3.6 enthalpy changes for solids and solutions

Question 1

in ionic bonding, the ions arrange themselves into a _____?

Answer 1

lattic
- so that the ions of opposite charge are next to one another

Question 2

enthalpy change of formation definition

Answer 2

• the energy transferred when 1 mole of the compound is formed from its elements under standard conditions (298K and 100KPa) with all the reactants and products being in their standard states

Question 3

is the enthalpy change of formation a positive or negative process?

Answer 3

negative

Question 4

enthalpy of atomisation definition

Answer 4

• the enthalpy change when 1 mole of gaseous atoms is formed from its own elements in its standard state

Question 5

is the enthalpy of atomisation a positive or negative process?

Answer 5

positive

Question 6

first ionisation enthalpy definition

Answer 6

the enthalpy change when one mole of gaseous atoms forms one mole of gaseous 1+ ions

Question 7

is the first ionisation enthalpy a positive or negative process?

Answer 7

positive

Question 8

first electron affinity definition

Answer 8

the enthalpy change that occurs when 1 mole of gaseous atoms forms 1 mole of gaseous 1- ions

Question 9

is the first electron affinity a positive or negative process?

Answer 9

negative

Question 10

second electron affinity definition

Answer 10

the enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ions

Question 11

is the second electron affinity a positive or negative process?

Answer 11

positive

Question 12

enthalpy of lattice formation definition

Answer 12

the standard enthalpy change when 1 mole of an ionic crystal lattice/solid ionic compound is formed from its constituent ions in gaseous form

Question 13

is the enthalpy of lattice formation a positive or negative process?

Answer 13

negative

Question 14

enthalpy of lattice dissociation definition

Answer 14

the standard enthalpy change when 1 mole of an ionic crystal lattice (ionic compound) is separated into its constituent ions in gaseous form

Question 15

is the enthalpy of lattice dissociation a positive or negative process?

Answer 15

positive

Question 16

categorise these:
1. Na (s) —> Na (g)
2. Na (g) —> Na+ (g) + e-
3. Na+ (g) + Cl- (g) —> NaCl (s)
4. O2 (g) —> 2O2 (g)

Answer 16

1. enthalpy of atomisation
2. first ionisation enthalpy
3. enthalpy of lattice formation
4. 2x enthalpy of atomisation

Question 17

what are born-haber cycles?

Answer 17

• thermochemical cycles that include all of the enthalpy changes involved in the formation of an ionic compound
• lattice enthalpy cant be calculated directly so a born-haber cycle can be used to calculate the lattice enthalpy by applying hess’s law

Question 18

drawing born-haber cycles:

Answer 18

• start with the elements in their standard states - these have zero enthalpy by definition
• the zero line will beer to be in the middle of the paper
• draw a line upwards from the zero line to the solid ionic compound made. this represents the enthalpy change of formation
• draw a lone upwards to usually atomise the elements individually
• positive (endothermic) changes are shown by an arrow pointing upwards
• negative (exothermic) changes are shown by an arrow pointing downwards

Question 19

• need to include state symbols in born-haber cycles *

Question 20

why is the second electron affinity positive while the first is negative?

Answer 20

because of the repulsion, so energy is required for the 2nd electron affinity a

Question 21

what are some factors affecting the strength of the lattice formation enthalpy?

Answer 21

• size of ions
• the charges on the ions

Question 22

how does the size of the ion affect the strength of the lattice formation enthalpy?

Answer 22

• the larger the ions, the less negative the enthalpies of lattice formation (weaker lattice)
• larger ions means the charges become further apart and have a weaker attractive force between them
• also larger ions means a smaller charge density

Question 23

how does the charges on the ion affect the strength of the lattice formation enthalpy?

Answer 23

• the bigger the charge of the ion, the greater the attraction between the ions so the stronger the lattice enthalpy (more negative values)
• this means more energy will be required to separate the ions

Question 24

what do born-haber cycles assume?

Answer 24

that an ionic compound is ‘perfectly ionic’:
- ions are treated as point charges, with the charge centred in the middle of the ion and 100% spherical
- attractions are assumed to be purely electrostatic

(the born-haber cycles use theoretical data)

Question 25

why might there be a difference between theoretical and experimental values of enthalpy?

Answer 25

- theoretical values assumes an ionic model

Question 26

how do you use the enthalpies of formation as a measure of a compound’s stability?

Answer 26

e.g calcium chloride - why does it have the formula CaCl2 and not CaCl or CaCl3 - need to calculate an enthalpy of formation for each case - the one with the most exothermic enthalpy of formation will be the one that forms as it will be the most thermodynamically stable

Question 27

do ionic solids dissolve in polar solvents?

Answer 27

yes

Question 28

what is the lattice enthalpy?

Answer 28

- in order to dissolve an ionic compound, the lattice must be broken up - this requires an input of energy - the lattice enthalpy

Question 29

what is hydration in terms of enthalpies of solution?

Answer 29

the separate ions (once the lattice is dissolved) are then solvated by solvent molecules, usually water and cluster around the ions - this is called hydration when the solvent is water

Question 30

enthalpy of hydration definition

Answer 30

the enthalpy change when one mole of gaseous ions become aqueous ions e.g Cu2+ (g) —> Cu2+ (aq)

Question 31

is the enthalpy of hydration a positive or negative process?

