Heats of Reactions Lecture 2: Enthalpy, Hess's Law, Entropy

HEATS OF
REACTIONS
Lecture 2

Recap from the first lecture
•All chemical reactions absorb heat to reach transition state
where the bonds of the reactants are being broken.
•All chemical reactions loose the absorbed heat –but to different
extents.
•Some reactions loose very little of this absorbed heat, and the
reaction would be referred to as endothermic reactions.
•Some reactions loose all of the absorbed heat and even more –
they referred to as exothermic reactions.

Reminder of energy diagrams
The change in heat energy (∆H) in a chemical reaction is called enthalpy
∆H ∆H
∆H = positive (+) = endothermic ∆H = negative (-) = exothermic
Endothermic exothermic
Enthalpy is an extensive property, meaning that its value is affected by the amount of material

Practice Question:
If the standard enthalpy change for phosphorus (P4) burning is ΔHreaction = -
3013 kJ.
How much heat is evolved when 266 g of white phosphorus (P4) burns in air?
Molar mass of P = 30.9737
Molar mass of P4= 123.896
Moles of P4= 266/123.896 = 2.15 moles
∆H = -3013 kJ for 1 mole
= -3013 ×2.15
= -6477.95kJ

•2 H2(g) + O2(g) →2 H2O(g) ΔH = –483.6 kJ
What is the enthalpy change when 178 g of H2O(g) are
produced?
Remember ΔH = –483.6 kJ is for 2 moles of H2O(g),
so ΔH for 1mol = –241.8 kJ
Molar mass or H2O(g) = 18 g/mol
Moles of H2O(g) = 178/18 = 9.88 moles
∆H = 9.88 x –241.8
= -2388,98 kJ