The characteristic properties of acids and bases- CIE iGCSE Chemistry Notes - New Syllabus
The characteristic properties of acids and bases for iGCSE Chemistry Notes
Core Syllabus
- Describe the characteristic properties of acids in terms of their reactions with:
(a) metals
(b) bases
(c) carbonates - Describe acids in terms of their effect on:
(a) litmus
(b) thymolphthalein
(c) methyl orange - State that bases are oxides or hydroxides of metals and that alkalis are soluble bases
Describe the characteristic properties of bases in terms of their reactions with:
(a) acids
(b) ammonium salts - Describe alkalis in terms of their effect on:
(a) litmus
(b) thymolphthalein
(c) methyl orange - State that aqueous solutions of acids contain H⁺ ions and aqueous solutions of alkalis contain OH⁻ ions
- Describe how to compare hydrogen ion concentration, neutrality, relative acidity and relative alkalinity in terms of colour and pH using universal indicator paper
- Describe the neutralisation reaction between an acid and an alkali to produce water, H⁺(aq) + OH⁻(aq) → H₂O(l)
Supplement Syllabus
- Define acids as proton donors and bases as proton acceptors
- Define a strong acid as an acid that is completely dissociated in aqueous solution and a weak acid as an acid that is partially dissociated in aqueous solution
- State that hydrochloric acid is a strong acid, as shown by the symbol equation, HCl(aq) → H⁺(aq) + Cl⁻(aq)
- State that ethanoic acid is a weak acid, as shown by the symbol equation, CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)
Definitions and Ion Theory of Acids, Bases, and Alkalis
Bases and Alkalis
- A base is a compound that reacts with acids to form a salt and water.
- An alkali is a base that is soluble in water, producing an aqueous solution that contains \( \text{OH}^- \) ions.
- All alkalis are bases, but not all bases are alkalis.
What counts as a base
- Metal oxides such as \( \text{CuO} \), \( \text{MgO} \), \( \text{CaO} \)
- Metal hydroxides such as \( \text{NaOH} \), \( \text{KOH} \), \( \text{Ca(OH)}_2 \)
What counts as an alkali (soluble base)
- Group 1 hydroxides: \( \text{NaOH} \), \( \text{KOH} \), \( \text{LiOH} \)
- \( \text{Ca(OH)}_2 \) is sparingly soluble but alkaline enough in water to be treated as an alkali
- \( \text{Ba(OH)}_2 \) is soluble and alkaline
- Aqueous ammonia, \( \text{NH}_3(aq) \), behaves as an alkali because it generates \( \text{OH}^- \) in water via \( \text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^- \)
Solubility snapshot
- Most metal hydroxides are insoluble, except \( \text{NaOH} \), \( \text{KOH} \), \( \text{Ba(OH)}_2 \) and the limited solubility of \( \text{Ca(OH)}_2 \)
- Most metal oxides are insoluble in water, so they are bases but not alkalis
Example
Classify each substance as a base, an alkali, or neither, and justify your choice: \( \text{MgO} \), \( \text{KOH} \), \( \text{Fe}_2\text{O}_3 \), \( \text{NH}_3(aq) \).
▶️Answer/Explanation
\( \text{MgO} \) is a base because it is a metal oxide and neutralises acids, but it is insoluble so it is not an alkali.
\( \text{KOH} \) is an alkali because it is a soluble metal hydroxide that gives \( \text{OH}^- \) in water.
\( \text{Fe}_2\text{O}_3 \) is a base because it is a metal oxide, but it is insoluble so it is not an alkali.
\( \text{NH}_3(aq) \) is an alkali because in water it produces \( \text{OH}^- \) via \( \text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^- \).
Ion Theory of Acids and Alkalis
The characteristic properties of acids and alkalis arise from the types of ions they produce in aqueous solution:
- Acids in water produce hydrogen ions, \( \text{H}^+ \) (also called protons). In solution, these protons are hydrated and often represented as \( \text{H}_3\text{O}^+ \) (hydronium ions).
- Alkalis in water produce hydroxide ions, \( \text{OH}^- \).
