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Aqa a level chemistry > Thermodynamics (Y13) > Flashcards

Thermodynamics (Y13) Flashcards

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1
Q

explain what happens to the size of the lattice energy when there is a highly charged cation and a small anion

A

-if the cation has a higher charge and the anion is smaller, there will be a stronger electrostatic force of attraction meaning more energy will be released upon forming a bond

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2
Q

explain why the lattice enthalphy for calcium flouride is much more exothermic than for potassium chloride

A

-F- has a smaller ionic radius than Cl-
-Ca2+ has larger chanrge than K+
-lattice enthaply is more exothermic for CaF
-greater attraction between Ca2+ and F-

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3
Q

which compound shows the greatest percentage difference between these 2 values

A

Lil

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4
Q

which equation represents the standard enthalphy of atomisation of iodine

A

1/2 I (s) –> I (g)

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5
Q

lattice energy

A

-the energy change when one mole of a solid ionic lattice is formed from its constituent gaseous ions (100kpa and 298K)

Na+ (g) + Cl- (g) –> NaCl (s)

-reactants must be gaseous

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6
Q

exothermic

A

Lattice energy is an exothermic process –> bonds are always being formed

Ionic compounds form strong giant lattices

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7
Q

lattice energy equation

A

e.g Mg2+ (g) + O2- (g) –> MgO

(must be ions + must be gaseous reactants)

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8
Q

factors which affect the size of lattice energy

A

-nuclear charge
-ionic radii

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9
Q

size of lattice energy

A

-Ionic radii –> the larger the radii of the ions, the less exothermic the lattice enthalphy

-Charges of ions –> the larger the charge of the ions, the more exothermic the lattice enthalphy

-increases across a period
-decreases down a group

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10
Q

electron affinity

A

Cl (g) + e- –> Cl- (g) (1st EA)

-1st electron affinity –> the energy change when one electron is added to each atom in 1 mol of gaseous atoms to form 1 mol of gaseous uninegative ions

-First electron affinity is an endothermic process since like charges of anion and electron repel (2nd EA)

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11
Q

enthalphy of atomisation

A

Na (s) –> Na (g) and 1/2Cl2 (g) —> Cl (g)

-Energy change when one mole of gaseous atoms is formed from its element in its standard states (100kpa and 298K) –> endothermic process

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12
Q

enthalphy of formation

A

-Na (s) + 1/2Cl2 (g) –> NaCl (s)

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13
Q

first ionisation energy

A

-the energy required to remove 1 mol of electrons from 1 mol of gaseous atoms to form 1 mol of positive ions in the gaeous state

Na (g) –> Na+ (g) + e-

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14
Q

exothermic vs endothermic in terms of electron affinity

A

-1st electron affinity of chlorine is exothermic as e- produces cl- ion with a more stable electron structure with lower energy

-2nd electron affinity of chlorine is more endothermic as gain a 2nd e- to cl- which requires energy to overcome the repulsion of like charges

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15
Q

formation

A

Formation

1 mol of compound

Na+ (g) + Cl- (g) –> NaCl (s)

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16
Q

atomisation

A

Atomisation

1 mol of gaseous atoms

Na (s) –> Na (g)

1/2Cl2 (g) –> Cl (g)

-for MgCl2 multiply enthalphy by 2

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17
Q

1st and 2nd ionisation

A

1st ionisation

1 mol of gaseous +1 ions

Na (g) –> Na+ (g) + e-

-for Na2O multiply enthalphy by 2

2nd ionisation

1 mol of gaseous 2+ ions

Mg+ (g) –> Mg2+ (g) + e-

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18
Q

1st and 2nd electron affinity

A

Na+ + Cl-

1st electron affinity

1 mol of gaseous 1- ions

Cl (g) + e- –> Cl- (g)

2nd electron affinity

1 mol of gaseous 2- ions

O- (g) + e- –> O2- (g)

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19
Q

lattice energy

A

Lattice

1 mol of ionic compound

Mg2+ (g) + 2Cl- (g) –> MgCl2 (s)

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20
Q

steps to calculate born haber cycle

A

1) standard enthalphy of formation of MgCl2

2) standard enthalphy of atomisation of magnesium

3) standard first ionisation energy of magnesium

4) second ionisation energy of magnesium

5) first electron affinity of chlorine

6) lattice energy of MgCl2 (s)

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21
Q

FAI(E)L

A

formation, atomisation (metal) , IE (metal) , electron affinity (non-metal) , lattice energy

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22
Q

arrows down

A

exothermic

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23
Q

which equation represents the process when the standard enthalphy of atomisation of iodine is measured

A

1/2I2 (s) –> I (g)

