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A Level Chemistry > Enthaplies A2 > Flashcards

Enthaplies A2 Flashcards

1
Q

Def of enthalpy of formation

A

Enthalpy change when 1 mol of a substance is formed from its elements with all substances in their standard states

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2
Q

Symbol and example of enthalpy of formation
(Use sodium chloride)

A

ΔfH⦵
E.g Na (s) + 1/2Cl2 (g) —> NaCl (s)

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3
Q

Def of the first ionisation enthalpy

A

Enthalpy change when 1 mol of electrons is removed from 1 mol of gaseous species

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4
Q

Def of enthaply of atomisation (element)

A

Enthalpy change when 1 mol of gaseous atoms is formed from its element in its standard state

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5
Q

Def enthalpy of atomisation (compound)

A

Enthalpy change when 1 mol of a compound in gas state is converted into separate gaseous ions

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6
Q

Def bond dissociation enthalpy

A

Enthalpy change when 1 mol of given covalent bond, of a compound in the gas states is broken

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7
Q

Def of first electron affinity enthalpy

A

Enthalpy change when 1 mol of electrons is added to 1 mol of gaseous species

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8
Q

Def of lattice enthalpy (also called lattice formation)

A

Enthalpy change of formation of 1 mol of an ionic lattice from its gaseous ions

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9
Q

Def of enthalpy of hydration

A

Enthalpy change when one mole of gaseous ions are diluted to give no further temperature change and one mole of hydrated ions

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10
Q

Def of enthalpy of solution

A

Enthalpy change when 1 mole of solute is dissolved in sufficient solvent to give no interaction between dissolved species

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11
Q

Symbol and example equation for the first ionisation enthalpy
(Use sodium)

A

Δ1st i.e H⦵

E.g Na (g) —> Na+ (g) + e-

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12
Q

Symbol and example equation for enthalpy of atomisation of an element
(Use sodium and chlorine)

A

ΔatH⦵

E.g Na(s) —> Na (g)
1/2 Cl2 (g) —> Cl2 (g)

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13
Q

Symbol and example equation for the enthalpy of atomisation of a compound
(Use methane)

A

ΔatH⦵

E.g CH4 —> C(g) + 4H(g)

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14
Q

Symbol and example equation for bond dissocation enthalpy
(Use methane)

A

ΔdissH⦵

E.g CH4 (g) —> CH3 (g) + H (g)

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15
Q

Symbol and example equation for first electron affinity enthalpy
(Use chlorine )

A

Δ1st eaH⦵

E.g Cl (g) + e- —> Cl- (g)

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16
Q

Symbol and example equation for lattice enthalpy
(Use sodium chloride)

A

ΔLEH⦵

E.g Na+ (g) + Cl- (g) —> NaCl(s)

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17
Q

Symbol and example equation of enthalpy of solution
( use sodium chloride)

A

ΔsolH⦵

E.g NaCl (s) —> Na+ (aq) + Cl- (aq)

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18
Q

Symbol and example equation of enthalpy of hydration
( use sodium)

A

ΔhydH⦵

E.g Na+ (g) + H2O (l) —> Na+ (aq)

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19
Q

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

A

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

20
Q

What is the effect of cation (positive) size (down the group) on
1) attraction
2) lattice energy
3) melting point
Na+ < K+ < Rb+

A

As ionic size increases
1- ionic radius increases
2- attraction between ions decreases
3- lattice energy less negative
4- melting point decreases

21
Q

What is the effect of increased cation (positive) charge on
1) attraction
2) lattice energy
3) melting point
Na+ vs Ca2+

A

As ionic charge increases
Attraction between ions increases
Lattice energy becomes more negative
Melting point increases

22
Q

What is the effect of cation (positive) size (down the group) on
1) attraction
2) hydration enthalpy

Na+ < K+ < Rb+

A

Ionic radius increases
Attraction between ion and water molecules decreases
Hydration energy less negative

23
Q

What is the effect of increased cation (positive) charge on
1) attraction
2)hydration energy

Na+ vs Ca2+

A

Ionic charge increases
Attraction with water molecules increases
Hydration energy becomes more negative

24
Q

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

A

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

25
Q

Def of entropy

A

A measure of the dispersal of energy in a system which is greater, the more disordered a system

26
Q

Why might lattice enthalpy be different to calculated values

A

Calculated values assume perfect ionic model whereas experimentally there can be some degree of covalent bonding in lattices

27
Q

What does perfect ionic model assume

A
  • all ions are perfectly spherical
  • charge is distributed evenly throughout the ion
  • the ions display no covalent charater
28
Q

Why is enthalpy of hydration always negative

A

Water molecules have + dipole regions that naturally attract negative ions

29
Q

What factors affect enthalpy of hydration and lattice enthalpy

A

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

30
Q

What is the symbol for entropy

A

S

31
Q

What effect does increasing temp have on entropy and why

A

Entropy increases as temperature increase
- particles gain energy and move faster and further apart, the particles become less ordered.

32
Q

Compare the entropy of solids liquids and gases

A

Solids have lowest entropy
Then liquids
Gases have the greatest entropy

33
Q

What happens what a lattice is dissolved in solution to the entropy

A

The entropy increases because the ions in lattice are being dissociated so can move, increasing the disorder

34
Q

Calc of entropy change
+ units for entropy

A

ΔS total = sum of ΔS products — sum of ΔS reactants

Entropy measured in J K-1 mol-1

35
Q

If entropy change is positive what does that mean

A

The products are more disordered that the reactants

E.g solid —> gas

36
Q

If entropy change is negative what does that mean?

A

The products are less disordered than the reactants

E.g liquid —> solid

37
Q

What is Gibbs’s free energy

A

Way to quantify the balance between entropy and enthalpy of a system acting as an idicator of reaction feasibility

38
Q

Calc of Δgibbs free energy
+ what are its units

A

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

39
Q

If ΔG (gibbs free-energy) is less than or equal to zero what happens

A

If ΔG negative or 0 than the reaction is feasible if given activation energy

40
Q

If ΔG (gibbs free-energy) is more that 0 what happens

A

If ΔG is positive the reaction is not feasible at the temperature even if activation energy given

41
Q

Calc of minimum temp reaction is feasible at using entropy

A

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

42
Q

Limitations of using ΔG for feasibility, in terms of kinetics

A

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

43
Q

Explain why the lattice enthalpy of NaBr is more exothermic than for KBr (2 marks)

A

Na has a smaller radius
So stronger attraction to Br

44
Q

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)

A

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

45
Q

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)

A

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

46
Q

State 2 large scale uses for the hydrogen produced

A
  • manufacture of margarine
    -to produce fuel cells
  • to manufacture ammonia
  • manufacture HCl
  • reduction of metal ores

A Level Chemistry (12 decks)

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