DP Chemistry HL R3.1 Proton Transfer Reactions

DP Chemistry HL R3.1 Proton Transfer Reactions - Full notes

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DP Chemistry HL | R3.1 Proton Transfer Reactions

DP CHEMISTRY · REACTIVITY

R3.1

Proton Transfer Reactions

Full notes for HL · IB-style

Contents

1	Brønsted-Lowry acids and bases
2	Conjugate acid-base pairs and amphiprotic species
3	The pH scale and the ionic product of water
4	Strong and weak acids and bases
5	Reactions of acids and neutralization
6	pH curves and indicators
7	Acid and base dissociation constants (HL only)
8	Buffer solutions and salt hydrolysis (HL only)
9	Brønsted–Lowry acids and bases in practice
10	Conjugate acid–base pairs
11	Reactions of acids
12	The pH scale and calculations
13	The ionic product of water
14	Strong and weak acids
15	Neutralization and salts
16	Titration curves
17	Indicators
18	pH of everyday substances
19	Acid deposition
20	Weak acid equilibria (HL)
21	Practical: preparing a soluble salt
22	Mixed review: worked problems
23	Additional exam-style practice
24	Buffers (introduction)
25	Common misconceptions and links
26	Exam-style question with markscheme
27	Check your understanding
28	Summary and topic map

These are full teaching notes. Read each section in order, try the Think about it question before you read on, and do every Quick check before moving to the next section.

Before you start

Rate yourself on each statement now, and again after you've worked through the notes. Shade R (not yet), A (getting there) or G (confident).

I can…	R	A	G
define Brønsted-Lowry acids and bases
identify conjugate acid-base pairs and amphiprotic species
calculate pH from [H+] and use Kw
distinguish strong and weak acids and bases using experimental evidence
write equations for reactions of acids with metals, bases and carbonates
interpret pH curves for strong acid-strong base titrations and choose indicators
calculate pH of weak acids using Ka and pKa
explain buffer action and the pH of salt solutions
interpret pH curves for weak acids and bases

Sections marked HL only (Sections 7 and 8) are Additional Higher Level content. Every other section is common to SL and HL.

Guiding questions

What happens when acids and bases react? How is pH used to describe acidity, and how can it be controlled?

How this topic is examined

R3.1 is examined through pH calculations, identification of conjugate pairs, comparisons of strong and weak acids, and pH curves. HL questions include Ka calculations, buffers and pH curves of weak acids.

Keep strength (extent of dissociation) and concentration (amount per volume) separate in all answers.

Command term	What the examiner wants
State / Identify	a brief answer, no explanation
Outline / Describe	a brief or detailed account, without reasons
Explain	a detailed account with causes, reasons or mechanisms
Calculate / Determine	a numerical answer with working, correct significant figures and units
Deduce / Predict	reach a conclusion from given information
Draw / Sketch	a clear labelled diagram, structure or graph
Evaluate / Discuss	weigh strengths and limitations; reach a judgement

How to use these notes

Each section teaches one part of the syllabus in full. You'll see these boxes throughout; each has one job:

Box	What it's for
Think about it	a question to get you thinking before a section; the answer comes in the section
Key point	the one idea you must not forget
Common mistake	an error examiners see every year, and how to avoid it
In the real world	where the chemistry shows up outside the classroom
Try this	something quick you can do yourself to test the idea
Worked example	full solutions with IB-style marks (M1 method, A1 answer, R1 reasoning)
Quick check	short questions to test yourself at the end of each section

1 Brønsted-Lowry acids and bases

Acids donate protons; bases accept them.

Think about it

Both vinegar and hydrochloric acid are acids, but you put one on chips. What makes an acid 'strong' or 'weak'?

In the Brønsted-Lowry theory, an acid is a proton (H+) donor and a base is a proton acceptor. Acid-base reactions are proton transfer reactions. For example, when hydrogen chloride dissolves in water: HCl(g) + H2O(l) → H3O+(aq) + Cl−(aq). HCl donates a proton (acid); water accepts it (base), forming the hydronium (oxonium) ion, H3O+. H+(aq) and H3O+(aq) are used interchangeably.

