Chapter 5 - Chemical Energetics: Thermochemistry & Thermodynamics

Question 1

What does energetically more stable mean?

Answer 1

If the energy level of the product is lower than that of the reactants, the products are energetically more stable than the reactants.

Question 2

Define activation energy.

Answer 2

Activation energy of a reaction is the minimum energy which the reacting particles must possess in order to overcome the energy barrier before becoming products.

Question 3

What are the standard conditions for thermochemical measurements? (3)

Answer 3

1) Temperature: 298K (25C)
2) Pressure: 1 bar
3) Concentration of any solution: 1 mol/dm3

Question 4

What does it mean by standard state?

Answer 4

A substance in its normal (most stable) physical state at 298K and 1 bar, where H=0 (zero enthalpy).

Question 5

Define standard enthalpy change of reaction. (ΔHr⊖)

Answer 5

The enthalpy change where molar quantities of reactants as specified by the chemical equation react to form products at 298K and 1 bar.

Question 6

Define standard enthalpy change of formation. (ΔHf⊖)

Answer 6

The enthalpy change when 1 mole of substance is formed from its constituent elements in their standard states at 298K and 1 bar.
(Measure of energetic stability of a substance relative to its constituent elements)

Question 7

Define standard enthalpy change of combustion. (ΔHc⊖)

Answer 7

The heat evolved when 1 mole of a substance is completely burnt in excess oxygen at 298K and 1 bar.
(Always exothermic)

Question 8

Define Hess’ law of constant heat summation.

Answer 8

The enthalpy change (ΔH) of a reaction is determined only by the initial and final states and is independent of the reaction pathway taken.

Question 9

Define standard enthalpy change of neutralisation. (ΔHneut⊖)

Answer 9

The heat evolved when 1 mole of water is formed in the neutralisation reaction between an acid and a base, at 298K and 1 bar.
(Always exothermic)

Question 10

What is the enthalpy change of neutralisation of a strong acid with a strong base?

Answer 10

-57.0kJ/mol

Question 11

What is the enthalpy change of neutralisation of a weak acid with a weak base? Why?

Answer 11

slightly less exothermic (less negative) than -57.0kJ/mol. Weak acids and bases do not ionise completely in dilute aqueous solution. During neutralisation, energy is absorbed to ionise the un-ionised weak acid/base. Thus, less energy is released and it is less endothermic.

Question 12

What is the equation for heat change of solution?

Answer 12

q=mcΔT or q=CΔT

Question 13

Define specific heat capacity (c).

Answer 13

The amount of heat required to raise the temperature of 1 g of the substance by 1 K. The unit is J g-1 k-1.

Question 14

Define heat capacity (C).

Answer 14

The amount of heat required to raise the temperature of the substance by 1 K (or 1C). The SI unit is J K-1.

Question 15

What are 2 assumptions made in thermochemistry calculations?

Answer 15

1) No heat loss to/gain from surrounding air

2) Heat capacity of the calorimeter is omitted

Question 16

What is the density of solutions assumed to be?

Answer 16

It is assumed to be same as water, at 1.00gm/cm3.

Question 17

Define bond dissociation energy.

Answer 17

It is the energy required to break 1 mole of a particular covalent bond in a specific molecule in the gaseous state.
(Always endothermic. When the same bond is formed, the same amount of energy is released, but BDE of the same type of bond in different molecules may differ)

Question 18

What is the relation between bond dissociation energy and strength of covalent bonds?

Answer 18

The more endothermic the bond dissociation energy, the stronger the covalent bond.

Question 19

Define bond energy.

Answer 19

It is the average energy required to break 1 mole of covalent bond in the gaseous state.
(Data in data booklet are average values derived from a large range of molecules containing that bond; can result in discrepancies)

Question 20

Define standard enthalpy change of atomisation for elements. (ΔHatom⊖)

Answer 20

It is the energy required when 1 mole of gaseous atoms is formed from the element at 298K and 1 bar.
(always endothermic)

Question 21

Define standard enthalpy change of atomisation for compounds. (ΔHatom⊖)

Answer 21

It is the energy required to convert 1 mole of the compound into gaseous atoms at 298K and 1 bar.
(always endothermic)

Question 22

How are gaseous diatomic elements’ enthalpy change of atomisation related to bond energy?

Answer 22

Bond energy is half of the standard enthalpy change of atomisation.

Question 23

What is first electron affinity?

Answer 23

It is the enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of singly charged gaseous anions.
(usually negative; always positive from 2nd onwards)

Question 24

Why is there a difference between experiment lattice energy and theoretical lattice energy?

Answer 24

Theoretical lattice energy is calculated with the assumption that the compound is completely ionic. However, there is no compound that is completely ionic.

Question 25

Why is there a discrepancy between experimental and theoretical lattice energy in a certain compound? (2 possible reasons)

Answer 25

1) When comparing cation: In MgCl2, Mg2+ has high charge and small ionic radius. With its high charge density, it has great polarising power, distorting the electron cloud of Cl- to a great extent, resulting in covalent character in MgCl2.
2) When comparing anion: In AgI, I- has a large ionic radius and is hence more polarisable. Thus Ag+ polarises the electron cloud of I- to a large extent, resulting in covalent character in AgI.

Question 26

Define standard enthalpy change of hydration. (ΔHhyd⊖)

Answer 26

It is the heat evolved when 1 mole of free gaseous ions is dissolved in an infinite volume of water at 298K and 1 bar.
(always exothermic as heat is evolved in forming ion-dipole interactions between ion&polar water molecules)

Question 27

What is the relationship between enthalpy change of hydration and charge density?

Answer 27

The higher the charge density, the stronger the ion-dipole interaction, the more exothermic enthalpy change of hydration.

