If a precipitate forms, name it.

Identify all potential products: Swap the cations and anions of the reactants.
- Silver ( $Ag+\text{Ag}^+$ ) combines with chloride ( $Cl−\text{Cl}^-$ ) to form silver chloride ( $AgCl\text{AgCl}$ ).
- Sodium ( $Na+\text{Na}^+$ ) combines with nitrate ( $NO3−\text{NO}_3^-$ ) to form sodium nitrate ( $NaNO3\text{NaNO}_3$ ).
Apply the solubility rules to the first product: Rule 3 states that common chlorides are soluble, except silver chloride and lead chloride. Therefore, silver chloride ( $AgCl\text{AgCl}$ ) is insoluble.
Apply the solubility rules to the second product: Rule 1 and Rule 2 state that all sodium salts and all nitrates are soluble. Therefore, sodium nitrate ( $NaNO3\text{NaNO}_3$ ) is soluble.
Formulate the final prediction: Since silver chloride is insoluble, a precipitate will indeed form. The name of the precipitate is silver chloride.

\text{AgNO}_3(aq) + \text{NaCl}(aq) \to \text{AgCl}(s) + \text{NaNO}_3(aq)

Exam technique

In the exam

When asked how to prepare a salt, always look at your solubility rules first. Your very first step in the exam should be to decide: Is the salt I am trying to make soluble or insoluble? This determines whether you use the excess base method, a titration, or precipitation.
Do not use a Bunsen burner directly to evaporate all of the water from hydrated copper sulfate crystals. If you heat them to dryness, they lose their water of crystallisation and turn into a dull white, anhydrous powder. Always use a water bath and heat gently until the crystallisation point is reached.
If writing a neutralisation equation, remember that the ionic equation is always $H+(aq)+OH−(aq)→H2O(l)\text{H}^+(aq) + \text{OH}^-(aq) \to \text{H}_2\text{O}(l)$ , provided it is a reaction between an acid and a soluble alkali.

Self review

Check yourself

State the difference between a strong acid and a concentrated acid.
Write the balanced chemical equation, including state symbols, for the reaction between calcium carbonate and hydrochloric acid.
Describe how you would test for the gas produced when magnesium ribbon is added to sulfuric acid.
Explain why we add an excess of copper oxide when preparing copper sulfate crystals, and how we ensure no excess copper oxide remains in our final crystal sample.

Acids

What you'll learn

How to define acids, bases, and alkalis using ions and pH.
Higher Tier: How concentration, strength, and pH are mathematically linked.
How to write equations for acid reactions and prepare pure, dry salts using three different techniques.

1. Acids, Alkalis, and Indicators

Before we can look at chemical reactions, we need to understand what is happening at the atomic level when these substances dissolve.

Definition

Acid

An acid is a substance that releases hydrogen ions ( $H+\text{H}^+$ ) when dissolved in water (aqueous solution).

For example, when hydrogen chloride gas ( $HCl\text{HCl}$ ) dissolves in water, it splits into its ions:

\text{HCl}(aq) \to \text{H}^+(aq) + \text{Cl}^-(aq)

Definition

Alkali

An alkali is a soluble base. In aqueous solution, alkalis release hydroxide ions ( $OH−\text{OH}^-$ ).

For example, sodium hydroxide ( $NaOH\text{NaOH}$ ) is a soluble ionic solid that releases hydroxide ions when dissolved in water:

\text{NaOH}(aq) \to \text{Na}^+(aq) + \text{OH}^-(aq)

Definition

Base

A base is any substance that reacts with an acid to form a salt and water only.

This means all alkalis are bases, but only soluble bases are alkalis!

Analogy

The Circle of Bases

The pH Scale and Indicators

The pH scale tells us how acidic or alkaline a solution is.

Acidic solutions have a pH less than 7.
Neutral solutions have a pH equal to 7.
Alkaline solutions have a pH greater than 7.

To find the pH of a solution, we can use chemical indicators. An indicator is a dye that changes colour depending on the pH. You must memorise the colours of these three key indicators:

Indicator	Colour in Acid	Colour in Neutral	Colour in Alkali
Litmus	Red	Purple	Blue
Methyl Orange	Red	Yellow	Yellow
Phenolphthalein	Colourless	Colourless	Pink

Common Mistake

Confusing Indicator Colours

2. pH and Hydrogen Ion Concentration (Higher Tier)

This section is Higher Tier only.

The pH of a solution is directly determined by the concentration of hydrogen ions ( $H+\text{H}^+$ ).

