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F11SEB Thermodynamics > Thermodynamics > Flashcards

Thermodynamics Flashcards

1
Q

What is heat?

A

Heat is energy transferred as a result of temperature differences.

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2
Q

What is work?

A

Work is done as a result of motion against an opposing force.

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3
Q

What are heat and work both methods of?

A

Heat and work are both methods of transferring energy to the system.

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4
Q

What is an isolated system?

A

A system not connected thermally or mechanically to the surroundings.

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5
Q

What is internal energy?

A

Internal energy is the capacity to do work (U).

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6
Q

State the molecular origins of internal energy (3).

A
  1. Kinetic energy from moving molecules.
  2. Potential energy possessed by molecules.
  3. Attractive and repulsive forces exerted by molecules on one another, related to electrostatic force between two charged particles.
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7
Q

Molecules possess potential energy as a result of…(3).

A
  1. Attractions between nuclei and electrons.
  2. Chemical bonds.
  3. Intermolecular attractions.
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8
Q

What is internal energy (U)?

A

The sum of all kinetic and potential energy contributions.

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9
Q

When will internal energy change?

A

Internal energy will change if its volume is changed, as this changes the average separation between molecules.

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10
Q

What is the relationship between work and internal energy?

A
  • If work is done BY a system, its internal energy DECREASES.
  • If work is done ON a system, its internal energy INCREASES.
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11
Q

Explain the meaning of the +/- signs of q and w.

A

w: POSITIVE means that work is done ON the system.
NEGATIVE means that work is done BY the system.
q: POSITIVE means that heat is supplied TO the system.
NEGATIVE means that heat is released FROM the system.

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12
Q

Give 3 ways of stating the first law of thermodynamics.

A
  1. The internal energy of an isolated system is constant ∆U = 0.
  2. Energy can neither be created nor destroyed, only interconverted between forms.
  3. The total amount of energy in the universe is constant.
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13
Q

What is expansion work? Give an example.

A

Expansion work involves a change in the volume of the system i.e. a reaction with gases.

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14
Q

What is non-expansion work? Give an example.

A

Non-expansion work doesn’t involve a change in volume i.e. chemical reaction done in a battery.

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15
Q

What are the consequences of the first law of thermodynamics? (2).

A
  1. Energy is conserved, heat + work are equivalent forms of energy.
  2. Chemical reactions at constant volume have no expansion work, therefore w = 0 and heat is the ONLY contribution to ∆U. Therefore, ∆U = qv.
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16
Q

What is meant by the term ‘state functions’?

A

Pressure, volume and temperature depend only on the state of the system and are independent of how the change was brought about.

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17
Q

What is meant by the term ‘path functions’?

A

Work and heat depend on the path between states and depend on how the change was brought about.

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18
Q

What is an exothermic reaction?

A

They give out energy/heat, the reacting system loses energy.
∆H < 0, heat is released from the system at constant pressure.

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19
Q

What is an endothermic reaction?

A

They take in energy/heat, reacting system gains energy.

∆H > 0, heat is absorbed by the system at constant pressure.

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20
Q

What is meant by the term molar enthalpy of melting?

A

The enthalpy change which occurs when 1 mol of a SOLID melts to form a liquid.

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21
Q

What is meant by the term molar enthalpy of vaporisation?

A

The enthalpy change which occurs when 1 mol of a LIQUID boils to form a gas.

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22
Q

State Hess’s Law.

A

∆H = ∆H(products) - ∆H(reactants)

∆H = ∆U + p∆V

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23
Q

What is meant by the term standard enthalpy of formation?

A

The standard reaction enthalpy for the formation of a compound from its elements in their most stable form.

24
Q

How do we calculate the standard reaction enthalpy?

A

∆rH = Zn∆fH - Zn∆fH

Z: sum of
n: number of moles of each reactant/product

25
Q

What is bond enthalpy?

A

Bond enthalpy is equal to the standard enthalpy change for the dissociation of a molecule in the gas phase in atoms.
XY (g) –> X (g) + Y (g)

26
Q

What are the features of a spontaneous reaction? (3)

A
  1. Once started will continue without any intervention.
  2. Will move towards their equilibrium state without any external influence.
  3. Can be exothermic, endothermic or neither.
27
Q

What is meant by the term entropy?

A

A measure of the amount of disorder.

Entropy increases for a spontaneous process, making ∆S positive.

28
Q

State the equation for ∆S.

A

∆S = Sf - Sin

f: final
in: initial
∆S: JK-1

29
Q

State the second law of thermodynamics. (2)

A
  1. The entropy of an isolated system increases during any spontaneous process.
  2. Spontaneous processes are those that increase the total entropy of the universe.
30
Q

Explain what the values of ∆S tell us.

