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Chemistry Y13 > Thermodynamics > Flashcards

Thermodynamics Flashcards

1
Q

Enthalpy changes

A
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2
Q

What is the definition for standard enthalpy of formation?

A

the enthalpy change when 1 mole of a substance is formed from it’s elements under standard with elements in their standard state at 100KPa

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3
Q

standard enthalpy of formation usually concerns forming what compounds?

A

usually concerned about forming ionic compounds

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4
Q

write standard enthalpy of formation for the following compounds:
1. calcium fluoride
2. sodium sulphide
3. lithium carbonate

A
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5
Q

During the formation of an ionic compounds, ionic..?

A

ionic bonds are formed

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6
Q

is bond making exothermic or endothermic?

A

Bond making is exothermic

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7
Q

Describe ionic bonding as you go down a group?

A
  • down a group ionic radii increases (ions get bigger)
  • so weaker electrostatic forces between oppositely charged ions
  • so weaker ionic bonding
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8
Q

the ion is stronger if the..?

A

Charge on the ion increases
e.g Na+ > Mg2+ > Al3+
Aluminium has stronger ionic bonding because it has a smaller ionic radii. Al3+ has a greater charge/size ratio, so stronger electrostatic attraction between oppositely charged ions, so stronger ionic bonding

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9
Q

For negative ions, when will the ionic bond be strongest?

A

the grater the charge, the stronger the ionic bonding

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10
Q

Which ionic compounds has stronger ionic bonds NaCl or KCl explain why?

A
  • NaCl has stronger ionic bonds
  • Na+ ions have a smaller ionic radii than K+ ions
  • so stronger electrostatic forces of attraction between Na+ and Cl- ions
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11
Q

Which ionic compounds have stronger ionic bonds MgO or MgCl2, explain why?

A
  • MgO has stronger ionic bonding
  • O2- ions have a greater negative charge than CL- ions
  • so stronger electrostatic attraction between Mg2+ and O2- ions
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12
Q

What is Definition for the Lattice Enthalpy Of Formation ΔH°LF? Give two examples.

A

The enthalpy change when 1 mole of an ionic lattice s formed from gaseous ions
e.g Na+ (g) + Cl- (g) —-> NaCl (s) ΔH°LF = -771 Kj mol-1
e.g Mg2+ (g) + o2- (g) —> MgO (s) ΔH°LF = -2500 Kj mol-1

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13
Q

What are the key ideas that all lattice formation enthalpies have?

A

All lattice formation enthalpies are large and negative, very exothermic as strong ionic bonds are formed

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14
Q

The stronger the ionic bond, the more..?

A

negative the lattice formation enthalpy

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15
Q

Explain why the ΔH°LF of BaO is more negative than the ΔH°LF of NaCl?

A
  • Ba2+ ions have a greater charge than Na+ and O2- have a greater charge than CL-,
  • so BaO has stronger ionic bonding,
  • as there are stronger electrostatic forces between Ba2+ and O2- ions
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16
Q

What is the definition for Lattice Dissociation enthalpy ΔH°LD?

A

enthalpy change when 1 mole of ionic lattice is converted into it’s gaseous ions

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17
Q

ΔH°LD is the opposite of..?

A

ΔH°LF.
ΔH°LD = - ΔH°LF

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18
Q

Lattice dissociation enthalpy ΔH°LD is always..?

A

Large, positive and endothermic as bonds are broken

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19
Q

Give an example

A

e.g NaCL (s) –> Na+ (g) + Cl- (g) ΔH°LD = +771 KJ mol-1

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20
Q

What is the definition for the standard enthalpy of atomisation ΔH°AT?
Give examples.

A

Enthalpy change when 1 mole of gaseous atoms are formed from it’s elements in their standard states
E.g Na (s) —> Na (g)
e.g Mg (s) –> Mg (g)
e.g 1/2 Cl2 (g) –> Cl (g)

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21
Q

this reaction is always, why…?

A

endothermic as heat energy is required to break bonds

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22
Q

Describe the trend in ΔH°AT down metal groups (1 and 2) ?

A
  • ΔH°AT decreases as metallic bonds are weaker and ionic radii increases
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23
Q

What is the definition for the standard Enthalpy of sublimation?
Give examples.

