Brønsted-Lowry Theory (HSC SSCE Chemistry): Revision Notes

Revisiting Neutralisation

What is neutralisation?

You learnt previously that a neutralisation reaction occurs when an acid and a base react together to form a salt and water. When we mix the correct amounts of acid and base, you might expect the resulting solution to always be neutral (neither acidic nor basic). However, this is only true when strong acids react with strong bases.

Understanding solution types

At $298 \text{ K}$, we can classify solutions based on the concentrations of hydronium ions ($\text{H}_3\text{O}^+$) and hydroxide ions ($\text{OH}^-$):

Acidic solution: A solution where the concentration of hydronium ions exceeds the concentration of hydroxide ions. This means $[\text{H}_3\text{O}^+] > [\text{OH}^-]$, which occurs when $[\text{H}_3\text{O}^+] > 10^{-7} \text{ mol L}^{-1}$.

Alkaline (basic) solution: A solution where the concentration of hydronium ions is less than the concentration of hydroxide ions. This means $[\text{H}_3\text{O}^+] < [\text{OH}^-]$, which occurs when $[\text{H}_3\text{O}^+] < 10^{-7} \text{ mol L}^{-1}$.

Neutral solution: A solution where the concentrations of hydronium ions and hydroxide ions are equal. This means $[\text{H}_3\text{O}^+] = [\text{OH}^-]$, which occurs when both species have a concentration of $10^{-7} \text{ mol L}^{-1}$.

Complete reactions in neutralisation

An important principle to understand is that when an acid reacts with a base (other than its conjugate base or water), the reaction will always proceed to completion. This happens provided the reactant quantities meet the required stoichiometric ratios.

For strong acids and strong bases, calculating the final pH is straightforward. For weak acids and bases, we need to use $K_a$ or $K_b$ values in our calculations.

Calculating pH after neutralisation

Let's look at how to calculate the pH of the final solution when a strong acid reacts with a strong base. Remember that sulfuric acid ($\text{H}_2\text{SO}_4$) is diprotic, meaning it can donate two protons. When it's the limiting reagent, all of its protons will react.

Salts: Not necessarily neutral

After an acid reacts with a base, the salt produced can create a solution that is acidic, basic, or neutral. Salts are ionic compounds formed when acids neutralise bases.

The Brønsted-Lowry theory helps us understand why salt solutions have different pH values. The conjugate acids or bases (the ions) formed during neutralisation may react with water to produce hydronium ions or hydroxide ions. This process is called hydrolysis.

Strength of conjugate acid-base pairs

The strength of a conjugate acid or base depends on the strength of the original acid or base. There's an inverse relationship:

Case 1: Strong acid + strong base → neutral salt (pH = 7)

When a strong acid reacts with a strong base, the resulting salt does not hydrolyse (react with water), so the solution remains neutral.

Example:

$\text{HCl}(\text{aq}) + \text{NaOH}(\text{aq}) \rightarrow \text{NaCl}(\text{aq}) + \text{H}_2\text{O}(\text{l})$

When sodium chloride dissolves in water:

$\text{NaCl}(\text{aq}) \rightarrow \text{Na}^+(\text{aq}) + \text{Cl}^-(\text{aq})$

The sodium ion ($\text{Na}^+$) does not react with water, so it doesn't affect the pH. Hydrochloric acid is a strong acid, producing chloride ion ($\text{Cl}^-$) as its conjugate base. This conjugate base is very weak and will not remove a proton from water. The resulting solution is neutral with pH = 7.

Case 2: Strong acid + weak base → acidic salt (pH < 7)

When a strong acid reacts with a weak base, the resulting solution will be acidic. The anion from the strong acid won't affect the pH, but the conjugate acid of the weak base (cation) will hydrolyse to produce $\text{H}_3\text{O}^+$ ions.

Example:

$\text{HCl}(\text{aq}) + \text{NH}_3(\text{aq}) \rightarrow \text{NH}_4\text{Cl}(\text{aq})$

The ammonium chloride salt dissociates:

$\text{NH}_4\text{Cl}(\text{aq}) \rightarrow \text{NH}_4^+(\text{aq}) + \text{Cl}^-(\text{aq})$

The ammonium ion reacts with water:

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

This produces hydronium ions, making the solution acidic (pH < 7).