A

∆S(total) > 0, the reaction is spontaneous.
∆S(total) < 0, the reaction is non-spontaneous.
∆S(total) = 0, the reaction is at equilibrium.

31
Q

State the equation for ∆S when heat (qrev) is added reversibly.

A

∆S = qrev/T

32
Q

How do we achieve reversible heat flow?

A

The temperature of the system needs to be at thermal equilibrium with the surroundings.

33
Q

State the equation for ∆S for a phase transition.

A

∆S = ∆H/T

34
Q

Outline the features of a phase transition (3).

A
  1. Temperature of a substance remains constant while a phase transition occurs.
  2. Heat is transferred reversibly.
  3. If the transition takes place at constant pressure, the heat supplied = enthalpy change ∆H.
35
Q

State the Boltzmann formula and what is calculates.

A

S = kB lnW

kB: boltzmann constant (1.381 x 10-23 JK-1)
W: number of ways of arranging molecules and their energies.

The Boltzmann formula calculates absolute S values.

36
Q

State the third law of thermodynamics.

A

The entropy of a perfect crystal is S=0 at T=0K.

37
Q

Explain why we cannot have a perfect crystal.

A

Residual entropy arises.
No thermal motion.
Entropy arises from the number of possible orientations a molecule can take up.

38
Q

State the equation for Gibbs free energy and what it measures.

A

∆G = ∆H - T∆S

A measure of the total entropy in the system and the surrounding.

39
Q

What is the significance of a negative ∆G value?

A

Implies the reaction occurs spontaneously, favoured by a positive ∆S and a negative ∆H.

40
Q

Draw the phase diagram and annotate.

A

See notes.

  • Triple point: all three phase boundaries intersect and all three phases coexist. This point is different for every molecule.
  • Critical point: above this, the molecule becomes a supercritical fluid.
41
Q

State the Clausius-Clapeyron equation and explain.

A

ln (p2/p1) = -(∆vapH)/R (1/T2 - 1/T1)

This equation gives the form of the vapour-liquid boundary, where ∆G = 0.

42
Q

State the term dynamic equilibrium.

A

When the rate of the forward reaction(s) equals that of the reverse reaction, with no net change in composition (∆G = 0).

43
Q

How do thermodynamics and kinetics relate to equilibrium?

A

Thermodynamics: calculates the position of equilibrium and whether products or reactants will eventually dominate.
Kinetics: used to assess how quickly equilibrium is achieved.

44
Q

Outline the relationship between equilibrium and ∆G.

A

-∆rG: more reactants than equilibrium mixture.
+∆rG: more products than equilibrium mixture.
∆rG = 0: at equilibrium.

When ∆G = 0, ∆rG = -RTlnK

45
Q

What is meant by the symbol K?

A

K: equilibrium constant which is equal to the reaction quotient for the equilibrium composition.

46
Q

Outline the relationship between K and ∆G.

A

-∆rG, K>1: PRODUCTS are favoured

+∆rG, K<1: REACTANTS are favoured

47
Q

State le Chatelier’s principle.

A

When a change is made to a system in dynamic equilibrium, the equilibrium responds so as to minimise the effect of change.

48
Q

State how the Haber process is arranged to give the maximum yield possible (2).

A
  1. Ammonia gas product is continuously removed by dissolving it in water to cause more products to form.
  2. Equilibrium responds to a pressure increase by decreasing the total amount of gas, this favours the products in the Haber process.
49
Q

State the equation for the Haber process.

A

N2(g) + 3H2(g) –> 2NH3(g)

50
Q

State the relationship between equilibrium and pressure (2).

A
  1. The value of the equilibrium constant does NOT depend on the overall pressure.
  2. The composition of the equilibrium mixture DOES depend on the overall pressure.
51
Q

State the relationship between equilibrium and temperature (2).

A
  1. For ENDOthermic reactions, increasing temp (T) favours products as heat is absorbed.
  2. For EXOthermic reactions, decreasing temp (T) favours products as heat is released.
52
Q

State the temperature dependence of equilibrium, van’t Hoff, equation.

A

ln (K2/K1) = -(∆rH)/R (1/T2 - 1/T1)

53
Q

How does this relate to the graph of a straight line?

A

ln K vs. 1/T is a straight line if ∆rH is constant over the required temperature range.
The slope: -∆rH/R

54
Q

How would this graph (ln K vs. 1/T) look for an EXOthermic reaction?

A

Positive slope.
K decreases as T increases.
Favours reactants.

55
Q

How would this graph (ln K vs. 1/T) look for an ENDOthermic reaction?

A

Negative slope.
K increases as T increases.
Favours products.

56
Q

How does a catalyst alter a reaction?

A

Increases the rate at which the reaction mixture reaches equilibrium, doesn’t alter the composition.
Lowers the temperature required for a significant ROR.

F11SEB Thermodynamics (2 decks)

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