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Edexcel Chemistry A level > Energetics II > Flashcards

Energetics II Flashcards

1
Q

Define lattice enthalpy of formation

A

The enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions

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2
Q

define entropy

A
  • measure of disorder
  • number of ways that energy can be shared between particles
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3
Q

Define enthalpy change of atomisation

A

the enthalpy change when 1 mole of gaseous ions is made from an element in its standard state.

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4
Q

Define 1st electron affinity

A

the enthalpy change when 1 mole of gaseous 1- ions are formed from 1 mole of gaseous atoms.

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5
Q

Name 2 factors that affect strength of ionic bonding

A
  • size of charge
  • ionic radius
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6
Q

Explain how Size of charge effects strength of ionic bond and melting and boiling point (2)

A
  • larger charge = stronger electrostatic attraction between ions
  • therefore, more energy needed to overcome these forces, leading to a higher melting and boiling point.
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7
Q

Describe and explain how ionic radius affects the strength of ionic bonds (2)

A
  • smaller ion = stronger electrostatic attraction between ions
  • can also pack closer together, more energy required to overcome the stronger forces = higher boiling and melting points.
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7
Q

Describe the structure of a born haber cycle from bottom to top. (enthalpies) and whether the steps are exothermic or endothermic. (5)

A
  1. original ionic compound (formation of which is exothermic)
  2. enthalpy of atomisation (for all elements if needed) (endothermic)
  3. ionisation energy of Cation (endothermic)
  4. 1st electron affinity of Anion (endothermic)
  5. (going right now) enthalpy of lattice (exothermic) formation.
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8
Q

what are the assumptions that theoretical values of lattice energy make? (2)

A
  • assume a perfect ionic model
  • 100%, and attractions are purely electrostatic.
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9
Q

Why might theoretical and actual (born haber) values be different? (3)

A
  • different if the compound shows covalent character - ie the cation distorts the anion slightly (kind of like they are sharing)
  • The higher the covalent character, the higher the larger the difference would be in the theoretical and Born Haber value.
  • the Born Haber value for a compound would be larger than the theoretical when the compound has some covalent character.
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10
Q

Define enthalpy of solution.

A
  • the enthalpy change when 1 mole of an ionic solid dissolves in a large enough amount of water to ensure they don’t interact.
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11
Q

Define enthalpy of lattice formation

A
  • the enthalpy change when 1 mole of an ionic crystal is formed from its constituent gaseous ions.
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12
Q

define enthalpy of hydration

A
  • when one mole of gaseous ions become hydrated such that further dilution causes no heat change.
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13
Q

How do we calculate enthalpy change of solution?

A

Enthalpy change of solution = -Enthalpy of lattice formation + sum of enthalpy of hydration.

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14
Q

What can be concluded about a substance from its enthalpy change of solution?

A
  • substance more likely soluble if exothermic
  • substance more likely to be insoluble if endothermic (as lattice energy much larger than hydration enthalpy.)
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15
Q

what happens to the entropy when a solid dissolves into ions?

A
  • entropy increases as there is more disorder, as solid changes to solution and number of particles increases.
  • this can make the total ΔSsystem can make the ΔStotal positive even if ΔH solution is endothermic, esp at higher temps.
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16
Q

what is a system?

A
  • your reactants and products
  • does not change temp or pressure
  • mass cannot be transferred to the surroundings
  • energy can be transferred to the surroundings.
17
Q

why do solids have lower entropies than liquids, which are lower than gasses?

A
  • as a solid is heated up, and turns into liquid, its particles vibrate more, meaning that its entropy increases
  • gasses are most disordered, so have the highest enthalpy.
18
Q

What will an increase in disorder and entropy lead to

A

a positive entropy change ( ΔS’ system = +ve)

19
Q

Give the equation to calculate ΔS system.

A

= ΣS products - ΣS reactants

20
Q

What is the unit for entropy?

21
Q

Give the equation to calculate ΔS surroundings
what are some things to keep in mind when using the equation?

A

-ΔH reaction/T
- covert ΔH to J/mol (x1000)
- convert degrees to Kelvin (+273)

22
Q

Give the equation to calculate total entropy change.

A

ΔS total = ΔS system + ΔS surrounding

23
Q

What must be positive for a reaction to be spontaneous?

24
What is Gibbs free energy?
- balance between entropy and enthalpy that determines the feasibility of a reaction.
25
What must be the condition for a reaction to be feasible?
the ΔG (gibbs free energy) must be LESS than 0.
26
Give the equation of Gibbs free energy.
ΔG = ΔH - TΔS system (gibbs free energy = enthalpy change - (Temp x entropy change of system).
27
What will happen if the reaction involves a decrease in entropy? (ΔG is +ve)
- less likely that ΔG is negative, and less likely that the reaction will occur.
28
What happens if the reaction involves an increase in entropy?
- more likely the reaction tales place and more likely that ΔG is -ve.
29
When y=mx+c is applied to the Gibbs energy equation. give the equation that results.
c = ΔH m=-ΔS
30
Which Kc value predicts that the reaction will not happen?
Kc < 10^-10
31
Which Kc value predicts that the reactants dominate in eq
Kc = 0.1
32
Which Kc value predicts that the reaction has equal amounts of reactants and products?
- Kc = 1
33
Which Kc value predicts that the products predominate in eq?
Kc = 10
34
Which Kc value predicts that the reaction goes to completion?
- Kc > 10^10
35
What can limit a reaction, causing it to not go to completion?
- if it has a high activation energy, kinetic factors limit the reaction.
36
What is the cycle that links enthalpy of hydration, solution, and Lattice enthalpy.
ΔsolutionH LiCl -----> Li+ + Cl- I ΔDissH I ΔhydH V V Li+ + Cl-
37
How to find the temp at what the reaction is feasible?
rearrange the Activation energy equation - LnK = -ΔG/RT
38
Explain, in terms of entropy, why the combustion of prop-2-en-1-ol is always feasible in the gaseous state.
ΔH is negative and ΔsSystem is positive - so ΔG is always negative
39

Edexcel Chemistry A level (18 decks)

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