Chemical Energetics

Exothermic & endothermic reactions

Some reactions warm the room; some cool it. That difference is the whole of Core energetics — the classification, the reason, and how to spot it from a temperature change.

Reading the temperature

If a reaction mixture rises in temperature during the reaction, energy has left the system — the reaction is exothermic. If the mixture falls in temperature, energy has been drawn in — the reaction is endothermic.

Where they show up

Reaction pathway diagrams

A pathway diagram is a picture of the energy story. Two horizontal lines — one for the reactants, one for the products — connected by a curve. Which line is higher tells you the reaction type immediately.

Exothermic and endothermic pathways

On an exothermic pathway the products sit lower on the y-axis than the reactants: the reaction releases energy to the surroundings and the mixture loses energy overall. On an endothermic pathway the products sit higher than the reactants: the reaction absorbs energy from the surroundings and the mixture ends up with more energy than it started with.

ExtendedThe sign of ΔH

Enthalpy change is written as ΔH (“delta H”). By convention its sign gives the direction of energy flow: ΔH < 0 is exothermic, ΔH > 0 is endothermic. A reaction with ΔH = −184 kJ releases 184 kJ per mole to the surroundings; one with ΔH = +103 kJ absorbs 103 kJ per mole from them. The size of the number tells you how much energy is transferred; the sign tells you in which direction.

ExtendedThe labelled diagram

A reaction must climb over an energy barrier — the activation energy Ea — before it can proceed. Ea is measured from the reactant line up to the peak of the curve; ΔH is measured from the reactant line to the product line. The endothermic diagram is the mirror image: products above reactants, the ΔH arrow pointing upwards, and Ea still measured from the reactant line to the peak, which is now above the products.

Bond breaking & forming

The sign of ΔH isn’t magic — it comes from bonds. Every reaction breaks the bonds in the reactants (costs energy) and forms new bonds in the products (releases energy). The net difference is ΔH.

ExtendedThe two directions

Breaking bonds is endothermic: every chemical bond holds atoms together, and pulling them apart requires energy from outside. Nothing about a bond breaks “for free”. Making bonds is exothermic: when atoms come together to form a bond they fall into a more stable, lower-energy state, and the energy they lose is released to the surroundings as heat.

ExtendedThe net outcome

A reaction is exothermic overall when more energy is released forming product bonds than is taken in breaking reactant bonds, so ΔH is negative. It is endothermic overall when more energy is taken in breaking reactant bonds than is released forming product bonds, so ΔH is positive. Every ΔH value you calculate comes from this single comparison — one number in, one number out, and the difference is what you feel.

Bond energy calculations

Given a table of bond energies, you can predict ΔH for any reaction. The method is a four-step accountancy: draw the bonds, sum the ones broken, sum the ones formed, subtract.

ExtendedWorked example: H₂ + Cl₂ → 2HCl

Bond energies (kJ/mol): H–H = 436, Cl–Cl = 242, H–Cl = 431. Bonds broken: 1 × H–H + 1 × Cl–Cl, so energy in = 436 + 242 = 678 kJ. Bonds formed: 2 × H–Cl, so energy out = 2 × 431 = 862 kJ. ΔH = 678 − 862 = −184 kJ/mol — negative, so the reaction is exothermic, as expected.

ExtendedWorked example: finding an unknown bond energy

Hydrogen bromide decomposes, 2HBr → H₂ + Br₂, endothermic with ΔH = +103 kJ/mol. Given H–Br = 366 and H–H = 436 kJ/mol, find the Br–Br bond energy. Bonds broken: 2 × H–Br = 732 kJ. Bonds formed: 436 + Br–Br. Substitute into ΔH = in − out: +103 = 732 − (436 + Br–Br), so Br–Br = 732 − 436 − 103 = +193 kJ/mol. Bond energies are always positive, and +193 kJ/mol is chemically reasonable for a halogen–halogen bond.