Thermodynamics 1

Question 1

what holds giant ionic lattices together?

Answer 1

• solid ionic substances form regular structure called giant ionic lattices
• the positively and negatively charges ions are held together by electrostatic interactions

Question 2

define standard lattice enthalpy of formation. ΔLEH°

Answer 2

The enthalpy change when 1 mole of an ionic lattice is formed from its gaseous ions under standard conditions

+ Endo
Na+ (g) + Cl- (g) –> NaCl (s) ΔLE(f)H° = -787kJmol-1

Question 3

define standard lattice enthalpy of dissociation ΔfH°

Answer 3

the enthalpy change when 1 mole of ionic lattice completely dissociates into its gaseous ions under standard conditions.

-the enthalpy of formation and the enthalpy of dissociation are opposites

• Exo
  NaCl (s) –> Na+ (g) + Cl- (g) ΔLE(d)H°= +787kJ kJmol-1

Question 4

how do we know the strength of the ionic bonds?

Answer 4

the more positive the lattice enthalpy of dissociation, the stronger the ionic boning, because it is harder to break up the lattice

Question 5

define standard enthalpy of formation, ΔfH

Answer 5

the enthalpy change when 1 mole of a compound is formed from its constituent elements in their standard states, under conditions

Na(s) + 1/2 Cl (g) –> NaCl (s) ΔfH° = -411 kJmol-1

This process is usually exothermic because bonds are being formed.

Question 6

define standard enthalpy of atomisation, ΔaH°

Answer 6

the enthalpy change when one mole of gaseous atoms are formed from an element in its standard state

Na (s) –> Na (g) ΔH° = + 107 kJ mol-1
1/2 Cl2 (g) –> Cl (g) ΔaH° = +121 kJ mol-1

This process is always endothermic because bonds are being broken

Question 7

define first ionisation energy, ΔI1H°

Answer 7

the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1 + ions

Na (g) –> Na+ (g) + e- define ionisation energy, define ionisation energy, ΔI1H° = +496 kJ mol-1

This process is always endothermic as energy needs to be put in to pull an electron away from a positively charged nucleus

Question 8

define second ionisation energy, ΔI2H°

Answer 8

the enthalpy change when 1 mole of gaseous 2+ ions are formed from 1 mole of gaseous 1+ ions

Mg + (g) –> Mg 2+ (g) + e- define ionisation energy, ΔI2H° = + 1450 kJ mol-1

its much more endothermic than 1st define ionisation energy, ΔI1H° as the atoms is already positively charges so the electrons is more strongly attracted to the nucleus

Question 9

define electron affinity, define ionisation energy, ΔEA1H°

Answer 9

the enthalpy change when 1 mole of gaseous 1- ions are formed from 1 mole of gaseous atoms

Cl- (g) + e- –> Cl- (g) define ionisation energy, ΔEA1H° = -346 kJ mol-1

The process is always exothermic because the electron is attracted to the positively charges nucleus of an atom.

Question 10

define electron affinity , define ionisation energy, ΔEA2H°

Answer 10

the enthalpy change when 1 mole of gaseous 2- ions are formed from one mole of gaseous 1- ions

O- (g) + e- –> O2 (g) define ionisation energy, ΔEA2H°= +790 kJ mol-1

This process is always endothermic because energy must be put in to overcome the repulsion between an electron and a negatively charged ion

Question 11

what’s Hess’ law?

Answer 11

the total enthalpy change is independent of the route taken

Question 12

why is it that the experimental lattice enthalpy value can differ from the calculated value?

Answer 12

• a purely ionic model of a lattice assumes that all of the ions are spherical, and have evenly distributed charge around them
• however the values differ as this is evidence that some ionic compounds have partially covalent character
• positive ions normally polarise neighbouring negative ions
• more polarisation= more covalent character
• small, positively charged ions ( like H+) are more polarising than large positive ions, and large negatively charged ions (like I-) are more polarisable than small negative ions (like F-)

Question 13

what are the 2 steps that are taken to dissolve an ionic compound so the ionic lattice must be broken up?

Answer 13

1- the bonds between the ions in the lattice break- this is an endothermic process

2-bonds between the ions and the water are made- this is exothermic, and is known as hydration

Question 14

define standard enthalpy of solution, ΔsolH°

Answer 14

the enthalpy change when 1 mole of solute dissolves completely in sufficient solvent under standard conditions to form a solution in which the molecules or ions are far enough apart not to interact with each other

E.g, when NaCl dissolves completely in water
NaCl (s) -> Na+ (aq) + Cl- (aq)
The associated enthalpy change is the enthalpy of solution

Question 15

explain the hydration of ions

Answer 15

This s where free ions are incorporated into the solution

• the interactions between the solvent and the solute need to be of a similar strength to the interactions between the positively and negatively charged ions in the lattice for the ions to dissolve
• ions dissolve well in polar solvents, like water, because of the favorable electrostatic interactions between the oppositely charged ions in the solvent and the ions

Question 16

define standard enthalpy of hydration, ΔhydH°

Answer 16

the enthalpy chnage when 1 mole of aqueous ions is formed from gaseous ions under standard conditions

Cl- (g) –> Cl- (aq) ΔhydH°= -363 kJmol-1

This is an exothermic process

Question 17

how can we calculate the standard enthalpy of solution using lattice enthalpy and hydration data?

Answer 17

Standard enthalpy of solution= lattice enthalpy + standard enthalpy of hydration

Question 18

what is meant by entropy?

Answer 18

its a measure of disorder or chaos
-the more disorder and chaos= increasing entropy

entropy can also determined in terms of the number of configurations the particles can be in, or as the number of ways to distribute energy. More configuration/ ways= greater entropy

Question 19

how do we calculate entropy changes?

Answer 19

ΔS = ΣS° (products) - ΣS°(reactants)
ΔS= change in the total entropy
ΣS° (products)= total standard entropy of the products
ΣS°(reactants)=total standard entropy of the reactants

Question 20

how can we predict entropy changes?

Answer 20

solid fewer particles -> fewer configurations -> lower entropy (homogenous reactants + all in same state)

Question 21

what the Gibbs free energy equation?

Answer 21

ΔG = ΔH - TΔS
ΔG= change in free energy of the system (kJ mol-1)
ΔH= change in enthalpy of the system (kJ mol-1)
T= temperature of the system (K)
ΔS= change in entropy of the system (K-1 mol-1)

Question 22

whats the importance of ΔG?

Answer 22

-ΔG allows us to predict whether a reaction is feasible
-reactions are only feasible if ΔG is negative
ΔG depends on temperature- some reactions may be feasible at 1 temp and not another

Question 23

what does it mean if ΔG=0?

Answer 23

there’s an equilibrium and the reaction is just about favourable

Question 24

in which conditions will a reaction not happen regardless of whether the reaction is thermodynamically favourable?

Answer 24

• if the activation energy is too high

- if the rate of reaction is very slow