Short Questions

Question 1

Define Compression Factor.

Answer 1

Z = PVₘ / RT
where
Z is the compression factor
P is pressure
Vₘ is molar volume
R is universal gas constant
T is temperature

Question 2

Define Critical Temperature.

Answer 2

The temperature at the critical point, above which the substance cannot be liquified by pressure alone. The substance exists in equilibrium between two phases.

Question 3

Define enthalpy.

Answer 3

Enthalpy is the total energy of a system.

It is the sum of the internal energy and the product of the pressure and volume of a system.
H = U + PV

Question 4

Define expansion work for a gas.

Question 5

Define free maximum work function (Helmholtz free energy).

Question 6

Define Gibbs Free energy.

Question 7

Define heat capacity.

Answer 7

Heat capacity is the amount of energy needed to increase the temperature of a system by 1K.

Question 8

Define molar heat capacity.

Answer 8

Molar heat capacity is the amount of energy needed to change the temperature of 1 mole by 1K.

Question 9

Define partial pressure.

Answer 9

The partial pressure of a gas in a mixture is the pressure which it would exert if it occupied the whole volume of the mixture on its own, at the same temperature.

Question 10

Define pH.

Answer 10

pH = -log[H⁺]

Question 11

Define the standard enthalpy of formation of a substance.

Answer 11

Standard enthalpy of formation is the change in enthalpy when 1 mole of a compound is made up from its elements.

Question 12

Define thermodynamic system.

Question 13

Explain how “a” and “b” in the van der Waals equation of state (pV = [RT/(V-nb)] - an²/V²) account for the inter-particle interactions in real gases.

Answer 13

“a” is a constant that explains the attraction between the molecules in the gas.
“b” is the volume occupied by Nₐ number of molecules

Question 14

Explain how enthalpy corresponds to the heat transferred between the system and the surroundings at constant pressure.

Question 15

Explain how heat capacity at constant volume differs from heat capacity at constant pressure for an ideal gas.

Answer 15

At constant volume, a system is constrained to the same volume, and/so all the energy added to the system is used to increase the temperature.

At constant pressure, the system is allowed to expand to maintain constant pressure, so some of the energy added may be used for expansion work, meaning more energy will be needed to increase the temperature.

As a result, the heat capacity at constant volume is smaller than the heat capacity at constant pressure because all the energy added is used for increasing the temperature, rather than expansion work.

Question 16

Explain how the absorption of gases on solids depends on temperature.

Answer 16

As temperature increases, adsorption decreases.

Question 17

Explain how the First Law of Thermodynamics operates in closed systems.

Answer 17

Closed systems allow only heat and work to be exchanged between the system and its surroundings.

Question 18

Explain how the First Law of Thermodynamics operates in isolated systems.

Answer 18

Isolated systems do not allow any energy exchange between the system and its surroundings, and as a result the energy remains constant.

Question 19

Explain how the First Law of Thermodynamics operates in open systems.

Answer 19

Open systems allow matter, heat and work to be exchanged between the system and its surroundings.

Question 20

Explain the reasons behind deviations from ideality for a gas.

Answer 20

The volume of molecules, although small, is not completely negligible.

Molecules have both repulsive and attractive forces operating between them.

Question 21

Explain the terms in the equation
η = (T₂-T₁)/T₂.

Answer 21

η is the Efficiency of the system
T₂ is the
T₁ is the

Question 22

Explain three main types of thermodynamic systems.

Answer 22

Open systems allow matter, heat and work to be exchanged between the system and its surroundings.

Closed systems allow only heat and work to be exchanged between the system and its surroundings.

Isolated systems do not allow any energy exchange between the system and its surroundings, and as a result the energy remains constant.

Question 23

Explain why the standard enthalpies of formation of oxygen (O₂) and ozone (O₃) are different.

Question 24

Give 4 examples of state functions.

Question 25

Give an example of an irreversible process.

Question 26

Give an example of electrolysis.

Question 27

Give two examples of a thermodynamics function that is not a state function.

Answer 27

Work and heat

Question 28

How can compression factor be used to describe the deviation of ideality?

Answer 28

If Z > 1, is a positive deviation from ideality, and the gas is harder to compress than an ideal gas.

If Z < 1, is a negative deviation from ideality, and the gas is easier to compress than an ideal gas.

Question 29

Name the constant, R in the ideal gas equation of state PV=nRT.

Answer 29

Universal Gas constant

Question 30

State Dalton’s Law of Partial Pressures.

Question 31

State Faraday’s Laws.

Question 32

State Gay-Lussac’s Law.

Question 33

State Kohlrausch’s Law of Independent Migration of Ions.

Question 34

State the First Law of Thermodynamics.

Answer 34

The total energy of a system and its surroundings must remain constant, but energy can change from one form to another.

Question 35

State the Second Law of Thermodynamics for an isolated system.

Question 36

State the Second Law of Thermodynamics.

Question 37

State the Zeroth Law of Thermodynamics.

Answer 37

If system A is in thermal equilibrium with system B, and system B is in thermal equilibirum with system C, then system A is in thermal equilibrium with system C.

Question 38

What is the change of Gibbs Free Energy for a spontaneous process taking place at constant temperature and pressure?

Question 39

What is the condition for thermodynamic equilibrium under constant temperature and pressure?

Question 40

What is the difference between metallic and ionic conductors/ electrolytes?