Case 3: Weak acid + strong base → basic salt (pH > 7)

When a weak acid reacts with a strong base, the resulting solution will be basic. The cation from the base won't affect the pH, but the conjugate base of the weak acid (anion) will hydrolyse to produce $\text{OH}^-$ ions.

Example:

$\text{HF}(\text{aq}) + \text{NaOH}(\text{aq}) \rightarrow \text{NaF}(\text{aq}) + \text{H}_2\text{O}(\text{aq})$

The salt dissociates:

$\text{NaF}(\text{aq}) \rightarrow \text{Na}^+(\text{aq}) + \text{F}^-(\text{aq})$

The sodium cation doesn't affect pH. However, the fluoride ion is the conjugate base of the weak acid HF. It reacts with water:

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

This produces hydroxide ions, making the solution basic (pH > 7).

Case 4: Weak acid + weak base → depends on Ka and Kb

When a weak acid reacts with a weak base, both the cation and anion of the resulting salt will react with water. Whether the solution is acidic or basic depends on which ion is stronger. We need to compare the Ka and Kb values to determine this.

Example:

$\text{HCOOH}(\text{aq}) + \text{NH}_3(\text{aq}) \rightarrow \text{NH}_4\text{HCOO}(\text{aq})$

The salt produces $\text{NH}_4^+$ (conjugate acid of weak base $\text{NH}_3$) and $\text{HCOO}^-$ (conjugate base of weak acid $\text{HCOOH}$). Both can hydrolyse:

$\text{NH}_4^+(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{NH}_3(\text{aq}) + \text{H}_3\text{O}^+(\text{aq}) \quad K_a = 5.6 \times 10^{-10}$

$\text{HCOO}^-(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{HCOOH}(\text{aq}) + \text{OH}^-(\text{aq}) \quad K_b = 6.25 \times 10^{-11}$

Summary table: pH of salt solutions

Acid Type	Base Type	Example Salt	Resulting pH
Strong acid	Strong base	$\text{NaCl}$	pH = 7 (neutral)
Strong acid	Weak base	$\text{NH}_4\text{Cl}$	pH < 7 (acidic)
Weak acid	Strong base	$\text{CH}_3\text{COONa}$	pH > 7 (basic)
Weak acid	Weak base	$\text{CH}_3\text{COONH}_4$	Depends on $K_a$ and $K_b$ values

Amphiprotic salts

Some salts, like some molecules, can act as both acids and bases. These are called amphiprotic salts. You already know that water is amphiprotic, but certain salt ions share this property.

Common examples

Consider sodium hydrogen carbonate ($\text{NaHCO}_3$) and potassium dihydrogen phosphate ($\text{KH}_2\text{PO}_4$). When these salts dissolve in water, they produce ions:

• $\text{NaHCO}_3 \rightarrow \text{Na}^+ + \text{HCO}_3^-$
• $\text{KH}_2\text{PO}_4 \rightarrow \text{K}^+ + \text{H}_2\text{PO}_4^-$

The anions ($\text{HCO}_3^-$ and $\text{H}_2\text{PO}_4^-$) are amphiprotic. The $\text{Na}^+$ and $\text{K}^+$ cations are spectator ions - they don't react with water.

Dual behaviour of amphiprotic ions

In aqueous solution, an amphiprotic substance acts as both an acid (donating a proton) and a base (accepting a proton) simultaneously, though only to a small extent.

Behaviour in acidic and basic solutions

When placed in strongly acidic or basic solutions, amphiprotic substances react much more completely, and one behaviour dominates.

In acidic solution, the hydrogen carbonate ion acts as a base, accepting a proton:

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

In basic solution, the hydrogen carbonate ion acts as an acid, donating a proton:

$\text{HCO}_3^-(\text{aq}) + \text{OH}^-(\text{aq}) \rightarrow \text{CO}_3^{2-}(\text{aq}) + \text{H}_2\text{O}(\text{l})$

Other amphiprotic ions

Other examples of amphiprotic substances include:

• Hydrogen sulfate ion ($\text{HSO}_4^-$)
• Hydrogen sulfite ion ($\text{HSO}_3^-$)
• Hydrogen sulfide ion ($\text{HS}^-$)