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24
Q

most covalent =

A

least polarisation = AlBr3

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25
Q

standard enthalphy of formation of BaCl2

A

Ba (s) + Cl2 (g) –> BaCl2 (s)

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26
Q

which way does final arrow face for enthalphy for lattice

A

down

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27
Q

suggest how the enthalphy of lattice formation of NaCl compares with that of NaF

A

lattice enthalphy for NaF is more exothermic than NaCl

F- is a smaller anion than Cl- so greater attraction between Na+ in NaF

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28
Q

what does it mean if the arrow for enthalphy of lattice faces upwards

A

lattice dissociation

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29
Q

hess cycle for enthalphy of solution

A

enthalphy for solution
solid ionic compound
(up arrow)
negative lattice enthalpy
gaseous ions
(down arrow)
enthalphy of hydration x2
aqueous ions

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30
Q

soluble vs insoluble

A

exothermic = soluble

endothermic = insoluble

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31
Q

solubility of ionic compound

A

When NaCl dissolves in water, the ions interact with water molecules and they become hydrated

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32
Q

negative lattice enthalphy

A

In order for a substance to dissolve the ionic lattice needs to be broken down and the ions need to be hydrated

-When the ionic lattice is broken down, gaseous ions are formed. This is called negative lattice enthalphy or lattice dissociation

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33
Q

enthalphy of hydration

A

-When ions become hydrated, excess water must be used. We call this enthalphy of hydration

-smaller ionic radii and high charge = more attracted to water molecules

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34
Q

dissolution

A

-the dissolution of NaCl can be pictured as 2 steps:

-enthalphy for solution = NaCl (s) –> Na+ (aq) + Cl- (aq) (ions are separated to form well-spaced ions in the gaseous state)

-lattice enthalphy = NaCl (s) –> Na+ (g) + Cl- (g)

-Enthalphy for hydration = Na+ (g) + aq –> Na+ (aq)

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35
Q

enthalphy of hydration

A

Enthalpy change of hydration = the enthalpy change when one mole of gaseous ions is dissolved in excess water

g dehydration of chlorine = Cl- (g) + aq –> Cl- (aq)

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36
Q

enthalphy of solution

A

Enthalphy change of solution = enthalphy change when one mole of an ionic substance dissolves in excess water

e.g solution of BaS = Bas (s) –> Ba2+ (aq) + S2- (aq)

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37
Q

hydration vs lattice

A

negative lattice energy = endothermic

-enthalphy for hydration = exothermic

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38
Q

negative enthalphy of solution

A

The enthalphy of a solution can either be exothermic or endothermic.

-A negative or small DHsol suggestes the ionic solid is soluble (sum of hydration is more negative than lattice energy)

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39
Q

true or false -> you always follow the arrows of lattice enthalphy

A

true

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40
Q

explain why the dissolving of magnesium sulfate in water is exothermic

A

Mg2+ has a large ionic charge = greater attraction

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41
Q

equation for enthalphy of solution of magnesium chloride

A

MgCl2 (s) –> Mg2+ (aq) + 2Cl- (aq)

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42
Q

state why there is a difference between the theoretical and experimental value

A

not purley ionic due to covalent character

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43
Q

hess cycle for enthalphy of solution

A

Negative LE + (hydration of Na+) + (hydration of Cl-) = enthalphy of solution

44
Q

trends in enthalphy of hydration

A

Solubility decreases as you go down the group. The lattice dissociation energy and hydration energy both decrease as you go down the group. The hydration energy decreases more than the lattice dissociation energy. Therefore the enthalpy of solution becomes more endothermic (or less exothermic).

decrease = Larger ions = weaker attraction to water molecules

45
Q

greater ionic charge and smaller radius

A

The greater the ionic charge and the smaller the radius the more exothermic the enthalphy of hydration

If the enthalphy of hydration becomes more exothermic then enthalphy of solution becomes more negative

46
Q

attraction between water molecules

A

If we increase ionic charge there will be an increase in the attraction between ions and water molecules = enthalphy of hydration is more exothermic (negative)

Lattice energy has a greater difference for small anions and large cations = more negative = more soluble

The more negative the enthalphy of solution the more soluble

47
Q

spontaneous reactions

A

-a spontaneous change is one that tends to occur of its own accord without being driven by an outisde force

-Two types of driving forces to a reaction: enthalphy or entropy

48
Q

exothermic vs endothermic entropy

A

-Exothermic reaction = spontaneous as it changes from a higher to lower energy state. (doesn’t require external heat energy) -> increases

-Endothermic reaction = energy is required as the products have a higher energy state and are less stable –> decreases

49
Q

entropy

A

-a measure of the disorder of substances

-Solid –> liquid = positive entropy change

-Solid = ordered, regular arrangement, low entropy

-Liquid = disordered, random arrangement, high entropy

50
Q

disordered

A

Disorder = number of ways of arranging the particles in a molecule. Water is more disordered since there are a greater number of ways arranging the particles.

