Standard Enthalpy Change of Neutralisation
What is Enthalpy and Neutralisation?
To understand $\Delta H_n^\circ$, we first need to break down the words. Enthalpy is a measure of the total heat energy in a chemical system. You can think of it like a battery's stored energy; in a reaction, this energy can be released or absorbed. The symbol $\Delta H$ represents the change in this enthalpy. If $\Delta H$ is negative, heat is released to the surroundings (an exothermic reaction), and if it's positive, heat is absorbed from the surroundings (an endothermic reaction).
Neutralisation is the chemical reaction between an acid and a base (alkali is a soluble base). The products are always a salt and water. For example, when you mix hydrochloric acid ($HCl$) with sodium hydroxide ($NaOH$), you get sodium chloride ($NaCl$), which is table salt, and water ($H_2O$).
The "standard" in $\Delta H_n^\circ$ means the measurement is taken under standard conditions: a pressure of $100\ kPa$ (which is close to normal atmospheric pressure) and a temperature of $298\ K$ ($25^\circ C$). The little $^\circ$ symbol reminds us of these conditions.
The Core Concept of $\Delta H_n^\circ$
The official definition is: The enthalpy change when one mole of water is formed from the reaction between an acid and an alkali under standard conditions. Let's unpack this.
It specifies "one mole of water". A mole is just a chemist's way of counting a very large number of particles ($6.02 \times 10^{23}$, to be exact). So, we are not looking at the heat for the entire reaction, but specifically for the amount of heat released when exactly one mole of water molecules is produced.
Since the reaction is exothermic, the value of $\Delta H_n^\circ$ is always negative. For reactions involving strong acids and strong bases, the value is remarkably constant at about $-57\ kJ\ mol^{-1}$. This is because the same net ionic reaction is always taking place.
Comparing Strong and Weak Acids and Bases
Not all acids and bases are created equal. Strong acids (like $HCl$, $HNO_3$, $H_2SO_4$) and strong bases (like $NaOH$, $KOH$) fully dissociate (break apart) into their ions in water. This means the neutralisation reaction is simply the combination of $H^+$ and $OH^-$ ions.
Weak acids (like ethanoic acid, $CH_3COOH$) and weak bases (like ammonia, $NH_3$) only partially dissociate. The neutralisation enthalpy is less exothermic (e.g., around $-55\ kJ\ mol^{-1}$ for a weak acid/strong base reaction) because some energy is used to first break the bonds within the weak acid or base molecules before the $H^+$ and $OH^-$ ions can combine.
| Acid Type | Base Type | Example Reaction | Approx. $\Delta H_n^\circ\ (kJ\ mol^{-1})$ |
|---|---|---|---|
| Strong | Strong | $HCl + NaOH \rightarrow NaCl + H_2O$ | $-57$ |
| Weak | Strong | $CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O$ | $-55$ |
| Strong | Weak | $HCl + NH_3 \rightarrow NH_4Cl$ | $-52$ |
Measuring $\Delta H_n^\circ$ in the Laboratory
How do scientists measure this energy change? A common school experiment involves a simple calorimeter, often just an insulated polystyrene cup. Let's see how we can find $\Delta H_n^\circ$ for the reaction between hydrochloric acid and sodium hydroxide.
Step 1: Measure out a known volume (e.g., $25\ cm^3$) of $1.0\ mol\ dm^{-3}$ $HCl$ and pour it into the insulated cup. Record its initial temperature.
Step 2: Measure an equal volume ($25\ cm^3$) of $1.0\ mol\ dm^{-3}$ $NaOH$. Record its initial temperature (it should be the same as the acid).
Step 3: Quickly add the alkali to the acid, stir gently with the thermometer, and record the highest temperature reached.
Calculations:
1. Temperature change, $\Delta T = T_{final} - T_{initial}$.
2. Total volume of solution = $50\ cm^3$. Assuming the density of the solution is $1\ g\ cm^{-3}$, the total mass, $m = 50\ g$.
3. Heat energy released, $q = m \times c \times \Delta T$, where $c$ (specific heat capacity of water) is $4.18\ J\ g^{-1}\ K^{-1}$.
4. Moles of water produced: Since we used $0.025\ dm^3$ of $1.0\ mol\ dm^{-3}$ solutions, we have $0.025$ moles of acid and alkali. The reaction produces $0.025$ moles of water.
5. Enthalpy change for the reaction: $\Delta H = -q / moles\ of\ water$. The negative sign is because heat is released. Convert the answer from Joules to kiloJoules ($kJ$).
Real-World Applications of Neutralisation Energy
The heat released during neutralisation isn't just a topic for textbooks; it has practical uses all around us.
1. Self-Heating Cans: Some food and beverage cans have a separate compartment at the bottom containing a solid salt and a sealed bag of water. When you push a button on the can, you break the seal, allowing the water to mix with a quicklime ($CaO$), which reacts exothermically with water (a type of base-forming reaction), heating the drink or meal without any external heat source.
2. Treating Indigestion: Indigestion is often caused by too much stomach acid ($HCl$). Antacid tablets contain weak bases like calcium carbonate ($CaCO_3$) or magnesium hydroxide ($Mg(OH)_2$). When you swallow one, a neutralisation reaction occurs in your stomach, producing salt and water and relieving discomfort. The exothermic nature of this reaction is why you sometimes feel a warming sensation.
3. Industrial Waste Treatment: Factories that use acids in their processes cannot simply dump acidic waste into rivers. They use large-scale neutralisation, often with cheap alkalis like slaked lime ($Ca(OH)_2$), to make the waste water safe before disposal. The heat released is a sign that the reaction is proceeding effectively.
Important Questions
The answer lies in the net ionic equation. For all strong acids and bases, the actual chemical change is identical: $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$. The spectator ions (like $Na^+$ and $Cl^-$) do not participate in the energy-changing step. Since the same bonds are being formed in every case (the O-H bonds in water), the energy released is very consistent.
Weak acids like ethanoic acid ($CH_3COOH$) exist mainly as whole molecules in solution. Only a tiny fraction are dissociated into $H^+$ ions. During neutralisation, the weak acid molecule must first be broken apart, which requires an input of energy (an endothermic step). This energy cost partially offsets the large amount of heat released when the $H^+$ and $OH^-$ ions combine. The net result is a less negative $\Delta H_n^\circ$.
If you use more concentrated solutions (e.g., $2.0\ mol\ dm^{-3}$ instead of $1.0\ mol\ dm^{-3}$), you will have more moles reacting in the same volume of water. This means more heat is released, leading to a larger temperature rise ($\Delta T$). However, the value of $\Delta H_n^\circ$ (the energy per mole of water) should remain roughly the same, as long as the solutions are not so concentrated that their properties change significantly.
Footnote
[1] $\Delta H_n^\circ$: The standard enthalpy change of neutralisation. The standard state is defined as a pressure of $100\ kPa$ and a temperature of $298\ K$.
[2] Enthalpy (H): A thermodynamic quantity equivalent to the total heat content of a system.
[3] Exothermic: A chemical reaction that releases energy by light or heat.
[4] Dissociation: The process in which a compound separates into particles such as atoms, ions, or molecules.
[5] Calorimeter: A device used to measure the heat involved in a chemical reaction or physical change.