22) Enthalpy and entropy

Question 1

Define lattice enthalpy

Answer 1

the enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions

Question 2

ΔLEH⦵ is an _ (exothermic / endothermic) change therefore, ΔH is _ (positive / negative)

Answer 2

exothermic

negative

Question 3

What is lattice enthalpy a measure of?

Answer 3

the strength of ionic bonding

Question 4

Lattice enthalpy cannot be measured directly therefore, what is required?

Answer 4

a Born-Haber cycle for indirect determination

Question 5

Define standard enthalpy change of atomisation ΔatH⦵

Answer 5

the enthalpy change that takes place when one mole of gaseous atoms forms from the element in its standard state

Question 6

ΔatH⦵ is always _ (exothermic / endothermic) ?

Answer 6

endothermic - bonds broken

Question 7

Define first ionisation energy ΔIEH⦵

Answer 7

the energy required to remove 1 e- from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions

Question 8

ΔIEH⦵ is _ (exothermic / endothermic) ?

Answer 8

endothermic - overcome nuclear attraction

Question 9

What is first ionisation energy a measure of?

Answer 9

ability to lose e-

Question 10

Define first electron affinity ΔEAH⦵

Answer 10

the enthalpy change that takes place when 1e- is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions

Question 11

What is first electron affinity a measure of?

Answer 11

ability to gain e-

Question 12

ΔEAH⦵ is _ (exothermic / endothermic) ?

Answer 12

exothermic - e- added is attracted towards nucleus

Question 13

What 2 things are important to remember about between each horizontal level of a Born-Haber cycle?

Answer 13

only one species has changed

all the species are balanced

Question 14

Give the alternative path for lattice enthalpy on a Born-Haber cycle

Answer 14

(gaseous ions) - first electron affinity - first ionisation energy (gaseous atoms) - enthalpy change of atomisation (elements in standard states) + enthalpy change of formation (ionic lattice)

Question 15

Define second electron affinity

Answer 15

the enthalpy change that takes place when 1e- is added to each ion in one mole of gaseous 1- ions to form one mole of gaseous 2- ions

Question 16

Second electron affinity is _ (exothermic / endothermic) ?

Answer 16

endothermic - negative ion repels e-

Question 17

When are successive electron affinities required?

Answer 17

when an anion has a greater charge than 1-

Question 18

What will you need to do if more than 1 atom of an element is involved throughout a Born-Haber cycle?

Answer 18

multiply anything involved by a factor of the no. of atoms

Question 19

Define enthalpy change of solution

Answer 19

the enthalpy change that takes place when one mole of a compound is completely dissolved in water under standard conditions

Question 20

How can enthalpy change of solution be determined experimentally?

Answer 20

using q = mcΔT
calculate moles dissolved
(if 1mol -> aqueous ions) 1 mole would gain energy q in kJ / mol
uses the mass of solution - not the mass of water alone!

Question 21

Give two features of the dissolving process, when a solid ionic compound dissolves in water

Answer 21

ionic lattice breaks up (opposite of ΔLEH⦵)

water molecules are attracted to and surround the ions

Question 22

Define enthalpy change of hydration

Answer 22

the enthalpy change that takes place when one mole of isolated gaseous ions is dissolved in water forming one mole of aqueous ions under standard conditions

Question 23

Give the alternative path for hydration on a Born-Haber cycle

Answer 23

(gaseous ions) + lattice enthalpy (ionic lattice) + enthalpy change of solution (aqueous ions)

Question 24

What can lattice enthalpy be a good indicator for?

Answer 24

melting point
(other factors e.g. packing of ions, needs to be considered also)

Question 25

How does increasing ionic radius affect lattice enthalpy and hydration, and therefore, melting point?

Answer 25

attraction between ions / ions and water molecules decreases lattice enthalpy and hydration energy is less negative melting point decreases

Question 26

How does increasing ionic charge affect lattice enthalpy and hydration, and therefore, melting point?

Answer 26

attraction between ions / ions and water molecules increases lattice enthalpy and hydration energy is more negative melting point increases

Question 27

What will happen if the sum of hydration enthalpies > magnitude of lattice enthalpy?

Answer 27

the overall energy change (enthalpy of solution) will be exothermic and the compound should dissolve

Question 28

How are many compounds with endothermic enthalpy changes of solution still soluble?

Answer 28

solubility also depends on temperature and entropy

Question 29

Define entropy S

Answer 29

term used for the dispersal of energy and disorder within the chemicals making up the chemical system

Question 30

As entropy increases, the dispersal of energy _ and the disorder _

Answer 30

increases | increases

Question 31

What happens if a system changes to be more random?

Answer 31

energy can be spread out more so ΔS is positive

Question 32

What happens if a system changes to be less random?

Answer 32

energy becomes more concentrated so ΔS is negative

Question 33

Give 2 changes that would result in a system becoming more random

Answer 33

solid -> liquid -> gas | increase in the number of gaseous molecules

Question 34

Define standard entropy S⦵

Answer 34

the entropy of one mole of a substance, under standard conditions (every substance has one, a positive value with the units JK-1 mol-1)

Question 35

entropy change of a reaction ΔS⦵ = ?

Answer 35

∑S⦵(products) - ∑S⦵(reactants)

Question 36

Define free energy change ΔG

Answer 36

the overall energy change during a chemical reaction; the balance between enthalpy, entropy and temperature

Question 37

What is the Gibbs' equation?

Answer 37

ΔG = ΔH - TΔS | free energy change = enthalpy change with surroundings - temp. in K x entropy change of system

Question 38

What should you remember about ΔS in the Gibbs' equation?

Answer 38

must be changed to kJ K-1 mol-1 by x 10^-3 to match units of ΔH

Question 39

What must there be for a reaction to be feasible?

Answer 39

a decrease in free energy | ΔG < 0

Question 40

ΔH has a _ than TΔS - however as temperature increases TΔS becomes _?

Answer 40

larger magnitude | more significant

Question 41

Feasibility is supported when ΔH is _ and ΔS is _ - but as T increases the significance of _?

Answer 41

negative positive ΔS

Question 42

Is a reaction feasible when ΔH is negative and ΔS is positive?

Answer 42

yes, ΔG is negative

Question 43

Is a reaction feasible when ΔH is negative, ΔS is negative and temperature is low?

Answer 43

yes, ΔG is negative

Question 44

Is a reaction feasible when ΔH is negative, ΔS is negative and temperature is high?

Answer 44

no, ΔG is positive

Question 45

Is a reaction feasible when ΔH is positive and ΔS is negative?

Answer 45

no, ΔG is positive

Question 46

Is a reaction feasible when ΔH is positive, ΔS is positive and temperature is low?

Answer 46

no, ΔG is positive

Question 47

Is a reaction feasible when ΔH is positive, ΔS is positive and temperature is high?

Answer 47

yes, ΔG is negative

Question 48

Give the two equations for minimum temperature for a reaction to be feasible

Answer 48

``` ΔG = ΔH - TΔS = 0 T = ΔH / ΔS ```

Question 49

What are the limitations of predictions made for feasibility? Why do many reactions have a negative ΔG and not seem to take place?

Answer 49

ΔG indicates thermodynamic feasibility but takes no account of kinetics or rate of reaction - so a reaction may not seem to take place if it has a very large Ea or very slow rate of reaction