General ion equations
- Acid dissociation: \( \text{HA}_{(aq)} \rightarrow \text{H}^+_{(aq)} + \text{A}^-_{(aq)} \). Hydrogen ions do not exist as bare protons in solution; they bond with water molecules to form \( \text{H}_3\text{O}^+ \).
- Example : \( \text{HCl}_{(aq)} \rightarrow \text{H}^+_{(aq)} + \text{Cl}^-_{(aq)} \)
- Alkali ionisation: \( \text{MOH}_{(aq)} \rightarrow \text{M}^+_{(aq)} + \text{OH}^-_{(aq)} \)
- Example : \( \text{NaOH}_{(aq)} \rightarrow \text{Na}^+_{(aq)} + \text{OH}^-_{(aq)} \)
Why these ions matter
- \( \text{H}^+ \) ions give acids their sour taste, low pH, and reactivity with metals, carbonates, and bases.
- \( \text{OH}^- \) ions give alkalis their slippery feel, high pH, and reactivity with acids and certain salts.
Example
Explain why pure hydrogen chloride gas is not acidic, but its aqueous solution is.
▶️Answer/Explanation
Acidity depends on the presence of free \( \text{H}^+ \) ions in solution. Pure \( \text{HCl}_{(g)} \) contains covalent H–Cl molecules and no free ions. When dissolved in water, H–Cl bonds break and \( \text{H}^+ \) ions are released, producing an acidic solution: \( \text{HCl}_{(aq)} \rightarrow \text{H}^+_{(aq)} + \text{Cl}^-_{(aq)} \).
Brønsted–Lowry Definitions of Acids and Bases
A Brønsted–Lowry acid is a proton donor — it donates an \( \text{H}^+ \) ion to another species.
- Example : \( \text{HCl} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{Cl}^- \).
Here \( \text{HCl} \) donates an \( \text{H}^+ \) to water, so \( \text{HCl} \) is the acid and water is the base (proton acceptor).
- Example : \( \text{HCl} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{Cl}^- \).
A Brønsted–Lowry base is a proton acceptor — it accepts an \( \text{H}^+ \) ion from another species.
- Example : \( \text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^- \)
Here \( \text{NH}_3 \) accepts an \( \text{H}^+ \) from water, so \( \text{NH}_3 \) is the base and water is the acid (proton donor).
- Example : \( \text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^- \)
Why this matters
This definition is more general than the Arrhenius definition. It works for both aqueous and non-aqueous systems and explains acid–base behaviour in terms of proton transfer.
Conjugate acid–base pairs
- When an acid loses a proton, it forms its conjugate base.
- When a base gains a proton, it forms its conjugate acid.
Example
Identify the acids, bases, conjugate acids, and conjugate bases in the following reaction:
\( \text{H}_2\text{SO}_4 + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{HSO}_4^- \)
▶️Answer/Explanation
\( \text{H}_2\text{SO}_4 \) is the acid (proton donor) and \( \text{H}_2\text{O} \) is the base (proton acceptor). After donating a proton, \( \text{H}_2\text{SO}_4 \) becomes \( \text{HSO}_4^- \) (conjugate base). After accepting a proton, \( \text{H}_2\text{O} \) becomes \( \text{H}_3\text{O}^+ \) (conjugate acid).
Example
Explain why \( \text{NH}_3 \) can be described as a Brønsted–Lowry base but not an Arrhenius base.
▶️Answer/Explanation
Arrhenius bases must contain \( \text{OH}^- \) in their formula, but \( \text{NH}_3 \) does not. However, in Brønsted–Lowry terms, \( \text{NH}_3 \) is a base because it accepts an \( \text{H}^+ \) from water to form \( \text{NH}_4^+ \) and \( \text{OH}^- \).
Strong and Weak Acids
Strong and Weak Acids
- A strong acid is one that is completely dissociated into ions when dissolved in water. This means every molecule of the acid releases a hydrogen ion (\( \text{H}^+ \)) in aqueous solution.
