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A Level Physical Chemistry > 1.08 Thermodynamics > Flashcards

1.08 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 its constituent elements in 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

Zero, by definition

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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 (excess) oxygen

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5
Q

Define standard enthalpy of atomisation

A

Enthalpy change when one mole of gaseous atoms is formed from, a compound in its standard state in standard conditions

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6
Q

Define first ionisation energy

A

Enthalpy change when one mole of electrons is removed from one mole of gaseous atoms to form one mole gaseous 1+ ions

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7
Q

Define second ionisation energy

A

Enthalpy change when one mole of electrons is removed 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

Enthalpy change when one mole of gaseous atoms gains one mole of electrons to form one mole gaseous 1- ions

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9
Q

Define second electron affinity

A

Enthalpy change when one mole 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

Enthalpy change when one mole of solid ionic lattice is formed from its constituent gaseous ions

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11
Q

Define lattice enthalpy of dissociation

A

Enthalpy change when one mole of solid ionic lattice is dissociated (broken into) into its gaseous ions

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12
Q

Define enthalpy of hydration

A

Enthalpy change when one mole of gaseous ions become hydrated/dissolved in water to infinite dilution (water molecules totally surround the ion)

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13
Q

Define enthalpy of solution

A

Enthalpy change when one mole of solute dissolves completely in a solvent to infinite dilution

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14
Q

Define mean bond dissociation enthalpy

A

Enthalpy change when one mole of (a certain type of) covalent bonds is broken, with all species in the gaseous state

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15
Q

Equation for standard enthalpy of formation

A

Mg(s) + 1/2O2(g) —> MgO (s)

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16
Q

Equation for standard enthalpy of combustion

A

CH4(g) +2O2(g) —> CO2(g) + 2H2O(g)

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17
Q

Equation for standard enthalpy of atomisation

A

1/2 I2 (g) —> I(g)

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18
Q

Equation for first ionisation energy

A

Li(g) —> Li+ (g) + e-

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19
Q

Equation for second ionisation energy

A

Mg+ (g) —> Mg2+ (g) + e-

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20
Q

Equation for first electron affinity

A

Cl(g) + e- —> Cl- (g)

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21
Q

Equation for second electron affinity

A

O-(g) + e- —> O2-(g)

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22
Q

Equation for lattice enthalpy of formation

A

Na+(g) + Cl-(g) —> NaCl(s)

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23
Q

Equation for
Lattice enthalpy of dissociation

A

NaCl(s) —> Na+ (g) + Cl-(g)

24
Q

Equation for
Enthalpy of hydration

A

Na+ (g) —> Na+ (aq)

25
Q

Equation for
Enthalpy of solution

A

NaCl (s) —> Na+ (aq) + Cl- (aq)

26
Q

Equation for
Mean bond dissociation enthalpy

A

Br2(g) —> 2Br (g)

27
Q

What is a Born-Haber cycle?

A

Thermochemical cycle showing all the enthalpy changes involved in the formation of an ionic compound. Start with elements in their standard states (enthalpy of 0)

28
Q

Draw a labelled Born Haber cycle for the formation of sodium chloride lattice, including ionisation, electron affinity, formation, lattice formation

A
29
Q

What factors affect the lattice enthalpy of an ionic compound?

A

Size of the ions and the charge on the ions

30
Q

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

A

Smaller ions, since the charge centres will be closer together. Increased charge, since the will be a greater electrostatic force of attraction between the oppositely charged ions.

31
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 these compounds would be thermodynamically stable. Was used to predict the existence of the first noble gas containing compound

32
Q

What actually happens when a solid is dissolved in terms of interactions of the ions with water molecules

A

lattice breaks and turns into gaseous ions; dissolve each gaseous ion water. The aqueous ions are surrounded by water molecules (which have a permanent dipole due to polar O-H bond)

33
Q

What is the perfect ionic model

A

Assumes that ions are perfectly spherical and that there is an even charge distribution (100% polar bonds). Act as point charges

34
Q

Why is the perfect ionic model often not accurate

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. Some lattices are not regular and the crystal structure can differ

35
Q

Which kind of bonds will be the most ionic? Why?

A

Between large positive ions and small negative ions e.g. CsF

36
Q

Define the terms spontaneous and feasible

A

If a reaction is spontaneous and feasible, it will take place of its own accord; does not take account of rate of reaction

37
Q

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

A

Negative- exothermic

38
Q

Define entropy

A

Randomness/disorder of a system.

Higher value of entropy = more disordered

39
Q

What units is entropy measured in?

A

JK-1mol-1

40
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

41
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

42
Q

Compare the general entropy values for solids, liquids and gases

A

Solids

43
Q

How would you calculate the entropy change for a reaction

A

Entropy change = sum of products’ entropy - sum of reactants’ entropy

44
Q

Define Gibbs free energy using an equation

A

ΔG = ΔH - TΔS

G=Gibbs free energy
H=enthalpy change
S=entropy change
T=temperature

45
Q

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

A

If G<0, reaction is feasible. If G=0, reaction is JUST feasible. If G>0, reaction is not feasible

46
Q

What is the significance of the temperature at which G=0

A

This is the temperature (in kelvin) at which the reaction becomes feasible

47
Q

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

A

Rearrange to T = (ΔH)/(ΔS) since G = 0

48
Q

What are the limitations of using G as an indicator of whether a reaction will occur?

A

Gibbs free energy only indicates if a reaction is feasible. It does not take into account the rate of reaction (the kinetics of the reaction). In reality, many reactions that are feasible at a certain temperature have a rate of reaction that is so slow that effectively no reaction is occurring

49
Q

If the reaction is exothermic and entropy increases, what is the value of G and what does this mean

A

G always negative so reaction is always feasible - product is favoured

50
Q

If the reaction is endothermic and entropy decreases, what is the value of G and what does this mean?

A

G always positive, so reaction is never feasible- reactant favoured

51
Q

If the reaction is exothermic and entropy decreases, what is the value of G and what does this mean?

A

Temperature dependent

52
Q

If the reaction is endothermic and entropy increases, what is the value of G and what does this mean?

A

Temperature dependent

53
Q

Why is entropy zero at 0K.

A

No disorder - molecules/atoms are not moving or vibrating and cannot be arranged in any other way. Maximum possible state of order

54
Q

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

A

Number of moles- more moles made —> increase in entropy

Going from solid -liquid/gas or liquid —> gas

55
Q

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

A

The heat that is given out escapes to the surroundings

A Level Physical Chemistry (12 decks)

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