A base must have a lone pair to accept a proton: ammonia is a base: NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH−(aq). An alkali is a base that dissolves in water to produce OH− ions (NaOH, KOH, NH3(aq)).

2 Conjugate acid-base pairs and amphiprotic species

Every acid has a conjugate base.

When an acid donates a proton, the species formed is its conjugate base; when a base accepts a proton, it forms its conjugate acid. A conjugate acid-base pair differs by one proton (H+).

In CH3COOH + H2O ⇌ CH3COO− + H3O+: the pairs are CH3COOH / CH3COO− and H3O+ / H2O.

Amphiprotic species can both donate and accept a proton: water (forms H3O+ or OH−), HCO3−, HSO4−, H2PO4−. Amphoteric is the broader term for substances that react with both acids and bases (such as Al2O3).

WORKED EXAMPLE 1 Identifying conjugate pairs

Identify the conjugate pairs in NH3 + H2O ⇌ NH4+ + OH−.

Solution

NH3 accepts a proton: base; conjugate acid NH4+A1

H2O donates a proton: acid; conjugate base OH−A1

Quick check

1. Define Brønsted-Lowry acid.

2. What is the conjugate base of H2SO4?

Answers: 1. proton donor 2. HSO4−

3 The pH scale and the ionic product of water

A logarithmic scale for [H+].

pH = −log10[H+], and [H+] = 10−pH. Because the scale is logarithmic, a change of one pH unit is a tenfold change in [H+]. Acidic solutions have pH < 7, neutral pH = 7, alkaline pH > 7 (at 298 K).

Water ionizes slightly: H2O(l) ⇌ H+(aq) + OH−(aq). The ionic product of water, Kw = [H+][OH−] = 1.00 × 10−14 at 298 K. In pure water [H+] = [OH−] = 1.00 × 10−7 mol dm−3, so pH = 7. Kw increases with temperature (the ionization is endothermic), so the pH of pure water falls below 7 at higher temperatures, though it is still neutral because [H+] = [OH−].

pH can be measured with a pH meter (precise) or universal indicator (approximate).

WORKED EXAMPLE 2 pH calculations

Calculate (a) the pH of 0.0100 mol dm−3 HCl; (b) [H+] in a solution of pH 3.40; (c) the pH of 0.0500 mol dm−3 NaOH.

Solution

(a) [H+] = 0.0100; pH = −log(0.0100) = 2.00A1

(b) [H+] = 10−3.40 = 3.98 × 10−4 mol dm−3A1

Common mistake

Thinking pH 2 is twice as acidic as pH 4. pH 2 has 100 times the [H+] of pH 4.

Quick check

1. What is Kw at 298 K?

2. By what factor does [H+] change from pH 5 to pH 3?

Answers: 1. 1.00 × 10−14 2. 100 times greater

4 Strong and weak acids and bases

Strength means extent of dissociation.

A strong acid dissociates completely in water: HCl, HNO3, H2SO4 (first proton). A weak acid dissociates only partially: an equilibrium is set up with mostly undissociated molecules, such as ethanoic acid (CH3COOH) and carbonic acid. Similarly, a strong base dissociates completely (NaOH, KOH, Ba(OH)2); a weak base reacts only partially with water (NH3, amines).

Strength is not the same as concentration. A dilute strong acid can have a higher pH than a concentrated weak acid.

Distinguishing strong and weak acids of equal concentration: the strong acid has a lower pH, higher electrical conductivity (more ions), and reacts faster with magnesium or carbonates (higher [H+]). Both neutralize the same amount of base, because the weak acid dissociates further as H+ is removed.

Test (0.10 mol dm−3)	HCl (strong)	CH3COOH (weak)
pH	1.0	about 2.9
Conductivity	high	low
Rate with Mg	fast	slow
Volume of NaOH to neutralize	same	same

Quick check

1. Distinguish strong and weak acids.

2. Why do equal amounts of strong and weak acid need the same amount of base?

Answers: 1. complete vs partial dissociation 2. weak acid keeps dissociating as H+ is neutralized

5 Reactions of acids and neutralization

Acids form salts.