Symbol Equation : \( \text{HA}_{(aq)} \rightarrow \text{H}^+_{(aq)} + \text{A}^-_{(aq)} \)
- A weak acid is one that is partially dissociated in aqueous solution. Only a small proportion of the acid molecules release hydrogen ions, and the rest remain undissociated.
Symbol Equation : \( \text{HA}_{(aq)} \rightleftharpoons \text{H}^+_{(aq)} + \text{A}^-_{(aq)} \)
Key Features
- • At the same concentration, a strong acid produces a higher hydrogen ion concentration (\( \text{[H}^+\text{]} \)) than a weak acid, giving it a lower pH.
- • Reactions involving \( \text{H}^+ \) (such as with metals, carbonates, or alkalis) occur faster for strong acids of the same concentration.
- • Strong acids generally conduct electricity better than weak acids at the same concentration due to more free-moving ions.
- • Acid strength is not the same as acid concentration — a concentrated weak acid can still be less acidic than a dilute strong acid depending on \( \text{[H}^+\text{]} \).
Examples of Strong and Weak Acids
• Strong acids: hydrochloric acid (\( \text{HCl} \)), sulfuric acid (\( \text{H}_2\text{SO}_4 \)), nitric acid (\( \text{HNO}_3 \))
• Weak acids: ethanoic acid (\( \text{CH}_3\text{COOH} \)), citric acid, carbonic acid (\( \text{H}_2\text{CO}_3 \))
Hydrochloric Acid as a Strong Acid
Hydrochloric acid (\( \text{HCl} \)) is a strong acid. This means that when it is dissolved in water, it undergoes complete ionisation, producing a high concentration of hydrogen ions (\( \text{H}^+ \)) in solution.
Chemical explanation:
When \( \text{HCl} \) is added to water, the covalent bond between hydrogen and chlorine breaks completely. The hydrogen atom is released as a proton (\( \text{H}^+ \)), and the chlorine atom becomes a chloride ion (\( \text{Cl}^- \)). This reaction is irreversible under normal conditions, meaning all \( \text{HCl} \) molecules dissociate.
Symbol equation:
\( \text{HCl(aq)} \rightarrow \text{H}^+(aq) + \text{Cl}^-(aq) \)
Key points:
- Complete dissociation — no undissociated \( \text{HCl} \) molecules remain in the solution.
- Produces a high concentration of \( \text{H}^+ \) ions → low pH (close to 0 for concentrated solutions).
- Highly reactive in acid–base reactions due to high availability of protons.
- Corrosive and must be handled with care.
Example
Explain why hydrochloric acid conducts electricity well in aqueous solution.
▶️ Answer/Explanation
Hydrochloric acid is a strong acid and dissociates completely into \( \text{H}^+ \) and \( \text{Cl}^- \) ions in water.
These ions are mobile and act as charge carriers, allowing electric current to pass through the solution efficiently.
Ethanoic Acid as a Weak Acid
Ethanoic acid (\( \text{CH}_3\text{COOH} \)) is a weak acid. This means that when it is dissolved in water, it undergoes partial ionisation, producing only a small concentration of hydrogen ions (\( \text{H}^+ \)) in solution.
Chemical explanation:
In aqueous solution, most ethanoic acid molecules remain undissociated. Only a small fraction ionises to produce \( \text{H}^+ \) and \( \text{CH}_3\text{COO}^- \) ions. This establishes a reversible equilibrium between the undissociated acid and its ions.
Symbol equation:
\( \text{CH}_3\text{COOH(aq)} \rightleftharpoons \text{H}^+(aq) + \text{CH}_3\text{COO}^-(aq) \)
Key points:
- Partial dissociation — a significant proportion of molecules remain as \( \text{CH}_3\text{COOH} \).
- Produces a low concentration of \( \text{H}^+ \) ions → higher pH than a strong acid of the same concentration.
- Reversible reaction — represented by the ⇌ sign.
- Less reactive than a strong acid of the same concentration.
Example
At \( 0.10\,\text{mol dm}^{-3} \), which solution has the lower pH: \( \text{HCl}_{(aq)} \) or \( \text{CH}_3\text{COOH}_{(aq)} \)? Explain using dissociation.