A

enthalpy change when 1 mole of a solid particles are converted to gaseous particles
e.g Na (s) —-> Na (g)
e.g Mg (s) —> (Mg (g)
e.g Cl2 (s) —> l2 (g)
e.g F2 (s) —> Cl 2 (g)

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24
Q

For metals and solids made up of atoms, the enthalpy of sublimation ΔH°Sub= …?

A

ΔH°at (atomisation)

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25
Q

For substances made up of molecules, The enthalpy of sublimation ≠..?

A

Does not equal the ΔH°at (atomisation)

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26
Q

What is the definition for the bond dissociation enthalpy?
Give examples.

A

energy change when 1 mole of a certain type of bond are broken with all species in the gaseous state
e.g Cl2 (g) —> 2 Cl(g)
e.g F2 (g) —-> 2 F (g)

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27
Q

How is it written as?

A

ΔH°D or ΔH°B.D

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28
Q

the enthalpy of bond dissociation = 2 x the bond enthalpy of…?
This is only true for what molecules?

A

ΔH°B.D = 2 ΔH°at (atomisation)
only true for diatomic molecules that are gases in their standard states

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29
Q

Give an example.

A

for bromine ΔH°B.D ≠ 2 ΔH°at as bromine is a liquid in t’s standard state

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30
Q

Calculate he ΔH°at for:
1. fluorine given that the ΔH°B.D = +196 KJ mol-1
2. Chlorine given then ΔH°B.D = +242 KJ mol -1

A
  1. ΔH°at = +98
  2. ΔH°at = +121
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31
Q

what is the definition for the Electron Affinity?

A

enthalpy change when 1 mole of gaseous atoms each gain one electron

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32
Q

What type of atoms is this for and why?

A

for non-metallic atoms as they form negative ions

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33
Q

Give examples?

A

e.g F (g) + e- –> F- (g) 1st e.a
e.g Cl (g) + e- –> Cl- (g)
e.g O (g) + e - –> O- (g) 1st e.a ΔH° e.a = -142 KJ mol-1
e.g O- + e- —> O2- (g) 2nd e.a ΔH° e.a - +242 KJ mol -1

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34
Q

why are all 1st Electron affinities (e.a) negative (exothermic)

A
  • as the electron added is attracted by the positive nucleus
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35
Q

Why is the second electron affinity positive (endothermic)?

A

due to repulsion between negative ion and the electron being added

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36
Q

What is the definition for the Ionisation enthalpy ΔH°I.E?

A

the enthalpy change when 1 mole of gaseous atoms each lose 1 electron to form gaseous ions

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37
Q

Give examples?

A

e.g Na (g) —> Na+ (g) + e-
e.g Mg (g) —-> Mg+ (g) + e - 1st I.E
e.g Mg + —> Mg 2+ (g) + e - 2nd I.E

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38
Q

Why does Ionisation energy decrease down a group?

A
  • due to larger atomic radii
  • so greater shielding
  • so the outer electron is attracted less strongly by the nucleus
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39
Q

why does the Ionisation energy increase across a period?

A
  • due to greater nuclear charge
  • smaller atomic radii
  • so outer electron is attracted more strongly by the nucleus
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40
Q

why is the 2nd IE greater than the 1st IE?

A
  • smaller ionic radii
  • so the nucleus of the positive ion attracts the electron more strongly than the neutral atom
41
Q

Draw out the Bon Haber Cycle for NaCl?

A
42
Q

Comparing Born-Haber Lattice Enthalpy with Theoretically calculated Lattice Enthalpy

A
43
Q

what is a perfect ionic model? - MSA

A
  • All ions are perfect charges/ spherical
  • All the attraction (bonding) is purely ionic (there is no covalent bonding)
  • All lattices are perfect (have no defects)
44
Q

Suggest two properties of ions that influence the value of a lattice enthalpy calculated using a perfect ionic model? - MSA

A
  • (Ionic) radius / distance between ions / size
  • (Ionic) charge / charge density
45
Q

The experimental Born Haber enthalpies for perfect ionic compounds are..?

A

ver close (within 2%) of the theoretical lattice enthalpies, these compounds have virtually 100% ionic bonding (perfect ionic model)

46
Q

The experimental Born Haber enthalpies for silver halides are..?

A
  • about 15% greater
  • so have stronger bonding than pure ionic bonding
  • have ionic bonding with covalent character
47
Q

what is a covalent character?

