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A-Level Edexcel Chemistry > Chapter 13 - Further Energetics > Flashcards

Chapter 13 - Further Energetics Flashcards

1
Q

What is standard lattice formation enthalpy?

A

Where one mole of an ionic solid is formed from its gaseous ions in standard conditions

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2
Q

What factors affect standard lattice enthalpy?

A

The charges on the ions

The radius of the ions

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3
Q

What is the relationship between standard lattice enthalpy and the magnitude of charges on the ions?

A

The greater the magnitude of charges on the ions, the more negative the value for the lattice enthalpy

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4
Q

What is the relationship between standard lattice enthalpy and the size of the ions?

A

The smaller the ions (the smaller the inter-ionic distance), the more negative the value for the lattice enthalpy

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5
Q

What is the inter-ionic distance?

A

The distance between the centre of two ions

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6
Q

What is standard enthalpy change of atomisation?

A

The enthalpy change when one mole of gaseous atoms are formed from an element in its standard state at standard conditions

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7
Q

Give an example of the atomisation of hydrogen

A

1/2 H2(g) -> H(g)

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8
Q

What is first electron affinity?

A

The enthalpy change when each atom in one mole of atoms in the gaseous state gain an electron to form one mole of -1 ions

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9
Q

Give an example of the first electron affinity of oxygen

A

O(g) + e- -> O-(g)

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10
Q

Is first electron affinity exothermic or endothermic?

A

Exothermic (negative)

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11
Q

What is the exception to first electron affinity being exothermic?

A

Noble gases first electron affinity is not exothermic

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12
Q

Is second electron affinity endothermic or exothermic?

A

Endothermic (positive)

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13
Q

Why is second electron affinity endothermic?

A

Because the electron is added to an ion which is already negative therefore it must overcome the repulsion

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14
Q

What assumptions are made when finding theoretical lattice energy?

A
  • The ions are in contact with one another
  • The ions are perfectly spherical
  • The charge on each ion is evenly distributed around
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15
Q

What causes differences between theoretical and experimental lattice energy?

A

Covalent character in the bonding of an ionic lattice

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16
Q

What causes covalent character?

A

The polarisation of the anion by the cation

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17
Q

What does the polarisation of the anion by the cation mean?

A

There is a distortion in the electron density within the anion, resulting in a higher electron density near the cation i.e. there is some electron density shared between the two ions (covalent character)

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18
Q

Why does the cation polarise the anion?

A

The positive cation will attract the negative electrons of the anion

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19
Q

What two factors affect the extent of covalent character?

A

The size of the ions

The charge of the ions

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20
Q

What will increase the extent of polarisation?

A
  • High charge and small cation (i.e high charge density)

- High charge and large anion

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21
Q

What is the ability of a cation to attract electrons from the anion known as?

A

Polarising power

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22
Q

What can be calculated to give an approximate value of the polarising power of the cation?

A

Charge density (charge/r^2)

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23
Q

What will be the effect on the difference between theoretical and experimental lattice energy of a cation with high polarising power?

A

High polarising power = greater covalent character = greater difference

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24
Q

What factors affect how easily an anion is polarised?

A

The larger the charge and the larger the size of the anion, the more easily it is polarised

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25
Q

What is the enthalpy change of solution?

A

The enthalpy change when one mole of an ionic solid dissolves in water to form an infinitely dilute solution

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26
Q

What is the enthalpy change of hydration?

A

The enthalpy change when one mole of an ion in its gaseous state is completely hydrated by water

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27
Q

What is the point of infinite dilution?

A

The point when further dilution has no measurable effect on the enthalpy of solution

28
Q

Can the value of enthalpy of solution be determined directly experimentally?

A

No it cannot, so it must be found by a process of extrapolation

29
Q

How do we know if something has been completely hydrated by water?

A

It is at infinite dilution, meaning further dilution has no measurable effect

30
Q

In the hydration of NaCl, what interactions occur with Na+?

A

Ion-dipole interactions between the delta negative oxygen and the positive Na+ ion

31
Q

In the hydration of NaCl, what interactions occur with Cl-?

A

Ion-dipole interactions as well as some hydrogen bonds between the Cl- and the hydrogen atoms

32
Q

What factors affect hydration enthalpy?

A

The larger charge on the ions and the smaller the size of the ionic radius, the more negative the value (the more energy released upon hydration decreases)

33
Q

What kind of process is enthalpy of hydration?

A

Exothermic

34
Q

What kind of process is lattice formation enthalpy?

A

Exothermic

35
Q

What kind of process is enthalpy of solution?

A

It can be exothermic or endothermic

36
Q

How can you deduce enthalpy of solution?

