Concentration Versus Strength (HSC SSCE Chemistry): Revision Notes

Concentration Versus Strength

Introduction

In everyday language, the terms 'concentration' and 'strength' are often used interchangeably. However, in chemistry, these terms have very different meanings. Understanding this distinction is crucial for working with acids and bases.

Ionisation and dissociation reactions

Before exploring concentration and strength, we need to understand how acids and bases react with water.

Ionisation of acids

When an acid reacts with water, it undergoes an ionisation reaction because ions are formed. Acids produce hydronium ions ($\text{H}_3\text{O}^+$) in aqueous solution, which is often simplified to $\text{H}^+$.

For example, hydrochloric acid ionises as follows:

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

This can be simplified to:

$\text{HCl(aq)} \rightarrow \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$

Dissociation of bases

When a base dissolves in water, it forms separate ions through a dissociation reaction. Bases typically produce hydroxide ions ($\text{OH}^-$) in aqueous solution.

For example:

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

$\text{K}_2\text{O(s)} + \text{H}_2\text{O(l)} \rightarrow 2\text{K}^+\text{(aq)} + 2\text{OH}^-\text{(aq)}$

Understanding concentration

Concentration refers to the amount of solute present in a specified volume of solution. While there are many ways to express concentration, the most convenient measure in chemistry is molarity.

Molarity

Molarity is measured in moles per litre and is represented as $\text{mol L}^{-1}$ or $M$.

A concentrated solution contains a high total concentration of solute. For example, $4~M$ or $4~\text{mol L}^{-1}$ would be considered concentrated.

A dilute solution contains a low total concentration of solute. For example, less than $2~M$ or $2~\text{mol L}^{-1}$ would be considered dilute.

Understanding strength of acids and bases

Strength refers to the degree to which an acid ionises or a base dissociates in aqueous solution. This can be determined by measuring the pH of solutions at the same concentration.

Strong acids

Strong acids are acids in which almost all of the acid molecules ionise to produce hydrogen ions. This means they produce a high concentration of hydronium ions in solution.

Common examples of strong acids include:

• Hydrochloric acid ($\text{HCl}$)
• Nitric acid ($\text{HNO}_3$)
• Sulfuric acid ($\text{H}_2\text{SO}_4$)

Weak acids

Weak acids are acids in which only some of the molecules ionise. This means they produce a much lower concentration of hydronium ions compared to a strong acid at the same concentration.

Common examples of weak acids include:

• Ethanoic acid ($\text{CH}_3\text{COOH}$)
• Carbonic acid ($\text{H}_2\text{CO}_3$)
• Hydrofluoric acid ($\text{HF}$)
• Citric acid ($\text{C}_6\text{H}_8\text{O}_7$)

Strong bases

Strong bases are bases in which almost all of the species have dissociated in aqueous solution, producing a high concentration of hydroxide ions.

Common examples of strong bases include:

• Sodium hydroxide ($\text{NaOH}$)
• Potassium hydroxide ($\text{KOH}$)
• Barium hydroxide ($\text{Ba(OH)}_2$)
• Sodium oxide ($\text{Na}_2\text{O}$)

Weak bases

Weak bases are bases in which only a small proportion of the species produce hydroxide ions in aqueous solution.

Common examples of weak bases include:

• Ammonia ($\text{NH}_3$)
• Methylamine ($\text{CH}_3\text{NH}_2$)

Writing equations for weak acids and bases

Because weak acids and bases only partially ionise or dissociate, they exist in equilibrium with their ions. Therefore, we must use the equilibrium arrow ($\rightleftharpoons$) when writing their equations.

For example, the ionisation of weak hydrofluoric acid is written as:

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

This equilibrium arrow shows that both the undissociated acid molecules and the ions coexist in solution.

Comparing concentration and strength

It is crucial to understand that concentration and strength are independent properties. You can have:

• A concentrated solution of a strong acid
• A dilute solution of a strong acid
• A concentrated solution of a weak acid
• A dilute solution of a weak acid

Effect on pH

When comparing solutions of a strong acid and a weak acid at the same concentration, the strong acid solution will have a lower pH than the weak acid solution. This is because:

• The strong acid produces a higher concentration of hydronium ions ($[\text{H}_3\text{O}^+]$) since almost all its molecules ionise
• The weak acid produces a lower concentration of hydronium ions since only some molecules ionise
• pH is a measure of hydronium ion concentration, so more ions means a lower pH

Effect on hydroxide ion concentration

Similarly, when comparing solutions of a strong base and a weak base at the same concentration:

• The strong base completely dissociates to produce a high concentration of hydroxide ions ($[\text{OH}^-]$)
• The weak base produces a much lower concentration of hydroxide ions due to limited dissociation
• The strong base solution will have a higher pH than the weak base solution

Polyprotic acids

Some acids can donate more than one proton (hydrogen ion) per molecule. Understanding these acids is important for accurate pH calculations and stoichiometry.

Monoprotic acids

Monoprotic acids donate one proton per molecule when they ionise. Examples include $\text{HCl}$, $\text{HNO}_3$, and $\text{HF}$.

pH of monoprotic acids

For a strong monoprotic acid, the concentration of hydronium ions equals the concentration of the acid because ionisation is complete.

For a weak monoprotic acid, only partial ionisation occurs, so $[\text{H}_3\text{O}^+]$ is much less than the acid concentration.

Diprotic acids

Diprotic acids can donate two protons per molecule. The term 'polyprotic' refers to the ability to donate multiple protons, not how readily they ionise.

pH of diprotic acids

Sulfuric acid ($\text{H}_2\text{SO}_4$) is an important example of a diprotic acid. A $0.1~\text{mol L}^{-1}$ solution of sulfuric acid has a pH of approximately $0.69$, not $1.0$ as might be expected. This lower pH indicates a more acidic solution than $0.1~\text{mol L}^{-1}$ hydrochloric acid.

The ionisation of sulfuric acid occurs in two steps:

First ionisation (strong acid):

$\text{H}_2\text{SO}_4\text{(l)} + \text{H}_2\text{O(l)} \rightarrow \text{HSO}_4^-\text{(aq)} + \text{H}_3\text{O}^+\text{(aq)}$

Second ionisation (weak acid):

$\text{HSO}_4^-\text{(aq)} + \text{H}_2\text{O(l)} \rightleftharpoons \text{SO}_4^{2-}\text{(aq)} + \text{H}_3\text{O}^+\text{(aq)}$

Triprotic acids

Triprotic acids can donate up to three protons per molecule. An example is phosphoric acid ($\text{H}_3\text{PO}_4$). Like diprotic acids, the second and third ionisation steps involve weak acids.

Calculating degree of ionisation

The fraction of a weak acid that has ionised can be calculated using:

$\text{Fraction ionised} = \frac{[\text{H}^+]}{\text{Total concentration of acid in solution}}$

This is useful for comparing the relative strengths of different weak acids.

Remember!