▶️ Answer/Explanation
\( \text{HCl} \) is a strong acid: \( \text{HCl}_{(aq)} \rightarrow \text{H}^+_{(aq)} + \text{Cl}^-_{(aq)} \). It dissociates completely, so \( \text{[H}^+\text{]} \approx 0.10\,\text{mol dm}^{-3} \), so its pH is very low (around 0).
\( \text{CH}_3\text{COOH} \) is a weak acid: \( \text{CH}_3\text{COOH}_{(aq)} \rightleftharpoons \text{H}^+_{(aq)} + \text{CH}_3\text{COO}^-_{(aq)} \). It dissociates only partially, so \( \text{[H}^+\text{]} \lt 0.10\,\text{mol dm}^{-3} \), so its pH is higher (around 2–3).
Therefore, \( \text{HCl}_{(aq)} \) has the lower pH at the same stated concentration.
Characteristic Reactions of Acids and Bases
Characteristic Reactions of Acids and Bases
Acids with Metals → Salt + Hydrogen
Acids react with reactive metals to produce a salt and hydrogen gas. This reaction is typically faster with more reactive metals such as magnesium and zinc. The hydrogen gas produced can be tested using a lighted splint, which produces a ‘pop’ sound.
General word equation: Acid + Metal → Salt + Hydrogen
General ionic equation: \( 2\text{H}^+(aq) + \text{M}(s) \rightarrow \text{M}^{2+}(aq) + \text{H}_2(g) \)
Example: \( \text{Mg}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g) \)
Acids with Carbonates → Salt + Carbon Dioxide + Water
Acids react with carbonates to produce a salt, carbon dioxide gas, and water. The presence of carbon dioxide can be confirmed by bubbling the gas through limewater, which turns milky.
General word equation: Acid + Carbonate → Salt + Carbon Dioxide + Water
General ionic equation: \( 2\text{H}^+(aq) + \text{CO}_3^{2-}(aq) \rightarrow \text{CO}_2(g) + \text{H}_2\text{O}(l) \)
Example: \( \text{CaCO}_3(s) + 2\text{HCl}(aq) \rightarrow \text{CaCl}_2(aq) + \text{CO}_2(g) + \text{H}_2\text{O}(l) \)
Example
Write balanced equations for:
- Zinc with dilute sulfuric acid
- Calcium carbonate with hydrochloric acid
▶️Answer/Explanation
1. \( \text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{H}_2(g) \)
2. \( \text{CaCO}_3(s) + 2\text{HCl}(aq) \rightarrow \text{CaCl}_2(aq) + \text{CO}_2(g) + \text{H}_2\text{O}(l) \)
Bases with Acids → Salt + Water (Neutralisation)
Bases, whether soluble (alkalis) or insoluble, react with acids to produce a salt and water. This reaction is called neutralisation. If the base is soluble, the reaction occurs in solution and is often used in titrations to determine concentrations. If the base is insoluble, it can be added in excess to ensure complete reaction, followed by filtration of the salt solution.
- All neutralisation reactions between an acid and a base can be simplified to the same net ionic equation, showing that it is the reaction between hydrogen ions from the acid and hydroxide ions from the base to form water.
General word equation: Base + Acid → Salt + Water
General ionic equation: \( \text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l) \)
Example:
In the reaction between hydrochloric acid and sodium hydroxide:
\( \text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) \)
When written in ionic form:
\( \text{H}^+(aq) + \text{Cl}^-(aq) + \text{Na}^+(aq) + \text{OH}^-(aq) \rightarrow \text{Na}^+(aq) + \text{Cl}^-(aq) + \text{H}_2\text{O}(l) \)
After cancelling the spectator ions (\( \text{Na}^+ \) and \( \text{Cl}^- \)), the net ionic equation remains:
\( \text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l) \)
Example
Write the net ionic equation for the reaction between nitric acid and potassium hydroxide.
