Theories of Acids and Bases

Learn and apply Theories of Acids and Bases in the published Chemistry course sequence.

GCE A-Level H1 Chemistry 8873-2027

Learning goals

show understanding of, and apply the Arrhenius theory of acids and bases
show understanding of, and apply the Brønsted-Lowry theory of acids and bases, including the concept of conjugate acids and conjugate bases
explain qualitatively the differences in behaviour between strong and weak acids and bases in terms of the extent of dissociation
explain the terms pH; Ka; Kb; Kw [the relationship Kw = KaKb is not required]
calculate [H+(aq)] and pH values for strong acids and strong bases
explain the choice of suitable indicators for acid-base titrations, given appropriate data, in terms of the strengths of the acids and bases
— explain how buffer solutions control pH
— describe and explain the uses of buffer solutions including the role of CO32− /HCO3− in maintaining pH in oceans, and how the rapid increase in atmospheric carbon dioxide gas contributed to ocean acidification

H1 Theories of Acids and Bases learning outcomes

H1 acids and bases joins theory, strong-solution calculations, indicator choice and qualitative buffer reasoning into one 8873 sequence. Follow the proton or electron evidence first, then calculate only after the reacting amounts and ion stoichiometry are clear.

Follow all six lessons for first learning. During revision, work out whether you need more help with model language, proton bookkeeping, logarithmic calculation, indicator evidence or equilibrium reasoning.

What You’ll Learn

Apply Arrhenius and Brønsted–Lowry theories using the evidence and vocabulary required by each model.
Identify conjugate acid–base pairs and demonstrate amphiprotic behaviour with equations.
Distinguish acid or base strength from concentration and explain the meanings of pH, Ka, Kb and Kw.
Calculate pH for strong acids and strong bases, including ion stoichiometry and dilution where required.
Choose acid–base indicators from acid/base strengths and supplied transition-range data.
Explain qualitative buffer action and apply carbonate equilibria to ocean buffering and ocean acidification.

Lessons (Recommended Order)

The sequence moves from acid–base models and conjugate pairs to strength, strong-solution pH, indicators and buffers.

Arrhenius and Brønsted–Lowry Theories
Choose the correct acid–base model and justify proton donation or acceptance from an equation.
Open lesson
Conjugate Acid–Base Pairs
Identify conjugate pairs and prove amphiprotic behaviour by tracking one proton and one charge unit.
Open lesson
Strong and Weak Acids and Bases
Separate strength from concentration and interpret pH, Ka, Kb and Kw without importing H2 calculations.
Open lesson
pH of Strong Acids and Strong Bases
Calculate strong-acid and strong-base pH after establishing ion stoichiometry and concentration.
Open lesson
Choosing Acid–Base Indicators
Select indicators from acid/base strengths and supplied transition ranges.
Open lesson
Buffer Solutions and Ocean Acidification
Explain qualitative buffer action, capacity and the carbonate chemistry of ocean acidification.
Open lesson

Quick Reference

Idea	Exam-ready reminder
Arrhenius	aqueous acid produces H⁺; aqueous base produces OH⁻
Brønsted–Lowry	acid donates H⁺; base accepts H⁺
Conjugate pair	members differ by exactly one H⁺ and their charges differ by one
Strong	essentially complete ionisation in water; not the same as concentrated
Strong-acid pH	find [H⁺] after any reaction or dilution, then use pH = −lg[H⁺]
Strong-base pH	find [OH⁻], then use Kw or pOH and pH + pOH = 14 at 25 °C
Indicator choice	the transition range must lie within the rapid pH change near equivalence
Buffer	conjugate base removes added H⁺; weak acid removes added OH⁻; capacity is finite
Ocean acidification	more dissolved CO₂ raises [H⁺], lowers pH and consumes CO₃²⁻

Name the model, identify the transferred species or relevant ions, complete any stoichiometry, then apply the equilibrium or logarithmic relationship allowed by the H1 scope.

Common Exam Traps

Importing Lewis acid–base theory into an H1 response when the question does not supply it as unfamiliar information.
Pairing species as conjugates when they differ by more than one H⁺.
Calling a dilute acid weak or a concentrated acid strong; strength and concentration describe different properties.
Using the weak-acid or weak-base pH workflows required in H2 instead of staying within the H1 strong-acid/base boundary.
Calculating strong-base pH directly from −lg[OH⁻]; that value is pOH.
Choosing an indicator because its colour is familiar rather than because its range lies in the rapid pH change.
Claiming a buffer keeps pH perfectly constant or has unlimited capacity.
Saying ocean acidification necessarily means seawater has pH below 7; it means its pH decreases.

