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R3.1Reactivity 3 — What are the mechanisms of chemical change?

Proton transfer reactions — IB Chemistry HL Revision Notes & Topical Questions

45 min read 62 parts Lesson 1 of 11 1 quiz

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R3.1 Proton transfer reactions — IB CHEMISTRY HL notes: Use Brønsted–Lowry theory, pH, strong and weak acids, pH curves, Ka and Kb, salt hydrolysis, indicators and buffers.Acid = a species that donates a proton (H+).Base = a species that accepts a proton using a lone pair of elec…

  • Water with HCl
  • Conjugate pair
  • Amphiprotic
  • Amphoteric
  • Relationship
  • Aluminium oxide

Part 1 of 10Brønsted–Lowry acids and bases

Acid
Acid a species that donates a proton (H+).
Base
Base a species that accepts a proton using a lone pair of electrons.
Scope
  • Scope the theory also applies in the gas phase, for example HCl(g) and NH3(g).
Water with HCl
  • HCl(g) + H2O(l) → H3O+(aq) + Cl−(aq), with HCl the acid and water the base.
Conjugate pair
Conjugate pair two species that differ by one H+, for example CH3COOH / CH3COO− and H3O+ / H2O.
Proton
  • Proton an H+ ion is a bare proton, as hydrogen has 1 proton, 1 electron and no neutrons.

Part 2 of 10Amphiprotic and amphoteric species

Amphiprotic
Amphiprotic able to act both as a proton donor and as a proton acceptor, such as water and H2PO4−.
Amphoteric
Amphoteric able to act as both an acid and a base.
Relationship
  • Relationship all amphiprotic species are amphoteric, but not all amphoteric species are amphiprotic.
Aluminium oxide
  • Aluminium oxide amphoteric, reacting with HCl and NaOH, but not amphiprotic.

Part 3 of 10The pH scale and Kw

1pH
  • pH pH = −log10[H+],
  • so [H+] = 10−pH.
    What the symbols mean: pH has no unit · [H+] = concentration of hydrogen ions (mol dm−3); each 1 pH unit is a ten-fold change in [H+]
2Logarithmic
  • Logarithmic each pH unit is a tenfold change in [H+].
3Acids
  • [H+] > [OH−], so pH < 7 at 298 K.
4Bases
  • [OH−] > [H+], so pH > 7.
5Measuring
  • Measuring a pH meter is most accurate, universal indicator paper less accurate.
6Ion product of water
  • Kw = [H+][OH−] = 1.00 × 10−14 at 298 K.
    What the symbols mean: Kw = ionic product of water (1.00 × 10−14 mol2 dm−6 at 298 K) · [H+], [OH−] = ion concentrations (mol dm−3)
7Temperature
  • Temperature ionisation of water is endothermic, so Kw rises and pH falls as temperature rises.
8Neutrality
  • Neutrality pure water is neutral when [H+] = [OH−] at any temperature, and pH 7 is neutral only at 298 K.

Part 4 of 10Strong and weak acids and bases

Strong acidWeak acid
DissociationAlmost completePartial, equilibrium established
Position of equilibriumRightLeft
[H+]HighLow
ExamplesHCl, HNO3, H2SO4 (first ionisation)Ethanoic acid, HCN, H2S, H2CO3
Strong and weak acids and bases
Strong bases
  • Group 1 hydroxides, fully dissociated.
Weak bases
  • NH3, amines and some transition metal hydroxides.
HX trend
  • HX trend acid strength increases down Group 17 as H–X bonds lengthen and weaken.
    • (HF < HCl < HBr < HI)
Conjugates
  • Conjugates strong acids have weak conjugate bases and strong bases have weak conjugate acids.
Strength versus concentration
  • Strength versus concentration strong and weak refer to dissociation;
    • dilute and concentrated refer to concentration.

Part 5 of 10Distinguishing strong and weak acids, and neutralisation

Distinguishing strong and weak acids, and neutralisation
pH
  • pH at equal concentration a stronger acid has lower pH.
Conductivity
  • Conductivity a stronger acid has more ions, so conducts better.
Reactivity
  • Reactivity a stronger acid reacts faster with metals, carbonates and oxides, with a larger temperature rise.
Neutralisation
  • Neutralisation acid + base → salt + water, with H+ + OH− → H2O.
Acid with metal
  • Acid with metal salt + hydrogen;
    • with metal oxide or hydroxide — salt + water;
    • with carbonate or hydrogencarbonate — salt + water + CO2.
Naming salts
  • Naming salts hydrochloric acid gives chlorides, nitric acid nitrates and sulfuric acid sulfates.
Enthalpy of neutralisation
  • Enthalpy of neutralisation about −57 kJ mol−1 for a strong acid and strong base.

Part 6 of 10pH curves

1Shape
  • Shape an S-shaped curve with a vertical section, the midpoint being the equivalence point.
2Information
  • Information initial pH, pH and volume at equivalence, and the pH range of the sharp change.
3Strong acid + strong base
  • Strong acid + strong base starts low, equivalence at pH 7.
4Weak acid + strong base
  • Weak acid + strong base starts near 3, has a buffer region, equivalence above 7;
    • at half equivalence pH = pKa.
5Weak base + strong acid
  • Weak base + strong acid starts near 11, has a buffer region, equivalence below 7;
    • at half equivalence pOH = pKb.
6Weak acid + weak base
  • Weak acid + weak base no steep section and equivalence about 7, so no suitable indicator and it is rarely done.

