Characteristic Properties of Acids and BasesCambridge IGCSE Chemistry: Revision notes
Section 1
What are the key reactions of acids?
Acids are substances that release H⁺ ions (hydrogen ions) in aqueous solution. They exhibit three characteristic reactions:
-
Reaction with metals: Acids react with reactive metals to produce a salt and hydrogen gas.
- General equation: acid + metal → salt + hydrogen gas
- Example: HCl + Mg → MgCl₂ + H₂↑
-
Reaction with bases: Acids neutralise bases to form a salt and water only.
- General equation: acid + base → salt + water
- Example: HCl + NaOH → NaCl + H₂O
-
Reaction with carbonates: Acids react with carbonates (and hydrogencarbonates) to produce a salt, water, and carbon dioxide gas.
- General equation: acid + carbonate → salt + water + carbon dioxide
- Example: 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂↑
These reactions are diagnostic tests for the presence of an acid.
Examiners expect you to write balanced equations for these reactions. Always show the three products clearly: the salt, and either water and CO₂, or hydrogen gas. Include state symbols in your answer (aq), (s), (l), (g).
Students often forget that only reactive metals (e.g. Mg, Zn, Fe) react with dilute acids. Copper and metals below it in the reactivity series do not react with dilute acids, so this is not a universal property of all acids.
Section 2
How do acids affect indicators?
Acids contain high concentrations of H⁺ ions, and their effect on different indicators is predictable:
| Indicator | Colour in Acid | Original Colour | Colour in Alkali |
|---|---|---|---|
| Litmus | Red | Purple/Blue | Blue |
| Thymolphthalein | Colourless | Colourless | Blue |
| Methyl orange | Red | Orange | Yellow |
| Universal indicator | Red (pH 1–3) | Green (pH 7) | Blue/Purple (pH 11–14) |
These colour changes occur because indicators are weak acids or bases themselves, and they change structure and colour when H⁺ ions are present or absent.
Using Universal Indicator Paper:
- Compare the colour of the indicator paper against the colour scale provided
- The pH scale ranges from 0–14, where:
- pH < 7 = acidic solution
- pH = 7 = neutral solution
- pH > 7 = alkaline solution
- The lower the pH, the higher the concentration of H⁺ ions
When answering indicator questions, always state the initial colour and the final colour after adding the acid. Simply saying 'it turns red' without reference to the starting colour will lose marks.
If you add hydrochloric acid to a solution of thymolphthalein that is initially blue, the solution becomes colourless because the high concentration of H⁺ ions causes the indicator to change structure and lose colour.
Section 3
What is the difference between bases and alkalis?
This is a crucial distinction that many students confuse:
Bases:
- Bases are oxides or hydroxides of metals
- Examples: CuO (copper oxide), Ca(OH)₂ (calcium hydroxide), Al₂O₃ (aluminium oxide)
- Bases are not necessarily soluble in water
- All bases neutralise acids, but not all bases are soluble
Alkalis:
- Alkalis are soluble bases
- They are specifically hydroxides of Group 1 metals (and some Group 2 metals)
- Examples: NaOH (sodium hydroxide), KOH (potassium hydroxide), Ca(OH)₂ (calcium hydroxide)
- Alkalis produce OH⁻ ions (hydroxide ions) in aqueous solution
- All alkalis are bases, but not all bases are alkalis
The key difference is solubility: only soluble bases are classified as alkalis.
Characteristic Properties of Bases and Alkalis:
-
Reaction with acids: Bases neutralise acids to form a salt and water
- General equation: base + acid → salt + water
-
Reaction with ammonium salts (alkalis only): Alkalis react with ammonium salts when heated, releasing ammonia gas
- General equation: alkali + ammonium salt → salt + ammonia + water
- Example: 2NaOH + (NH₄)₂SO₄ → Na₂SO₄ + 2NH₃↑ + 2H₂O
- The ammonia gas can be detected by its pungent smell
Examiners frequently test the base vs. alkali distinction. When asked 'Is X a base or an alkali?', you must consider solubility. If soluble in water, it is an alkali; if insoluble, it is a base only.
Think of bases as the broader category (like 'vehicles') and alkalis as a specific subset (like 'cars'). All alkalis are bases, but only some bases are alkalis.
Section 4
How do alkalis affect indicators?
Alkalis contain high concentrations of OH⁻ ions, and their effect on different indicators is the opposite to acids:
| Indicator | Colour in Acid | Original Colour | Colour in Alkali |
|---|---|---|---|
| Litmus | Red | Purple/Blue | Blue |
| Thymolphthalein | Colourless | Colourless | Blue |
| Methyl orange | Red | Orange | Yellow |
| Universal indicator | Red (pH 1–3) | Green (pH 7) | Blue/Purple (pH 11–14) |
Effect of pH on Alkalinity:
- Solutions with pH > 7 are alkaline
- The higher the pH above 7, the greater the concentration of OH⁻ ions
- For example, pH 13 represents a much stronger alkali than pH 8
- Universal indicator paper is used to compare relative alkalinity by matching the colour to the pH scale
Assessing Relative Acidity and Alkalinity:
- Use universal indicator paper by placing a drop of solution on the paper
- Compare the colour to the colour scale provided
- A pH of 3 is more acidic than pH 6
- A pH of 11 is more alkaline than pH 9
- This method allows you to rank solutions by acidity or alkalinity
When comparing the alkalinity of multiple solutions, state the pH value of each from the universal indicator scale, then rank them. For example, 'Solution A (pH 10) is less alkaline than Solution B (pH 12)' shows clear reasoning.