A

By forming a Born-Haber cycle, and using the values of the enthalpy of hydration and the lattice formation enthalpy to find the value of enthalpy of solution

37
Q

What is a spontaneous process?

A

A process that takes place without continuous intervention from us

38
Q

What is the relationship between entropy and disorder?

A

Increasing disorder = Increasing entropy

39
Q

What is entropy?

A

A property of matter that is associated with the degree of disorder/randomness of the particles

40
Q

What is the order of entropy from solids -> liquids -> gases

A

Solid < Liquid < Gas

Solid has the lowest entropy, gas has the highest

41
Q

What two components make up the total entropy change?

A

Entropy change of the system

Entropy change of the surroundings

42
Q

What is the total entropy change?

A

The sum of the entropy change of the system and the entropy change of the surroundings

43
Q

For a reaction to be spontaneous, what must total entropy change be?

A

Total entropy change must be positive for a reaction to be spontaneous

44
Q

How do you calculate the entropy change of a system?

A

Sum of entropy of products - Sum of entropy of reactants

45
Q

How do you calculate the entropy of surroundings?

A
  • (Change in enthalpy / Temperature in kelvin)
46
Q

What mustn’t you forget when calculating entropy of surroundings?

A

That there is a - sign before the fraction!!!

47
Q

For exothermic reactions, what will entropy of surroundings always be?

A

Positive i.e. entropy has increased

48
Q

For endothermic reactions, what will entropy of surroundings always be?

A

Negative i.e. entropy has decreased

49
Q

What is the effect of temperature on entropy?

A

Entropy increases with increased temperature

50
Q

What is the effect of an increased temperature on the entropy of a cold object compared to an already hot object?

A

The increase in heat energy (temp) will increase the entropy of the cold object to a greater degree than the already hot object, as in the hot object the molecules are already moving around vigorously

51
Q

How does a change in the number of moles from reactants to products affect entropy of system?

A

More moles = more particles = more ways in which the particles can be arranged = increased entropy

52
Q

What will the value of entropy of system be when the number of moles increases from reactants to products?

A

Entropy of system will be positive (increased)

53
Q

What will the value of entropy of system be when the number of moles decreases from reactants to products?

A

Entropy of system will be negative (decreased)

54
Q

What is the formula for Gibbs energy?

A

Gibbs energy = change in enthalpy - (temperature x change in entropy of system)

55
Q

What does it mean for a reaction to be thermodynamically feasible?

A

It is spontaneous

56
Q

What does it mean for the value of Gibbs energy to be less than 0 (negative)?

A

The reaction is thermodynamically feasible at that temperature

57
Q

What does it mean for the value of Gibbs energy to be more than 0 (positive)?

A

The reaction is not thermodynamically feasible at that temperature

58
Q

What does it mean for the value of Gibbs energy to be 0?

A

The reaction is in equilibrium

59
Q

How do you rearrange the Gibbs energy formula to find the minimum temperature at which a reaction is thermodynamically feasible?

A

Change in enthalpy / Change in entropy of system = Minimum temperature

60
Q

What is the equation that relates Gibbs energy and equilibrium constant?

A

Gibbs = - (Gas constant x Temp x ln(equilibrium constant)

61
Q

How can you rearrange the Gibbs energy equation to find the equilibrium constant?

A

e^ - (Gibbs/Gas constant x Temp)

62
Q

What does it mean for the value of Gibbs energy to be less than 0 (negative) on the value of the equilibrium constant?

A

The equilibrium constant will be greater than 1, meaning the products will be favoured

63
Q

What does it mean for the value of Gibbs energy to be more than 0 (positive) on the value of the equilibrium constant?

A

The equilibrium constant will be less than 1, meaning the reactants will be favoured

64
Q

If the value of Gibbs energy is positive, does this mean the reaction has not occured?

A

Not necessarily. It just means the position of equilibrium would favour the reactants side, meaning it is possible for some of the reactants to be converted into products, just not all of them

65
Q

What assumptions do you make when finding the minimum temperature at which a reaction is thermodynamically feasible?

A

That enthalpy of reaction and entropy of system do not change with temperature

66
Q

What are the limitations of using Gibbs energy?

A
  • The reactants may be kinetically stable, meaning a the activation energy for the reaction is high, so very few particles result in reaction
  • It is calculated only for standard conditions (100kPa 298K and 1 mol dm-3)
67
Q

Why is CaCO3 thermodynamically stable at room temperature, but not higher temperatures?

A

Because the total enthalpy change is positive.
Gibbs = Total enthalpy change - Entropy of system x Temp. Given that we know the entropy change of system is positive, if total enthalpy was negative, this would make the decomposition of CaCO3 occur at all temperatures, therefore not stable at room temp.

A-Level Edexcel Chemistry (24 decks)

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