Acids, Bases and Alkalis

Acids

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

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

Common acids:

Bases

A base is a substance that neutralises an acid. Bases are usually metal oxides or metal hydroxides.

Examples: CuO, MgO, NaOH, Ca(OH)₂

Alkalis

An alkali is a soluble base — a base that dissolves in water. Alkalis produce hydroxide ions ($\text{OH}^-$) in solution.

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

Key relationship: All alkalis are bases, but not all bases are alkalis. A base that doesn't dissolve (like CuO) is a base but not an alkali.

The pH scale measures how acidic or alkaline a substance is, ranging from 0 to 14:

What pH Really Measures

pH is a measure of the concentration of hydrogen ions ($\text{H}^+$) in a solution:

• Lower pH = higher $\text{H}^+$ concentration = more acidic
• Higher pH = lower $\text{H}^+$ concentration = more alkaline
• pH 7 = neutral (equal $\text{H}^+$ and $\text{OH}^-$)

Indicators

An indicator is a substance that changes colour depending on pH.

Universal indicator gives a range of colours and can estimate pH. It comes as a solution or paper.

pH Meter

A pH meter (or pH probe) gives a precise numerical pH reading. It is more accurate than indicators.

Strong Acids

A strong acid completely ionises (dissociates) in water — every acid molecule breaks apart:

$\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$

Strong acids: hydrochloric acid (HCl), sulfuric acid (H₂SO₄), nitric acid (HNO₃)

Weak Acids

A weak acid only partially ionises in water — only a small fraction of molecules release H⁺:

$\text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+$

The ⇌ symbol shows this is a reversible reaction — an equilibrium.

Weak acids: ethanoic acid, citric acid, carbonic acid

Comparing Strong and Weak Acids at the Same Concentration

Important: "Strong" and "concentrated" are different! Strong = fully ionised. Concentrated = lots of acid dissolved in solution.

When an acid reacts with a base, they neutralise each other:

$\text{acid} + \text{base} \rightarrow \text{salt} + \text{water}$

The essential ionic equation for all neutralisation reactions:

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

Types of Neutralisation

Acid + alkali: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

Acid + metal oxide: $2\text{HCl} + \text{CuO} \rightarrow \text{CuCl}_2 + \text{H}_2\text{O}$

Acid + metal hydroxide: $\text{H}_2\text{SO}_4 + \text{Cu(OH)}_2 \rightarrow \text{CuSO}_4 + 2\text{H}_2\text{O}$

The salt produced depends on the acid and the base:

The metal (or ammonium) comes from the base.

Examples

• HCl + NaOH → sodium chloride + water
• H₂SO₄ + MgO → magnesium sulfate + water
• HNO₃ + KOH → potassium nitrate + water

• Know the ions produced: acids → H⁺; alkalis → OH⁻
• To name salts: acid tells you the second part (chloride/sulfate/nitrate), base gives the metal
• For strong vs weak: mention "fully ionised" vs "partially ionised"
• The ionic equation $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$ applies to ALL acid-alkali neutralisations

• Acids produce $\text{H}^+$ ions; alkalis produce $\text{OH}^-$ ions
• pH scale: 0–14 (acid → neutral → alkaline)
• Strong acids fully ionise; weak acids partially ionise
• Neutralisation: acid + base → salt + water
• Salt name = metal (from base) + ending (from acid: chloride/sulfate/nitrate)