Acid + metal → salt + hydrogen: Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g). Acid + base (metal oxide or hydroxide) → salt + water: CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l). Acid + carbonate or hydrogencarbonate → salt + water + carbon dioxide: CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g).

Neutralization is the reaction between an acid and a base, forming a salt and water; it is exothermic. The ionic equation for any strong acid with a strong base is H+(aq) + OH−(aq) → H2O(l). Uses include treating acidic soils with lime, antacids for indigestion and treating acidic lakes.

6 pH curves and indicators

Following a titration with a pH meter.

A pH curve shows how pH changes as a base is added to an acid (or vice versa) during a titration. For a strong acid with a strong base: the pH starts low (about 1); rises slowly; then rises very steeply near the equivalence point (where moles of acid = moles of base), which is at pH 7; then levels off at a high pH (about 13).

An acid-base indicator is a weak acid whose undissociated form (HIn) and conjugate base (In−) have different colours. The indicator changes colour over a pH range of about 2 units. A suitable indicator must change colour within the steep part of the pH curve. For strong acid-strong base titrations, both phenolphthalein (8.3 to 10.0) and methyl orange (3.1 to 4.4) are suitable because the vertical section is very long.

Figure 1. pH curve for a strong acid titrated with a strong base: steep rise around pH 7 at the equivalence point.

Indicator	pH range	Colour change (acid to alkali)
Methyl orange	3.1 to 4.4	red to yellow
Bromothymol blue	6.0 to 7.6	yellow to blue
Phenolphthalein	8.3 to 10.0	colourless to pink

Quick check

1. What is the pH at the equivalence point for strong acid and strong base?

2. Why is phenolphthalein suitable for this titration?

Answers: 1. 7 2. its range lies within the steep part of the curve

7 Acid and base dissociation constants (HL only)

Quantifying the strength of weak acids and bases.

HL only

This section is Additional Higher Level content. SL students do not need it.

For a weak acid HA ⇌ H+ + A−, the acid dissociation constant is Ka = [H+][A−] ÷ [HA]. A larger Ka means a stronger acid. pKa = −log Ka: a smaller pKa means a stronger acid. For a weak base B + H2O ⇌ BH+ + OH−, Kb = [BH+][OH−] ÷ [B]. For a conjugate pair, Ka × Kb = Kw and pKa + pKb = 14 at 298 K.

pH of a weak acid: assume [H+] = [A−] and that dissociation is small, so [HA] at equilibrium ≈ initial concentration c. Then [H+] = √(Ka × c).

WORKED EXAMPLE 3 pH of a weak acid

Calculate the pH of 0.100 mol dm−3 ethanoic acid (Ka = 1.74 × 10−5).

Solution

[H+] = √(1.74 × 10−5 × 0.100) = √(1.74 × 10−6)M1

= 1.32 × 10−3 mol dm−3M1

pH = 2.88A1

8 Buffer solutions and salt hydrolysis (HL only)

Resisting changes in pH.

HL only

This section is Additional Higher Level content. SL students do not need it.

A buffer solution resists changes in pH when small amounts of acid or base are added. An acidic buffer contains a weak acid and its conjugate base (such as ethanoic acid and sodium ethanoate). Added H+ reacts with the conjugate base: CH3COO− + H+ → CH3COOH. Added OH− reacts with the weak acid: CH3COOH + OH− → CH3COO− + H2O. Because large reservoirs of both are present, pH changes very little. A basic buffer contains a weak base and its conjugate acid (NH3 and NH4Cl).

The pH of a buffer: pH = pKa + log([A−] ÷ [HA]). When [HA] = [A−], pH = pKa. Blood is buffered at pH 7.4 mainly by the carbonic acid/hydrogencarbonate system.

Salt hydrolysis: salts of a strong acid and strong base (NaCl) give neutral solutions; salts of a weak acid and strong base (sodium ethanoate) are alkaline, because the anion accepts protons from water; salts of a strong acid and weak base (NH4Cl) are acidic.

pH curves for weak acids: a weak acid titrated with a strong base starts at a higher pH, rises more gradually, has a buffer region, and an equivalence point above 7. At the half-equivalence point, pH = pKa. Phenolphthalein is suitable; methyl orange is not.