A

this means that negative ions are non-spherical/ distorted

48
Q

State what the term covalent character means and explain why AL203 is ionic with a covalent character

A

a covalent character is:
- distorted/ non-spherical (negative O2- ions)
- caused by small highly charged Al3+ ion
- causes some sharing of electron density

49
Q

Dissolving Ionic lattices

A
50
Q

when a solid ionic substance dissolves in water which two things happen?

A
  1. the bonds between the ions break (lattice dissociation enthalpy) - endothermic
  2. Bonds between the ions and water are made (hydration enthalpy) - exothermic
51
Q

why do the ions react with water?

A

the charges on the ion are attracted to the charges on the water molecule

52
Q

What is the definition for the standard enthalpy of solution?

A

the enthalpy change when 1 mole of an ionic lattice is completely dissolved in excess water to form aqueous ions

53
Q

Give examples of Standard enthalpy of solution?

A
  • NaCl (s) —> Na+ (aq) + Cl- (aq)
  • CaBR2 (s) —#> Ca2+ (aq) + 2BR- (aq)
54
Q

Are all Enthalpy of solution values exothermic or endothermic?

A

All ΔH°solution are small. Can be both exothermic and endothermic

55
Q

What is the standard enthalpy of hydration?

A

the enthalpy change when 1 mole of a gaseous ion is converted into aqueous ions.

56
Q

Give examples of this?

A
  • Na+ (g) —> Na+ (aq)
  • Cl- (g) —> Cl- (aq)
57
Q

Are ΔH°Hydration values exothermic or endothermic?

A

All ΔH° Hydration values are exothermic as hydration involves forming bonds with water

58
Q

Explain why the standard enthalpy of Na+ is more negative than K+?

A
  • Na+ has a smaller ionic radius, so greater electrostatic attraction between the Na+ and δ- on the polar water molecule
59
Q

Explain why the ΔH°hydration for Mg2+ is more negative than Na+?

A

Mg2+ has geater nuclear charge (and a smaller ionic radius), so greater electrostatic forces of attraction between the Mg2+ and the δ− in the polar water molecule

60
Q

How can we calculate the ΔH° solution?

A

using enthalpy cycle and
ΔH° solution = ΔH° solutionL.D + ΔH° solution Hydration

61
Q

Draw an enthalpy cycle and calculate the standard enthalpy of solution for NaCl

A
62
Q

Calculating Harder bond dissociation enthalpies

A
63
Q

In order to calculate bond dissociation enthalpies, what do we need to do in steps?

A
  1. Hess’s law cycle - to calculate the formation
  2. then from then you calculate the bond dissociation
64
Q

Calculate the ΔH for the following reaction:
C2H4 (g) + H2O (g) —> CH3CH2OH (g)

A

ΔH = -42kj mol-1

65
Q

Why might the bond enthalpy values change slightly?

A

Because all values are mean bond enthalpies calculated from a range of compounds containing that bond

66
Q

Entropy

A
67
Q

what helps us understand why some reactions occur spontaneously?

A

the concept of entropy can help us understand why

68
Q

What is entropy?

A

a measure of how much disorder there is

69
Q

What do substances like, and what do particles do?

A

Substances like disorder and the particles move to try and increase entropy

70
Q

what are three ways entropy can be increased?

A
  • change in physical state
  • dissolution
  • number of particles
71
Q

Why does Entropy increase when there is a change in state from a solid to a liquid and then a larger increase from a liquid to a gas?

A
  • there is a large increase in entropy when a solid turns into a liquid
  • Vibrating solid particles turn into randomly moving liquid particles, so there is a large increase in disorder
  • Even larger increase in entropy at the boiling point as randomly moving liquid particles turn into choatically and randomly moving gas particles
72
Q

The entropy of vaporisation…?

A

is greater than the entropy of melting

73
Q

Describe why dissolving a solution increases the entropy?

A
  • The particles become smaller
  • The particles can move freely, they are not fixed in position
74
Q

Why does entropy increase as the number of particles increase?

A
  • more particles mean more entropy
  • e.g 2O3 (g) –> 3O2 (g)
  • the Right Hand Side contains more moles of gas than the Left Hand Side
75
Q

What is the total entropy change?

