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Chemistry > 1.8 Thermodynamics > Flashcards

1.8 Thermodynamics Flashcards

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1
Q

What is ∆H the symbol for?

A

Enthalpy change.

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2
Q

What is enthalpy change?

A

Enthalpy change is the heat energy transferred in a reaction at constant pressure.

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3
Q

What does ∆H* mean?

A

The enthalpy change was measured under standard conditions (298K and 100kPa)

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4
Q

Do exothermic reactions have a negative or positive ∆H value?

A

Negative, because heat energy is given out.

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5
Q

Do endothermic reactions have a negative or positive ∆H value?

A

Positive, because heat energy is absorbed.

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6
Q

Define the Enthalpy change of formation. (∆fH)

A

The standard enthalpy change of formation of a compound is the energy transferred when 1 mole of the compound is formed from its elements under standard conditions, all reactants and products being in their standard states.

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7
Q

Define the Enthalpy change of atomisation of an element. (∆atH)

A

The enthalpy of atomisation of an element in the enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state.

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8
Q

Define the First ionisation energy. (∆ie1H)

A

The first ionisation enthalpy is the enthalpy change required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a +1 charge.

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9
Q

Define the First electron affinity. (∆ea1H)

A

The first electron affinity is the enthalpy change that occurs when 1 mole of gaseous atoms gain 1 mole of electrons to form 1 mole of gaseous ions with a -1 charge.

The first electron affinity is exothermic for atoms that normally form negative ions. This is because the ion is more stable than the atom, and there is an attraction between the nucleus and the electron.

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10
Q

Define the Enthalpy change of hydration. (∆hydH)

A

The enthalpy change when 1 mole of gaseous ions become 1 mole of aqueous ions.

This is always an exothermic reaction because bonds are made between the ions and the water molecules.

(see page 1 in the chemrevise revision guide for an example equation)

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11
Q

Define the enthalpy of lattice dissociation enthalpy. (∆dissH)

A

The standard enthalpy change when 1 mole of an ionic crystal lattice form is separated into its constituent ions in gaseous form.

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12
Q

Define the Enthalpy change of atomisation of a compound. (∆atH)

A

The enthalpy change when 1 mole of a compound in its standard state is converted to gaseous atoms.

(see page 54 in the revision guide for an example equation)

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13
Q

Define the Second ionisation energy. (∆ie2H)

A

The second ionisation energy is the enthalpy change to remove 1 mole of electrons from 1 mole of gaseous 1 + ions to produce 1 mole of gaseous 2+ ions.

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14
Q

Define the Second electron affinity. (∆ea2H)

A

The second electron affinity is the enthalpy change when 1 mole of gaseous 1- ions gains 1 electron per ion to produce gaseous 2- ions.

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15
Q

Define the Enthalpy change of solution. (∆solutionH)

A

The enthalpy of solution is the standard enthalpy change when 1 mole of an ionic solid is dissolved in enough solvent that no further enthalpy change occurs on further dilution.

(see page 54 in the revision guide for an example equation)

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16
Q

Define the Lattice enthalpy of formation. (∆lattH)

A

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

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17
Q

Define the Lattice enthalpy of dissociation. (∆lattH)

A

The lattice enthalpy of dissociation is the standard enthalpy change when 1 mole of a solid ionic compound is completely dissociated into its constituent gaseous ions under standard conditions.

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18
Q

Draw a model Born-Haber cycle.

A

(see page 57 in the revision guide)

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19
Q

Calculate the lattice enthalpy of NaCl.

(/\fH) = -411kJ mol-1
(/\atH) of Cl = +122kJ mol-1
(/\atH) of Na = +107kJ mol-1
(/\ieH) of Na = +496kJ mol-1
(/\eaH) of Cl = -349kJ mol-1
A

(see page 57 in the revision guide)

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20
Q

Calculate the atomisation enthalpy of Mg.

