Topic 6 and 15 – Energetics
Bond enthalpy: The average enthalpy change of breaking a bond between gaseous atoms into its constituent gaseous atoms.
Born-Haber cycle: Energy cycles for the formation of ionic compounds. If there is little agreement between the theoretical and experimental values, this could indicate a degree of covalent character. ΔHf = ΔHat(1) + ΔHat(2) + ΔHIE(1) + ΔHEa(2) – ΔHLatt
Calorimetry: A device for measuring enthalpy changes for reactions. In a simple calorimeter all the heat evolved in an exothermic reaction is used to raise the temperature of a known mass of water.
Electron affinity: Enthalpy change when an e- is added to an isolated atom in the gaseous state.
Endothermic: A reaction in which energy is absorbed. ΔH is +. Reactants more stable than products.
Enthalpy: The internal energy stored in the reactants. Only changes in enthalpy can be measured.
Entropy: A measure of the disorder of a system. The absolute entropy can be found in terms of the probability of a state existing. Things causing entropy to increase: 1) increase of number of moles of gaseous molecules; 2) change of state from solid to liquid or liquid to gas
Exothermic: A reaction in which energy is evolved. ΔH is –. Products more stable than reactants.
Gibb’s free energy: Must be –ve for reaction to be spontaneous. ΔG = ΔH – TΔS
Heat: A measure of the total energy in a given amount of substance.
Hess’ law: Enthalpy change for a reaction depends only on difference between enthalpy of products and enthalpy of reactants independent of pathway.
Lattice enthalpy: The endothermic process of converting a crystalline solid into its gaseous ions, or the reverse exothermic process. The lattice enthalpy increases with decreasing size of the ions and increasing charge.
Spontaneous: A reaction that has a natural tendency to occur.
Standard conditions: 298 K and 1 atm.
Temperature: A measure of the average kinetic energy.