3 Hess Law

Hess’s law states that total enthalpy change for a reaction is independent of the route by which the chemical change takes place Hess’s law is a version of the first law of thermodynamics, which is that energy is always conserved. 2H (g) + 2Cl(g) H2 + Cl2 2HCl (g) a b ΔH On an energy level diagram the directions of the arrows can show the different routes a reaction can proceed by. In this example one route is arrow ‘a’ The second route is shown by arrows ΔH plus arrow ‘b’ So a = ΔH + b And rearranged ΔH = a – b H+ (g) + Br – (g) H+ (aq) + Br – (aq) H (g) + Br (g) HBr (g) a c d ΔH Interconnecting reactions can also be shown diagrammatically. In this example one route is arrow ‘a’ plus ΔH The second route is shown by arrows ‘c’ plus arrow ‘d’ So a+ ΔH = c + d And rearranged ΔH = c + d – aThis Hess’s law is used to work out the enthalpy change to form a hydrated salt from an anhydrous salt. This cannot be done experimentally because it is impossible to add the exact amount of water and it is not easy to measure the temperature change of a solid. Often Hess’s law cycles are used to measure the enthalpy change for a reaction that cannot be measured directly by experiments. Instead alternative reactions are carried out that can be measured experimentally. Instead both salts are dissolved in excess water to form a solution of copper sulphate. The temperature changes can be measured for these reactions. ∆H reaction = Σ ∆fH products – Σ ∆fH reactants Mean Bond energies Definition: The Mean bond energy is the enthalpy needed to break the covalent bond into gaseous atoms, averaged over different molecules These values are positive because energy is required to break a bond. The definition only applies when the substances start and end in the gaseous state. We use values of mean bond energies because every single bond in a compound has a slightly different bond energy. E.g. In CH4 there are 4 CH bonds. Breaking each one will require a different amount of energy. However, we use an average value for the C-H bond for all hydrocarbons. In general (if all substances are gases) ∆H = Σ bond energies broken – Σ bond energies made ∆H values calculated using this method will be less accuate than using formation or combustion data because the mean bond energies are not exact Reaction profile for an EXOTHERMIC reaction Reaction profile for an ENDOTHERMIC reaction. Enthalpies of Combustion in a Homologous Series When comparing the enthalpies of combustion for successive members of a homologous series such as alkanes or alcohols there is a constant rise in the size of the enthalpies of combustion as the number of carbon atoms increases. As one goes up the homologous series there is a constant amount and type of extra bonds being broken and made e.g. 1C-C, 2C-H and 1.5 O=O extra bonds broken and 2 C=O and 2 O-H extra bonds made, so the enthalpy of combustion increases by a constant amount. Mr of alcohol experimental calculated mol combustion kJ Enthalpy of -1 If the results are worked out experimentally using a calorimeter the experimental results will be much lower than the calculated ones because there will be significant heat loss. There will also be incomplete combustion which will lead to less energy being released. Remember that calculated values of enthalpy of combustions will be more accurate if calculated from enthalpy of formation data than if calculated from average bond enthalpies. This is because average bond enthalpy values are averaged values of the bond enthalpies from various compounds.