R1.2 Energy cycles in reactionsIB Chemistry HL: Flashcards

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ΔH from bond enthalpies?

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ΔH from bond enthalpies?: Σ(bonds broken) − Σ(bonds formed), all species gaseous.
State Hess's law.: The enthalpy change is independent of the pathway between the initial and final states.
Define standard enthalpy of formation.: Enthalpy change when one mole of a compound forms from its elements in their standard states under standard conditions.
Define standard enthalpy of combustion.: Enthalpy change when one mole of a substance burns completely in oxygen under standard conditions.
ΔHf⦵ of O₂(g)?: Zero: it is an element in its standard state.
Formula using ΔHf⦵ data?: ΔH⦵ = Σ ΔHf⦵(products) − Σ ΔHf⦵(reactants).
Formula using ΔHc⦵ data?: ΔH⦵ = Σ ΔHc⦵(reactants) − Σ ΔHc⦵(products).
Why is ΔHf⦵ of propane found indirectly?: Carbon and hydrogen do not react directly to give only propane.
Define lattice enthalpy (IB convention).: Enthalpy change when one mole of an ionic solid separates into gaseous ions; endothermic.
Define first electron affinity.: Enthalpy change when one mole of gaseous atoms each gains an electron: X(g) + e⁻ → X⁻(g).
Why is the second electron affinity of oxygen endothermic?: The electron is added to a negative ion and must overcome repulsion.
Steps in the Born–Haber cycle for MgCl₂?: Atomise Mg and 2Cl, IE₁ + IE₂ of Mg, 2 × EA of Cl, then lattice formation.
Why is the lattice enthalpy of MgCl₂ much larger than that of a hypothetical MgCl?: Mg²⁺ has a higher charge and smaller radius, so attraction to Cl⁻ is stronger.