5. Thermo Chemistry

The amount of energy or heat content of a substance. The energy is stored in chemical bonds (including kinetic and potential energy)

The difference between the enthalpy of the products and the enthalpy of the reactants at 298k and 1.00x105Pa dH = p - r

The temperature of the surroundings increases and dH is negative, energy is given out • combustion and neutralization reactions • products from an exothermic reaction have less stored energy, are more stable and have stronger bonds

The temperature of the surroundings decreases as dH is positive, energy is absorbed • the higher the nethalpy the less stable the substance

The enthalpy change that results when 1 mole of a compound reacts with oxygen at 298K and 1.00 x 105Pa. All reactants and products are in their standard states dH = r - p

Q = mcdT m = mass in g (volume in cm3) c = specific heat capacity of the substance dT = change in temperature in K dH = Q / n Q = heat energy change dH = heat energy change per mole

The technique used to measure enthalpy change Assumptions and sources of error • that no heat is lost to or gained from the surroundings • significant errors are associated with reactions • heat rises and dT under these conditions are reported to be less than the literature stipulates • Thermometers often have precision uncertainty of +-0.1C or greater

In a chemical reaction the total change in chemical potential energy must be equal to the energy lost or gained by the reaction system dH2 = 2dH1 - dH3

The enthalpy change that results when 1 mole of a compound is formed from its elements at 298K and 1.00 x 105 Pa, in the standard states dHf = p - r

The amount of energy required to break 1 mole of bonds in the gaseous state averaged across a range of compounds containing that bond • breaking bonds = endothermic • making bonds = exothermic dH = r - p

Ozone is decomposed more easily than oxygen • O2 requires 498kJ/mol to break • O3 requires 364kJ/mol to break

The energy change required to change 1 mole of atoms from their standard state to a gaseous state M(s) –> M(g) 0.5X2(g) –> X(g) dH > 0

The standard enthalpy change that occurs by the removal of 1 mole of electrons from 1 mole of atoms or positive ions in the gaseous phase IE1: M(g) –> M+(g) + e- IE2: M+(g) –> M2+(g) + e- dHIE > 0

The energy change that occurs when 1 mole of electrons joins to one mole of atoms in the gaseous phase X(g) + e- –> X-(g) dHEA < 0

The energy required to convert 1 mole of the solid compound into gaseous ions (association/ dissociation) MX(s) –> M+(g) + X-(g) dH > 0 Magnitude of Lattice Enthalpy • charge on the ions • greater charge, greater the attraction, greater lattice energy • size of the ions • larger the ions the further apart they are and the more separate the charges, the smaller the lattic energy

Diagrammatic representation of the formation of ionic compounds, starting from atoms, moving to gaseous atoms to gaseous ions to a solid compound dHf = dHatm + dHIE + dHEA + dHlat

The energy change that occurs when 1 mole of substance dissolves in excess solvent MX(s) –> M+(aq) + X-(aq)

The energy change that occurs when 1 mole of gaseous ions are added to water to form a dilute solution • value is always negative • if another solvent other than water is used it is enthalpy of solvation M+(g) –> M+(aq) X-(g) –> X-(aq)

The measure of the distribution of total available energy between the particles. The greater the disorder the particles, the great the entropy G > L > S • The entropy of a substance increases with temperature dS = entropy of p - entropy of r An increase in disorder produces a positive dS • decomposition reactions • changes in state • dissolution reactions • increases in temperature

The energy available to do work dG = dH - TdS • negative dG = spontaneous reaction • positive dG = not spontaneous reaction • For any element in its standard state dG = 0