3.1.8 Thermodynamics

Question 1

Define enthalpy change of formation (Δ_fH)

Answer 1

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

Question 2

Write an equation representing the enthalpy change of formation of ethanol

Answer 2

Question 3

Define bond dissociation enthalpy (Δ_dissH)

Answer 3

Enthalpy change when one mole of a (covalent) bond is broken into (two) gaseous atoms (or free radicals)

Question 4

Write an equation representing the bond dissociation enthalpy of chlorine

Answer 4

Question 5

Define enthalpy change of atomisation of an element (Δ_atH)

Answer 5

Enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state

Question 6

Write an equation representing the enthalpy change of
atomisation of chlorine

Answer 6

Question 7

Define enthalpy change of atomisation of a compound (Δ_atH)

Answer 7

Enthalpy change when 1 mole of a compound in its standard state in converted to gaseous atoms

Question 8

Write an equation representing the enthalpy change of atomisation of sodium chloride

Answer 8

Question 9

Define first ionisation energy (Δ_ie1H)

Answer 9

Enthalpy change when 1 mole of gaseous 1+ ions is formed from 1 mole of gaseous atoms

Question 10

Write an equation representing the first ionisation energy of magnesium

Answer 10

Question 11

Define second ionisation energy (Δ_ie2H)

Answer 11

Enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions

Question 12

Write an equation representing the second ionisation energy of magnesium

Answer 12

Question 13

Define first electron affinity (Δ_ea1H)

Answer 13

Enthalpy change when 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms

Question 14

Write an equation representing the first electron affinity of oxygen

Answer 14

Question 15

Define second electron affinity (Δ_ea2H)

Answer 15

Enthalpy change when 1 mole of gaseous 2- ions is formed from 1 mole of gaseous 1- ions

Question 16

Write an equation representing the second electron affinity of oxygen

Answer 16

Question 17

Define enthalpy change of hydration (Δ_hydH)

Answer 17

Enthalpy change when 1 mole of aqueous ions is formed from 1 mole of gaseous ions

Question 18

Write an equation representing the enthalpy change of hydration of sodium

Answer 18

Question 19

Define enthalpy change of solution (Δ_solutionH)

Answer 19

Enthalpy change when 1 mole of solute is dissolved in enough solvent that no further enthalpy change occurs on further dilution

Question 20

Write an equation representing the enthalpy change of solution of sodium chloride

Answer 20

Question 21

Define lattice enthalpy of formation (Δ_latticeH)

Answer 21

Enthalpy change when 1 mole of solid ionic compound is formed from its gaseous ions under standard conditions

Question 22

Write an equation representing the lattice enthalpy of formation of magnesium chloride

Answer 22

Exothermic

Question 23

Define lattice enthalpy of dissociation (Δ_latticeH)

Answer 23

Enthalpy change when 1 mole of solid ionic compound is completely dissociated into its gaseous ions under standard conditions

Question 24

Write an equation representing the lattice enthalpy of dissociation of magnesium chloride

Answer 24

Endothermic

Question 25

What can you use to calculate lattice enthalpy since you can't calculate it directly?

Answer 25

Born-Haber cycle

Question 26

Calculate the lattice enthalpy of formation of sodium chloride given that the enthalpy of formation of sodium chloride is -411 kJ mol^-1

Answer 26

Question 27

Calculate the atomisation enthalpy of magnesium given that the enthalpy of formation of magnesium oxide is -548 kJ mol^-1 and lattice enthalpy of formation is -3791 kJ mol^-1

Answer 27

Question 28

Calculate the lattice enthalpy of formation of aluminium oxide given that the enthalpy of formation of aluminium oxide is -1669 kJ mol^-1

Answer 28

Question 29

What does the purely ionic model of a lattice assume?

Answer 29

* All ions are spherical & their charge is evenly distributed around them * Only ionic bonding

Question 30

Why are experimental lattice enthalpy values different to theoretical lattice enthalpy values?

Answer 30

* Most ionic compounds have some partial covalent bonding * +ve/-ve ions in lattice aren't usually spherical * +ve ions polarise neighbouring -ve ions to different extents

Question 31

Explain why the differences between experimental and lattice enthalpies are much bigger for magnesium halides than sodium halides

Answer 31

* Shows bonding in magnesium halides is stronger than ionic model predicts * ∴ bonds are strongly polarised & have quite a lot of covalent character * Bonding in sodium halides is similar to production of ionic model = compounds closer to being purely ionic

Question 32

When a solid ionic lattice dissolves in water, what 2 things occur?

