Thermodynamics

Question 1

Define the first ionization energy Δie1H:

Answer 1

It is the enthalpy change when 1 mole of gaseous 1+ ions formed from 1 mole of gaseous atoms.
Mg(g) —> Mg+ (g) + e-

Question 2

Define the second ionization energy Δie2H:

Answer 2

It is the enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions.
Mg+ (g) —> Mg2+ (g) + e-

Question 3

Define the first electron affinity Δea1H:

Answer 3

It is the enthalpy change when 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms.
O (g) + e- —> O- (g)

Question 4

Define the second electron affinity Δea2H:

Answer 4

It is the enthalpy change when 1 mole of gaseous 2- ions is formed from 1 mole of gaseous 1- ions.
O- (g) + e- —> O2- (g)

Question 5

Define bond dissociation enthalpy ΔdissH:

Answer 5

It is the enthalpy change when all the bonds of the same type in 1 mole of gaseous molecules are broken.
Cl2 (g) —> 2Cl (g)

Question 6

What are the two types of lattice enthalpy ΔlatticeH?

Answer 6

Lattice enthalpy can be defined as either enthalpy of lattice dissociation or enthalpy of lattice formation.

Question 7

Define lattice enthalpy of formation:

Answer 7

It is the enthalpy change when 1 mole of solid ionic compound is formed from its gaseous ions under standard conditions.
Na+ (g) + Cl- (g) —> NaCl (s) = -787 exothermic
Mg2+ (g) + 2Cl- —> MgCl2 (s) = -2526 exothermic

Question 8

Define lattice enthalpy of dissociation:

Answer 8

It is the enthalpy change when 1 mole of a solid ionic compound is completely dissociated into its gaseous ions under standard conditions.
NaCl (s) —> Na+ (g) + Cl (g)= +787 (endothermic)
MgCl2 (s) —> Mg2+ (g) + 2Cl (g)= +2526 (endothermic)

Question 9

Define enthalpy of atomization of an element ΔatH:

Answer 9

It is the enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard states.
1/2Cl2(g)—> Cl(g)

Question 10

Define enthalpy of atomization of a compound ΔatH:

Answer 10

It is the enthalpy change when 1 mold of a compound is in its standard state is converted to gaseous atoms.
NaCl(s)—> Na(g) + Cl (g)

Question 11

Define enthalpy change of formation ΔfH:

Answer 11

It is the enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions.
2C(s) + 3H2(g) + 1/2O2(g)—> C2H5OH(l)

Question 12

Define enthalpy change of hydration ΔhydH:

Answer 12

It is the enthalpy change when 1 mole of aqueous ions is formed from 1 mole of gaseous ions.
Na+(g)—-> Na+(aq)

Question 13

Define enthalpy change of solution ΔsolutionH:

Answer 13

It is the enthalpy change when 1 mole of solute is dissolved in enough solvent that no further enthalpy change occurs on further dilution.
NaCl(s)—-> NaCl(aq)

Question 14

Define the perfect ionic model:

Answer 14

It contains perfect spheres and electrostatic forces.

Question 15

Define the standard entropy of a substance ΔS⦵:

Answer 15

The standard entropy of a substance (S⦵) is the entropy of 1 mole of that substance under standard conditions (100kpa and 298K).

Question 16

What is entropy (S) and how does it increase?

Answer 16

It is a measure of the number of ways that particles can be arranged and the number of ways that the energy can be shared out between particles. The more disordered the particles are, the higher the entropy is. A large positive value of entropy shows high level of entropy. The more particles you have the more ways their energy can be arranged so entropy increases.

Question 17

How does physical states affect entropy?

Answer 17

Solid particles contain fixed positions meaning there is hardly any disorder so have the lowest entropy. Whereas gas particles are freely moving particles and have the most disordered arrangement of particles so have the highest entropy.

Question 18

What is the formula of standard entropy ΔS?

Answer 18

ΔS= S(products)-S(reactants)

Question 19

What is free energy change ΔG?

Answer 19

It is a measure used to predict whether a reaction is feasible.
If ΔG is negative or equal to zero then the reaction may happen itself.

Question 20

What is the formula of Gibbs free energy change ΔG?

Answer 20

ΔG = ΔH − TΔS

• ΔH= enthalpy change (ΔfH= products-reactants)
• ΔS= entropy change (S(products)-S(reactants)

Question 21

How does feasibility of reactions depend on temperature?

Answer 21

• If reaction is exothermic (negative ΔH and positive ΔS) then ΔG is always positive because ΔG = ΔH − TΔS is feasible at any temperatures.
• If reaction is endothermic (positive ΔH and negative ΔS) then ΔG is always positive and these reactions are not feasible at any temperatures.
• If the ΔH is positive (endo) and ΔS is positive then reaction won’t be feasible at some temperatures but will be at high enough temperatures.
• If ΔH is negative (exo) and ΔS is negative the reaction will be feasible at lower temperatures but not higher temperatures.

Question 22

When is a reaction feasible?

Answer 22

When ΔG = 0 or negative and then you can find the temperature by rearranging Gibbs free energy equation.

Question 23

How can you work out temperature using Gibbs free energy equation?

Answer 23

ΔT= ΔH (x1000)/ΔS

Question 24

Why are theoretical lattice different from experimental values?

Answer 24

You can work out a theoretical lattice enthalpy using calculations from purely (perfect) ionic model of lattice which assumes that ions are spherical and have evenly distributed charge. But when working lattice enthalpy experimentally the value is different and shows that most ionic compounds have covalent characters; positive and negative ions usually aren’t spherical and positive ions polarize negative ions so the more polarization, the more covalent bonding there is.