Enthaplies A2 Flashcards
Def of enthalpy of formation
Enthalpy change when 1 mol of a substance is formed from its elements with all substances in their standard states
Symbol and example of enthalpy of formation
(Use sodium chloride)
ΔfH⦵
E.g Na (s) + 1/2Cl2 (g) —> NaCl (s)
Def of the first ionisation enthalpy
Enthalpy change when 1 mol of electrons is removed from 1 mol of gaseous species
Def of enthaply of atomisation (element)
Enthalpy change when 1 mol of gaseous atoms is formed from its element in its standard state
Def enthalpy of atomisation (compound)
Enthalpy change when 1 mol of a compound in gas state is converted into separate gaseous ions
Def bond dissociation enthalpy
Enthalpy change when 1 mol of given covalent bond, of a compound in the gas states is broken
Def of first electron affinity enthalpy
Enthalpy change when 1 mol of electrons is added to 1 mol of gaseous species
Def of lattice enthalpy (also called lattice formation)
Enthalpy change of formation of 1 mol of an ionic lattice from its gaseous ions
Def of enthalpy of hydration
Enthalpy change when one mole of gaseous ions are diluted to give no further temperature change and one mole of hydrated ions
Def of enthalpy of solution
Enthalpy change when 1 mole of solute is dissolved in sufficient solvent to give no interaction between dissolved species
Symbol and example equation for the first ionisation enthalpy
(Use sodium)
Δ1st i.e H⦵
E.g Na (g) —> Na+ (g) + e-
Symbol and example equation for enthalpy of atomisation of an element
(Use sodium and chlorine)
ΔatH⦵
E.g Na(s) —> Na (g)
1/2 Cl2 (g) —> Cl2 (g)
Symbol and example equation for the enthalpy of atomisation of a compound
(Use methane)
ΔatH⦵
E.g CH4 —> C(g) + 4H(g)
Symbol and example equation for bond dissocation enthalpy
(Use methane)
ΔdissH⦵
E.g CH4 (g) —> CH3 (g) + H (g)
Symbol and example equation for first electron affinity enthalpy
(Use chlorine )
Δ1st eaH⦵
E.g Cl (g) + e- —> Cl- (g)
Symbol and example equation for lattice enthalpy
(Use sodium chloride)
ΔLEH⦵
E.g Na+ (g) + Cl- (g) —> NaCl(s)
Symbol and example equation of enthalpy of solution
( use sodium chloride)
ΔsolH⦵
E.g NaCl (s) —> Na+ (aq) + Cl- (aq)
Symbol and example equation of enthalpy of hydration
( use sodium)
ΔhydH⦵
E.g Na+ (g) + H2O (l) —> Na+ (aq)
Draw a Born-Haber cycle to calculate the ΔleH of LiF
If
Δ1st ea H(f) = -328 KJ mol-1
Δ atH(Li) = +159 KJ mol -1
Δ1st ie H(Li) = +520 KJ mol -1
ΔatH(F) = +79 KJ mol-1
ΔfH(LiF) = -616 KJ mol-1
See notes for cycle
ΔleH (LiF) = - Δ1st ea H(f) - Δ atH(Li) - Δ1st ie H(Li) - ΔatH(F) +ΔfH(LiF)
ΔleH (LiF) = -(-328) - (+79) - (+520) - (+159) + (-616)
ΔleH (LiF) = -1046 KJ mol-1
What is the effect of cation (positive) size (down the group) on
1) attraction
2) lattice energy
3) melting point
Na+ < K+ < Rb+
As ionic size increases
1- ionic radius increases
2- attraction between ions decreases
3- lattice energy less negative
4- melting point decreases
What is the effect of increased cation (positive) charge on
1) attraction
2) lattice energy
3) melting point
Na+ vs Ca2+
As ionic charge increases
Attraction between ions increases
Lattice energy becomes more negative
Melting point increases
What is the effect of cation (positive) size (down the group) on
1) attraction
2) hydration enthalpy
Na+ < K+ < Rb+
Ionic radius increases
Attraction between ion and water molecules decreases
Hydration energy less negative
What is the effect of increased cation (positive) charge on
1) attraction
2)hydration energy
Na+ vs Ca2+
Ionic charge increases
Attraction with water molecules increases
Hydration energy becomes more negative
How do you calculate enthalpy of solution of LiF from lattice enthalpies and enthalpies of hydration using Hess cycle or Born-Haber cycle
If
Lattice enthalpy = -1031
Hydration of Li+ = -520
Hydration of F- = -524
Enthalpy of solution (LiF) = - lattice enthalpy + enthalpy of hydration( Li+) + enthalpy of Hydration (F-)
Enthalpy of solution (LiF) = -(-1031) + ((-520) + (-524))
Enthalpy of solution (LiF) = -13 KJ mol-1
Def of entropy
A measure of the dispersal of energy in a system which is greater, the more disordered a system
Why might lattice enthalpy be different to calculated values
Calculated values assume perfect ionic model whereas experimentally there can be some degree of covalent bonding in lattices
What does perfect ionic model assume
- all ions are perfectly spherical
- charge is distributed evenly throughout the ion
- the ions display no covalent charater
Why is enthalpy of hydration always negative
Water molecules have + dipole regions that naturally attract negative ions
What factors affect enthalpy of hydration and lattice enthalpy
hydration
- attraction are stronger with smaller ions
- with greater change.