A

the total entropy change is the sum of the entropy changes of the system and the surroundings

76
Q

What is the equation used to calculate the total entropy change?

What are the units

A

ΔS total =ΔS system + ΔS surroundings

units: J K-1 Mol-1

77
Q

How do we calculate the enthalpy of a system ΔS system?

A

ΔS products - ΔS reactants

78
Q

How do we calculate the enthalpy of surroundings?

A

ΔS surroundings = -ΔH / T
ΔH = enthalpy change
T = temperature in Kelvins

79
Q

Example - Calculate the total entropy change for the reaction of Ammonia and hydrogen chloride under standard conditions
NH3 () + HCl () —> NH4CL () . Where ΔH = -315 KJMol-1 at 298K.

A

Total entropy = + 7772.5 JK Mol-1

80
Q

What is entropy?

A

entropy is a measure of how disordered a system is?

81
Q

The higher the temperature, the..?

A

greater the disorder, so the greater the entropy of a substance

82
Q

At 0 Kelvins, what is the entropy of all subatnces, and why?

A

At 0 kelvins, all substances have an entropy of zero, because all substances are in their most ordered arrangement

83
Q

What are three general rules of entropy of a system?
ΔS system = ΔS products - ΔS reactants

A
  1. if the ΔS system is large and positive - there is a large increase in entropy and disorder, often when there is more moles of gaseous product
  2. If ΔS system is large and negative - then there is a large decrease in entropy and disorder, often when there is less moles of gaseous product
  3. if S system is small and slightly positive or negative - there is little change in entropy and disorder, when there are equal moles of gaseous reactants and products
84
Q

Describe entropy solid<Liquid<Gas

A

Entropy and disorder increases

85
Q

the larger and more complex the molecule and formula, the..?

A

Greater the entropy

86
Q

Describe the entropy change for the following:
1. H2O (l) –> H2O (g)
2. H20 (g) –> H2O (l)
3. C (diamond) (s)—-> C (graphite) (s)
4. C (s) + 1/2O2 (g) –> Co (g)
5. C (s) + O2 (g) –> CO2 (g)
6. CO (g) + 1/2 O2 –> CO2 (g)

A
  1. More moles of gaseous product, so large and positive increase in entropy
  2. less moles of gaseous product, so large and negative decrease in entropy
  3. equal moles of gas on both sides of equation, so small increase in entropy. Value is very low as both graphite and diamond are very ordered solids
  4. more moles of gaseous product, so large and positive increase in entropy
  5. equal moles on both sides of the equation, so small increase in entropy
  6. less moles of gaseous product, so large and negative decrease in entropy.
87
Q

Gibbs Free Energy

A
88
Q

what is the change in Gibbs free energy ΔG?

A

is a measure used to predict whether a reaction is feasible or not

89
Q

Gibbs free energy an explain why some..? reactions can happen spontaneously?

A

endothermic

90
Q

What is the equation used to calculate Gibbs free energy change?

A

ΔG = ΔH - TΔS

ΔG = Gibbs free energy change Kj mol-1
ΔH = enthalpy change
T = temperature in Kelvins
ΔS = enthalpy change of a system

91
Q

For a reaction to be feasible, Gibbs free energy change must be?

A

ΔG ≤ 0

92
Q

What does Feasible mean?

A

ΔG ≤ 0 but also means spontaneous, but even if a reaction is feasible the activation energy might be so high that the reaction rate is very slow, so you cannot observe it happening

93
Q

Calculate ΔG for the following reaction: C (s) + O2 (g) —> CO2 (g)

A

-350 KJ mol-1

94
Q

When are reactions feasible in terms of Gibbs free energy?

A

When Gibbs free energy is Negative - all reactions are feasible

95
Q

What is y=mx + c , using the ΔG = ΔH - TΔS

A

y =ΔG
m = ΔS
x = T
c = -ΔH

96
Q

Calculating the temperature a reaction becomes feasible?

A
  • when ΔG = 0, a reaction is just feasible
  • ΔG = ΔH - TΔS
  • When ΔG = 0, ΔG/TΔS
  • so T = ΔH/ ΔS
97
Q

Calculate the minimum temperature that the reaction becomes feasible (example 1)

A
98
Q

Calculate the minimum temperature that the reaction becomes feasible (example 2)

A
99
Q

DONE!!

A

Chemistry Y13 (13 decks)

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