Key Idea

The pH-Ion Relationship

The higher the concentration of hydrogen ions ( $H+\text{H}^+$ ) in an acidic solution, the lower the pH.
The higher the concentration of hydroxide ions ( $OH−\text{OH}^-$ ) in an alkaline solution, the higher the pH.

How are they mathematically linked? The pH scale is logarithmic. This means every single step on the pH scale represents a tenfold (factor of 10) change in ion concentration.

Key Idea

The Factor of 10 Rule

As the hydrogen ion concentration ( $H+\text{H}^+$ ) in a solution increases by a factor of 10, the pH of the solution decreases by 1.

Let's look at this mathematically. If we change the hydrogen ion concentration by a factor of $10^x$ , the pH changes by $x$ units in the opposite direction:

If $[H+][\text{H}^+]$ increases by a factor of $10$ , pH decreases by 1.
If $[H+][\text{H}^+]$ increases by a factor of $100$ (which is $10^2$ ), pH decreases by 2.
If $[H+][\text{H}^+]$ decreases by a factor of $1000$ (which is $10^3$ ), pH increases by 3.

Example

Calculating pH changes after dilution

A sample of hydrochloric acid has a pH of 2.0. The acid is diluted with water until its volume is increased by a factor of 100. Calculate the new pH of the solution.

Identify the change in concentration: Increasing the volume of a solution by a factor of 100 dilutes it, meaning the concentration of hydrogen ions ( $H+\text{H}^+$ ) decreases by a factor of 100.
Express the concentration change as a power of 10: A decrease by a factor of 100 is equivalent to a decrease of $10^2$ .
Apply the pH rule: Since the concentration decreased by a power of $2$ , the pH must increase by 2 units.
Calculate the final pH:

\text{New pH} = 2.0 + 2 = 4.0

3. Strong vs. Weak, Dilute vs. Concentrated (Higher Tier)

This section is Higher Tier only.

It is extremely common to confuse "strong" with "concentrated" in everyday language, but in GCSE Chemistry they mean completely different things.

Strong vs. Weak Acids

This describes how easily the acid molecules split apart into ions when dissolved in water. This splitting process is called dissociation (or ionisation).

Definition

Strong Acid

An acid that completely dissociates (splits up) into hydrogen ions ( $H+\text{H}^+$ ) in aqueous solution.

Examples of strong acids include hydrochloric acid ( $HCl\text{HCl}$ ), sulfuric acid ( $H2SO4\text{H}_2\text{SO}_4$ ), and nitric acid ( $HNO3\text{HNO}_3$ ).

\text{HCl}(aq) \to \text{H}^+(aq) + \text{Cl}^-(aq)

(Note the single irreversible arrow $→\to$ , showing all reactant molecules turn into ions.)

Definition

Weak Acid

An acid that only partially dissociates into hydrogen ions ( $H+\text{H}^+$ ) in aqueous solution.

Examples of weak acids include ethanoic acid ( $CH3COOH\text{CH}_3\text{COOH}$ ) and citric acid.

\text{CH}_3\text{COOH}(aq) \rightleftharpoons \text{CH}_3\text{COO}^-(aq) + \text{H}^+(aq)

(Note the reversible arrow $⇌\rightleftharpoons$ , showing that the reaction is in equilibrium and most molecules remain unsplit.)

Dilute vs. Concentrated

This describes the actual amount of the acid solute dissolved in a specific volume of water, regardless of whether it has dissociated or not.

Concentrated solution: Contains a large amount of solute dissolved in a given volume of solution.
Dilute solution: Contains a small amount of solute dissolved in a given volume of solution.

Common Mistake

Strength vs Concentration

4. Reactions of Acids and Chemical Tests

Acids react with different types of bases to form ionic compounds called salts. You must learn these general equations, including their state symbols!

The General Reactions

Acid + Metal $→\to$ Salt + Hydrogen
- Example: $Mg(s)+2HCl(aq)→MgCl2(aq)+H2(g)\text{Mg}(s) + 2\text{HCl}(aq) \to \text{MgCl}_2(aq) + \text{H}_2(g)$
Acid + Metal Oxide (Base) $→\to$ Salt + Water
- Example: $CuO(s)+H2SO4(aq)→CuSO4(aq)+H2O(l)\text{CuO}(s) + \text{H}_2\text{SO}_4(aq) \to \text{CuSO}_4(aq) + \text{H}_2\text{O}(l)$
Acid + Metal Hydroxide (Base/Alkali) $→\to$ Salt + Water
- Example: $HCl(aq)+NaOH(aq)→NaCl(aq)+H2O(l)\text{HCl}(aq) + \text{NaOH}(aq) \to \text{NaCl}(aq) + \text{H}_2\text{O}(l)$
Acid + Metal Carbonate $→\to$ Salt + Water + Carbon Dioxide
- Example: $CaCO3(s)+2HCl(aq)→CaCl2(aq)+H2O(l)+CO2(g)\text{CaCO}_3(s) + 2\text{HCl}(aq) \to \text{CaCl}_2(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$

