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Chemistry revision > Thermodynamics > Flashcards

Thermodynamics Flashcards

1
Q

What does Hess’s Law state?

A

The enthalpy change for a reaction is independent of the route taken

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2
Q

Define standard enthalpy of formation.

A

The enthalpy change when one mole of a compound is formed from its elements in their standard states, under standard conditions

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3
Q

What is the standard enthalpy of an element?

A

Zero, by definition.

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4
Q

Define standard enthalpy of combustion

A

The enthalpy change when one mole of a substance is completely burnt in (excess) oxygen

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5
Q

Define standard enthalpy of atomisation

A

Enthalpy change when one mole of gaseous atoms is formed from elements in their standard state.

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6
Q

Define first ionisation energy

A

Enthalpy change when one mole of gaseous atoms is converted into 1 mole of gaseous 1+ ions

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7
Q

Define second ionisation energy.

A

Enthalpy change when one mole of electrons is removed from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions

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8
Q

Define first electron affinity

A

Enthalpy change when 1 mole of gaseous atoms forms 1 mole of gaseous 1- ions

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9
Q

Define second electron affinity.

A

Enthalpy change when one mole of gaseous 1- ions gains one mole of electrons to form one mole of gaseous 2- ions

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10
Q

Define lattice enthalpy of formation

A

Enthalpy change when one mole of solid ionic lattice is formed from its constituent gaseous ions.

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11
Q

Define lattice enthalpy of dissociation.

A

Enthalpy change when one mole of solid ionic lattice is dissociated (broken into) into its gaseous ions

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12
Q

Define enthalpy of hydration.

A

Enthalpy change when one mole of gaseous ions become hydrated/dissolved in water to infinite dilution [water molecules totally surround the ion]

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13
Q

Define enthalpy of solution

A

Enthalpy change when one mole of solute dissolves completely in a solvent to infinite dilution.

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14
Q

Define mean bond dissociation enthalpy

A

Enthalpy change when one mole of (a certain type of) covalent bonds is broken, with all species in the gaseous state

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15
Q

Write an example equation for standard enthalpy of formation

A

Mg (s) + ½ O2 (g) → MgO (s)

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16
Q

Write an example equation for standard enthalpy of combustion

A

CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)

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17
Q

Write an example equation for standard enthalpy of atomisation

A

½ I2 (g) → I (g)

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18
Q

Write an example equation for first ionisation energy

A

Li (g) → Li+ (g) + e-

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19
Q

Write an example equation for second ionisation energy

A

Mg+ (g) → Mg2+ (g) + e-

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20
Q

Write an example equation for first electron affinity

A

Cl (g) + e- → Cl - (g)

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21
Q

Write an example equation for second electron affinity

A

O- (g) + e- → O2- (g)

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22
Q

Write an example equation for lattice enthalpy of formation

A

Na+ (g) + Cl- (g) → NaCl (s)

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23
Q

Write an example equation for lattice enthalpy of dissociation

A

NaCl (s) → Na+ (g) + Cl- (g)

24
Q

Write an example equation for enthalpy of hydration

A

Na+ (g) → Na+ (aq)

25
Q

Write an example equation for enthalpy of solution

A

NaCl (s) → Na+ (aq) + Cl- (aq)

26
Q

Write an example equation for mean bond dissociation enthalpy

A

Br2 (g) → 2Br (g)

27
Q

What is a Born-Haber cycle?

A

Thermochemical cycle showing all the enthalpy
changes involved in the formation of an ionic compound.
Start with elements in their standard states (enthalpy of 0)

28
Q

What factors affect the lattice enthalpy of an ionic

compound?

A

Size of the ions, charge on the ions

29
Q

How can you increase the lattice enthalpy of a compound?

Why does this increase it?

A

Smaller ions, since the charge centres will be closer together.
Increased charge, since there will be a greater electrostatic attraction between the oppositely charged ions.
N.B. Increasing the charge on the anion has a much smaller effect than increasing the charge on the cation, since increasing anion charge also has the effect of increasing ionic size.

30
Q

How can Born-Haber cycles be used to see if compounds could theoretically exist?

