top of page

Energetics II

​

  • -  Standard Lattice Energy: Enthalpy change when 1 mole of an ionic solid is

    formed from its gaseous ions under standard conditions of 298 K and 100kPa [1

    mol of ionic solid formed from gaseous ions]

  • -  Lattice dissociation: Enthalpy change when one mole of a solid ionic compound

    is completely dissociated into its gaseous constituent ions under standard

    conditions (ENDOTHERMIC)

  • -  Lattice formation: Enthalpy change when one mole of solid ionic compounds

    formed from its gaseous constituent ions under standard conditions

  • -  Enthalpy of Hydration: Enthalpy change when 1 mole of gaseous ions dissolves in water to form one mole of aqueous ions under standard condition [1 mole of gaseous ions dissolve fully to form 1 mol of aqueous ions]

  • -  Enthalpy of Solution: Enthalpy change when 1 mole of ionic solid is dissolved to infinite dilution so that the ions no longer interact under standard conditions [1 mol of ionic solid infinitely diluted so the ions no longer interact]

  • -  Enthalpy of Atomisation: Enthalpy change when 1 mole of gaseous atoms is formed from its constituent elements in its standard state under standard conditions [1 mol of gaseous atoms from elements in standard state]

  • -  Ionic bond strength is related to lattice enthalpy. The higher the enthalpy the stronger the bond strength

  • -  Electron Affinity: The energy change when 1 electron is added to each atom in 1 mole of gaseous atoms to form an ion. (typically occurs after atomising.

  • -  Lattice Energy:

    • -  Standard enthalpy of formation is shown going down from the eqauton to

      the product.

    • -  A series of upward energies:

      - Atomisation - Ionisation
      - Atomisation

- Electron Affinity

- THESE ARE ALL TO MAKE THE IONS WHICH REACT TO FORM THE LATTICE

- There will be the lattice energy which you may be asked to calculate front he other data and this will be going down from the gaseous ions under state to the ionic solid and so everything must be made similar to this

  • -  The value calculculated is merely theoretical and it is not the true lattice energy value. Often the experimental value will be more negative

  • -  This occurs due to the polarisation of ions and this means that the experimental values which ar calculated using the purely ionic model cannot function as in actuality the bonding id not so definitively ionic as is assumed

  • -  The positive ions may attract the outer most electrons of a negative ion pulling these electrons into a region between the ions and creating a degree of

   

covalency in the bond. This contribution from covalent bonding leads to the actual bond enthalpy being more negative as the bond is even stronger than assumed in the pure ionic model

  • -  The degree to which this occurs is based on the polarising power - the ability of a positive ion to distort the negative electron cloud of a negative ion. As well as the polarisability - the egress to which the electron cloud in a molecule can be distorted by the positive charge.

  • -  Polarising power is greater when the charge is larger and the radius of the ion is smaller

  • -  Polarisability is greater when the ionic radius is larger and there are mre electrons

  • -  The closer to the theoretical value the more closely the perfect ionic model is followed

    Difficult to make assumptions about solubility:

  • -  The enthalpy change of solution is usually a very small difference between two values

    for the enthalpy changes. So even a small error can lead to major flaw in estimating the

    trend

  • -  Both lattice energy and hydration enthalpies are affected in the same way which is

    changes in bonding type etc which means clear trends are difficult

  • -  With a small enthalpy value especially the sign is not a clear indicator of whether a

    reaction is feasible

    Entropy:

  • -  Feasible reactions go naturally

  • -  Exothermic reactions are often feasible but some endothermic reactions can be feasible

    too

  • -  As a result the use of enthalpy alone is limited

  • -  Random changes tend to occur in the direction which increases the number of ways in

    which the order or arrangement of molecules can be

  • -  Entropy is a measure of disorder or randomness

  • -  EQUATION: delta S total = delta S system + delta S surroundings

  • -  Entropy rises as temperature rises as the more KE the more ways to arrange

  • -  EQUATION: Delta S system = Sum of S products - Sum of S reactants

  • -  EQUATION: Delta S surroundings = - delta H/T

  • -  Positive delta S means the reaction is feasible

    • -  So most exothermic are feasible because the surroundings value is more positive than the system value so overall positive

    • -  As tempt increases it may become feasible

  • -  Gibbs free energy:

    • -  EQUATION: DeltaG = Delta H -TdeltaS (system)

    • -  If value is negative then feasible

    • -  UNITS would be kJ K-1 mol-1 so convert entropy to kJ

    • -  The delta S used is that of the system

  • -  We can also work out/derive from this the relationship between DeltaG and delta S total

  • -  Delta G =-TdeltaS (total)

  • -  EQUATION: Delta G = -RT lnK

  • -  Enables us to infer rate by calculating the eq constant -

    Relationship between hydration, solution and formation

  • -  A+(g) + B- (g) → AB(s) Lattice Enthalpy

  • -  A+ (g) + B- (g) → A+ (aq) + B- (aq) Hydration enthalpy

  • -  AB(s) → A+(aq) + B- (aq)

  • -  NOW: enthalpy of solution = enthalpy of hydration - lattice enthalpy

  • -  Create a Table -

Enthalpy Change

Value

Lattice Enthalpy

Always -ve (making bonds which always gives out energy)

Hydration enthalpy

Always -ve (making intermolecular bonds)

Enthalpy of Solution

Can be +ve or -ve as entirely dependent on the ionic salt (often group 1 in Qs as very varied behaviour)

If the hydration enthalpy ‘overtakes’ the lattice enthalpy then it will dissolve really easily

E.G.

LiCl - -37 kJ/mol NaCl - +4 kJ/mol KCl - + 26 kJ/mol

We can see it goes from being exothermic to endothermic and

Hydration enthalpy is more exothermic than the lattice enthalpy to make it a negative value (for LiCl) we can explain this using delta H sol = Hydration - (-Lat)
NaCl and KCl have more shells in the metal so there is a lower charge density and the lattice enthalpy increases and thus is greater than the hydration enthalpy so the overall value is positive

FOR ALL entropy will be positive as going from a solid to a liquid/solution/aqueous

Assuming the same temperature then the more or less exothermic that the delta H is the G will vary

As a result the temperature can be increased but only to the point of the boiling point of Water

bottom of page