51
Q

entropy of values

A

Symbol for entropy = S and units are joules per Kelvin per mol

Values for entropy are always positive

52
Q

exothermic

A

-movement of particles in surroundings increases as there is more kinetic energy

-Entropy of surroundings = positive

-enthalpy of system = negative

53
Q

endothermic

A

-movement of particles in surroundings decreases as there is less kinetic energy

-entropy of surroundings = negative

-enthalpy of system = positive

54
Q

temperature for entropy

A

If the temperature of the surroundings is already hot the energy released only has a small effect on particle motion so entropy is small

If the temperature of the surroundings is cold then the energy released has a large effect on particle motion so entropy is big

Entropy -> enthalpy = multiply enthalpy by 1000

55
Q

equations for entropy

A

-(enthalphy change) / temperature = entropy change

Total entropy = system + surroundings

Entropy total = entropy of system – -(change in enthalpy / time)

56
Q

what do spontaneous reactions depend on

A

-enthalphy change of the system

-entropy change of the system

-temperature

57
Q

Gibbs free energy

A

The change in quantity defined a Gibbs free energy provides a measure of whether a reaction is spontaneous or not

Δ G = Δ H − T Δ S ‍ (must remember equation)

-If Gibbs is less than 0 then the reaction is spontaneous

-If Gibbs is equal to zero then the reaction is at equilibirum

When calculating entropy using Gibbs the entropy values must be divided by 1000

58
Q

positive entropy and positive enthalphy

A

Depends on T (gets more favourable with bigger T)

59
Q

negative entropy and positive enthalpy

A

never spontaneous

60
Q

positive entropy and negative enthalphy

A

always spontaneous

61
Q

negative entropy and negative enthalphy

A

Depends on T (less favourable with a bigger T)

62
Q

Kp

A

p (c) x p (d) / p (a) x p (b)

63
Q

mole fraction

A

no of moles of gas / total number of moles of gas

64
Q

partial pressure =

A

mole fraction x total pressure

65
Q

what is used to calculate kp

A

partial pressure

only gaseous compounds

66
Q

what affects Kp

A

temperature only not catalysts and pressure

if equilibrium shifts right due to change in temperature Kp increases

if equilibrium shifts left then Kp decreases

67
Q

thermodynamics

A

formation = exo 2C(s) + 2H2 (g) -> C2H (g)

ionisation = endo Na (g) -> Na+ (g)

atomisation = endo 1/2F2 -> F (g)

affinity = exo O (g) -> O- (g)

68
Q

remember when calculating

A

if atomisation of an element is X2 and you need 1/2X2 then divide by 2

69
Q

theoretical vs experimental lattice enthalphy values

A

theoretical = purely ionic

experimental = covalent character
positive ion polarises negative ion
higher lattice enthalpy = larger distortion of negative ion = more polarisation

70
Q

enthalphy change of solution

A

left arrow up = dissociation
right arrow up = hydration
accross at top = solution

71
Q

positive entropy value

A

entropically feasible

units = JK-1 mol-1

more moles produced = entropy increases

more disorder e.g gas = higher entropy

72
Q

Gibbs = feasible

A

negative value or zero

73
Q

dissociation vs lattice enthalohy

A

dissociation = positive
lattic = negatvie

74
Q

The enthalpy of hydration of Ca2+(g) is –1650 kJ mol–1
Suggest why this value is less exothermic than that of Mg2+(g)

A

-Ca2+ has a larger charge to size ratio
-weaker attraction to the O- in water

75
Q

why is the standard entropy value for CO2 greater than carbon

A

CO2 is more disordered

76
Q

enthalphy lattice dissociation

A

The lattice dissociation enthalpy is the enthalpy change needed to convert 1 mole of solid crystal into its scattered gaseous ions. Lattice dissociation enthalpies are always positive

77
Q

suggest one reason why the first electron affinity of oxygen is exothermic

A

large nuclear charge
strong electrostatic forces of attraction

78
Q

By describing the nature of the attractive forces involved, explain why the value for the enthalpy of hydration of the chloride ion is more negative than that for the bromide ion

A

During enthalpy of hydration gaseous ions form aqueous ions by forming bonds with water molecules which are polar.

If the value is more negative, then more energy is released so the bond must be stronger.

Cl- ion is smaller than Br- ion (smaller ionic radius)

Forces of attraction between Cl- ion and water are therefore stronger than between Br- ion and water.

