Unit 1: Section 6 - Thermodynamics

The heat energy transferred in a reaction at constant pressure

ΔfH - the enthalpy change when 1 mole of a substance is formed from its constituent elements with all substances in their standard states - exothermic for most substances 2Na(s) + 1/2O2(g) -> Na2O(s)

ΔcH - When 1 mole of a substance undergoes complete combustion in oxygen with all substances in their standard states - exothermic H2 (g) + 1/2O2(g) -> H2O(l)

ΔneutH - when 1 mole of water is formed in a reaction betweeen an acid and alkali under standard conditions - exothermic 1/2H2SO4(aq) + NaOH(aq) -> 1/2Na2SO4(aq) + H2O(l)

ΔieH 1st ionisation enthalpy - enthalpy change whe each atom in 1 mole of gaseous atoms loses 1 electron to form 1 mole of gaseous 1+ ions - endothermic Mg(g) -> Mg+(g) +e- 2nd ionisation enthalpy - enthalpy change when each ion in 1 mole of gasous 1+ ions loses 1 electron to form 1 mole of gaseous 2+ ions - endothermic Mg+(g) -> Mg2+ +(g) +e-

ΔeaH 1st electron affinity - enthalpy change when each atom in 1 mole of gaseous atoms gains 1 electron to form 1 mole of gaseous 1- ions - exothermic for many non-metals O(g) + e- -> O-(g) 2nd electron affinity - enthalpy change when each ion in 1 mole of gasous 1- ions gains 1 electron to form 1 mole of gasoues 2- ions - endothermic O-(g) + e- -> O2-(g)

ΔatH - when 1 mole of gaseous atoms is produced from an element in its standard state - endothermic 1/2I2(s) -> I(g)

ΔhydH When 1 mole of gaseous atoms dissolve in water (become hydrated) - exothermic Mg2+(g) + aq -> Mg2+(aq)

ΔsolH When 1 mole of an ionic solid dissolved in an infinitely dilute solution so that the dissolved ions are well separated and don’t interact with each other - varies MgCl2(s) + aq -> Mg2+(aq) + 2Cl-(aq)

ΔdisH When 1 mole of covalent bonds is broken in the gaseous state - endothermic I2(g) -> 2I (g) 2 x enthalpy of atomisation

ΔLEFH When 1 mole of a solid ionic compound is formed from its consistuent ions in the gas phase - exothermic Mg2+(g) + 2Cl-(g) -> MgCl2(s)

When 1 mole of a solid ionic compound is broken up into its constituent ions in the gas phase - endothermic MgCl2(s) -> Mg2+(g) + 2Cl-(g)

ΔvapH Whe 1 mole of a liquid is turned into a gas - endothermic H2O(l) -> H2O(g)

ΔfusH When 1 mole of a solid is turned into a liquid - endothermic Mg(s) -> Mg(l)

Used to calculate lattice enthalpies

Assumes that all the ions are spherical, and have their charge evenly distributed around them

• Ions are not perfectly spherical • Polarisation often occurs when small positive ions or large negative ions are involved • Some lattices are not regular and the crystal structure can differ

• The lattice enthalpy found experimentally is often not the theoretical lattice enthalpy • This is evidence that most ionic compounds have some covalent character

• The positive and negative ions in a lattice aren’t exactly spherical • Positive ions polarise neighbouring negative ions to different extents • The more polarisation there is, the more covalent the bonding will be

• Solubility - makes the compound less soluble • General conductivity

• The bonds between the ions break to give free ions - endothermic • Bonds between the ions and the water are made - exothermic The enthalpy change of solution is the overall effect on the enthalpy of these 2 things

If the energy released is roughly the same, or greater than the energy taken in

• A measure of the number of ways that particles can be arranged and the number of ways the energy can be shared out between the particles in a system • The more disordered the particles are, the higher the entropy is

Physical state - solids have the lowest entropy as they just vibrate about a fixed position but gases have the highest entropy as they have the most disordered arrangements More particles - more ways they can be arranged and the more energy there is