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AQA A Level Chemistry > thermodynamics [P2] PAPER 1 > Flashcards

thermodynamics [P2] PAPER 1 Flashcards

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1
Q

gibbs equation and units?

A

• ΔG = ΔH – TΔS

• ΔG = Gibbs free energy change, kJmol-1
• ΔH = enthalpy change, kJmol-1
• T = temperature, K
• ΔS = entropy change, JK-1mol-1

⚠️ ΔS is in JOULES per kelvin per mole, while the rest are KILOJOULES, so need to convert to kJ when doing calculations

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2
Q

how is the feasibility of a reaction determined using the value of ΔG?

A

if ΔG is negative, it is feasible

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3
Q

State the meaning of the term enthalpy change.

A

heat energy change at constant pressure

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4
Q

Define the term enthalpy of lattice formation.

A

enthalpy change when one mole of ionic lattice is formed from its constituent gaseous ions under standard conditions

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5
Q

define the term electron affinity

A

the enthalpy change when one mole of gaseous atoms form one mole of gaseous negative ions under standard conditions

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6
Q

Define the term enthalpy of lattice dissociation

A

enthalpy change when separating one mole of ionic lattice into its constituent gaseous ions under standard conditions

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7
Q

Lattice enthalpy can be calculated theoretically using a perfect ionic model. Explain the meaning of the term perfect ionic model.

A

• ions are point charges
• ions are perfect spheres
• only electrostatic attraction between ions
• no covalent character
• only ionic bonding

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8
Q

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

A

• ion size
• charge density

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9
Q

lattice enthalpy can be defined as…

A

either enthalpy of lattice dissociation or enthalpy of lattice formation

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10
Q

define enthalpy of hydration

A

enthalpy change when one mole of gaseous ions forms one mole of aqueous ions under standard conditions

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11
Q

define enthalpy of solution

A

enthalpy change when one mole of ionic lattice is dissolved in water to infinite dilution, so that the ions no longer interact under standard conditions

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12
Q

define enthalpy of atomisation

A

enthalpy change when one mole of gaseous atoms are formed from the element in its standard state

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13
Q

why is enthalpy of hydration almost always negative?

A

• water is polar, with H δ+ and O δ+
• the negative ion attracts the H in water molecules
• the positive ion attracts the O in water molecules

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14
Q

calculating enthalpy of solution?

A

ΔH(solution) = ΔH(lattice dissociation) + Σ(ΔHhydration)

OR

ΔH(solution) = – ΔH(lattice formation) + Σ(ΔHhydration)

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15
Q

how to calculate at which temperature a reaction becomes feasible using the Gibbs free energy equation?

A

• ΔG = ΔH – TΔS

• make ΔG = 0 [as a reaction is feasible when ΔG <0], then solve for T

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16
Q

what is entropy?

A

measure of disorder

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17
Q

calculating entropy change, ΔS?

A

ΔS = S products – S reactants

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18
Q

ΔG = ΔH – TΔS as y = mx + c, with y-axis as ΔG in kJ/mol, and x-axis as temperature in K

A

ΔG = ΔH – TΔS
⇒ ΔG = – TΔS + ΔH
• ΔG = y
• ΔH = c
• T = x
• –ΔS = m

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19
Q

solid, liquid and gas states of a substance in order of increasing entropy, S?

A

S solid < S liquid < S gas

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20
Q

value of ΔG when finding a melting/boiling point, provided you have ΔH and ΔS?

A

use ΔG = 0, then rearrange Gibbs free energy equation for T

21
Q

why may a reaction not be feasible at very low/high temperatures?

A

substances may change state

22
Q
for Mg2+(g), enthalpy of hydration = –1920kJ/mol
A

• Ca2+ ion is bigger than Mg2+
• weaker attraction to O δ+ of water

23
Q

state the temperature at which the standard entropy value of __ is 0 J/K/mol

A

0K

24
Q

Why may there be a difference between theoretical (based on the perfect ionic model) and experimental values for enthalpies of formation?