Practice and Check Your Understanding

H1 Acids and Bases Topic Practice
Course-specific 8873 questions across theory, strong-solution pH, indicators and qualitative buffer action.
Start topic practice
Extended Acids and Bases Practice
Eighteen additional questions for consolidation and transfer.
Open extended practice
H1 Acids and Bases: Check Your Understanding
Check each reasoning step, read the targeted feedback, then try a fresh question independently.
Practice onlyCheck understanding

Practise

Work through questions with marking and feedback as you learn.

About 10 minutes

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Hydrogen chloride dissolved in water behaves as an acid, but hydrogen chloride dissolved in methylbenzene does not. What is the Arrhenius explanation?

In water HCl ionises to give H⁺(aq); in methylbenzene it stays as molecules and releases no H⁺ ions.
In methylbenzene HCl becomes a base.
Methylbenzene neutralises the H⁺ ions as they form.
HCl is an acid only when OH⁻ ions are present to accept its proton.

CH₃COO⁻(aq) + H₂O(l) ⇌ CH₃COOH(aq) + OH⁻(aq). Which species acts as a Brønsted–Lowry base in the forward reaction?

CH₃COO⁻
H₂O
OH⁻
CH₃COOH

HNO₂(aq) + H₂O(l) ⇌ NO₂⁻(aq) + H₃O⁺(aq). Which are the two conjugate acid–base pairs?

HNO₂/NO₂⁻ and H₃O⁺/H₂O
HNO₂/H₂O and H₃O⁺/NO₂⁻
HNO₂/H₃O⁺ and H₂O/NO₂⁻
HNO₂/NO₂⁻ and H₂O/OH⁻

Why does 1.0 mol dm⁻³ ethanoic acid conduct electricity much less well than 1.0 mol dm⁻³ hydrochloric acid?

Ethanoic acid is only partially dissociated, so its solution contains far fewer ions.
Ethanoic acid is more dilute than hydrochloric acid.
Ethanoic acid molecules are larger, so they move more slowly and carry less current.
Hydrogen chloride is ionic, whereas ethanoic acid is covalent.

At 298 K, Ka of HCN is 4.9 × 10⁻¹⁰ mol dm⁻³ and Ka of HNO₂ is 4.5 × 10⁻⁴ mol dm⁻³. Which statement about 0.10 mol dm⁻³ solutions of each is correct?

HNO₂ is the stronger acid: its larger Ka means a greater extent of dissociation, a higher [H⁺] and a lower pH.
HCN is the stronger acid, because a smaller Ka means it holds its proton less tightly.
Both solutions have the same pH, because they have the same concentration.
HNO₂ has the larger Ka, because its solution is more concentrated.

What is the pH of 0.025 mol dm⁻³ barium hydroxide, Ba(OH)₂, at 298 K? Kw = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K.

12.70
12.40
1.30
12.10

0.10 mol dm⁻³ ethanoic acid is titrated with 0.10 mol dm⁻³ sodium hydroxide. Indicator ranges: methyl orange pH 3.1–4.4; phenolphthalein pH 8.3–10.0. Which choice is correct?

Phenolphthalein, because the ethanoate ion makes the equivalence point alkaline and the steep part of the curve lies above pH 7.
Methyl orange, because the acid being titrated is ethanoic acid.
Methyl orange, because the equivalence point is at pH 7.
Either indicator, because the steep section runs from pH 3 to pH 11.

A buffer contains aqueous ammonia and ammonium chloride. Which equation shows how it removes a small amount of added sodium hydroxide?

NH₄⁺ + OH⁻ → NH₃ + H₂O
NH₃ + OH⁻ → NH₂⁻ + H₂O
NH₃ + H⁺ → NH₄⁺
Cl⁻ + OH⁻ → ClOH⁻

River water brings small amounts of acid into the sea, yet seawater stays at about pH 8. Which explanation is correct?

Seawater contains CO₃²⁻ and HCO₃⁻; added H⁺ is removed by CO₃²⁻ + H⁺ → HCO₃⁻, so the pH changes very little.
The dissolved sodium chloride neutralises the added acid.
The sea is so large that added H⁺ ions are diluted to zero.
Added H⁺ is removed by HCO₃⁻ + H⁺ → CO₃²⁻ + H₂.