WORKED EXAMPLE 4 Buffer pH

Calculate the pH of a buffer containing 0.100 mol dm−3 ethanoic acid (pKa = 4.76) and 0.200 mol dm−3 sodium ethanoate.

Solution

pH = pKa + log([A−]/[HA]) = 4.76 + log(2)M1

= 5.06A1

Quick check

1. What is in an acidic buffer?

2. Why is a solution of sodium ethanoate alkaline?

3. What is pH at half-equivalence for a weak acid?

Answers: 1. weak acid and its conjugate base 2. ethanoate accepts H+ from water, forming OH− 3. pKa

9 Brønsted–Lowry acids and bases in practice

Proton donors and acceptors.

A Brønsted–Lowry acid donates a proton (H+); a base accepts a proton. In every acid–base reaction, a proton is transferred from the acid to the base.

WORKED EXAMPLE 5 Identifying (standard)

Identify the acid and base in each direction: NH3 + H2O ⇌ NH4+ + OH−.

Solution

Forward: H2O acid, NH3 baseA1

Reverse: NH4+ acid, OH− baseA1

Quick check

1. Identify the acid in HCO3− + OH− → CO32− + H2O.

2. Why is water described as amphiprotic?

Answers: 1. HCO3− 2. it can donate or accept a proton

10 Conjugate acid–base pairs

Differ by one proton.

A conjugate pair differs by one H+: the conjugate base of HCl is Cl−; the conjugate acid of NH3 is NH4+. The stronger the acid, the weaker its conjugate base.

Acid	Conjugate base
HCl	Cl−
H2SO4	HSO4−
CH3COOH	CH3COO−
NH4+	NH3
H2O	OH−
H3O+	H2O

WORKED EXAMPLE 6 Pairs (standard)

Identify both conjugate pairs in HNO3 + H2O → H3O+ + NO3−.

Solution

HNO3/NO3−A1

H3O+/H2OA1

Quick check

1. Give the conjugate base of HSO4−.

2. Give the conjugate acid of HPO42−.

Answers: 1. SO42− 2. H2PO4−

11 Reactions of acids

Typical reactions.

Acids react with: metals (salt + hydrogen); metal oxides and hydroxides (salt + water); carbonates and hydrogencarbonates (salt + water + carbon dioxide); ammonia (ammonium salt).

WORKED EXAMPLE 7 Equations (standard)

Write equations for HCl with (a) zinc; (b) calcium carbonate; (c) copper(II) oxide.

Solution

(a) Zn + 2HCl → ZnCl2 + H2A1

(b) CaCO3 + 2HCl → CaCl2 + H2O + CO2A1

Quick check

1. Write the equation for sulfuric acid with sodium hydroxide.

2. Name the salt formed from nitric acid and magnesium oxide.

Answers: 1. H2SO4 + 2NaOH → Na2SO4 + 2H2O 2. magnesium nitrate

12 The pH scale and calculations

pH = −log[H+].

pH = −log10[H+] and [H+] = 10−pH. A change of 1 pH unit is a tenfold change in [H+]. Neutral water at 298 K has pH 7.

WORKED EXAMPLE 8 pH (basic)

Find the pH of 0.0050 mol dm−3 HCl and [H+] at pH 3.6.

Solution

pH = −log(0.0050) = 2.30A1

[H+] = 10−3.6 = 2.5 × 10−4 mol dm−3A1

Quick check

1. How many times more acidic is pH 2 than pH 5?

2. Find the pH of 0.020 mol dm−3 H2SO4 assuming full dissociation.

Answers: 1. 1000 times 2. [H+] = 0.040: pH 1.40

13 The ionic product of water

Kw = [H+][OH−].

Water self-ionizes: H2O ⇌ H+ + OH−. At 298 K, Kw = [H+][OH−] = 1.00 × 10−14. In any aqueous solution, [H+] = Kw/[OH−], and pH + pOH = 14 at 298 K.

