Acids and Bases

Acids, Bases and Alkalis: What Is the Difference?

Before learning any reactions, the three terms need to be separated carefully. They describe related ideas, but they do not mean the same thing.

What is an acid?

An acid is a substance that produces H⁺ ions when dissolved in water. We therefore describe an acidic solution in terms of the H⁺ ions present in the aqueous solution.

Common acids encountered at O-Level include:

• hydrochloric acid, HCl
• sulfuric acid, H₂SO₄
• nitric acid, HNO₃
• ethanoic acid, CH₃COOH

What is a base?

A base is a substance that neutralises an acid. Many bases are metal oxides, metal hydroxides or metal carbonates.

Examples include:

• copper(II) oxide, CuO
• magnesium oxide, MgO
• sodium hydroxide, NaOH
• calcium carbonate, CaCO₃

Notice that a base does not have to contain an –OH group. Copper(II) oxide, CuO, contains no hydroxide ions in its formula, yet it is a base because it neutralises acids.

What is an alkali?

An alkali is a soluble base. When it dissolves in water, it produces OH⁻ ions in the aqueous solution.

Examples include:

• sodium hydroxide solution, NaOH(aq)
• potassium hydroxide solution, KOH(aq)
• calcium hydroxide solution, Ca(OH)₂(aq)

Indicators and the pH Scale

Definitions tell us what an acid or alkali is, but in the laboratory we need a safe way to identify them. We do not taste or touch an unknown solution. Instead, we use indicators or measure its pH.

Litmus tells us whether a solution is acidic or alkaline

• An acid turns blue litmus paper red.
• An alkali turns red litmus paper blue.
• A neutral solution produces no such change.

Litmus is useful for a quick classification, but it cannot tell us how acidic or alkaline the solution is.

Universal indicator gives an approximate pH

Universal indicator contains a mixture of indicators and changes through a range of colours. The observed colour is compared with a reference chart to estimate the pH.

What does pH actually tell us?

For the usual O-Level treatment:

• pH below 7 means the solution is acidic.
• pH 7 means the solution is neutral.
• pH above 7 means the solution is alkaline.

As the concentration of H⁺ ions increases, the pH decreases. A very acidic solution therefore has a high concentration of H⁺ ions.

Other indicators

Strong and Weak Acids Are Not the Same as Concentrated and Dilute Acids

This distinction is one of the most important pieces of language in acid–base chemistry. The words strong and concentrated answer completely different questions.

Strong and weak describe ionisation

A strong acid ionises almost completely in water. A high proportion of the dissolved acid particles therefore produces H⁺ ions.

Hydrochloric acid is an example:

A weak acid ionises only partially in water. Only a fraction of the dissolved acid particles forms ions at any one time.

Ethanoic acid, CH₃COOH, is a common weak acid.

Concentrated and dilute describe how much solute is present

Concentration tells us how much dissolved substance is present in a given volume of solution.

• A concentrated acid solution contains a relatively large amount of acid per unit volume.
• A dilute acid solution contains a relatively small amount of acid per unit volume.

Predicting the Characteristic Reactions of Acids

Acid reactions become much easier when you learn them as patterns. The acid supplies H⁺ ions, while the substance reacting with the acid determines whether a gas is formed and what other products appear.

Acid + reactive metal → salt + hydrogen

Reactive metals above hydrogen in the reactivity series can react with dilute hydrochloric acid or dilute sulfuric acid to produce a salt and hydrogen gas.

For example, magnesium reacts with dilute hydrochloric acid:

The essential ionic change is:

You would expect to observe:

• effervescence as hydrogen gas is produced;
• the metal gradually becoming smaller as it reacts; and
• the mixture becoming warmer in many cases.

Test the gas with a lighted splint. Hydrogen gives a squeaky pop.

Acid + metal oxide or hydroxide → salt + water

A base neutralises an acid. For a metal oxide or hydroxide base, the products are a salt and water.

For example, copper(II) oxide reacts with sulfuric acid:

Black CuO disappears as it reacts, forming a blue copper(II) sulfate solution. If excess CuO is added after all the acid has been neutralised, the extra black solid remains unreacted.

Acid + alkali → salt + water

Because an alkali is simply a soluble base, it also neutralises an acid.

If we remove the spectator ions, the essential ionic equation for neutralisation between an acid and an alkali is:

Why may there be no obvious visible change?

Hydrochloric acid, sodium hydroxide solution and sodium chloride solution are all colourless. Their reaction can therefore occur without bubbles, a precipitate or a striking colour change. An indicator is often used when we need to detect the point at which neutralisation has occurred.

Acid + carbonate → salt + water + carbon dioxide

Carbonates behave differently from metal oxides and hydroxides because the carbonate ion ultimately produces carbon dioxide as well as water.

For example:

The ionic equation is:

Effervescence is observed because carbon dioxide gas is produced. To confirm the gas, bubble it through limewater. The limewater turns milky.

Comparing the three important acid reaction patterns

Alkalis and Ammonium Salts

Alkalis have another important reaction that appears frequently in qualitative analysis. When an ammonium salt is warmed with an alkali, ammonia gas is released.