Answer 31

negative (it gives out energy (exothermic) because bonds are made between the ions and the water molecules)

Question 32

the higher the charge density, the (greater/lower) the hydration enthalpy

Answer 32

greater - as the ions attract the water molecules more strongly

Question 33

e.g compare the enthalpies of hydration of Na+ and Mg2+

Answer 33

- Mg2+ has stronger attraction - more exothermic - bc bond making is exothermic ??? i think

Question 34

e.g compare the enthalpies of hydration of Ca2+ and Mg2+

Answer 34

- Ca2+ less exothermic - Ca2+ has lower charge density (bc bigger) - weaker attraction between water and Ca2+

Question 35

enthalpy of solution definition

Answer 35

- the standard enthalpy change when 1 mole of an ionic solid dissolves in a large enough amount of water to ensure that the dissolved ions are well separated and do not interact with one another e.g NaCl (s) —> Na+ (aq) + Cl- (aq)

Question 36

you can link the enthalpy of solution, the enthalpy of hydration and lattice enthalpy in a Hess’s law cycle

Question 37

we can think of dissolving an ionic compound as the sum of 3 processes:

Answer 37

1. breaking the ionic lattice up to give gaseous ions 2. hydrating the positive ions 3. hydration the negative ions

Question 38

therefore, what can the equation be for dissolving an ionic compound?

Answer 38

∆H soln = -∆H latt (of formation) + ∆H hyd

Question 39

exam questions tend to focus on you working out the lattice formation/dissociation enthalpy in a Hess’s law cycle and then use that value to calculate solution/hydration enthalpies

Question 40

e.g calculate the enthalpy of solution of NaCl given that the lattice enthalpy of formation of NaCl is -771 kJmol^-1 and the enthalpies of hydration of sodium and chloride ions are -406 and -364 kJmol^-1 respectively

Answer 40

∆H soln = -∆H latt + ∆H hyd = - (-771) + (-406 + - 364) = 1 kJmol^-1

Question 41

what does ∆H solution tell us?

Answer 41

- in general, the substance is more likely to be soluble if the ∆H solution is exothermic - however its not always the case as e.g NaCl = 1 and can dissolve in water - (bc whether a compound is soluble or not also depends on other thermodynamic quantities such as entropy change - when a solid dissolves into ions the entropy increases - this positive entropy can make ∆G negative even if ∆H solution is endothermic, especially at higher temperatures )

Question 42

what type of enthalpy change is taking place in the following reaction? Na+ (g) —> Na+ (aq)

Answer 42

enthalpy change of hydration

Question 43

what type of enthalpy change is taking place in the following reaction? Mg2+ (g) + 2Cl- (g) —> MgCl2 (s)

Answer 43

lattice enthalpy of formation

Question 44

what type of enthalpy change is taking place in the following reaction? 1/2Cl2 (g) —> Cl (g)

Answer 44

enthalpy of atomisation

Question 45

what affects the solubility of ionic compounds in water?

Answer 45

- the balance between the hydration enthalpy of the ions in the compound and the lattice dissociation enthalpy of the compound

Question 46

what factors affect enthalpy of hydration?

Answer 46

- depends on the amount of attraction between the ions and the water molecules - the bigger the charge and the smaller the ion, the larger the enthalpy of hydration

Question 47

what is Hess’s law?

Answer 47

the enthalpy change of a reaction is independent of the route taken

Question 48

how can the enthalpy change of a solution be calculated using an energy cycle?

Answer 48

- lattice breaking enthalpy and enthalpy change of hydration can be used in an energy cycle to calculate the enthalpy change of solution using hess’s law

Question 49

give the equations for the first electron affinity of oxygen and the first ionisation energy of magnesium:

Answer 49

first electron affinity of oxygen: O (g) + e- —> O- (g) first ionisation energy of magnesium: Mg (g) —> Mg+ (g) + e-

Question 50

what is the perfect ionic model?

Answer 50

- the perfect ionic model assumes that all ions are perfectly spherical and that the bonds have no covalent character

Question 51

why are theoretical lattice enthalpies often different from experimental ones?

Answer 51

- theoretical lattice enthalpies use calculations which are based on a perfect ionic model of the lattice - the experimental values often differ bc most ionic compounds have some covalent character - the more polarisation in the ionic bond, the more covalent character of the compound

Question 52

how does exothermicity or endothermicity of ∆fH (enthalpy change of formation) act as a qualitative indicator of a compound’s stability?

Answer 52

- for an exothermic reaction the products are more stable than the reactants - for an endothermic reaction the reactants are more stable than the products

Question 53

- the lattice enthalpies of sodium fluoride, sodium chloride and magnesium chloride are shown below: - sodium fluoride = -918 kJmol^-1 - sodium chloride = -780 kJmol^-1 - magnesium fluoride = -2957 kJmol^-1 - explain the differences between these lattice enthalpies

Answer 53

- Na+ F- - Na+ Cl- - Mg2+ F- - Mg2+ charge —> more attraction between ions - Cl- larger than F- —> lower charge density - less attraction between ions in NaCl than NaF