Part 7 of 10pOH, K<sub>a</sub> and K<sub>b</sub> (HL)

pOH
  • pOH pOH = −log[OH−], and pH + pOH = pKw (14 at 298 K).
Acid dissociation
  • Ka = [A−][H+] / [HA].
Base dissociation
  • Kb = [BH+][OH−] / [B].
p-values
  • p-values pKa = −log Ka, pKb = −log Kb.
Strength
  • Strength larger Ka or smaller pKa means a stronger acid; larger Kb means a stronger base.
Conjugate pair
  • Ka × Kb = Kw and pKa + pKb = 14.
Assumptions
  • [HA] ≈ initial concentration, [A−] = [H+], negligible ionisation of water and 298 K.

Part 8 of 10Salt hydrolysis and indicators (HL)

Salt hydrolysis and indicators (HL)
Strong acid + strong base
  • Strong acid + strong base the ions do not react with water, so the solution is neutral.
Strong acid + weak base
  • NH4+ + H2O ⇌ H3O+ + NH3 makes the solution acidic.
Weak acid + strong base
  • CH3COO− + H2O ⇌ CH3COOH + OH− makes it basic.
Weak acid + weak base
  • Weak acid + weak base compare Ka of the cation with Kb of the anion.
Metal ions
  • Metal ions small, highly charged 3+ ions such as [Al(H2O)6]3+ release H+ and lower the pH;
    • 1+ and 2+ ions do not.
Indicator
  • Indicator a weak acid HIn whose conjugate base In− has a different colour.
End point
  • [HIn] = [In−], so pKa = pH; colour changes over about pKa ± 1.
Choosing
  • Choosing match the indicator range to the pH at equivalence using Section 18: methyl red or phenolphthalein for strong–strong, phenolphthalein for weak acid–strong base, methyl red for strong acid–weak base.
End point versus equivalence point
  • End point versus equivalence point the end point is the colour change;
    • the equivalence point is equal amounts of acid and base.

Part 9 of 10Buffer solutions (HL)

1Buffer
Buffer a solution that resists changes in pH when small amounts of acid or alkali are added.
2Acidic buffer
  • Acidic buffer a weak acid and its conjugate base (salt), such as ethanoic acid and sodium ethanoate.
3Basic buffer
  • Basic buffer a weak base and its conjugate acid, such as NH3 and NH4Cl.
4Adding H+
  • Adding H+ reacts with the conjugate base (CH3COO− + H+ → CH3COOH).
5Adding OH−
  • Adding OH− removes H+, so the weak acid dissociates to replace it.
6Limit
  • Limit only small additions are resisted.
7Acidic buffer pH
  • Acidic buffer pH pH = pKa + log10 .
    • ([salt] / [acid])
8Basic buffer pOH
  • Basic buffer pOH pOH = pKb + log10 , and pH = 14 − pOH.
    • ([salt] / [base])
9Dilution
  • Dilution the ratio is almost unchanged so pH barely changes, but buffer capacity falls.
10Temperature
  • Temperature it changes Ka and Kb, so a constant temperature is needed, for example for blood.

Part 10 of 10More from the specification

1Analytical sensitivity and calibration scale
  • Analytical sensitivity and calibration scale constructing multi-point standard absorbance curves enables reliable sample quantification across the full analytical measurement scale.
2Bronsted-Lowry acid-base theory
  • Defines an acid as a proton (H+) donor and a base as a proton acceptor in aqueous solutions.
3Conjugate acid-base pairs
  • Pairs of species related by the transfer of a single proton (e.g., NH4+ / NH3 or H2O / OH−).
4Electrical conductivity of solutions
  • Strong acids and bases completely dissociate in aqueous solution, yielding high ion concentrations and high electrical conductivity.
5Acid-base titration and equivalence point
  • Quantitative neutralization using standardized solutions to determine unknown molar concentrations.
6Aqueous
  • Dissolved in water.
7Titration
  • Method for finding the volume of one solution that reacts exactly with another.
8Carbon dioxide
  • A gas that turns limewater cloudy.

Key terms

TermMeaning
Brønsted–Lowry acidA proton donor.
Conjugate acid–base pairTwo species differing by one H+.
AmphiproticAble to donate and accept protons.
pH−log10[H+].
KwIon product of water, 1.00 × 10−14 at 298 K.
KaAcid dissociation constant.
BufferA solution that resists pH change.
Equivalence pointThe point where acid and base are in equal amounts.

Avoid these mistakes

1
Confusing amphiprotic with amphoteric.
2
Saying a strong acid is a concentrated acid.
3
Saying pure water is only neutral at pH 7 at all temperatures.
4
Confusing the end point with the equivalence point.
5
Using the buffer equation with the concentration of the acid in the numerator.

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