Energy changes in reactions — enthalpy, exo/endothermic reactions, Hess's law, calorimetry and bond energies — exam-focused for the BIEK / Sindh Board paper. Read it, or open the interactive lecture and run the calorimetry simulator.
1 — System & surroundings
- System — the part of the universe under study (e.g. the reaction mixture).
- Surroundings — everything else that can exchange energy/matter with it.
| System | Exchanges with surroundings |
|---|
| Open | energy AND matter (a beaker) |
| Closed | energy only (a sealed flask) |
| Isolated | neither (a thermos/vacuum flask) |
2 — Internal energy, enthalpy & state functions
- Internal energy (U) — the total energy (kinetic + potential) stored in a system.
- Enthalpy (H) — the heat content of a system at constant pressure: H = U + PV.
- State function — a property that depends only on the present state, not the path taken (U, H, T, P all are).
We measure the change: the enthalpy change ΔH = H(products) − H(reactants).
3 — Exothermic & endothermic reactions
| Exothermic | Endothermic |
|---|
| Heat | released to surroundings | absorbed from surroundings |
| ΔH sign | negative (−) | positive (+) |
| Temperature | rises | falls |
| Example | combustion, neutralisation | photosynthesis, thermal decomposition |
4 — Enthalpy (energy) diagrams
On an enthalpy diagram, the vertical axis is enthalpy. For an exothermic reaction the products lie below the reactants (energy is given out); for an endothermic reaction the products lie above (energy is taken in).
The size of the arrow from reactants to products is |ΔH|.
5 — Standard enthalpy changes
Measured at standard conditions (298 K, 1 atm, 1 mol). Common ones:
| ΔH° | Defined as the enthalpy change when… |
|---|
| Formation (ΔH°f) | 1 mol of a compound forms from its elements |
| Combustion (ΔH°c) | 1 mol of a substance burns completely in O₂ |
| Neutralisation | 1 mol of water forms from acid + base (≈ −57 kJ) |
| Solution | 1 mol of a solute dissolves |
6 — Hess's law
- Hess's law — the total enthalpy change for a reaction is the same whatever route is taken, provided the initial and final states are the same.
This follows from H being a state function. It lets us find a ΔH that is hard to measure directly by adding known steps.
Adding stepsΔH(overall) = ΔH₁ + ΔH₂ + ΔH₃ …
7 — Calorimetry
Heat is measured with a calorimeter. The heat gained or lost by the water is:
Heatq = m c ΔT
m = mass · c = specific heat (water = 4.18 J g⁻¹ °C⁻¹) · ΔT = temperature change
calorimetry
100 g of water rises by 6 °C. How much heat was released?
q = mcΔT = 100 × 4.18 × 6 = 2508 J = 2.51 kJ
8 — Bond energy & ΔH from bonds
Breaking a bond absorbs energy (endothermic); making a bond releases energy (exothermic).
From bond enthalpiesΔH = Σ(bonds broken) − Σ(bonds formed)
If more energy is released forming new bonds than is used breaking old ones, ΔH is negative (exothermic).
9 — First law of thermodynamics
- First law (conservation of energy) — energy can neither be created nor destroyed, only transferred or transformed.
Change in internal energyΔU = q + w
q = heat added to the system · w = work done on the system
10 — Worked numericals
heat released
250 g of water cools by 8 °C. Heat lost?
q = 250 × 4.18 × 8 = 8360 J = 8.36 kJ
Hess's law
If A→B is −100 kJ and B→C is −50 kJ, find ΔH for A→C.
ΔH = −100 + (−50) = −150 kJ
ΔH from bonds
Bonds broken = 680 kJ, bonds formed = 800 kJ. Find ΔH.
ΔH = 680 − 800 = −120 kJ (exothermic)
11 — Applications
- Fuels & food — enthalpy of combustion gives the energy content (calorific value).
- Hand-warmers (exothermic) and cold-packs (endothermic).
- Designing reactions that release or store energy efficiently.
12 — Exam recap
- System/surroundings; open, closed, isolated.
- Internal energy, enthalpy, state functions; ΔH.
- Exothermic (−ΔH) vs endothermic (+ΔH); enthalpy diagrams.
- Standard enthalpy changes; Hess's law.
- Calorimetry q = mcΔT; ΔH from bond energies.
- First law: ΔU = q + w.