For example:

The ionic equation is:

Ammonia is a colourless gas. The confirmatory test is to expose damp red litmus paper to the gas. The paper turns blue.

Why must the litmus paper be damp?

Ammonia must dissolve in water before its alkaline behaviour is observed. The moisture on the litmus paper provides the water needed for the gas to form an alkaline solution.

How Acids Determine the Names of Salts

Many acid reactions produce salts. Rather than memorising the name of every possible salt, recognise that the acid supplies the negative ion in the salt.

The other reactant normally supplies the positive ion.

Magnesium + hydrochloric acid gives magnesium chloride.
Magnesium + sulfuric acid gives magnesium sulfate.

A reliable method for predicting the salt formula

Suppose sulfuric acid reacts with sodium hydroxide.

Sodium hydroxide supplies Na⁺ ions, while sulfuric acid supplies SO₄²⁻ ions. Two Na⁺ ions are needed to balance one SO₄²⁻ ion, so the salt formula is Na₂SO₄.

Classifying Acidic, Basic, Amphoteric and Neutral Oxides

The acid–base idea can also be used to classify oxides. The most reliable method is not simply to ask whether the element is a metal or non-metal. Instead, ask what the oxide reacts with.

Acidic oxides

An acidic oxide reacts with a base or alkali. Many non-metal oxides are acidic.

Examples include:

• carbon dioxide, CO₂
• sulfur dioxide, SO₂
• sulfur trioxide, SO₃

For example:

Basic oxides

A basic oxide reacts with an acid to form a salt and water. Many metal oxides behave in this way.

Examples include MgO and CuO.

Amphoteric oxides

An amphoteric oxide reacts with both acids and alkalis. This means it can display acidic behaviour in one reaction and basic behaviour in another.

Important examples include:

• aluminium oxide, Al₂O₃
• zinc oxide, ZnO

These are important exceptions to the convenient shortcut that “metal oxides are basic”.

Neutral oxides

A neutral oxide does not normally react with acids or bases to form salts under the conditions studied at O-Level.

Examples include:

• carbon monoxide, CO
• nitrogen monoxide, NO
• water, H₂O

Using Neutralisation to Control Soil pH

Acid–base chemistry has an important agricultural application. Different crops grow best within particular ranges of soil pH. If a soil becomes too acidic, plant growth can be affected because the availability of mineral ions and other soil conditions change.

Farmers can reduce excessive soil acidity by adding a basic material. This process is commonly described as liming.

Suitable materials include:

• calcium carbonate, CaCO₃
• calcium oxide, CaO
• calcium hydroxide, Ca(OH)₂

These substances react with acids in the soil and raise the soil pH.

Connecting Observations to Chemical Conclusions

Examination questions often give you observations rather than telling you which reaction has occurred. The skill is to move from observation → substance produced → reaction type.

Example 1: A metal and an acid

A grey metal is added to dilute hydrochloric acid. Effervescence occurs and the gas gives a squeaky pop with a lighted splint.

The squeaky pop identifies the gas as hydrogen. This strongly suggests the familiar acid + reactive metal reaction.

Example 2: An unknown carbonate

A white solid is added to dilute acid. A colourless gas is produced and turns limewater milky.

The limewater result confirms carbon dioxide. In this context, the observation points to an acid reacting with a carbonate.

Example 3: An ammonium salt

An unknown salt is warmed with sodium hydroxide solution. A gas is produced that turns damp red litmus paper blue.

The gas is ammonia, showing that the original substance contained NH₄⁺ ions.

Common Acids and Bases Mistakes

Mistake 1: Saying every base is an alkali

An alkali must be soluble in water. CuO is a base because it neutralises acids, but it is not an alkali because it is insoluble in water.

Mistake 2: Using “strong” to mean “concentrated”

Strength refers to the extent of ionisation. Concentration refers to the amount of dissolved acid per unit volume. They are independent ideas.

Mistake 3: Saying that pH directly measures acid strength

pH tells us about the acidity of a particular solution and is related to its H⁺ concentration. Both acid strength and concentration can affect the pH.

Mistake 4: Assuming every metal reacts with dilute acid

Check the reactivity series. Metals below hydrogen, such as copper and silver, do not normally displace hydrogen from dilute hydrochloric acid or dilute sulfuric acid.

Mistake 5: Confusing the gas tests

Hydrogen gives a squeaky pop with a lighted splint. Carbon dioxide turns limewater milky. Ammonia turns damp red litmus paper blue.

Mistake 6: Forgetting water in an acid–carbonate reaction

Acid + carbonate gives salt + water + carbon dioxide. Carbon dioxide is not the only non-salt product.

Mistake 7: Guessing the salt name

Hydrochloric acid forms chlorides, sulfuric acid forms sulfates and nitric acid forms nitrates. Use the ions supplied by the reactants to construct the salt formula.

Mistake 8: Classifying an oxide only by whether its element is a metal

The metal/non-metal pattern is useful, but it has exceptions. Al₂O₃ and ZnO are amphoteric, while CO and NO are neutral oxides. Reaction behaviour gives the more reliable classification.

A Practical Checklist for Acids and Bases Questions

When you meet an acids-and-bases question, work from the particles and reaction pattern rather than trying to recall an isolated fact.