Enthalpy of formation
The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.
1st ionisation enthalpy =
the enthalpy change when one mole of gaseous 1 + ions is made from one mole of gaseous atoms.
2nd ionisation enthalpy=
the enthalpy change when one mole of gaseous 2+ ions is formed from one mole of 1+ ions.
Enthalpy of atomisation
the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state
the enthalpy change when 1 mole of gaseous atoms is formed from a compound in its standard states
Bond Dissociation enthalpy
The enthalpy change when all the bonds of the same type in one mole of gaseous molecules are broken.
1st electron affinity =
if the enthalpy change when one mole of gaseous 1- ions is made from one mole of gaseous atoms
2nd electron affinity=
the enthalpy change when one mole of gaseous 2- ions is made from 1 mole of gaseous 1- ions.
Lattice enthalpy of formation
the standard enthalpy change when one mole of solid ionic compound is formed from gaseous ions
Enthalpy of hydration
the enthalpy change when one mole of aqueous ions is formed from one mole of gaseous ions.
Enthalpy change of solution
The enthalpy change when one mole of solute is dissolved in sufficient solvent that no further enthalpy change occurs of further dilution.
Born haber cycles.
- These follow Hess's law that the enthalpy change from reactant to product is the same, no matter the direction taken.
- Requires understanding of the definitions already outlined.
- Used to calculate lattice enthalpy of formation.
differences between exp values and theoretical val
- most experimental values via born haber cycle agrees with theoretical values.
- shows that the correct model of ionic bonding has been decided upon.
- However, some compounds have large discrepancy between the two values
- due to the bond in question having some covalent character.
- e.g. Zinc Selenide, ZnSe, the theoretical values is 10% lower than experimental value.
- the greater experimental values implies some extra bonding is present.
- can be explained as follows:
- Zn2+ = relatively small with high positive charge
- Se2- = relatively large with high neg charge.
- small Zn2+ able to approach Se2- and distorts electron clouds.
- Se2- electron clouds are distorted easily.
- this distortion causes more electrons than expected conc between the Zn and nuclei, represents a degree of electron sharing or covalance which occurs for the lattice enthalpy discrepancy.
- Se2- ions is polarised
calculation enthalpies of solution from enthalpies
- this is done using an enthalpy cycle
- Need to know:
- lattice dissociation of the compound
- enthalpies of hydration of the ions.
- connect ionic lattice and dissolved ions using enthalpy change of solution. This is the direct route
- connect ionic lattice to gaseous ions by the lattice enthalpy of dissociation. This is a +ve no.
- connect gaseous ions to dissolved ions by hydration enthalpies of each ions.
- Use Hess's law to calculate enthalpy change of solution.
using mean bond enthalpies to calculate approx ent
- enthalpy change for a reaction = sum of bonds broken - sum of enthalpies of bonds formed.
- if you need more energy to break bonds than is released when bonds are made, its an endothermic reaction and the enthalpy change is positive
- if more energy is released than is taken in, it's exothermic and enthalpy is negative
- sketch out molecules
- add up mean bond enthalpies for reactant bonds broken
- add up the mean bond enthalpies for bond formed in products
- enthalpy change = bonds broken - bonds formed.
Mean bond calculations only approximations
- e.g. mean bond enthalpy for every N-H bond is not the same for every N-H bond.
- a given type of bond will vary in strength from compounds and can vary within a compound.
- mean bond enthalpies are the averages. Only the bond enthalpies of diatomic molecules such and HCl will always be the same
- calculations using mean bond enthalpies will never be perfectly accurate.
- more exact results obtained from specific compounds.
why mean bond enthalpies differ from enthalpy cycl
- theoretical lattice enthalpy assumes you are working with a purely ionic model of a lattice. assumes all ions are spherical and charge is evenly distributed around them.
- however, most ions aren't usually spherical.
- positive ions can polarise neighbouring negative ions. The higher the amount of polarisation the more covalent the molecules are.
- If the two methods produce the same value, or similar, than the lattice is purely ionic.
enthalpy does not explain spontaneous change
'Feasible' or 'spontaneous' is the term used to describe reactions that take place of its own accord.
terms take no account of rate of reaction.
enthalpy change takes no explanation why endothermic reactions are spontaneous.
Entropy is a measure of how much disorder there is
- measure of the number of ways the particles can be arranged and the number of ways the energy can be spread between the particles.
- the higher the disorder the more energetically stable substances are.
- Particles normally attempt to increase entropy
e.g. gas particles diffusing across a room.
Physical states and entropy
Solid (Ordered) --> Liquid (some disoorder) --> Gas ( random)
Dissolving a solid increases entropy --> evaporating a liquid increases entropy.
chemical changes effect entropy
C6H8O7 + 3NaHCO3 --> Na3C6H5O7 + 3H2O +3CO2
Produces more gas molecules so increasing entropy
Increasing disorder more molecules
Dissolution: The process by which a solid or liquid forms a homogenous mixture with a solvent. (breakdown of a crystal lattice into individual atoms)
calculating entropy values from absolute entropy v
The total energy change includes system and surroundings.
during a reaction, there's an entropy change between the reactants and products - the entropy change of the system.
The entropy of the surroundings changes too.
Entropy (S) total = S system + S surroundings.
S system = S products - S reactants
S surroundings = -(enthalpy/Temp in kelvin)
Free energy change (g) = enthalpy - (temp*entropy)
For spontaneous reactions G must be -ve or 0.
Feasibility of reaction depends on temp
- exothermic reaction and +ve entropy change
- G = -ve
- reactions feasible at any temp
- endothermic and -ve entropy change
- G = +ve
- reactions not feasible at any temp
- For any other combinations, temp has a effect.
- If enthalpy is positive and S system is positive
- reaction not feasible at some temps but will be at higher temps
- enthalpy -ve and S system -ve
- reaction will be feasible at some temps but won't be at higher temp.
- Temp at which the reaction is feasible = enthalpy/S system