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Alevel Chemistry Edexcel > Topic 13 - Energetics 2 > Flashcards

Topic 13 - Energetics 2 Flashcards

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1
Q

What is lattice enthalpy?

A

the energy change when one mole of an ionic
solid is formed from its gaseous ions
-always -ve

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2
Q

What is enthalpy change of atomisation?

A

the energy change to form a gaseous atom from its element in its standard state

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3
Q

What is first electron affinity?

A

the energy released when one electron is added to one mole of gaseous ions
-forming negative ions

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4
Q

What is a Born-Haber cycle used for?

A

to calculate an unknown enthalpy change

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5
Q

What do arrows going up in a Born-Haber cycle indicate?

A

energy required

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6
Q

What do arrows going down in a Born-Haber cycle indicate?

A

energy released

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7
Q

What factors affect lattice enthalpy?

A
  • charge (higher charge = less -ve lattice enthalpy)

- ionic radius (smaller ionic radius = less -ve lattice enthalpy)

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8
Q

What is the difference between theoretical and experimental lattice enthalpy values?

A

theoretical value -> based upon mathematical model (electrostatic theory) and assumes point charges
-less -ve value (more exothermic)

experimental value -> based upon experimental data (from a Born-Haber cycle)
-takes into account distortion of anion’s electron density (distortion makes bond more covalent + stronger -would release more energy)

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9
Q

What is enthalpy change of solution?

A

the energy used when one mole of a solid compound is dissolved in sufficient solvent to give an infinitely dilute solution
(whole dissolving process)

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10
Q

ΔHsol =

A

-ΔHlatt + 2ΔHhyd

can be used on a Hess cycle

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11
Q

What is enthalpy change of hydration?

A

the energy used when one mole of gaseous ions is dissolved in sufficient solvent to give an infinitely dilute solution

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12
Q

What happens when an ionic solid dissolves?

A

negative lattice enthalpy/lattice disassociation

  • lattice breaks up into separate gaseous ions
  • endothermic

hydration enthalpy

  • solid is surrounded by water molecules
  • attraction between ions (in solid) and dipoles (in water)
  • exothermic
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13
Q

What does the magnitude of lattice enthalpy depend on?

A

the strength of the forces acting on the ions

  • ionic charges (higher charge = greater lattice enthalpy)
  • ionic radii (smaller radii = greater lattice enthalpy)
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14
Q

What impact does ionic charge have on lattice enthalpy?

A

the greater the charge, the greater the lattice enthalpy

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15
Q

What impact does ionic radii have on lattice enthalpy?

A

the larger the radii, the smaller the lattice enthalpy

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16
Q

What does the magnitude of hydration enthalpy depend on?

A

the strength of the force between the ion and surrounding water molecules

  • ionic charge (higher charge = great force)
  • ionic radius (smaller radius = greater force)
17
Q

What impact does ionic charge have on hydration enthalpy?

A

the greater the charge, the greater the hydration enthalpy

18
Q

What impact does ionic radii have on hydration enthalpy?

A

the larger the radius, the smaller the hydration enthalpy

19
Q

What is entropy?

A

a measure of disorder

  • symbol = S
  • units = JK-¹mol-¹
  • naturally increases
20
Q

What are the three rules of thermodynamics?

A

1) energy can not be created or destroyed
2) every spontaneous process must result in an increase in disorder
- the entropy of the most perfect crystalline solid at 0K (absolute 0) would have an entropy of 0JK-¹mol-¹

21
Q

What factors affect entropy?

A
  • state (solid→gas = increasing entropy)

- complexity (the higher the MMR of the compound, the more electrons it has so the greater the disorder)

22
Q

ΔSsys =

A

ΣΔSprods - ΣSreacts

23
Q

Why would an entropy change of system be negative?

A

there’s a decrease in disorder

eg. 2 mols of ordered solid and 1 mol of highly disordered gas are being converted into 2 mols of ordered solid

24
Q

Why would an entropy change of system be positive?

A

there’s an increase in disorder

eg. 1 mol of ordered solid is being converted into 1 mol of ordered solid and 1 mol of highly disordered gas

25
Q

ΔSsurr =

A

-ΔH
____
T

NB.
ΔH is measured in kJmol-1 however ΔS is measured in JK-1mol-1 so ΔH must be multiplied by 1000 (to convert into Jmol-1)
temp in Kelvin

26
Q

ΔStot =

A

ΔSsys + ΔSsurr

27
Q

When is a reaction feasable?

A

when ΔStot is positive

28
Q

Why does a reaction not happen even though it has a positive total entropy?

A
  • reactants are thermodynamically unstable (likely to turn into products)
  • however… reactants are thermodynamically stable (high activation energy doesn’t allow them to react)
29
Q

ΔG =

A

ΔH-TΔSsys

NB.
ΔG and ΔH are measured in kJmol-1 but ΔS is measured in JK-1mol-1 so ΔS must be divided by 1000 (to convert into kJK-1mol-1)

30
Q

What does a negative ΔS say about a reaction’s feasibility?

A

rxn is not feasible

31
Q

What does a positive ΔS say about a reaction’s feasibility?

A

rxn is feasible

32
Q

What does a negative ΔG say about a reaction’s feasibility?

A

rxn is feasible

33
Q

What does a positive ΔG say about a reaction’s feasibility?

A

rxn is not feasible

34
Q

How do you calculate the minimum temperature to make a non-feasible reaction feasible USING ΔS?

A

Use ΔSurr = -ΔH/T rearranged to T=ΔH/ΔSsurr

use the opposite +/- to ΔSsys as ΔSsurr

35
Q

How do you calculate the minimum temperature to make a non-feasible reaction feasible USING ΔG?

A

For a feasible rxn ΔG<0 so use ΔG=ΔH-TΔSsys to find minimum temp by replacing ΔG with 0 so that ΔH-TΔSsys>0

36
Q

At equilibrium, what does a -ve ΔStot/a +ve ΔG mean?

A

backwards reaction is favoured

equilibrium lies to the left

37
Q

ΔG =

in terms of equilibrium

A

RTlnK

rearranged:
K = e ^(-ΔG/RT)

38
Q

ΔStot =

in terms of equilibrium constant

A

RlnK

Alevel Chemistry Edexcel (19 decks)

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