Cl- ions attract the delta positive H of water molecule more strongly

79
Q

enthaply of atomisation

A

The enthalpy of atomisation of an element is the enthalpy change when
1 mole of gaseous atoms is formed from the element in its
standard state

80
Q

diatomic molecules =

A

multipy by 2 for atomisation and electron affinity

1/2Cl2 –> Cl (g)
Cl2 –> 2Cl (g)

81
Q

first ionisation energy

A

The first ionisation enthalpy is the enthalpy change required to
remove 1 mole of electrons from 1 mole of gaseous atoms to form
1 mole of gaseous ions with a +1 charge

82
Q

first electron affinity

A

The first electron affinity is the enthalpy change that occurs when 1
mole of gaseous atoms gain 1 mole of electrons to form 1 mole of
gaseous ions with a –1 charge

The first electron affinity is exothermic for atoms that normally
form negative ions. This is because the ion is more stable than the
atom, and there is an attraction between the nucleus and the
electron

83
Q

enthalphy of lattice dissociation

A

The enthalpy of lattice dissociation is the standard enthalpy
change when 1 mole of an ionic crystal lattice form is separated
into its constituent ions in gaseous

84
Q

enthalphy of solution

A

The enthalpy of solution is the standard enthalpy change
when one mole of an ionic solid dissolves in a large enough
amount of water to ensure that the dissolved ions are well
separated and do not interact with one another.

85
Q

where does the arrow point for negative lattice dissociation

A

up

86
Q

lattice

A

formation - (atomisation + ionisation + electron affinity)

87
Q

first and second ionisation energy MgCl2

A

Mg2+ (g) + 2e- + Cl2 (g)

Mg2+(g) + 2e- + 2Cl (g)

88
Q

why is second electron affinity for oxygen endothermic

A

Notice the second electron affinity for
oxygen is endothermic because it
take energy to overcome the
repulsive force between the
negative ion and the electron

89
Q

theoretical vs experimental lattice enthalpy

A

The Born Haber lattice enthalpy is the real experimental value.
When a compound shows covalent character, the theoretical and the born
Haber lattice enthalpies differ. The more the covalent character the bigger
the difference between the values.

90
Q

why does gas have largest entropy

A

Solids have lower entropies than liquids, which are lower than gases.
When a solid increases in temperature its entropy increases as the
particles vibrate more.
There is a bigger jump in entropy with boiling than that with melting.
Gases have large entropies as they are much more disordered.

91
Q

increase in entropy =

A

An increase in disorder and entropy will lead to a positive entropy change ∆S˚ = +ve = increase in temperature = gibbs is negative = more likely for reaction to occur

92
Q

change in S =

A

S products - reactants

93
Q

T =

A

Make ∆G = 0 in the following equation ∆G = ∆H - T∆S
0 = ∆H - T∆S
So T = ∆H / ∆S
T= 0.94 / (10.3÷1000)
T= 91 K

94
Q

enthalphy of solution =

A

lattice dissociation + hydration

or negative lattice formation

95
Q

why are hydration enthalpies exothermic

A

Hydration enthalpies are exothermic as energy is given
out when water molecules bond to the metal ions.
The negative ions are electrostatically attracted to the δ+ hydrogens on the polar water molecules and the positive
ions are electrostatically attracted to the δ
- oxygen on the
polar water molecules

96
Q

equation for enthalphy of solution of MgCl2

A

MgCl2 (s) –> Mg2+ (aq) + 2Cl- (aq)

97
Q

why is the enthalphy of hydration of calcium less exothermic than magneisum

A

Ca2+ has lower charge to size ratio
-weaker attraction to partial negative oxygen in water

98
Q

units of entropy

A

J K-1 mol -1

99
Q

units of enthalphy

A

KJ mol -1

100
Q

total entropy =

A

change in entropy - change in enthalphy / time

101
Q

explain why the free-energy change for the dissolving of potassium chloride in water is negative even though enthalphy change is positive

A

entropy change is positive
-number of particles increases
Therefore T∆S > ∆H and ∆G becomes less than zero

102
Q

complete the born haber cycle for sodium flouride

A

Na (g) –> 1/2F2 (g)

Na+ (g) + e- + 1/2F2(g)

103
Q

F2(g) –> 2F =

A

divide by 2

104
Q

if sodium is in excess =

A

divide answer by 2

105
Q

use the equation and the data in the table above to calculate the minimum temperature in K at which this reaction becomes feasible

A

G = H- T?S

Find H (products - reactants)
Find S (products - reactants)

H/S to get answer

Aqa a level chemistry (19 decks)

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