WORKED EXAMPLE 9 Alkali pH (standard)

Find the pH of 0.025 mol dm−3 NaOH at 298 K.

Solution

[H+] = 10−14/0.025 = 4.0 × 10−13M1

pH = 12.40A1

Quick check

1. Find the pH of 0.0010 mol dm−3 Ba(OH)2.

2. Find [OH−] at pH 9.0.

Answers: 1. [OH−] = 0.0020: pH 11.30 2. 1.0 × 10−5 mol dm−3

14 Strong and weak acids

Degree of dissociation.

Strong acids (HCl, HNO3, H2SO4) dissociate completely in water. Weak acids (CH3COOH, H2CO3) dissociate only partially, establishing an equilibrium. At the same concentration, a weak acid has higher pH, lower conductivity and reacts more slowly with metals and carbonates.

WORKED EXAMPLE 10 Distinguishing (exam level)

Describe two experiments to distinguish 0.10 mol dm−3 HCl from 0.10 mol dm−3 CH3COOH.

Solution

Measure pH: HCl pH 1, CH3COOH pH about 2.9R1

Conductivity higher for HCl; or faster bubbling with MgR1

Common mistake

Confusing strength with concentration. A dilute strong acid can have a higher pH than a concentrated weak acid.

Quick check

1. Why do equal volumes of 0.10 mol dm−3 HCl and CH3COOH neutralize the same amount of NaOH?

Answers: 1. both contain the same amount of acid; the weak acid dissociates further as H+ is removed

15 Neutralization and salts

Acid + base.

The ionic equation for neutralization of any strong acid by a strong base is H+(aq) + OH−(aq) → H2O(l), ΔH ≈ −57 kJ mol−1. Salts are named from the metal (or ammonium) and the acid: chloride, sulfate, nitrate, ethanoate.

WORKED EXAMPLE 11 Neutralization enthalpy (standard)

Explain why ΔHneut is similar for HCl + NaOH and HNO3 + KOH.

Solution

Both are strong acid + strong baseR1

The same reaction occurs: H+ + OH− → H2OR1

16 Titration curves

pH against volume.

A pH curve plots pH against volume of titrant. For a strong acid titrated with a strong base, the pH rises slowly, then very steeply around the equivalence point (pH 7), then levels off.

Figure 2. Titration of 25.0 cm3 of 0.100 mol dm−3 HCl with NaOH.

WORKED EXAMPLE 12 Equivalence point (standard)

Find the equivalence volume and pH in the figure.

Solution

25.0 cm3A1

pH 7A1

Quick check

1. Find the pH after adding 10.0 cm3 of NaOH.

Answers: 1. n(H+) left 1.5 × 10−3 in 35 cm3: [H+] = 0.0429: pH 1.37

17 Indicators

Colour changes at end points.

An indicator is a weak acid whose conjugate base has a different colour. The colour change happens over a pH range; choose an indicator whose range lies within the steep part of the titration curve.

Indicator	pH range	Colour change
methyl orange	3.1–4.4	red to yellow
bromothymol blue	6.0–7.6	yellow to blue
phenolphthalein	8.3–10.0	colourless to pink

WORKED EXAMPLE 13 Choosing (standard)

Which indicator suits the titration of ethanoic acid with NaOH?

Solution

Equivalence point above 7 (weak acid, strong base)R1

PhenolphthaleinA1

18 pH of everyday substances

From battery acid to bleach.

Substance	Approximate pH
stomach acid	1–2
lemon juice	2
coffee	5
pure water	7
blood	7.4
baking soda solution	8.5
household ammonia	11
bleach	12.5

In the real world

Normal rain has a pH of about 5.6 because dissolved CO2 forms carbonic acid; acid rain has a pH below 5.6 due to dissolved sulfur and nitrogen oxides.

Quick check

1. Why is unpolluted rain slightly acidic?

Answers: 1. dissolved CO2 forms carbonic acid

19 Acid deposition

Causes and effects.