(/\fH) of MgO = -548kJ mol-1
(/\atH) of O = +249kJ mol-1
(/\ieH) of Mg = +738kJ mol-1
(/\ie2H) of Mg = +1451kJ mol-1
(/\eaH) of O = -141kJ mol-1
(/\ea2H) of O = +798kJ mol-1
(/\latticeH) of formation = -3791kJ mol-1
A

(see page 58 in the revision guide)

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21
Q

Why are theoretical lattice enthalpies often different from experimental values?

A

(see page 59 in the revision guide)

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22
Q

How can comparing the different values of the theoretical and experimental lattice enthalpy tell you ‘how ionic’ an ionic lattice is?

A

(read page 59 in the revision guide)

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23
Q

What happens when a solid ionic lattice dissolves in water?

A

(see page 60 in the revision guide)

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24
Q

Calculate the enthalpy change of solution for sodium chloride.

(/\dissH) = +787kJ mol-1
(/\hydH) of Cl- = -364kJ mol-1
(/\hydH) of Na+ = -406kJ mol-1

A

(see page 60 in the revision guide)

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25
Q

Calculate the enthalpy change of solution for silver chloride.

(/\dissH) = +905kJ mol-1
(/\hydH) of Cl- = -364kJ mol-1
(/\hydH) of Ag+ = - 464kJ mol-1

A

(see page 61 in the revision guide)

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26
Q

What is Entropy? (S)

A

The number of ways that particles can be arranged and the number of ways that the energy can be shared out between the particles.

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27
Q

What does a large, positive value of entropy show?

A

A high level of disorder.

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28
Q

What two things affect entropy?

A
  • Physical state
  • Number of particles

(see page 62 in the revision guide)

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29
Q

How do you calculate the change in entropy for a reaction?

A

(see page 62 in the revision guide)

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30
Q

Answer the two questions on page 62.

A

Get them right?

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31
Q

Answer the two questions on page 63.

A

Get them right?

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32
Q

When is a reaction feasible?

A

When /\G<0 or = 0

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33
Q

How do you calculate /\G?

A

/\G = /\H - T/\S

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34
Q

How do you calculate the temperature in which a reaction becomes feasible/

A

T =

/\H / /\S

35
Q

How do you calculate the temperature in which a reaction becomes feasible/

A

T =
/\H / /\S

(see page 65 in the revision guide)

36
Q

Give the ∆fH equation for NaCl.

A

Na(s) + 1/2Cl2(g) -> NaCl(s)

37
Q

Give the ∆atH equation for Na.

A

Na(s) -> Na(g)

38
Q

Give the ∆atH equation for O.

A

1/2O2(g) -> O(g)

39
Q

Give the ∆ie1H equation for Mg.

A

Mg(g) -> Mg+(g) + e-

40
Q

Give the ∆ea1H equation for O.

A

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

41
Q

Give the ∆dissH equation for NaCl.

A

NaCl(s) -> Na+(g) + Cl-(g)

42
Q

Give the ∆ie2H equation for Mg.

A

Mg(g) -> Mg2+(g) + e-

43
Q

Give the ∆ea2H equation for O.

A

O-(g) + e- -> O2-(g)

44
Q

Give the ∆solH equation for NaCl.

A

NaCl(s) + aq -> Na+ (aq) + Cl-(aq)

45
Q

Give the ∆lattH formation equation for NaCl.

A

Na+(g) + Cl-(g) -> NaCl(s)

46
Q

Give the ∆lattH dissociation equation for NaCl.

A

NaCl(s) -> Na+(g) + Cl-(g)

47
Q

What two factors does the strength of an enthalpy of lattice formation depend on? Explain both points.

A
  1. The size of ions.
    The larger the ions, the less negative the enthalpies of lattice formation (e.g. a weaker lattice). As the ions are larger the charges become further apart and so have a weaker attractive force between them.
  2. The charges on the ion.
    The bigger the charge of the ion, the greater the attraction between the ions so the stronger the lattice enthalpy (more negative values)
48
Q

What is defined as the perfect ionic model?

