IB Chemistry HL 100113

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.

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.

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.

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

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

The pH scale can be used to describe the H⁺ of a solution:

$\text{pH} = -\log_{10}{[H^+]}$; $[H^+] = 10^{-\text{pH}}$

Perform calculations involving the logarithmic relationship between pH and H⁺.

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.

The ion product constant of water, $K_\text{w}$, shows an inverse relationship between [H⁺] and [OH^-]. $K_\text{w}$ = [H⁺] [OH^-]

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

The equation for $K_\text{w}$ and its value at 298 K are given in the data booklet.

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.

HCl, HBr, HI, HNO₃ and H₂SO₄ 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.

Acids react with bases in neutralization reactions.

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

Identify the parent acid and base of different salts.

Bases should include ammonia, amines, soluble carbonates and hydrogencarbonates; acids should include organic acids.

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.

Interpretation should include the intercept with the pH axis and equivalence point.

Only monoprotic neutralization reactions will be assessed.

The pOH scale describes the [OH^-] of a solution.

$\text{pOH} = -\log_{10}[OH^-]$; $[OH^-] = 10^{-{\text{pOH}}}$

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

The equations for pOH are given in the data booklet.

The strengths of weak acids and bases are described by their $K_\text{a}$, $K_\text{b}$, $\text{p}K_\text{a}$ or $\text{p}K_\text{b}$ values.

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

For a conjugate acid–base pair, the relationship $K_\text{a}$ × $K_\text{b}$ = $K_\text{w}$ can be derived from the expressions for $K_\text{a}$ and $K_\text{b}$.

Solve problems involving these values.

The use of quadratic equations is not expected in calculations.

The pH of a salt solution depends on the relative strengths of the parent acid and base.

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

Examples should include the ammonium ion NH₄⁺, the carboxylate ion RCOO^-, the carbonate ion CO₃^2-, and the hydrogencarbonate ion HCO₃^-.

The acidity of hydrated transition element ions and (aq) is not required.

pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features.

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

Interpretation should include: intercept with the pH axis, equivalence point, buffer region, points where $\text{pH} = \text{p}K_\text{a}$ or $\text{pOH} = \text{p}K_\text{b}$.

Acid–base indicators are weak acids, where the components of the conjugate acid–base pair have different colours. The pH of the end point of an indicator, where it changes colour, approximately corresponds to its $\text{p}K_\text{a}$ value.

Construct equilibria expressions to show why the colour of an indicator changes with pH.

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.

An appropriate indicator for a titration has an end point range that coincides with the pH at the equivalence point.

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

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

A buffer solution is one that resists change in pH on the addition of small amounts of acid or alkali.

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

The pH of a buffer solution depends on both:

• the $\text{p}K_\text{a}$ or $\text{p}K_\text{b}$ 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.

Include explanation of the effect of dilution of a buffer.