Sulfur dioxide (from burning sulfur-containing fuels) and nitrogen oxides (from high-temperature combustion in engines) form sulfuric and nitric acids in the atmosphere. Acid deposition damages limestone buildings, leaches nutrients from soils, harms aquatic life and corrodes metals.

WORKED EXAMPLE 14 Equations (standard)

Write equations for SO2 forming sulfuric acid.

Solution

2SO2 + O2 → 2SO3A1

SO3 + H2O → H2SO4A1

Quick check

1. Write the equation for acid rain reacting with limestone.

Answers: 1. CaCO3 + H2SO4 → CaSO4 + H2O + CO2

20 Weak acid equilibria (HL)

Ka and pKa.

For HA ⇌ H+ + A−, Ka = [H+][A−]/[HA] and pKa = −log Ka. For a weak acid, [H+] ≈ √(Kac).

WORKED EXAMPLE 15 Weak acid pH (exam level)

Find the pH of 0.100 mol dm−3 ethanoic acid (Ka = 1.8 × 10−5).

Solution

[H+] = √(1.8 × 10−6) = 1.34 × 10−3M1 A1

pH = 2.87A1

21 Practical: preparing a soluble salt

Two methods.

Soluble salts are made by (1) reacting an acid with an excess of an insoluble base, metal or carbonate, filtering off the excess, and crystallizing the filtrate; or (2) titrating an acid with an alkali using an indicator to find the volumes, then repeating without indicator and crystallizing.

WORKED EXAMPLE 16 Copper(II) sulfate (standard)

Describe how to prepare crystals of copper(II) sulfate from copper(II) oxide.

Solution

Add excess CuO to warm dilute H2SO4, stirringR1

Filter off excess CuOR1

Heat to concentrate, cool to crystallize, filter and dryR1

Common mistake

Using the titration method with an insoluble base. Titration needs two solutions; insoluble bases are added in excess and filtered.

Quick check

1. Which method would you use to make sodium chloride?

2. Why is excess base used in method (1)?

Answers: 1. titration (both reactants soluble) 2. to ensure all the acid reacts

22 Mixed review: worked problems

Bringing the ideas together.

WORKED EXAMPLE 17 Mixed problem (exam level)

20.0 cm3 of 0.150 mol dm−3 HCl is mixed with 30.0 cm3 of 0.0800 mol dm−3 NaOH. Find the pH of the mixture.

Solution

n(H+) = 3.00 × 10−3; n(OH−) = 2.40 × 10−3A1

Excess H+ 6.0 × 10−4 in 50.0 cm3: 0.0120 mol dm−3M1

pH = 1.92A1

Quick check

1. A solution's pH increases from 3 to 5. By what factor has [H+] changed?

Answers: 1. decreased by a factor of 100

23 Additional exam-style practice

Paper 1 and Paper 2.

WORKED EXAMPLE 18 Paper 1 question (standard)

Which species is amphiprotic? A. Cl− B. HSO4− C. NH4+ D. SO42−

Solution

B: can donate or accept H+A1

WORKED EXAMPLE 19 Paper 2 question (exam level)

(a) Calculate the pH of 0.0040 mol dm−3 HNO3. (b) The solution is diluted tenfold. State the new pH. (c) Explain why diluting a weak acid tenfold changes pH by less than 1.

Solution

(a) 2.40A1

(b) 3.40A1

24 Buffers (introduction)

Resisting pH change.

A buffer contains a weak acid and its conjugate base (or a weak base and its conjugate acid) in similar amounts. Added acid reacts with the conjugate base; added base reacts with the weak acid, so pH changes very little. Blood is buffered at pH 7.4 by the H2CO3/HCO3− system.

WORKED EXAMPLE 20 Buffer action (standard)

Explain how an ethanoic acid/ethanoate buffer resists added acid.