A

Use known data to predict certain values of theoretical compounds, and then see if these compounds would be thermodynamically stable.
Was used to predict the existence of the first noble gas containing compound.

31
Q

What actually happens when a solid is dissolved in terms of interactions of the ions with water molecules?

A

Break lattice → gaseous ions; dissolve each gaseous ion in water.
The aqueous ions are surrounded by water molecules (which have a permanent dipole due to polar O-H bond)

32
Q

What is the perfect ionic model?

A

Assumes that ions are perfectly spherical and
that there is an even charge distribution (100%
polar bonds).
Act as point charges.

33
Q

Why i the perfect ionic model often not accurate?

A

Ions are not perfectly spherical.
Polarisation often occurs when small positive ions or large negative ions are involved, so the ionic bond gains covalent character. Some lattices are not regular and the crystal structure can differ.

34
Q

Which kind of bonds will be the most ionic? Why?

A

Between large positive ions and small negative ions e.g. CsF

35
Q

Define the terms spontaneous and feasible

A

If a reaction is spontaneous and feasible, it will take place of its own accord; does not take account of rate of reaction.

36
Q

Is a reaction with a positive or negative enthalpy change more likely to be spontaneous?

A

Negative - exothermic

37
Q

Define entropy

A

Randomness/disorder of a system.

Higher value for entropy = more disordered

38
Q

What units is entropy measured in?

A

JK-1mol-1

39
Q

What is the second law of thermodynamics?

A

Entropy (of an isolated system) always increases, as it is overwhelmingly more likely for molecules to be disordered than ordered

40
Q

Is a reaction with positive or negative entropy change more likely to be spontaneous?

A

Positive - reactions always try and increase the amount of disorder

41
Q

Compare the general entropy values for solids,

liquids and gases

A

Solids < liquids < gases

42
Q

How would you calculate the entropy change for a

reaction?

A

Entropy change = sum of products’ entropy - sum of reactants’ entropy
ΔS = ΣS(products) - ΣS(reactants)

43
Q

Define Gibbs free energy using an equation

A 
ΔG = ΔH - TΔS 
G = Gibbs free energy
H = enthalpy change
S = entropy change
T =temperature
44
Q

What does the value for Gibbs free energy for a

reaction show?

A

If G < 0, reaction is feasible. If G = 0, reaction is JUST feasible.
If G > 0, reaction is not feasible.

45
Q

What is the significance of the temperature at which

G = 0?

A

This is the temperature (in Kelvin) at which the reaction becomes feasible.

46
Q

How would you calculate the temperature at which a reaction becomes feasible?

A

Rearrange to T = (ΔH)/(ΔS) since G = 0

47
Q

What are the limitations of using G as an indicator of

whether a reaction will occur?

A

Gibbs free energy only indicates if a reaction is feasible.
It does not take into account the rate of reaction (the kinetics of the reaction).
In reality, many reactions that are feasible at a certain temperature have a rate of reaction that is so slow that effectively no reaction is occurring.

48
Q

If the reaction is exothermic and entropy increases, what is the value of G and what does this mean?

A

G always negative, so reaction is always feasible

- product favoured

49
Q

If the reaction is endothermic and entropy decreases, what is the value of G and what does this mean?

A

G always positive, so reaction is never feasible - reactant favoured

50
Q

If the reaction is exothermic and entropy decreases, what is the value of G and what does this mean?

A

Temperature dependent

51
Q

If the reaction is endothermic and entropy increases, what is the value of G and what does this mean?

A

Temperature dependent

52
Q

Why is entropy zero at 0K?

A

No disorder - molecules/atoms are not moving or vibrating and cannot be arranged in any other way.
Max possible state of order

53
Q

What are the two key things to look out for to decide if entropy increases/decreases/stays relatively constant?

A

Number of moles - more moles made → increase in entropy

Going from solid → liquid/gas or liquid → gas

54
Q

How is it possible for the temperature of a substance

undergoing an endothermic reaction to stay constant?

A

The heat that is given out escapes to the surroundings

55
Q

What is enthalpy?

A

Heat measured at constant pressure

56
Q

What is mean bond enthalpy?

A

The enthalpy to break a bond, averaged over different molecules

Chemistry revision (46 decks)

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