energetics

• heat energy is released to surroundings - products will have less energy than the reactants

• heat energy is absorbed from the surroundings - products will have more energy than the reactants

• the amount of heat given out/taken in during a reaction carried out at constant pressure • symbol: ∆H • units: kJmol-1

• the heat released during the complete combustion of one mole of that substance • always a -ve value

• a measure of the average bond strength for a particular covalent bond • measured in compounds in the gaseous state so the values only apply in calculations where all reactants and all products are in the gaseous state • ΔH = [sum of energy breaking bonds] + [sum of energy forming bonds]

• to break a covalent bond between 2 atoms, energy must be supplied • bond breaking is an endothermic process (+ve value)

• when a covalent bond between 2 atoms is made, energy is released • bond forming is an exothermic process (-ve value)

• bond energies are average values so values calculated using them are not specific to the molecules used • thus, the actual value is likely to have a difference from the calculated value of over 10%. • calculating the enthalpy change can only be done for reactions which take place entirely in the gaseous state • to obtain more accurate values for the ΔH, calorimetry experiments can be used

Eh = cmΔT • Eh = energy gained by water (in kJ) • c = specific heat capacity of water (4.2) • m = mass of water (in kg) 1l of water = 1kg of water • ΔT = change in temp of water ΔH =Eh/n (n=no. of moles), must have a +/- sign

• a measured volume of a known conc of solution is placed in a polystyrene cup (an insulator so less heat is lost/gained to/from the surroundings) • the temp of the solution is taken • a measured volume of a known conc of another solution is added • the mixture is stirred and the highest temp rise is found. -ΔT, can then be calculated and used in the formula

• at room temp and pressure, the density of water is 1.00g cm−3 and the density of a solution is assumed to be the same even though it will not be quite the same • the specific heat capacity is particular to the substance being heated. As most of an aqueous mixture is water, it can be assumed that the specific heat capacity of a solution is the same as that for pure water. • as water is the major component, it can be assumed that only the water has absorbed any heat and that the calorimeter doesn’t absorb any energy

• incomplete combustion: especially true when larger molecules are used as the fuel • heat loss to surroundings is more likely in an open setting: heat can be lost before it reaches the calorimeter, and if the calorimeter not insulated • heat is lost to the calorimeter: since it is a conductor, some heat will be absorbed by the calorimeter, (this error can be removed by calculating the energy change of the calorimeter and the water and adding these values.