▶️Answer/Explanation
Step 1: Full equation – \( \text{HNO}_3(aq) + \text{KOH}(aq) \rightarrow \text{KNO}_3(aq) + \text{H}_2\text{O}(l) \)
Step 2: Ionic equation – \( \text{H}^+(aq) + \text{NO}_3^-(aq) + \text{K}^+(aq) + \text{OH}^-(aq) \rightarrow \text{K}^+(aq) + \text{NO}_3^-(aq) + \text{H}_2\text{O}(l) \)
Step 3: Cancel spectator ions – \( \text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l) \)
Bases with Ammonium Salts → Ammonia + Salt + Water
Bases react with ammonium salts upon heating to release ammonia gas, along with the formation of a salt and water. Ammonia has a characteristic pungent smell and turns damp red litmus paper blue due to its basic nature.
General word equation: Base + Ammonium Salt → Ammonia + Salt + Water
Example: \( \text{NaOH}(aq) + \text{NH}_4\text{Cl}(aq) \rightarrow \text{NH}_3(g) + \text{NaCl}(aq) + \text{H}_2\text{O}(l) \)
Example
Complete and balance the equations for:
- Magnesium oxide with nitric acid
- Sodium hydroxide with ammonium sulfate
▶️Answer/Explanation
1. \( \text{MgO}(s) + 2\text{HNO}_3(aq) \rightarrow \text{Mg(NO}_3\text{)}_2(aq) + \text{H}_2\text{O}(l) \)
2. \( 2\text{NaOH}(aq) + (\text{NH}_4)_2\text{SO}_4(aq) \rightarrow 2\text{NH}_3(g) + \text{Na}_2\text{SO}_4(aq) + 2\text{H}_2\text{O}(l) \)
Indicators and pH
Effect of Acids on Indicators
Indicators are substances that change colour depending on whether they are in an acidic or alkaline environment. Different indicators have different colour changes.
In acidic solutions:
- Litmus: Turns red in acid due to the high concentration of \( \text{H}^+ \) ions.
- Thymolphthalein: Colourless in acid as the \( \text{H}^+ \) ions suppress the indicator’s ionisation.
- Methyl orange: Turns red in acid because the indicator’s acidic form is dominant.
Example
State the colours observed when each of the following indicators is added to hydrochloric acid:
▶️Answer/Explanation
• Litmus – red
• Thymolphthalein – colourless
• Methyl orange – red
Effect of Alkalis on Indicators
Alkalis are soluble bases that produce \( \text{OH}^- \) ions in aqueous solution. These ions affect the ionisation equilibrium of indicators, producing characteristic colour changes.
In alkaline solutions:
- Litmus: Turns blue because \( \text{OH}^- \) ions shift the indicator to its basic form.
- Thymolphthalein: Turns blue in alkalis as the basic form of the indicator is strongly coloured.
- Methyl orange: Turns yellow because the basic form of the indicator is dominant.
Example
State the colours observed when each of the following indicators is added to sodium hydroxide solution:
▶️Answer/Explanation
• Litmus – blue
• Thymolphthalein – blue
• Methyl orange – yellow
Universal Indicator and pH Scale
The universal indicator is a mixture of several different indicators, designed to produce a gradual colour change over a wide pH range. This allows it to show not only whether a solution is acidic or alkaline but also the degree of acidity or alkalinity.
Interpretation:
- pH 0–3: Strongly acidic (red)
- pH 4–6: Weakly acidic (orange to yellow)
- pH 7: Neutral (green)
- pH 8–10: Weakly alkaline (blue)
- pH 11–14: Strongly alkaline (purple)
The pH value is directly related to the concentration of hydrogen ions \( \text{[H}^+ \text{]} \) in solution. Lower pH values indicate higher \( \text{[H}^+ \text{]} \) and stronger acidity, while higher pH values indicate lower \( \text{[H}^+ \text{]} \) and stronger alkalinity.
Example
A student tests three solutions using universal indicator paper:
• Solution A turns the paper red
• Solution B turns the paper green
• Solution C turns the paper purple
▶️Answer/Explanation
• Solution A – strongly acidic (pH ~1)
• Solution B – neutral (pH 7)
• Solution C – strongly alkaline (pH ~13)