Acids and Bases (Grade 12 NSC Matric Physical Sciences): Revision Notes

Acids and Bases

What are acids and bases?

Acids and bases are substances we encounter regularly in our daily lives. Many common household items contain these important chemical compounds.

Acids typically have a sour taste and are found in various foods and products. Examples include:

• Vinegar (contains acetic acid)
• Lemon juice (contains citric acid)
• Wine (contains tartaric acid)

Bases usually have a bitter taste and feel slippery to touch. Common examples include:

• Sodium hydroxide (found in cleaning products)
• Ammonia (household cleaner)
• Magnesium hydroxide (antacid tablets)

Some acids like hydrochloric acid, sulfuric acid, and nitric acid are primarily used in laboratories and industry rather than in household products.

Models for acids and bases

Scientists have developed different models to define and understand acids and bases. These models help us predict how these substances will behave in chemical reactions.

Arrhenius model

The Arrhenius model was one of the earliest definitions for acids and bases. This model focuses specifically on what happens when acids and bases dissolve in water.

Arrhenius acid: A substance that increases the concentration of hydronium ions ($\text{H}_3\text{O}^+$) when dissolved in water.

Arrhenius base: A substance that increases the concentration of hydroxide ions ($\text{OH}^-$) when dissolved in water.

For example, when hydrochloric acid dissolves in water: $\text{HCl}(aq) + \text{H}_2\text{O}(l) \rightarrow \text{H}_3\text{O}^+(aq) + \text{Cl}^-(aq)$

When sodium hydroxide dissolves in water: $\text{NaOH}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Na}^+(aq) + \text{OH}^-(aq) + \text{H}_2\text{O}(l)$

Brønsted-Lowry model

The Brønsted-Lowry model provides a broader definition that applies to reactions beyond those in water. This model focuses on the transfer of protons (hydrogen ions, $\text{H}^+$).

Brønsted-Lowry acid: A substance that donates protons ($\text{H}^+$) in a reaction. We call this a proton donor.

Brønsted-Lowry base: A substance that accepts protons ($\text{H}^+$) in a reaction. We call this a proton acceptor.

Let's examine some examples to understand how this works:

In the reaction: $\text{HCl}(aq) + \text{NH}_3(g) \rightarrow \text{NH}_4^+(aq) + \text{Cl}^-(aq)$

• HCl donates a proton, so it acts as the acid
• NH₃ accepts a proton, so it acts as the base

In the reaction: $\text{CH}_3\text{COOH}(aq) + \text{H}_2\text{O}(l) \rightarrow \text{H}_3\text{O}^+(aq) + \text{CH}_3\text{COO}^-(aq)$

• CH₃COOH donates a proton, so it acts as the acid
• H₂O accepts a proton, so it acts as the base

Amphoteric substances

Some substances can act as either acids or bases depending on the reaction conditions. We call these amphoteric substances.

Amphoteric substance: A substance that can act as an acid in one reaction or as a base in another reaction.

Other examples of amphoteric substances include ammonia (NH₃), zinc oxide (ZnO), and beryllium hydroxide (Be(OH)₂).

Conjugate acid-base pairs

The Brønsted-Lowry model introduces an important concept called conjugate acid-base pairs. Understanding these pairs helps us track proton transfer in reactions.

Conjugate acid-base pair: Two compounds that differ only by one hydrogen ion ($\text{H}^+$) and a charge of +1.

When an acid donates a proton, it forms its conjugate base. When a base accepts a proton, it forms its conjugate acid.

In the reaction above:

• HCl (acid) and Cl⁻ (conjugate base) form one conjugate pair
• NH₃ (base) and NH₄⁺ (conjugate acid) form another conjugate pair

Another example shows water acting as an acid:

• H₂O (acid) and OH⁻ (conjugate base) form one conjugate pair
• NH₃ (base) and NH₄⁺ (conjugate acid) form another conjugate pair

Strong and weak acids and bases

Not all acids and bases behave the same way in solution. We classify them based on how completely they ionise or dissociate.

Strong acids and bases

Strong acid: An acid that almost completely dissociates to form ions in solution.

Strong base: A base that almost completely dissociates to form ions in solution.

When we say "almost completely," we mean that a large percentage of the molecules break apart into ions. For example, if we add 100,000 molecules of HCl to water, about 99,000 might form ions.

$\text{HCl}(g) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_3\text{O}^+(aq) + \text{Cl}^-(aq)$

The unequal arrows show that the equilibrium strongly favours ion formation.

Weak acids and bases

Weak acid: An acid where only a small percentage of molecules dissociate to form ions in solution.

Weak base: A base where only a small percentage of molecules dissociate to form ions in solution.

For weak acids and bases, only a small fraction of molecules form ions. For example, if we add 100,000 molecules of HF to water, only about 100 might form ions.

$\text{HF}(g) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_3\text{O}^+(aq) + \text{F}^-(aq)$

The unequal arrows show that the equilibrium does not favour ion formation.

An example of a weak base is Mg(OH)₂, which only partially dissociates into Mg²⁺ and OH⁻ ions.

Dilute and concentrated solutions

It's important not to confuse strong/weak with concentrated/dilute. These terms describe different properties of solutions.

Concentrated solution: A solution with a high ratio of dissolved substance (acid or base) to solvent.

Dilute solution: A solution with a low ratio of dissolved substance to solvent.

You can have:

• A concentrated solution of a weak acid (lots of weak acid molecules, few ions)
• A dilute solution of a strong acid (few strong acid molecules, but most form ions)
• Any other combination

The electrical conductivity of a solution depends on the concentration of mobile ions present. A concentrated solution of a strong acid or base will have high electrical conductivity, while a dilute solution of a weak acid or base will have low electrical conductivity.

Ka and Kb equilibrium constants

The equilibrium constants Ka (for acids) and Kb (for bases) provide a quantitative way to measure the strength of acids and bases.

For acids, we use the general expression: $K_a = \frac{[\text{products}]}{[\text{reactants}]}$

Strong acids

Strong acids have very large Ka values (typically greater than $1 \times 10^9$). This indicates that the equilibrium lies far to the right, meaning almost complete ionisation occurs.

For hydrobromic acid: $\text{HBr}(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_3\text{O}^+(aq) + \text{Br}^-(aq)$

$K_a = \frac{[\text{H}_3\text{O}^+(aq)][\text{Br}^-(aq)]}{[\text{HBr}(g)]}$

The very large Ka value (approximately $1 \times 10^9$) confirms that HBr is a strong acid.

Weak acids

Weak acids have much smaller Ka values (typically less than 1). This indicates that only a small fraction of molecules ionise.

For ethanoic acid: $\text{CH}_3\text{COOH}(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_3\text{O}^+(aq) + \text{CH}_3\text{COO}^-(aq)$

$K_a = \frac{[\text{H}_3\text{O}^+(aq)][\text{CH}_3\text{COO}^-(aq)]}{[\text{CH}_3\text{COOH}(aq)]}$

The small Ka value ($1.7 \times 10^{-5}$) confirms that ethanoic acid is weak.

Bases and Kb values

For bases, we use similar principles. Strong bases like NaOH have very large Kb values and dissociate almost completely:

$\text{NaOH}(aq) \rightarrow \text{Na}^+(aq) + \text{OH}^-(aq)$

Weak bases like NH₃ have small Kb values and only partially ionise:

$\text{NH}_3(g) + \text{H}_2\text{O}(l) \rightleftharpoons \text{NH}_4^+(aq) + \text{OH}^-(aq)$

$K_b = \frac{[\text{NH}_4^+(aq)][\text{OH}^-(aq)]}{[\text{NH}_3(g)]}$