This explains why hydration enthalpies decrease as you move down a group
lattice
- increases in magnitude with decreasing ionic radius
- and increasing change
this is due to ions form stronger attractions so more energy required for dissociation
What is the symbol for entropy
S
What effect does increasing temp have on entropy and why
Entropy increases as temperature increase
- particles gain energy and move faster and further apart, the particles become less ordered.
Compare the entropy of solids liquids and gases
Solids have lowest entropy
Then liquids
Gases have the greatest entropy
What happens what a lattice is dissolved in solution to the entropy
The entropy increases because the ions in lattice are being dissociated so can move, increasing the disorder
Calc of entropy change
+ units for entropy
ΔS total = sum of ΔS products — sum of ΔS reactants
Entropy measured in J K-1 mol-1
If entropy change is positive what does that mean
The products are more disordered that the reactants
E.g solid —> gas
If entropy change is negative what does that mean?
The products are less disordered than the reactants
E.g liquid —> solid
What is Gibbs’s free energy
Way to quantify the balance between entropy and enthalpy of a system acting as an idicator of reaction feasibility
Calc of Δgibbs free energy
+ what are its units
ΔG = ΔH — T ΔS
Gibbs free-energy measured in KJ mol-1
ensure that when using this calculation ΔH and T ΔS are in the same units as ΔS is J not KJ so needs converting
If ΔG (gibbs free-energy) is less than or equal to zero what happens
If ΔG negative or 0 than the reaction is feasible if given activation energy
If ΔG (gibbs free-energy) is more that 0 what happens
If ΔG is positive the reaction is not feasible at the temperature even if activation energy given
Calc of minimum temp reaction is feasible at using entropy
If ΔG = ΔH - T ΔS
And reaction is feasible when ΔG = 0
0 = ΔH - T ΔS
Rearranged give
TΔS = ΔH
Therefore the temp at which reaction becomes feasible
T = ΔH
——
ΔS
Limitations of using ΔG for feasibility, in terms of kinetics
Although reaction might be thermodynamically feasible it may not be able to occur due to
kinetic factors
- such a slow rate of reaction it appears not to happen at all
Or
To large activation energy that isn’t provided by reaction conditions
Explain why the lattice enthalpy of NaBr is more exothermic than for KBr (2 marks)
Na has a smaller radius
So stronger attraction to Br
The exothermic reaction below occurs spontaneously at low temperatures but not at very high temperatures
2SO2(G) + O2 (G) —> 2SO3(G)
Explain why
(2 marks)
If exothermic reaction ΔH is a negative value so
1) with high temperatures the T ΔS is more negative
2) at high temperatures ΔH -T ΔS > 0
When dissolved in water, the enthalpy change of solution of the salt KF, is -15kJ mol-1
The salt RbF has an enthalpy change of solution in water of -24kJ mol-1
Suggest reasons for the difference between the enthalpy changes of solution between KF & RbF
(4 marks)
K+ is smaller so stronger attraction to F
Lattice enthalpy of KF more negative than RbF
Δhydration H of K+ is more negative than Rb+
From data given in the question ΔH is affected more by lattice enthalpy than by hydration enthalpy
State 2 large scale uses for the hydrogen produced
- manufacture of margarine
-to produce fuel cells - to manufacture ammonia
- manufacture HCl
- reduction of metal ores
