Energetics (AS Level) flashcards

Topic 4: Energetics

Enthalpy change, ΔH, and standard conditions

ΔH is the heat energy change measured at constant pressure. Standard conditions: 100 kPa and a stated temperature, e.g. ΔH⦵₂₉₈ for 298 K.

Standard enthalpy of combustion, ΔcH⦵

The enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions, all reactants and products in their standard states.

Standard enthalpy of formation, ΔfH⦵

The enthalpy change when 1 mole of a compound is formed from its elements under standard conditions, all reactants and products in their standard states.

Hess’s law

The enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same.

Mean bond enthalpy

The enthalpy change to break 1 mole of a stated covalent bond in gaseous molecules, averaged over a range of compounds.

Calorimetry: the two equations, with units

q = mcΔT (m in g of water or solution, c in J g⁻¹ K⁻¹, ΔT in K; q in J). ΔH = −q ÷ n, with q in kJ and n the moles of the limiting reactant.

Hess’s law shortcuts with formation and with combustion data

Formation: ΔH = ΣΔfH(products) − ΣΔfH(reactants). Combustion: ΔH = ΣΔcH(reactants) − ΣΔcH(products).

Required practical 2: why extrapolate the cooling line back to the time of mixing?

To estimate the maximum temperature change, correcting for heat lost to the surroundings while the temperature was rising.

Why does a bond-enthalpy ΔH differ from a Hess’s law value?

Mean bond enthalpies are averages over many compounds, not the values in these molecules; and the method assumes all species are gases, ignoring enthalpy changes of state.

What is ΔfH⦵ of O₂(g), and of Br₂(g)?

O₂(g): zero, as an element in its standard state. Br₂(g): not zero, because bromine’s standard state is liquid.