A

has covalent character, as well as ionic bonding

25
Q

The experimental value of the enthalpy of lattice formation of an ionic substance is very close to the theoretical value, obtained using the perfect ionic model. What does this indicate about the bonding in this substance?

A

it is almost perfectly ionic

26
Q

A calculation of the enthalpy of lattice formation of silver iodide based on a perfect ionic model gives a smaller numerical value than the value calculated experimentally. Explain this difference.

A

• AgI contains covalent character
• forces holding the lattice together are stronger

27
Q
A

not possible to prevent some dissolving

28
Q
A
29
Q

Explain why the standard entropy value for carbon dioxide is greater than that for carbon.

A

• carbon dioxide is a gas and carbon is solid
• gas is more disordered than solid

30
Q
A

• ion size increases
• charge density of ion decreases
• attraction between metal ion and O δ– of water decreases

31
Q

general Born-Haber cycle?

A
i.e. 1 = 2 + 3 + 4 + 5
32
Q

Covalent characters arise due to distorted ions. As positive ions get ____ /____ , they are very good at distorting (i.e. polarising) negative ions. As negative ions get ____ / ____ , are easier to distort (i.e. they are polarisable).

A

• positive ions : smaller / more highly charged

• negative ions: larger / more highly charged

33
Q

lattice strength depends on…

A

• charge density
• ion size

34
Q

Born-Haber cycle for CaO?

A
35
Q

Born-Haber cycle for MgBr2?

A
36
Q
A

• At 0K, particles are stationary
• no disorder

37
Q

Explain, in terms of molecules, why the first part of the graph in Figure 1 is a line that slopes up from the origin.

A

• as temperature increases, particles start to move
• disorder increases

38
Q

In terms of the behaviour of molecules, explain why L2 is longer than L1 in Figure 1.

A

• L2 corresponds to boiling, and L1 corresponds to melting
• Bigger change in disorder for L2 compared with L1

39
Q
A

• ΔG = ΔH – TΔS
• ΔG = c, and –ΔS = m

40
Q

Explain why the magnitude of ΔG decreases as T increases in this reaction.

A

because ΔS is positive

41
Q

Explain why the lattice dissociation enthalpy of magnesium chloride is greater than that of calcium chloride.

A

• magnesium ion is smaller
• stronger attraction to chloride ion

42
Q

Explain why the lattice dissociation enthalpy of magnesium oxide is greater than that of magnesium chloride.

A

• oxide ion has a greater charge density
• attracts the magnesium in more strongly

43
Q

Energy is released when a magnesium ion is hydrated because magnesium ions attract water molecules.
Explain why magnesium ions attract water molecules.

A

• O on water has a δ– charge
• magnesium ion attracts the negative O on a water molecule

44
Q

Suggest why a value for the enthalpy of solution of magnesium oxide is not found in any data books.

A

magnesium oxide reacts with water to form Mg(OH)2

45
Q

Explain the significance, for this reaction, of temperatures below the temperature value where the line crosses the temperature axis.

A

• ΔG becomes negative
• reaction becomes feasible

46
Q

N2 + O2 → 2NO

Explain why the free-energy change for the reaction to form nitrogen monoxide stays approximately constant at different temperatures.

A

• no change in the number of moles of gas
• so, entropy change stays approximately constant

47
Q

The freezing of water is an exothermic process. Give one reason why the temperature of a sample of water can stay at a constant value of 0 °C when it freezes.

A

the heat given out escapes

48
Q
A

• ΔH is negligible as making and breaking the same number and types of bonds (bonds have similar enthalpies)
• increase in entropy (i.e. entropy is positive) as increasing the number of particles in solution (from 4 to 7 particles)
• ΔG = ΔH – TΔS
• ΔG is negative (for the forwards reaction) as -TΔS is negative, since ΔS is positive (and T is >0 )

AQA A Level Chemistry (36 decks)

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