A

The ions are 100% ionic and spherical and the attractions are purely electrostatic.

49
Q

What 4 factors causes a tendency towards covalent characters in an ionic substance?

A
  1. The positive ion is small
  2. The positive ion has multiple charges
  3. The negative ion is large
  4. The negative ion has multiple negative charges
50
Q

Why is the theoretical value usually smaller than a Born-Haber experimental value?

A

When a compound has some covalent character - it tends towards a giant covalent so the lattice is stronger than if it was 100% ionic.

51
Q

Explain the difference between 100% ionic ions and ions with covalent characters.

A

When ions are 100% ionic they are spherical.

When ions have a covalent character The charge cloud is distorted.

  • The negative ions become distorted and more covalent, it becomes polarised. The metal cation is called polarising if it polarises the negative ion.
  • This happens because electrons are moving around constantly, it cannot be predicted where the electrons will be at a certain time.
52
Q

Explain the difference between theoretical and Born-Haber (experimental) lattice enthalpies.

A

The Born-Haber lattice enthalpy is the real experimental value. When a compound shows covalent character, the theoretical and the Born-Haber lattice enthalpies differ. The more the covalent character the bigger the different between the values.

53
Q

Give the definition for a spontaneous process.

A

A spontaneous process (e.g diffusion) will proceed on its own without any external influence.

54
Q

Describe entropy.

A

Entropy is a description of the number of ways atoms can share quanta of energy. If the number of ways of arranging the energy (W) is high, then the system is disordered and the entropy (S) is high.

Substances with more ways of arranging their atoms and energy (more disordered) have a higher entropy.

Solids have lower entropies than liquids, which are lower than gases. When a solid increases in temperature its entropy increases as the particles vibrate more. There is a bigger jump in entropy with boiling than that with melting.

Gases have large entropies as they are much more disordered.

55
Q

Draw a generalised graph for the entropy change between a solid, liquid and gas. Draw temperature on the x axis and the entropy on the y axis.

A

(see page 5 in the chemrevise revision guide)

56
Q

Will an increase in disorder and entropy result in a -ve or +ve entropy change?

A

An increase in disorder and entropy will result in a +ve entropy change.

57
Q

Just by looking at an equation of a reaction, what two things can we use to tell if there is a significant increase in entropy?

A
  • If there is a change of state from solid or liquid to gas.

- If there is a significant increase in number of molecules between products and reactants.

58
Q

Predict whether this reaction shows a -ve or +ve entropy:

NH4Cl (s) -> HCl (g) + NH3 (g)

A

∆Sº = +ve

  • Change from cold reactant to gaseous products.
  • Increase in number of molecules.
59
Q

What is the symbol for entropy?

A

Sº

60
Q

Predict whether this reaction shows a -ve or +ve entropy:

Na(s) + 1/2Cl2(g) -> NaCl(s)

A

∆Sº = -ve

  • Change from gaseous and solid reactant to solid.
  • Decrease in number of molecules.
61
Q

What is the problem with ∆H?

A

A reaction that is exothermic will result in products that are more thermodynamically stable than the reactants. This is a driving force behind many reactions and causes them to be spontaneous.

62
Q

What is the unit of entropy?

A

J K-1 mol-1

63
Q

Give the equation for how to calculate the entropy change.

A

∆Sº = (∑ Sº products) - (∑ Sº reactants)

64
Q

When do substances have 0 entropy?

A

Elements in their standard states do not have zero entropy.

Only perfect crystals at absolute zero (T = 0K) will have zero entropy.

65
Q

Answer the example question labelled 1 on page 6 in the chemrevise revision guide.

A

correct?

66
Q

Give the equation used to calculate Gibbs free energy change (∆G)

What are all the units?

A

∆G = ∆H - T∆S

∆G =  KJ mol-1
∆H = KJ mol-1
T = K
∆S = J K-1 mol-1
67
Q

When a reaction has increasing energy and is exothermic, will ∆G be +ve or -ve? Will the reaction be feasible?

