Colourful Solutions > Proton transfer reactions

Introduction

Protons are hydrogen atoms that have lost their outer electron. Acids are said to contain aqueous protons, H+(aq). In reality, the proton is far too reactive to have independent existence and will be attached to the lone pair on the oxygen of a water molecule, H3O+. This is known as the hydronium or hydroxonium ion.

Syllabus reference R3.1.1
Syllabus reference R3.1.2
Syllabus reference R3.1.3
Syllabus reference R3.1.4
Syllabus reference R3.1.5
Syllabus reference R3.1.6
Syllabus reference R3.1.7
Syllabus reference R3.1.8
Syllabus reference R3.1.9
Syllabus reference R3.1.10
Syllabus reference R3.1.11
Syllabus reference R3.1.12
Syllabus reference R3.1.13
Syllabus reference R3.1.14
Syllabus reference R3.1.15
Syllabus reference R3.1.16
Syllabus reference R3.1.17

Reactivity 3.1.1 - Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor.

  • Deduce the Brønsted–Lowry acid and base in a reaction.

Guidance

  • A proton in aqueous solution can be represented as both H+(aq) and H3O+(aq).
  • The distinction between the terms “base” and “alkali” should be understood.

Tools and links

  • Nature of science, Reactivity 3.4 - Why has the definition of acid evolved over time?

Reactivity 3.1.2 - A pair of species differing by a single proton is called a conjugate acid–base pair.

  • Deduce the formula of the conjugate acid or base of any Brønsted–Lowry base or acid.

Guidance

Tools and links

  • Structure 3.1 - What is the periodic trend in the acid–base properties of metal and non-metal oxides?
  • Structure 3.1 - Why does the release of oxides of nitrogen and sulfur into the atmosphere cause acid rain?

Reactivity 3.1.3 - Some species can act as both Brønsted–Lowry acids and bases.

  • Interpret and formulate equations to show acid–base reactions of these species.

Guidance

Tools and links

  • Structure 3.1 - What is the periodic trend in the acid–base properties of metal and non-metal oxides?
  • Structure 3.1 - Why does the release of oxides of nitrogen and sulfur into the atmosphere cause acid rain?

Reactivity 3.1.4 - The pH scale can be used to describe the [H+] of a solution:

  • pH = –log10[H+]; [H+] = 10–pH
  • Perform calculations involving the logarithmic relationship between pH and [H+].

Guidance

  • Include the estimation of pH using universal indicator, and the precise measurement of pH using a pH meter/probe.
  • The equations for pH are given in the data booklet.

Tools and links

  • Tool 1, Tool 2, Tool 3 - What is the shape of a sketch graph of pH against [H+]?
  • Nature of science, Tool 2 - When are digital sensors (e.g. pH probes) more suitable than analogue methods (e.g. pH paper/solution)?

Reactivity 3.1.5 - The ion product constant of water, Kw, shows an inverse relationship between [H+] and [OH-]. Kw = [H+][OH-]

  • Recognize solutions as acidic, neutral and basic from the relative values of [H+] and [OH–].

Guidance

  • The equation for Kw and its value at 298 K are given in the data booklet.

Tools and links

  • Reactivity 2.3 - Why does the extent of ionization of water increase as temperature increases?

Reactivity 3.1.6 - Strong and weak acids and bases differ in the extent of ionization.

  • Recognize that acid–base equilibria lie in the direction of the weaker conjugate.

Guidance

  • HCl, HBr, HI, HNO3 and H2SO4 and are strong acids, and group 1 hydroxides are strong bases.
  • The distinction between strong and weak acids or bases and concentrated and dilute reagents should be covered.

Tools and links

  • Reactivity 2.3 - How would you expect the equilibrium constants of strong and weak acids to compare?
  • Reactivity 1.1 - Why does the acid strength of the hydrogen halides increase down group 17?
  • Tool 1, Inquiry 2 - What physical and chemical properties can be observed to distinguish between weak and strong acids or bases of the same concentration?

Reactivity 3.1.7 - Acids react with bases in neutralization reactions.

  • Formulate equations for the reactions between acids and metal oxides, metal hydroxides, hydrogencarbonates and carbonates.

Guidance

  • Identify the parent acid and base of different salts.
  • Bases should include ammonia, amines, soluble carbonates and hydrogencarbonates; acids should include organic acids.