Practise after feedback

After a check, practise the skills it showed you need to work on.

About 10 minutes

Questions are picked at random each time you start. You'll see the answer after each question. It's for practice only and doesn't count towards mastery.

Recent attempts

History is stored only in this browser.

No completed attempts are saved yet.

This check needs JavaScript to record answers. The reviewed questions are listed below for study.

Hydrogen chloride dissolved in water behaves as an acid, but hydrogen chloride dissolved in methylbenzene does not. What is the Arrhenius explanation?

In water HCl ionises to give H⁺(aq); in methylbenzene it stays as molecules and releases no H⁺ ions.
In methylbenzene HCl becomes a base.
Methylbenzene neutralises the H⁺ ions as they form.
HCl is an acid only when OH⁻ ions are present to accept its proton.

CH₃COO⁻(aq) + H₂O(l) ⇌ CH₃COOH(aq) + OH⁻(aq). Which species acts as a Brønsted–Lowry base in the forward reaction?

CH₃COO⁻
H₂O
OH⁻
CH₃COOH

HNO₂(aq) + H₂O(l) ⇌ NO₂⁻(aq) + H₃O⁺(aq). Which are the two conjugate acid–base pairs?

HNO₂/NO₂⁻ and H₃O⁺/H₂O
HNO₂/H₂O and H₃O⁺/NO₂⁻
HNO₂/H₃O⁺ and H₂O/NO₂⁻
HNO₂/NO₂⁻ and H₂O/OH⁻

Why does 1.0 mol dm⁻³ ethanoic acid conduct electricity much less well than 1.0 mol dm⁻³ hydrochloric acid?

Ethanoic acid is only partially dissociated, so its solution contains far fewer ions.
Ethanoic acid is more dilute than hydrochloric acid.
Ethanoic acid molecules are larger, so they move more slowly and carry less current.
Hydrogen chloride is ionic, whereas ethanoic acid is covalent.

At 298 K, Ka of HCN is 4.9 × 10⁻¹⁰ mol dm⁻³ and Ka of HNO₂ is 4.5 × 10⁻⁴ mol dm⁻³. Which statement about 0.10 mol dm⁻³ solutions of each is correct?

HNO₂ is the stronger acid: its larger Ka means a greater extent of dissociation, a higher [H⁺] and a lower pH.
HCN is the stronger acid, because a smaller Ka means it holds its proton less tightly.
Both solutions have the same pH, because they have the same concentration.
HNO₂ has the larger Ka, because its solution is more concentrated.

What is the pH of 0.025 mol dm⁻³ barium hydroxide, Ba(OH)₂, at 298 K? Kw = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K.

12.70
12.40
1.30
12.10

0.10 mol dm⁻³ ethanoic acid is titrated with 0.10 mol dm⁻³ sodium hydroxide. Indicator ranges: methyl orange pH 3.1–4.4; phenolphthalein pH 8.3–10.0. Which choice is correct?

Phenolphthalein, because the ethanoate ion makes the equivalence point alkaline and the steep part of the curve lies above pH 7.
Methyl orange, because the acid being titrated is ethanoic acid.
Methyl orange, because the equivalence point is at pH 7.
Either indicator, because the steep section runs from pH 3 to pH 11.

A buffer contains aqueous ammonia and ammonium chloride. Which equation shows how it removes a small amount of added sodium hydroxide?

NH₄⁺ + OH⁻ → NH₃ + H₂O
NH₃ + OH⁻ → NH₂⁻ + H₂O
NH₃ + H⁺ → NH₄⁺
Cl⁻ + OH⁻ → ClOH⁻

River water brings small amounts of acid into the sea, yet seawater stays at about pH 8. Which explanation is correct?

Seawater contains CO₃²⁻ and HCO₃⁻; added H⁺ is removed by CO₃²⁻ + H⁺ → HCO₃⁻, so the pH changes very little.
The dissolved sodium chloride neutralises the added acid.
The sea is so large that added H⁺ ions are diluted to zero.
Added H⁺ is removed by HCO₃⁻ + H⁺ → CO₃²⁻ + H₂.

Check what I know

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About 8 minutes

Answer 9 short questions. It shows what to work on next and doesn't count towards mastery.

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Which substance is a base according to the Arrhenius theory?