A

A reaction that has increasing entropy (+ve ∆S) and is exothermic (-ve ∆H) will make the ∆G be negative and will always be feasible.

If the ∆G is negative there is still a possibility, however, that the reaction will not occur or will occur so slowly that effectively it doesn’t happen. If the reaction has a high activation energy the reaction will not occur.

68
Q

How do you convert J K -1 mol-1 into KJ mol-1?

A

J K-1 mol-1 / 1000

69
Q

Answer the example question labelled 2 on page 6 in the chemrevise revision guide.

A

correct?

70
Q

Answer the example question labelled 3 on page 6 in the chemrevise revision guide.

A

correct?

71
Q

Answer the example question labelled 4 on page 6 in the chemrevise revision guide.

A

correct?

72
Q

For any spontaneous change, will the ∆G be +ve or -ve?

A

For any spontaneous change, ∆G will be -ve.

73
Q

If the reaction involves an increase in entropy (∆S is +ve) then what will increasing the temperature do to the value of ∆G?

A

If the reaction involves an increase in entropy (∆S is +ve) then increasing temperature will make it more likely that ∆G is -ve and more likely that the reaction occurs.

e.g. NaCl + aq -> Na+(aq) + Cl-(aq)

74
Q

If the reaction has a ∆S close to zero then what will increasing the temperature do to the value of ∆G?

A

If the reaction has a ∆S close to zero then temperature will not have a large effect on the feasibility of the reaction as -T∆S will be small and ∆G won’t change much.

e.g. N2(g) + O2(g) -> 2NO(g)

75
Q

If the reaction involves a decrease in entropy (∆S is -ve) then what will increasing the temperature do to the value of ∆G?

A

If the reaction involves a decrease in entropy (∆S is -ve) then increasing the temperature will make it more less likely that ∆G is -ve and less likely for the reaction to occur.

e.g. HCl(g) + -> NH3(g) -> NH4Cl(s)

76
Q

Read the section on the graph on page 7 in the chemrevise revision guide.

A

understand?

77
Q

What is the equation for the relationship between ∆solH, ∆LdissH and ∆hydH?

A

∆solH = ∆LdissH + ∑ ∆hydH

see page 7 in the chemrevise revision guide

78
Q

What is the equation for the relationship between ∆solH, -∆LformH and ∆hydH?

A

∆solH = -∆LformH + ∑ ∆hydH

see page 7 in the chemrevise revision guide

79
Q

Answer the example question labelled 1 on page 8 in the chemrevise revision guide.

A

correct?

80
Q

Are hydration enthalpies exo or endothermic?

A

Hydration enthalpies are exothermic as energy is given out when water molecules bond to the metal ions.

The higher the charge density the greater the hydration enthalpy (e.g. smaller ions or ions with larger charges) as the ions attract the water molecules more strongly.
e.g. Fluoride ions have more negative hydration enthalpies than chloride ions.
Magnesium ions have a more negative hydration enthalpy than barium ions.

81
Q

Describe how ions dissolve in water.

A

The -ve ions are attracted to the partially +ve hydrogens on the polar water molecules and the +ve ions are attracted to the partially -ve oxygen on the polar water molecules.

(see page 8 in the chemrevise revision guide for a diagram)

82
Q

What does ∆solH tell us?

A

Generally ∆solH is not very eco or endothermic so the hydration enthalpy is about the same as a lattice enthalpy.
In general the substance is more likely to be soluble if the ∆solH is exothermic.
If a substance is insoluble it is often because the lattice enthalpy is much larger than the hydration enthalpy and it is not energetically favourable to break up the lattice, making ∆solH endothermic.

83
Q

When a solid dissolves into ions does the entropy increase or decrease?

A

When a solid dissolves into ions the entropy increases as there is more disorder as solid changes to solution and the number of particles increases.

84
Q

What happens to the ∆hydH of salts when it is endothermic and when it is exothermic?

A

(see bottom of page 8 in the chemrevise revision guide)

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