Neutralisation

\text{H}^+(aq) + \text{OH}^-(aq) \to \text{H}_2\text{O}(l)

Key Gas Tests

When acids react with metals or carbonates, gases are produced. You must know how to test for them:

Test for Hydrogen ( $H2\text{H}_2$ ): Hold a lighted splint near the mouth of the test tube. Hydrogen burns rapidly with a squeaky pop sound.
Test for Carbon Dioxide ( $CO2\text{CO}_2$ ): Bubble the gas through limewater (calcium hydroxide solution). The limewater turns cloudy (or milky) due to the formation of an insoluble white precipitate of calcium carbonate.

5. Core Practical: Investigating pH Changes (3.6)

The Method

Measure $cm3150\text{ cm}^3$ of dilute hydrochloric acid into a beaker.
Measure and record the initial pH of the acid using a pH probe or universal indicator paper.
Weigh out $g0.3\text{ g}$ of calcium hydroxide powder.
Add the powder to the beaker, stir thoroughly, and record the new pH once the reaction is complete.
Repeat steps 3 and 4 multiple times until the pH no longer changes (showing the base is in excess and the solution is fully neutralised).
Plot a graph of pH (y-axis) against the total mass of calcium hydroxide added (x-axis).

6. Solubility Rules and Preparing Salts

To prepare a pure, dry sample of a salt, you must first check whether the salt you want to make is soluble or insoluble in water.

The Solubility Rules

You must memorise these rules:

Compound Type	Solubility Rules
Sodium, Potassium, Ammonium	All common salts are soluble
Nitrates	All are soluble
Chlorides	Common chlorides are soluble (except silver chloride and lead chloride)
Sulfates	Common sulfates are soluble (except lead sulfate, barium sulfate, and calcium sulfate)
Carbonates & Hydroxides	Common ones are insoluble (except sodium, potassium, and ammonium ones)

Method A: Preparing a Soluble Salt from an Insoluble Reactant (Core Practical 3.17)

Use this method when making a soluble salt (like copper sulfate) from an acid and an insoluble base (like copper oxide).

Preparation of copper sulfate crystals

Why do we do these steps?

Add the insoluble reactant in excess: This ensures that all of the acid reacts completely and is used up, so your final salt isn't contaminated with unreacted acid.
Filter the mixture: To remove the leftover, unreacted excess solid base.
Heat gently to evaporate some water: Heat the filtrate in an evaporating basin over a water bath until crystals just begin to form at the edges (the crystallisation point).
Leave to crystallise: Let it cool slowly so large, pure hydrated crystals form. Filter the crystals, rinse with cold distilled water, and pat dry with filter paper.

Method B: Preparing a Soluble Salt from a Soluble Reactant (Titration)

Titration setup

The Method

Use a volumetric pipette to measure an exact volume (usually $cm325.0\text{ cm}^3$ ) of the alkali into a conical flask. Add a few drops of indicator (e.g., phenolphthalein).
Fill a burette with the acid and note the starting volume.
Place the conical flask on a white tile so you can see the colour change clearly.
Slowly run the acid from the burette into the flask, swirling constantly, until the indicator changes colour permanently (the end point). Record the final volume of acid added.
Repeat the experiment without any indicator, adding the exact volume of acid determined in step 4 to the same volume of alkali. This leaves you with only salt and water in the flask.
Heat the solution gently in an evaporating basin to crystallise the salt.

Method C: Preparing an Insoluble Salt (Precipitation)

When two soluble solutions react to form an insoluble solid, that solid is called a precipitate. You can prepare a pure, dry sample of an insoluble salt by mixing two soluble salt solutions.

Definition

Precipitation Reaction

A reaction in which soluble substances react to form an insoluble product (a precipitate).

The Method

Mix the two soluble salt solutions together in a beaker to form the precipitate.
Filter the mixture; the insoluble precipitate will be trapped as the residue on the filter paper.
Wash the residue thoroughly with distilled water while still in the filter paper to wash away any traces of the soluble reaction products.
Leave the pure residue to dry in a warm oven or pat dry between dry filter papers.

