22. enthalpy and entropy
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lattice enthalpy
A LE H* (always a - value as it is exothermic) enthalpy change that accompanies the formation of 1 mole of an ionic compound from its gaseous ions under standard conditions a measure of the strength of ionic bonding in a giant ionic lattice
lattice enthalpy cannot be directly measured so how is it measured
born harbor cycle uses known energy changes in an energy cycle
standard enthalpy change of formation
A f H* enthalpy change that takes place when 1 mole of a compound is formed from its elements under standard conditions in there standard states exothermic
standard enthalpy change of atomisation
A at H* enthalpy change that takes place for the formation of 1 mole of gaseous atoms from the element in its standard state under standard conditions always endothermic as bonds are broken for already gaseous atoms it is bond enthalpy of bond being broken e.g. 1/2 Cl2(g) —> Cl(g)
1st ionisation energy
A IE H* enthalpy change required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions endothermic as energy is required to overcome the attraction between a negative electron and positive
1st electron efinity
A EA H* enthalpy change that takes place when 1 electron is added to each atom of 1 mole of gaseous atoms to form 1 mole of gaseous 1- ions opposite of ionisation energy exothermic second electron affinities are endothermic e- being gained by a - ion
standard enthalpy change of solution
A sol H* enthalpy change that takes place when 1 mole of a solute dissolves in 1 mole of a solvent if solvent is water, ions from the ionic lattice finish up surrounded with water molecules as aqueous ions e.g Na+Cl- + aq(excess water) —> Na+(aq) + Cl-(aq) exo or endothermic
what is m in q=mcAT
mass what is changing temp could be solution
what happens when an ionic compound dissolves in water
ionic lattice breaks up forming separate gaseous ions (opposite of lattice enthalpy) gaseous ions interact with the polar water molecules forming hydrated aqueous ions, water molecules are attracted to and surround the ions enthalpy change of hydration
enthalpy change of hydration
A hyd H enthalpy change that accompanies the dissolving of gaseous ions in water to form 1 mol of aqueous ions
general properties of ionic compounds
high melting and boiling points soluble in polar solvents conduct electricity when molten or in aqueous solution
what affects lattice enthalpy
ionic size ionic charge
as ionic size increases A LE H*
ionic radius increases attraction between ions decreases lattice enthalpy less negative melting point decreases
as ionic charge increases A LE H*
attraction between ions increases lattice energy becomes more negative melting point increases
ionic charge and size across period 3
Na + -> Mg 2+ -> Al 3+ increasing charge more attraction decreasing size more attraction S 2- <- Cl - increasing charge more attraction increasing size less attraction
factors affecting hydration enthalpies
ionic size ionic charge
as ionic size increases A hyd H*
ionic radius increases attraction between ion and water molecules decrease hydration energy less negative
as ionic charge increases A hyd H*
attraction with water molecules increases hydration energy becomes more negative
predicting solubility
if sum of hydration enthalpies is larger than the magnitude of lattice enthalpy the overall enthalpy change (enthalpy change of solution) will be exothermic and the compound should dissolve but some with endothermic can dissolve to
entropy (S)
measures the dispersal of energy within chemicals making up a chemical system disorder JK-1mol-1 per kelvin per mol
greater the entropy
greater the dispersal of energy and disorder
rank the states of matter from lowest entropy to highest
solids, liquids, gasses mostly
entropy value at 0K
no energy so 0
effect of gaseous products on entropy
production of a gas increases entropy if products have less moles of gas than reactants entropy decreases