Revise: Energy Changes in Chemical Reactions
Energy changes, exothermic and endothermic reactions, enthalpy diagrams, standard enthalpy, and using Hess's law to calculate enthalpy changes indirectly.
Exothermic: releases heat, ΔH < 0. Endothermic: absorbs heat, ΔH > 0.
H(products)=320, H(reactants)=500 → ΔH=−180 kJ/mol, exothermic.
q = mcΔT measures heat from a temperature change.
100 g water, ΔT=20°C → q = 8360 J.
Products drawn below reactants = exothermic; above = endothermic.
Combustion diagrams always show products well below reactants.
Total ΔH is the same regardless of the reaction pathway.
C→CO→CO₂ two-step ΔH sums to the same −393.5 kJ/mol as C→CO₂ directly.
Chemical bonds store energy; reactions convert it to other forms.
A glow stick converts chemical energy directly to light.
Enthalpy (H) is heat content at constant pressure — measured via calorimetry.
50 g water, ΔT=12°C → q = 2508 J.
ΔH = H(products) − H(reactants).
Photosynthesis absorbs energy — endothermic, ΔH positive.
The vertical gap between reactant/product levels represents |ΔH|.
A downward arrow on the diagram = exothermic reaction.
ΔH° is measured under standard conditions (1 atm, stated temperature) for fair comparison.
Fuels are compared by their standard enthalpy of combustion.
Enthalpy is a state function — depends only on start/end, not the path.
Two-step and one-step combustion of carbon give the same total ΔH.
Reversing a reaction flips the sign of ΔH; scaling it scales ΔH by the same factor.
Combining 3 known combustion reactions gives ΔHf(CH₄) = −74.8 kJ/mol.