Weak acids have a lower enthalpy of neutralisation because energy is required to fully dissociate the acid before it can react with a base, consuming some of the heat that would otherwise be released.
Understanding Enthalpy of Neutralisation
The enthalpy of neutralisation ($\Delta H{neut}$) is defined as the heat energy released when one mole of water is formed from the reaction between an acid and a base under standard conditions. This reaction is always exothermic, meaning it releases heat, so $\Delta H{neut}$ values are negative.
For strong acids and strong bases, the reaction is essentially:
H$^+$ (aq) + OH$^-$ (aq) → H$_2$O (l)
This reaction consistently yields an enthalpy of neutralisation of approximately -57.3 kJ/mol because both the acid and base are fully dissociated in solution, meaning hydrogen ions (H$^+$) and hydroxide ions (OH$^-$) are readily available to react and form water.
Why Weak Acids Differ
The key difference with weak acids lies in their partial dissociation. Unlike strong acids, which dissociate completely in water, weak acids only partially ionize. This means that at any given time, only a fraction of their molecules release hydrogen ions (H$^+$) into the solution.
Here's a breakdown of why this leads to a lower (less negative) enthalpy of neutralisation:
- Incomplete Dissociation: Weak acids are partially dissociated, which means not all hydrogen ions are free to react with hydroxide ions immediately. For the neutralisation to proceed, the remaining undissociated weak acid molecules must dissociate further.
- Energy Input Required for Further Dissociation: Some energy is required to break the chemical bonds in weak acid molecules to produce hydrogen ions. This process is often endothermic or requires a small energy input, effectively consuming a portion of the heat energy that is simultaneously being released by the H$^+$ + OH$^-$ → H$_2$O reaction.
- Reduced Net Heat Release: Because energy is absorbed to drive the dissociation of the weak acid, the overall net release of heat energy during the neutralisation reaction is less than what would be observed with a strong acid. Consequently, the enthalpy of neutralisation for weak acids is numerically smaller (less negative, e.g., -50 kJ/mol instead of -57.3 kJ/mol).
Let's visualize this:
- Strong Acid Neutralisation: All H$^+$ ions are pre-formed and ready to react, leading to maximum heat release from H$^+$ + OH$^-$ → H$_2$O.
- Weak Acid Neutralisation: The reaction of already dissociated H$^+$ with OH$^-$ releases heat. This heat then drives the further dissociation of the weak acid (HA → H$^+$ + A$^-$), which absorbs some of the released heat. The net effect is a lower overall heat release.
Comparison Table: Strong vs. Weak Acids
| Feature | Strong Acid Neutralisation | Weak Acid Neutralisation |
|---|---|---|
| Dissociation in Water | Complete | Partial |
| H$^+$ Availability | All H$^+$ ions are free | Not all H$^+$ ions are free; further dissociation needed |
| Energy for Dissociation | Negligible (already dissociated) | Required (endothermic step) |
| Primary Reaction | H$^+$ (aq) + OH$^-$ (aq) → H$_2$O (l) | HA (aq) + OH$^-$ (aq) → A$^-$ (aq) + H$_2$O (l) |
| Enthalpy of Neutralisation | Approximately -57.3 kJ/mol | Less negative (e.g., -50 kJ/mol to -55 kJ/mol) |
| Example Acid | Hydrochloric acid (HCl), Nitric acid (HNO$_3$) | Acetic acid (CH$_3$COOH), Formic acid (HCOOH) |
Practical Insights
- Heat Generation: When performing neutralisation reactions in a laboratory setting, the temperature increase observed when neutralising a strong acid with a strong base will generally be greater than that observed for a weak acid with a strong base (given equimolar amounts).
- Buffer Systems: The partial dissociation of weak acids is crucial for their role in buffer solutions, where they can absorb added H$^+$ or OH$^-$ ions to maintain pH stability. Learn more about buffer solutions from resources like Khan Academy.
Understanding the fundamental difference in dissociation between strong and weak acids is key to explaining the variation in their enthalpy of neutralisation values.