Tools and links

  • Tool 1, Structure 1.1 - How can the salts formed in neutralization reactions be separated?
  • Reactivity 1.1 - Neutralization reactions are exothermic. How can this be explained in terms of bond enthalpies?
  • Reactivity 3.2 - How could we classify the reaction that occurs when hydrogen gas is released from the reaction between an acid and a metal?

Reactivity 3.1.8 - pH curves for neutralization reactions involving strong acids and bases have characteristic shapes and features.

  • Sketch and interpret the general shape of the pH curve.

Guidance

  • Interpretation should include the intercept with the pH axis and equivalence point.
  • Only monoprotic neutralization reactions will be assessed.

Tools and links

  • Structure 1.4 - Why is the equivalence point sometimes referred to as the stoichiometric point?
  • Tool 1 and Tool 3, Structure 1.3 - How can titration be used to calculate the concentration of an acid or base in solution?

Reactivity 3.1.9 - The pOH scale describes the [OH-] of a solution. pOH = –log10[OH-]; [OH-] = 10–pOH (HL)

  • Interconvert [H+], [OH-], pH and pOH values.

Guidance

  • The equations for pOH are given in the data booklet.

Tools and links

Reactivity 3.1.10 - The strengths of weak acids and bases are described by their Ka, Kb, pKa or pKb values. (HL)

  • Interpret the relative strengths of acids and bases from these data.

Guidance

Tools and links

Reactivity 3.1.11 - For a conjugate acid–base pair, the relationship Ka × Kb = Kw can be derived from the expressions for Ka and Kb. (HL)

  • Solve problems involving these values.

Guidance

  • The use of quadratic equations is not expected in calculations.

Tools and links

  • Reactivity 2.3 - How can we simplify calculations when equilibrium constants Ka and Kb are very small?

Reactivity 3.1.12 - The pH of a salt solution depends on the relative strengths of the parent acid and base. (HL)

  • Construct equations for the hydrolysis of ions in a salt, and predict the effect of each ion on the pH of the salt solution.

Guidance

  • Examples should include the ammonium ion NH4+, the carboxylate ion RCOO–, the carbonate ion CO32–, and the hydrogencarbonate ion HCO3–.
  • The acidity of hydrated transition element ions and Al3+(aq) is not required.

Tools and links

Reactivity 3.1.13 - pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features. (HL)

  • Interpret the general shapes of pH curves for all four combinations of strong and weak acids and bases.

Guidance

  • Interpretation should include: intercept with the pH axis, equivalence point, buffer region, points where pH = pKa or pOH = pKb.

Tools and links

  • Tool 1 - When collecting data to generate a pH curve, when should smaller volumes of titrant be added between each measurement?

Reactivity 3.1.14 - Acid–base indicators are weak acids, where the components of the conjugate acid–base pair have different colours. (HL)

  • The pH of the end point of an indicator, where it changes colour, approximately corresponds to its pKa value.
  • Construct equilibria expressions to show why the colour of an indicator changes with pH.

Guidance

  • The generalized formula HInd(aq) can be used to represent the undissociated form of an indicator.
  • Examples of indicators with their pH range are given in the data booklet.
  • Include universal indicator as a mixture of many indicators with a wide pH range of colour change.

Tools and links

  • Tool 1, Inquiry 2, Reactivity 3.2 - What are some of the similarities and differences between indicators used in acid–base titrations and in redox titrations?

Reactivity 3.1.15 - An appropriate indicator for a titration has an end point range that coincides with the pH at the equivalence point. (HL)

  • Identify an appropriate indicator for a titration from the identity of the salt and the pH range of the indicator.

Guidance

  • Distinguish between the terms “end point” and “equivalence point”.

Tools and links

Reactivity 3.1.16 - A buffer solution is one that resists change in pH on the addition of small amounts of acid or alkali. (HL)

  • Describe the composition of acidic and basic buffers and explain their actions.

Guidance

Tools and links

  • Reactivity 2.3 - Why must buffer solutions be composed of weak acid or base conjugate systems, not of strong acids or bases?

Reactivity 3.1.17 - The pH of a buffer solution depends on both: (HL)

  • • the pKa or pKb of its acid or base
  • • the ratio of the concentration of acid or base to the concentration of the conjugate base or acid.
  • Solve problems involving the composition and pH of a buffer solution, using the equilibrium constant.

Guidance

  • Include explanation of the effect of dilution of a buffer.

Tools and links

  • Reactivity 2.3 - How does Le Chatelier’s principle enable us to interpret the behaviour of indicators and buffer solutions?