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Chemistry > 1.8 Thermodynamics > Flashcards

1.8 Thermodynamics Flashcards

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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 it constituent elements under standard conditions, with all products and reactants in their standard states
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3
Q

What is the standard enthalpy of an element?

A
  • 0
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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 oxygen under standard conditions, with all products and reactants in their standard states.
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5
Q

Define standard enthalpy of atomisation.

A
  • the enthalpy change when one moles of gaseous atoms is formed from a compound in its standard state, under standard conditions
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6
Q

Define first ionisation energy.

A
  • the enthalpy change when one mole of electrons is removed from one mole of gaseous atoms to form one mole of gaseous 1+ ions
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7
Q

Define second ionisation energy.

A
  • the 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
  • the enthalpy change when one mole of gaseous atoms gains one mole of electrons to form one mole of gaseous 1- ions
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9
Q

Define second electron affinity.

A
  • the enthalpy change when one moles 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
  • the enthalpy change when one mole of solid ionic lattice is formed from from its constituent gaseous ions
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11
Q

Define lattice enthalpy of dissociation.

A
  • the enthalpy change when one mole of solid ionic lattice is dissociated into its gaseous ions
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12
Q

Define enthalpy of hydration.

A
  • the enthalpy change when one mole of aqueous ions is formed from gaseous ions
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13
Q

Define enthalpy of solution.

A
  • the enthalpy change when 1 mole of an ionic substance dissolves in enough solvent to form an infinitely dilute solution
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14
Q

Define bond dissociation enthalpy.

A
  • the enthalpy change when all the bonds of the same type in 1 mole of gaseous molecules are broken
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15
Q

Write example equations for:
- Standard enthalpy of formation
- Standard enthalpy of combustion
- Standard enthalpy of atomisation
- First ionisation energy
- Second ionisation energy
- First Electron affinity

A
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16
Q

Write example equations for:
- Second electron affinity
- Lattice enthalpy of formation
- Lattice enthalpy of dissociation
- Enthalpy of hydration
- Enthalpy of solution
- Bond dissociation enthalpy

A
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17
Q

What is a Born-Haber cycle?

A
  • thermochemical cycle showing all the enthalpy changes involved in the formation of an ionic compound
  • start will elements in their standard states (enthalpy of 0)
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18
Q

Draw a labelled Born-Haber cycle for the formation of sodium chloride lattice.

A
19
Q

What factors effect the lattice enthalpy of an ionic compound?

A
  • size of the ions
  • charge on the ions
20
Q

How can you increase the lattice enthalpy of a compound? Why does this increases it?

A
  • smaller ions, since the charge centres will be closer together
  • increased charge, since there will be a greater electrostatic force of attraction between the oppositely charged ions
    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
21
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 theses compounds would be thermodynamically stable.
    (was used to predict the existence of the first noble gas containing compound)
22
Q
  • What actually happens when a solid is dissolved in terms of interaction of the ions with water molecules?
A
  • break bonds in lattice to give gaseous ions (lattice dissociation enthalpy)
  • dipole of polar O-H bond in water allows the partially negative oxygen to weakly bond with the positive ion and the partially positive hydrogen’s to weakly bond with the negative ion
23
Q

What is the perfect ionic model?

A
  • assumes that ions are perfectly spherical and that there is an even charge distribution
24
Q

Why is the perfect ionic model often inaccurate?

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
25
Q

Which kind of bonds will be the most ionic?

A
  • ones between large positive ions and small negative ions
    e.g. CsF
26
Q

Define the terms spontaneous and feasible.

A
  • if a reaction is spontaneous and feasible, it will take place of its own accord and does not take account of rate of reaction
27
Q

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

A
  • negative - exothermic
28
Q

Define Entropy.

A
  • measure of the disorder of a system
  • higher entropy value = more disorder
29
Q

What units is entropy measured in?

A
  • JK‾¹mol‾¹
30
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
31
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
32
Q

Compare the general entropy values for solids, liquids and gases…

A
  • solids < liquids < gases
33
Q

How would you calculate the entropy change for a reaction?

A
34
Q

Define Gibbs free energy using an equation.

A

where G = gibbs free energy, H = enthalpy change, S = entropy change and T = temperature

35
Q

What does the value for Gibbs free energy of a reaction show?

A
  • if G < 0 reaction is feasible
  • if G = 0 reaction is just feasible
  • if G > 0 reaction is not feasible
36
Q

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

A
37
Q

If ΔH is negative (exothermic) and ΔS is positive, what is the value of G and what does this mean?

A
  • G is always negative, so reaction is always feasible
  • product favoured
38
Q

If ΔH is positive (endothermic) and ΔS is negative, what is the value of G and what does this mean?

A
  • G is always positive, so reaction is never feasible
  • reactant favored
39
Q

If ΔH is negative (exothermic) and ΔS is negative, what is the value of G and what does this mean?

A
  • it means the reaction is temperature dependent
40
Q

If ΔH is positive (endothermic) and ΔS is positive, what is the value of G and what does this mean?

A
  • it means the reaction is temperature dependent
41
Q

Why is entropy zero at 0K?

A
  • no disorder as molecules/atoms are not moving or vibrating and cannot be arranged in any other way.
  • maximum possible state of order
42
Q

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

A
  • more moles made = increase in entropy
  • going from solid to liquid/gas or liquid to gas
43
Q

How is it possible for the temperature of a substance undergoing an endothermic reaction to stay constant?

A
  • heat that is given out escapes to the surroundings

Chemistry (24 decks)

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