Chemical energetics

Enthalpy Changes and Standard DefinitionsSign up

know that the enthalpy change, delta H, is the heat energy change measured at constant pressure and that standard conditions are 100 kPa and a specified temperature, usually 298 K · know that, by convention, exothermic reactions have a negative enthalpy change and endothermic reactions have a positive enthalpy change · be able to construct and interpret enthalpy level diagrams, showing exothermic and endothermic enthalpy changes · know the definition of standard enthalpy change of: (i) reaction, delta_rH; (ii) formation, delta_fH; (iii) combustion, delta_cH; (iv) neutralisation, delta_neutH; (v) atomisation, delta_atH

Calorimetry and Hess's LawSign up

be able to use experimental data to calculate: (i) energy transferred in a reaction recalling and using the expression: energy transferred (J) = mass (g) x specific heat capacity (J g^-1 °C^-1) x temperature change (°C); (ii) enthalpy change of the reaction in kJ mol^-1. This will be limited to experiments where substances are mixed in an insulated container and combustion experiments using a suitable calorimeter. · know Hess's Law and be able to apply it to: (i) constructing enthalpy cycles; (ii) calculating enthalpy changes of reaction using data provided, or data selected from a table or obtained from experiments · CORE PRACTICAL 2: Determination of the enthalpy change of a reaction using Hess's Law · be able to evaluate the results obtained from experiments and comment on sources of error and uncertainty and any assumptions made in the experiments. Students will need to consider experiments where substances are mixed in an insulated container and combustion experiments using, for example, a spirit burner and be able to draw suitable graphs and use cooling curve corrections.

Bond Energies and Enthalpy CalculationsSign up

understand the terms 'bond enthalpy' and 'mean bond enthalpy', and be able to use bond enthalpies to calculate enthalpy changes, understanding the limitations of this method · be able to calculate mean bond enthalpies from enthalpy changes of reaction · understand that bond enthalpy data gives some indication about which bond will break first in a reaction, how easy or difficult it is and therefore how rapidly a reaction will take place at room temperature

Lattice Energy, Born-Haber Cycles and Enthalpy of SolutionSign up

be able to define the terms: (i) standard enthalpy change of atomisation, delta_atH; (ii) electron affinity; (iii) lattice energy (as the exothermic process for the formation of one mole of an ionic solid from its gaseous ions) · be able to construct Born-Haber cycles and carry out related calculations · understand that a comparison of the experimental lattice energy value (from a Born-Haber cycle) with the theoretical value (obtained from electrostatic theory) in a particular compound indicates the degree of covalent bonding · understand that polarisation of anions by cations leads to some covalency in an ionic bond, based on evidence from the Born-Haber cycle · be able to define the terms 'enthalpy change of solution, delta_solH' and 'enthalpy change of hydration, delta_hydH of an ion' · be able to use energy cycles and energy level diagrams to calculate the enthalpy change of solution of an ionic compound, using enthalpy change of hydration and lattice energy · understand the effect of ionic charge and ionic radius on the values of enthalpy change of hydration and the lattice energy of an ionic compound · be able to use entropy and enthalpy changes of solution values to predict the solubility of ionic compounds and discuss trends in the solubility of ionic compounds covered in Unit 2

Entropy and Gibbs Free EnergySign up

understand that, since endothermic reactions can occur spontaneously at room temperature, enthalpy changes alone do not control whether reactions occur · understand entropy as a measure of disorder of a system in terms of the random dispersal of molecules and of energy quanta between molecules · understand that the entropy of a substance increases with temperature, that entropy increases as solid -> liquid -> gas and that perfect crystals at zero kelvin have zero entropy · be able to interpret the natural direction of change as being in the direction of increasing total entropy (positive entropy change), including gases spread spontaneously through a room · understand why entropy changes occur during: (i) changes of state; (ii) dissolving of a solid ionic lattice; (iii) reactions in which there is a change in the number of moles from reactants to products · understand that the total entropy change of any reaction is the sum of the entropy change of the system and the entropy change of the surroundings, summarised by the expression: delta S_total = delta S_system + delta S_surroundings · be able to calculate the entropy change of the system for a reaction, delta S_system, given the entropies of the reactants and products · be able to calculate the entropy change in the surroundings, and hence delta S_total, using the expression delta S_surroundings = -delta H / T · understand that the feasibility of a reaction depends on: (i) the balance between delta S_system and delta S_surroundings, so that even endothermic reactions can occur spontaneously at room temperature; (ii) temperature, as higher temperatures decrease the magnitude of delta S_surroundings so its contribution to delta S_total is less. Students should be able to calculate the temperature at which a reaction is feasible. Students may also use delta G = delta H - T delta S_system in answers, although this approach is not a requirement of the specification. · understand that reactions can occur as long as delta S_total is positive even if one of the other entropy changes is negative · understand and distinguish between the concepts of thermodynamic stability and kinetic stability