Module 6 Crash Course

Want to revise Module 6 briefly or get a feel for what is to come? This short summary of acid-base reactions will be perfect - these are the notes I took when learning this content in Year 10.

Earlier Theories

• The Lavoisier theory states that acidity is caused by the presence of oxygen
• The Davy theory states it arises from replaceable hydrogen
• Arrhenius posits that acids produce hydrogen ions in water, bases produce hydroxide

Bronsted-Lowry

Bronsted-Lowry theory is the theory used in HSC Chemistry, and states that acids donate protons to bases, which accept them.
Conjugate pair: Reactions of acids and bases can be generalised:
$$
\text{HA} + \text{B} \rightleftharpoons \text{HB}^+ + \text{A}^-
$$
Acid-base pairs are called conjugate pairs, and differ due to the presence or absence of $\text{H}^+$. A conjugate acid remains when a base accepts a proton, and a conjugate base is what remains when an acid donates a proton.

Lewis Theory

While not explicitly in the syllabus, it is worth knowing the currently accepted theory of acids and bases. The Lewis theory states that acids are electron pair acceptors, and bases are electron pair donors.

pH Scale

pH: pH stands for ‘hydrogen power’, and is based on the concentration of hydrogen ions/hydronium $H_3O^+$ in solution. It can be calculated by:
$$
\text{pH} = -\log_{10}[\text{H}^+]
$$
A pH of 7 is neutral, $<7$ is acidic, and $>7$ is alkaline.

Strength vs. Concentration

Ionisation reaction: An ionisation reaction involves an acid reacting with water, producing hydronium. Strong acids/bases dissociate completely in water, and weak acids/bases only partly dissociate. Weak acid/base dissociation should be written as an equilibrium reaction.
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Mono/di/polyprotic: Monoprotic acids will only donate one $\text{H}^+$ ion per molecule, whereas polyprotic acids will donate multiple (diprotic acids donate two).

Water Self-Ionisation

Self-ionisation: Since water is amphiprotic, it can react with itself:
$$
\text{H}_2\text{O} + \text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{OH}^-
$$
Water is not included in the equilibrium expression (negligible change). The self-ionisation constant ($K_w$) can be calculated:
$$
K_w = [\text{OH}^-][\text{H}_3\text{O}^+]
$$
This formula can be used to calculate pH of alkaline solutions.

pOH

Same as pH, except measures hydroxide concentration:
$$
\text{pOH} = -\log_{10}[\text{OH}^+]
$$
Clearly, the sum of pH and pOH will equal 14.

Percentage Ionisation

For weak acids, percentage ionisation represents the amount of the original acid that has ionised:
$$
\text{\%} = \frac{[\text{A}^-]}{[\text{HA}]} \times 100
$$
Where $\text{A}^-$ is the conjugate base, and $\text{HA}$ is the original acid.

Polyprotic acids

Separate equations are written, and the $K_a$ values are usually written as $K_{a1}$, $K_{a2}$, etc.

p$K_a$

p$K_a$ can be calculated:
$$
\text{p}K_a = -\log_{10}{K_a}
$$

Neutralisation

Salts produced from neutralisation reactions are not necessarily neutral, such as the case of a strong acid and weak base reaction. Conjugate acid/base of a weak acid/base is strong, and can react with water to produce hydronium or hydroxide ions.

The reaction between a weak acid and a weak base will also produce a non-neutral solution, and its pH will depend on the $K_a$ and $K_b$ of the salt ions.