KOH, because it dissociates in water to give OH⁻(aq) ions.
CH₃OH, because its formula contains an OH group.
NH₃(g), because it reacts with HCl(g) to form a salt.
NaCl, because it dissolves in water to give ions.

In concentrated acid, HNO₃ + H₂SO₄ → H₂NO₃⁺ + HSO₄⁻. Which statement is correct?

HNO₃ acts as a Brønsted–Lowry base, because it accepts a proton from H₂SO₄.
HNO₃ acts as a Brønsted–Lowry acid, because nitric acid is always an acid.
H₂SO₄ acts as a Brønsted–Lowry base, because it forms an anion.
No acid–base reaction occurs, because both reactants are acids.

What is the conjugate base of H₂PO₄⁻?

HPO₄²⁻
H₃PO₄
PO₄³⁻
It has no conjugate base, because it is already an anion.

Separate 0.10 mol dm⁻³ solutions of hydrochloric acid and ethanoic acid are compared at 298 K. Which statement is correct?

Hydrochloric acid has the lower pH and reacts faster with magnesium, because HCl is fully dissociated and gives a higher [H⁺].
Both solutions have the same pH, because the acids have the same concentration.
Ethanoic acid reacts faster, because each molecule contains four hydrogen atoms.
Hydrochloric acid gives a larger total volume of hydrogen with excess magnesium.

What is Kw, and what is its value at 298 K?

Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶
Kw = [H⁺] + [OH⁻] = 2.0 × 10⁻⁷ mol dm⁻³
Kw = [H⁺][OH⁻] = 1.0 × 10⁻⁷ mol² dm⁻⁶
Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol dm⁻³

10.0 cm³ of 0.50 mol dm⁻³ nitric acid is diluted with water to 250 cm³. What is the pH of the diluted solution?

1.70
0.30
2.30
12.30

0.10 mol dm⁻³ aqueous ammonia is titrated with 0.10 mol dm⁻³ hydrochloric acid. Indicator ranges: methyl orange pH 3.1–4.4; bromothymol blue pH 6.0–7.6; phenolphthalein pH 8.3–10.0. Which indicator is most suitable?

Methyl orange, because the equivalence point of a weak base–strong acid titration is acidic and the steep part of the curve lies in the acidic range.
Phenolphthalein, because ammonia is a base.
Bromothymol blue, because every titration has its equivalence point at pH 7.
Any of them, because the pH changes sharply from about 3 to 11.

A buffer contains ethanoic acid and sodium ethanoate. How does it resist a change in pH when a small amount of hydrochloric acid is added?

The large amount of CH₃COO⁻ removes the added H⁺: CH₃COO⁻ + H⁺ → CH₃COOH, so [H⁺] changes very little.
CH₃COOH removes the added H⁺: CH₃COOH + H⁺ → CH₃COOH₂⁺.
The added H⁺ is neutralised by OH⁻ ions from the water.
The pH stays exactly constant, because Ka changes to compensate.

Rising atmospheric carbon dioxide is causing ocean acidification. Which description of the chemistry is correct?

More CO₂ dissolves and forms H⁺ and HCO₃⁻; the H⁺ reacts with CO₃²⁻ to form HCO₃⁻, so CO₃²⁻ is used up and the pH of seawater slowly falls.
Dissolved CO₂ reacts with OH⁻, so seawater becomes more alkaline.
The CO₃²⁻/HCO₃⁻ buffer keeps the pH exactly constant however much CO₂ dissolves.
Dissolved CO₂ forms extra CO₃²⁻ ions, which raise the pH.

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Theories of Acids and Bases

H1 Theories of Acids and Bases learning outcomes

What You’ll Learn

Lessons (Recommended Order)

Arrhenius and Brønsted–Lowry Theories

Conjugate Acid–Base Pairs

Strong and Weak Acids and Bases

pH of Strong Acids and Strong Bases

Choosing Acid–Base Indicators

Buffer Solutions and Ocean Acidification

Quick Reference

Practice and Check Your Understanding

H1 Acids and Bases Topic Practice

Extended Acids and Bases Practice

H1 Acids and Bases: Check Your Understanding

Practise

Practice complete

Your answers

Practise after feedback

Practice complete

Your answers

Check what I know

Starting check complete

Your answers

Check my progress

Progress check complete

Your answers

Check again

Progress check complete

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Review

Review complete

Your answers