Solution

CH3COO− + H+ → CH3COOHA1

Added H+ is removed, so [H+] changes littleR1

25 Common misconceptions and links

Students often think…	Actually…
strong means concentrated	strong means fully dissociated
pH 2 is twice as acidic as pH 4	100 times the [H+]
weak acids neutralize less base	same amount for equal moles
buffers keep pH exactly constant	they resist change for small additions

How R3.1 connects to the rest of the course

Topic	Connection
R2.3 How far?	equilibrium in weak acids
R2.1 How much?	titration calculations
S3.1 Periodic table	acidic and basic oxides
R3.4 Electron-pair sharing	Lewis acids (HL)

26 Exam-style question with markscheme

Try this under exam conditions (about 15 minutes) before reading the markscheme.

Figure 3. pH curve for 25.0 cm3 of HCl titrated with NaOH.

(a) Calculate the pH of 0.0250 mol dm−3 HNO3. [1]

(b) State the pH at the equivalence point and suggest a suitable indicator. [2]

(d) Identify the conjugate pairs in HCO3− + H2O ⇌ CO32− + H3O+. [2]

(e) Calculate the pH of 0.050 mol dm−3 HCOOH (Ka = 1.8 × 10−4). [2]

Markscheme

Part	Answer	Marks
(a)	pH = 1.60	A1
(b)	pH 7	A1
	phenolphthalein or methyl orange (any within steep section)	A1
(c)	pH meter: HCl lower pH	R1 R1
	conductivity / rate with Mg: HCl higher / faster	R1 R1
(d)	HCO3− / CO32−	A1
	H3O+ / H2O	A1
(e)	[H+] = √(9.0 × 10−6) = 3.0 × 10−3	M1
	pH = 2.52	A1

Examiner's comments. In (c), many candidates said 'the strong acid is more concentrated'. Both have the same concentration.

27 Check your understanding

Section A: multiple choice

1. A Brønsted-Lowry base is

A. a proton acceptor B. a proton donor C. an electron acceptor D. a hydroxide only

2. pH of 0.001 mol dm−3 HCl is

A. 1 B. 0.001 C. 3 D. 11

3. The conjugate acid of NH3 is

A. NH2− B. NH4+ C. H3O+ D. OH−

4. A weak acid

A. fully dissociates B. is always dilute C. cannot neutralize bases D. partially dissociates

5. Acid + carbonate gives

A. salt and hydrogen B. salt, water and CO2 C. salt only D. water only

6. Kw at 298 K is

A. 1.00 × 10−14 B. 1.00 × 10−7 C. 14 D. 7

7. A smaller pKa means

A. a weaker acid B. a stronger base C. a higher pH D. a stronger acid

8. At the half-equivalence point of a weak acid titration

A. pH = 7 B. pH = pKw C. pH = pKa D. pH = 14

Section B: structured

9. Explain the difference between a strong acid and a concentrated acid. [2]

10. Calculate the pH of 0.020 mol dm−3 Ba(OH)2. [2]

11. Describe the shape of a strong acid-strong base pH curve. [3]

12. Explain how an ethanoic acid/ethanoate buffer works. [4]

Answers

1. A 2. C 3. B 4. D 5. B 6. A 7. D 8. C

9. Strong: fully dissociated; concentrated: many moles per dm3.

10. [OH−] = 0.040; pOH 1.40; pH 12.60.

11. Starts ~1; gradual rise; steep vertical at pH 7; levels at ~13.

12. Weak acid and conjugate base reservoirs; added H+ reacts with ethanoate; added OH− reacts with acid; pH barely changes.

28 Summary and topic map

Idea	Remember
Brønsted-Lowry	acid donates H+; base accepts
Conjugate pairs	differ by one H+; amphiprotic species
pH	−log[H+]; Kw = 10−14
Strong vs weak	complete vs partial dissociation
Reactions	metals, bases, carbonates; neutralization
pH curves	steep at equivalence; indicators
HL	Ka, pKa, Kb; weak acid pH; buffers; salt hydrolysis

Before the exam, make sure you…

•separate strength from concentration
•use Kw for alkalis
•choose indicators from pH ranges
•identify conjugate pairs by one proton

Proton Transfer on one page

Cover this page with a sheet of paper and try to rebuild it from memory. Whatever you can't recall is what to revise next.

Figure 4. Topic map for Proton Transfer.

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