If you add sodium hydroxide solution to methyl orange indicator (originally orange), the solution turns yellow because the high concentration of OH⁻ ions causes the indicator to change colour.
Section 5
What is neutralisation and how do acid and base strengths differ?
The Neutralisation Reaction:
Neutralisation is the reaction between an acid and a base (or alkali) in which they react completely to form a salt and water only. The ionic equation for all neutralisation reactions is:
H⁺(aq) + OH⁻(aq) → H₂O(l)
This equation shows that neutralisation is fundamentally about the H⁺ ions from the acid combining with the OH⁻ ions from the base to form water.
Strong Acids:
- A strong acid is completely dissociated in aqueous solution
- This means 100% of the acid molecules break down to release H⁺ ions
- Example: Hydrochloric acid is a strong acid
- Equation: HCl(aq) → H⁺(aq) + Cl⁻(aq) (single arrow indicates complete dissociation)
- In a solution of HCl, there are virtually no HCl molecules; only H⁺ and Cl⁻ ions
Weak Acids:
- A weak acid is only partially dissociated in aqueous solution
- Only a small percentage of the acid molecules break down to release H⁺ ions
- Example: Ethanoic acid is a weak acid
- Equation: CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq) (equilibrium arrow indicates partial dissociation)
- In a solution of ethanoic acid, most molecules remain as CH₃COOH; only a few break down into H⁺ and CH₃COO⁻ ions
- The equilibrium favours the undissociated molecules
Key Differences:
| Property | Strong Acid | Weak Acid |
|---|---|---|
| Dissociation | Complete (100%) | Partial (typically < 5%) |
| Equation arrow | Single arrow (→) | Equilibrium arrow (⇌) |
| H⁺ ion concentration | High | Low |
| pH | Lower (e.g. pH 1–2) | Higher (e.g. pH 3–6) |
| Examples | HCl, H₂SO₄, HNO₃ | CH₃COOH, citric acid |
Both strong and weak acids produce H⁺ ions and can neutralise bases, but strong acids produce more H⁺ ions per molecule than weak acids.
Use the correct arrow in equations: single arrow (→) for complete dissociation and equilibrium arrow (⇌) for partial dissociation. The mark scheme specifically checks this distinction.
Students often think weak acids don't dissociate at all. This is wrong—weak acids do dissociate, just not completely. There are always some H⁺ ions present in a weak acid solution, which is why weak acids still turn litmus red and conduct electricity.
Section 6
What are acids and bases in terms of proton transfer?
The Brønsted–Lowry definition provides a more modern understanding of acids and bases based on proton (H⁺ ion) transfer:
Acids are proton donors:
- An acid is a substance that donates a proton (H⁺ ion) to another substance
- When HCl dissolves in water: HCl donates a proton to water
- When ethanoic acid dissociates: CH₃COOH donates a proton to water
Bases are proton acceptors:
- A base is a substance that accepts a proton (H⁺ ion) from another substance
- When NaOH dissolves in water, OH⁻ ions can accept protons from acids
- When ammonia (NH₃) reacts with an acid, it accepts a proton
Application to Neutralisation: In the neutralisation reaction H⁺(aq) + OH⁻(aq) → H₂O(l):
- The H⁺ ion (from the acid) is the proton donor
- The OH⁻ ion (from the base) is the proton acceptor
- The proton is transferred from the acid to the base, forming water
This definition explains why acids and bases react with each other—because proton donation and acceptance are complementary processes. It is broader than the previous definition because it can apply to non-aqueous solutions and explains the behaviour of substances like ammonia.
The Brønsted–Lowry definition is explicitly in the spec and may be tested in structured or extended-response questions. Be able to identify acids and bases in reactions using this definition, and explain neutralisation using proton transfer language.
In the reaction CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O, the ethanoic acid donates a proton to the hydroxide ion. The OH⁻ ion accepts this proton, showing that the acid is a proton donor and the base is a proton acceptor.
Must Know
- Acids react with metals, bases, and carbonates producing hydrogen gas (or water + CO₂ with carbonates); this is diagnostic of acids
- H⁺ ions in acids turn red: litmus red, methyl orange red, thymolphthalein colourless; OH⁻ ions in alkalis turn blue: litmus blue, methyl orange yellow, thymolphthalein blue
- Bases are metal oxides or hydroxides; alkalis are soluble bases that release OH⁻ ions in water—all alkalis are bases, but not vice versa
- Neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l); this ionic equation applies to all acid–base reactions
- Strong acids completely dissociate (single arrow: HCl → H⁺ + Cl⁻); weak acids partially dissociate (equilibrium arrow: CH₃COOH ⇌ H⁺ + CH₃COO⁻)—weak acids have fewer H⁺ ions and higher pH than strong acids
- Use universal indicator paper to compare pH: pH < 7 is acidic (H⁺ ions), pH = 7 is neutral, pH > 7 is alkaline (OH⁻ ions); lower pH = more acidic, higher pH = more alkaline
- Acids are proton donors; bases are proton acceptors—this explains why they react: the acid donates H⁺ to the base
That's the notes covered.
Carry on to the next subtopic.