Example

Predicting a precipitate using solubility rules

Predict whether a precipitate will form when aqueous solutions of silver nitrate ( $AgNO3\text{AgNO}_3$ ) and sodium chloride ( $NaCl\text{NaCl}$ ) are mixed. If a precipitate forms, name it.

Identify all potential products: Swap the cations and anions of the reactants.
- Silver ( $Ag+\text{Ag}^+$ ) combines with chloride ( $Cl−\text{Cl}^-$ ) to form silver chloride ( $AgCl\text{AgCl}$ ).
- Sodium ( $Na+\text{Na}^+$ ) combines with nitrate ( $NO3−\text{NO}_3^-$ ) to form sodium nitrate ( $NaNO3\text{NaNO}_3$ ).
Apply the solubility rules to the first product: Rule 3 states that common chlorides are soluble, except silver chloride and lead chloride. Therefore, silver chloride ( $AgCl\text{AgCl}$ ) is insoluble.
Apply the solubility rules to the second product: Rule 1 and Rule 2 state that all sodium salts and all nitrates are soluble. Therefore, sodium nitrate ( $NaNO3\text{NaNO}_3$ ) is soluble.
Formulate the final prediction: Since silver chloride is insoluble, a precipitate will indeed form. The name of the precipitate is silver chloride.

\text{AgNO}_3(aq) + \text{NaCl}(aq) \to \text{AgCl}(s) + \text{NaNO}_3(aq)

Exam technique

In the exam

When asked how to prepare a salt, always look at your solubility rules first. Your very first step in the exam should be to decide: Is the salt I am trying to make soluble or insoluble? This determines whether you use the excess base method, a titration, or precipitation.
Do not use a Bunsen burner directly to evaporate all of the water from hydrated copper sulfate crystals. If you heat them to dryness, they lose their water of crystallisation and turn into a dull white, anhydrous powder. Always use a water bath and heat gently until the crystallisation point is reached.
If writing a neutralisation equation, remember that the ionic equation is always $H+(aq)+OH−(aq)→H2O(l)\text{H}^+(aq) + \text{OH}^-(aq) \to \text{H}_2\text{O}(l)$ , provided it is a reaction between an acid and a soluble alkali.

Self review

Check yourself

State the difference between a strong acid and a concentrated acid.
Write the balanced chemical equation, including state symbols, for the reaction between calcium carbonate and hydrochloric acid.
Describe how you would test for the gas produced when magnesium ribbon is added to sulfuric acid.
Explain why we add an excess of copper oxide when preparing copper sulfate crystals, and how we ensure no excess copper oxide remains in our final crystal sample.

Acids

What you'll learn

1. Acids, Alkalis, and Indicators

The pH Scale and Indicators

2. pH and Hydrogen Ion Concentration (Higher Tier)

3. Strong vs. Weak, Dilute vs. Concentrated (Higher Tier)

Strong vs. Weak Acids

Dilute vs. Concentrated

4. Reactions of Acids and Chemical Tests

The General Reactions

Neutralisation

Key Gas Tests

5. Core Practical: Investigating pH Changes (3.6)

The Method

6. Solubility Rules and Preparing Salts

The Solubility Rules

Method A: Preparing a Soluble Salt from an Insoluble Reactant (Core Practical 3.17)

Why do we do these steps?

Method B: Preparing a Soluble Salt from a Soluble Reactant (Titration)

The Method

Method C: Preparing an Insoluble Salt (Precipitation)

The Method

Recap questions

You've reached the end

Acids Revision Guide

Acids

What you'll learn

1. Acids, Alkalis, and Indicators

The pH Scale and Indicators

2. pH and Hydrogen Ion Concentration (Higher Tier)

3. Strong vs. Weak, Dilute vs. Concentrated (Higher Tier)

Strong vs. Weak Acids

Dilute vs. Concentrated

4. Reactions of Acids and Chemical Tests

The General Reactions

Neutralisation

Key Gas Tests

5. Core Practical: Investigating pH Changes (3.6)

The Method

6. Solubility Rules and Preparing Salts

The Solubility Rules

Method A: Preparing a Soluble Salt from an Insoluble Reactant (Core Practical 3.17)

Why do we do these steps?

Method B: Preparing a Soluble Salt from a Soluble Reactant (Titration)

The Method

Method C: Preparing an Insoluble Salt (Precipitation)

The Method

Recap questions

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Acids Revision Guide