Question 28

Define standard enthalpy change of solution. (ΔHsol⊖)

Answer 28

The enthalpy change when 1 mole of solute is completely dissolved in an infinite volume of solvent at 298K and 1 bar.
(Solvent usually an ionic compound, solvent usually water)

Question 29

How is enthalpy change of solution used to determine solubility?

Answer 29

If it is highly positive, it is likely insoluble.
if it is negative, it is likely soluble.
(there are exceptions due to entropy change)

Question 30

What is the relationship between lattice energy, enthalpy change of hydration and enthalpy change of solution?

Answer 30

ΔHsol⊖ = -L.E. + sum of ΔHhyd⊖

Question 31

Define a spontaneous process.

Answer 31

It is a change that occurs without a need for continuous input of energy from outside the system.

Question 32

Define entropy (S).

Answer 32

It is a measure of randomness or disorder in a system, reflected in the number of ways that the energy of a system can be distributed through the motion of its particles. Its unit is J mol-1 K-1.
If a reaction or process results in more ways to disperse or distribute the energy, entropy increases and ΔS > 0.

Question 33

How does the entropy for solid, liquid and gas differ?

Answer 33

For the same amount of substance,

solid < liquid &laquo_space;gas

Question 34

What is the equation for ΔS?

Answer 34

ΔS = ∑ΔS products - ∑ΔS reactants

Question 35

What are the 4 types of processes that involves change is entropy? (ΔS)

Answer 35

1) Change in temperature
2) Change in phase
3) Change in number of particles
4) Mixing of particles
5) Dissolution of ionic solid

Question 36

How does change in temperature affect ΔS?

Answer 36

As temperature increases, the average kinetic energy of the particles and the range of energies increase. There are more ways to disperse the energy among particles, thus increasing entropy. (ΔS > 0)

Question 37

How does change in phase affect ΔS? (2)

Answer 37

When temperature increases, entropy increases gradually as kinetic energy of particles increase.

1) When a solid melts, the particles move freely in the liquid state and become more disordered. Hence, there is an abrupt increase in entropy (ΔS > 0) as there are more ways to distribute the particles and their energy in the liquid state.
2) During vaporisation, the liquid converts into a gas where the particles are able to move even more freely. Hence there is a large increase in entropy (ΔS&raquo_space; 0) as there are more ways to distribute the particles and their energy in the gaseous state.

Question 38

How does change in the number of particles affect ΔS?

Answer 38

When a chemical reaction results in an increase in the number of gas particles, there is a large increase in entropy, since particles in gas are the most disordered. The number of ways that the particles and the energy can be distributed increase greatly.
If there is no change in the number of gas particles, entropy may increase or decrease but ΔS will be relatively small numerically.

Question 39

How does mixing of particles affect ΔS? (4 scenarios)

Answer 39

1) When 2 equimolar gases are mixed by removing a barrier between them at constant pressure, the gas expands to occupy the whole container. Volume of each gas doubles, and the pressure is halved (but total pressure remains constant. As the volume available for each gas is increased, there are more ways to distribute the particles and hence their energy. Thus, entropy increases (ΔS > 0).
2) When a gas expands, the volume available for distribution of particles increase. Entropy increases as there are more ways that the particles and the energy can be distributed.
3) When gases are mixed at constant volume (remove barrier than compress gas into original volume), the volume available to distribute each gas particle is the same. Hence, the entropy does not change. (ΔS=0)
4) When liquids with similar polarities are mixed together, entropy increases. This is because total volume increases and hence there are more ways to distribute the particles and therefore their energy.

Question 40

How is ΔS determined during the dissolution of an ionic solid?

Answer 40

Two entropy terms operate:
1) Entropy increases because ions in the solid are free to move in solution
2) entropy decreases because water molecules that were originally free to move become restricted in motion as they arrange themselves around the ions.
Overall ΔS depends on which factor is more significant.

Question 41

Define standard Gibbs free energy change of reaction (ΔG⊖)

Answer 41

It is the change in Gibbs free energy needed to convert reactants into products at 1 bar and at constant temperature (usually 298K).

Question 42

What is the equation for ΔG?

Answer 42

ΔG = ΔH - TΔS

it is temperature dependent

Question 43

How does ΔG affect spontaneity?

Answer 43

1) ΔG < 0: Reaction is feasible and spontaneous.
2) ΔG = 0: The system is at equilibrium and no net reaction. (eg melting and boiling at mp&bp)
3) ΔG > 0: The reaction is not feasible and cannot take place spontaneously.

Question 44

How is spontaneity affected when ΔH is negative and ΔS is positive?

Answer 44

Reactions are spontaneous at all temperatures since ΔG < 0 at all temperatures.

Question 45

How is spontaneity affected when ΔH is positive and ΔS is negative?

Answer 45

Reactions are not spontaneous as ΔG > 0 at all temperatures.

Question 46

How is spontaneity affected when both ΔH and ΔS is negative?

Answer 46

Reactions are spontaneous at low temperatures.

Question 47

How is spontaneity affected when both ΔH and ΔS are positive?

Answer 47

These endothermic reactions, which may not be spontaneous at room temperature, become spontaneous if temperature is sufficiently raised.

Question 48

What are 2 limitations of using ΔG⊖ to predict spontaneity of a reaction?

Answer 48

1) Under non-standard conditions, it must be calculated as ΔG⊖ can only be used to predict spontaneity of a reaction under standard conditions.
2) Kinetics consideration: ΔG⊖ does not take into account the kinetics of the reaction (rate of reaction). They may have a large activation energy to overcome. Some reactions are thermodynamically (energetically) favourable but not kinetically favourable.