Topic 6.6 Notes – Introduction to Enthalpy of Reaction

1. What the Enthalpy of Reaction Is

Enthalpy (H) is the heat content of a system at constant pressure.
In AP Chemistry, reactions are assumed to occur at constant pressure, so:

$$\Delta H = q_{\text{reaction}}$$

That means the enthalpy change (ΔH) equals the heat absorbed or released.

Important: you can’t measure absolute enthalpy. Only changes in enthalpy.

Interpreting the Sign of ΔH

If ΔH is negative, heat is released. If positive, heat is absorbed. On tests, they often describe a temperature increase in the solution and expect you to recognize that the reaction was exothermic.

2. Where ΔH Comes From at the Particle Level

Reactions rearrange atoms by breaking and forming bonds.

• Breaking bonds requires energy → always endothermic.
• Forming bonds releases energy → always exothermic.

The overall ΔH depends on the balance:

• Energy released forming bonds > energy required to break bonds → exothermic
• Energy required to break bonds > energy released forming bonds → endothermic

Why Temperature Changes

Reactants and products have different chemical potential energy because their bonding is different.

That energy difference turns into a change in kinetic energy of particles.

• More kinetic energy → higher temperature
• Less kinetic energy → lower temperature

So in an exothermic reaction, the particles in the surroundings move faster. In an endothermic reaction, they slow down.

You do not need to distinguish enthalpy from internal energy on the AP exam. At constant pressure, just treat ΔH as the heat of reaction.

3. Molar Enthalpy of Reaction

The molar enthalpy of reaction (ΔH_rxn) is the heat released or absorbed per mole of reaction as written in the balanced equation.

Units: kJ/mol

That “per mole” refers to the entire balanced reaction.

Example idea:

$$\ce{N2(g) + 3H2(g) -> 2NH3(g)} \qquad \Delta H = -92\text{ kJ/mol}$$

This means:

• When 1 mol of reaction occurs (1 mol of $\ce{N2}$ reacts with 3 mol of $\ce{H2}$), 92 kJ is released.
• If 2 mol of $\ce{N2}$ react, twice as much heat is released.

If you multiply the whole equation by 3, ΔH also multiplies by 3. Students forget this constantly.

4. Calculating Heat Using $q = n\Delta H$

This is the math skill they love testing.

$$q = n \times \Delta H$$

• $q$ = heat (kJ or J)
• $n$ = moles of reaction (usually limiting reactant)
• $\Delta H$ = kJ/mol

Step-by-Step

1. Write and balance the equation.
2. Convert given quantities to moles.
3. Use mole ratios if needed.
4. Multiply moles by ΔH.
5. Interpret the sign.

Quick Example

$$\ce{2SO2(g) + O2(g) -> 2SO3(g)} \qquad \Delta H = -198\text{ kJ/mol}$$

Suppose 0.75 mol $\ce{SO2}$ reacts.

From the equation:
2 mol $\ce{SO2}$ corresponds to 1 mol reaction.

$$0.75 \text{ mol SO2} \times \frac{1 \text{ mol rxn}}{2 \text{ mol SO2}} = 0.375 \text{ mol rxn}$$

$$q = 0.375 \times (-198) = -74.3 \text{ kJ}$$

Negative means 74.3 kJ released.

On quizzes, mistakes usually happen in the mole ratio step, not the multiplication.

You might also solve for moles:

$$n = \frac{q}{\Delta H}$$

5. Heat Flow and Thermal Equilibrium

After a reaction, the products and surroundings may be at different temperatures.

Energy flows until thermal equilibrium is reached.

• Exothermic: products initially hotter → heat flows outward.
• Endothermic: products initially cooler → heat flows inward.

Key idea:
ΔH measures total energy transferred, not just temperature change. A small temperature change could still involve a large ΔH if a large amount of substance reacts.