Answer 32

1. Bonds between ions break to give free ions * Endothermic 2. Bonds between ions and water are made * Exothermic

Question 33

Explain why water molecules can bond to ions

Answer 33

* ∵ oxygen is more electronegative than hydrogen ∴ it draws electrons towards itself = creating a dipole * Dipole means +ve charged H atoms can form bonds with -ve ions & -ve charged O atoms can form bonds with +ve ions

Question 34

Substances generally only dissolve if the energy...

Answer 34

released is roughly the same, or greater than the energy taken in

Question 35

Souble substances tend to have ______ enthalpies of solution

Answer 35

exothermic

Question 36

Calculate the enthalpy of solution for sodium chloride given that the lattice dissociation enthalpy is 787 kJ mol^-1 and the enthalpies of hydration of sodium and chloride ions are -406 and -364 kJ mol^-1 respectively.

Answer 36

Question 37

Calculate the enthalpy of solution for silver chloride given that the lattice dissociation enthalpy is 905 kJ mol^-1 and the enthalpies of hydration of silver and chloride ions are -464 and -364 kJ mol^-1 respectively.

Answer 37

Question 38

Calculate the enthalpy of solution for magnesium chloride given that the lattice formation enthalpy is -2493 kJ mol^-1 and the enthalpies of hydration of magnesium and chloride ions are -1920 and -364 kJ mol^-1 respectively.

Answer 38

Question 39

What is entropy (*S*)?

Answer 39

Measure of no. of ways that particles can be arranged & no. of ways that energy can be shared out between particles

Question 40

More disordered the particles are, the ____ (more +ve the value) entropy is

Answer 40

higher

Question 41

Name and describe 2 factors that affect entropy

Answer 41

* Physical State * Solid = hardly any disorder = lowest entropy * Gas = most disordered arrangements = highest entropy * More Particles (means more entropy) * More particles you've got = more ways they + their energy can be arranged

Question 42

Why do substances always tend towards disorder?

Answer 42

They're more energetically stable when there's more disorder ∴ particles will move to increase their entropy

Question 43

Explain why the reaction of sodium hydrogencarbonate with hydrochloric acid is feasible even though it's an endothermic reaction

Answer 43

* Feasible due to the increase in entropy as the reaction produces CO₂ gas and water * Liquids and gases are more disordered that solids and so have a higher entropy * & more moles are being produced * This increase in entropy overcomes the change in enthaply * *T*Δ*S*\>Δ*H* * ΔG\<0

Question 44

During a reaction, there's an _____ \_\_\_\_ (Δ*S*) between reactants and products

Answer 44

Entropy change

Question 45

State the equation you would use to calculate Δ*S*

Answer 45

Question 46

State the units of entropy

Answer 46

J K^-1 mol^-1

Question 47

Define standard entropy of a substance (S^⦵)

Answer 47

Entropy of 1 mole of that substance under standard conditions

Question 48

Answer 48

Question 49

Answer 49

Question 50

What is free energy change (ΔG)?

Answer 50

Measure used to predict whether a reaction is feasible

Question 51

Free Energy Change When is a reaction considered feasible?

Answer 51

If ΔG = -ve or 0

Question 52

State the equation for calculating ΔG (free energy change). Include the units.

Answer 52

Question 53

Even if ΔG shows that a reaction is theoretically feasible, it might have a ... or be ... that you wouldn't notice it happening at all

Answer 53

Even if ΔG shows that a reaction is theoretically feasible, it might have a _very high activation energy_ or be _so slow_ that you wouldn't notice it happening at all

Question 54

Answer 54

Question 55

Feasibility depends on \_\_\_\_\_

Answer 55

Temperature

Question 56

When is ΔG always negative (i.e. reactions feasible at any temperature)? Name the 2 conditions

Answer 56

If reaction is exothermic (ΔH = -ve) & has +ve entropy change

Question 57

When is ΔG always postive (i.e. reactions are NOT feasible at any temperature)? Name the 2 conditions

Answer 57

If reaction is endothermic (ΔH = +ve) & has -ve entropy change

Question 58

What happens to the reaction's feasibility if ΔH = positive & ΔS = positive?

Answer 58

Reaction won't be feasible at some temps but will be at high enough temperature

Question 59

What happens to the reaction's feasibility if ΔH = negative & ΔS = negative?

Answer 59

Reaction will be feasible at lower temperatures but won't be feasible at higher temperatures

Question 60

Describe how you can calculate the temperature at which a reaction becomes feasible

Answer 60

Can find temperature when ΔG = 0

Question 61

State the equation for calculating the temperature at which a reaction becomes feasible. Include the units.

Answer 61

Question 62

Answer 62

Question 63

Suggest why the electron affinity of chlorine is an exothermic change (2)

Answer 63

* Net attraction between the chlorine nucleus and the extra electron * Energy is released (when the electron is gained)

Question 64

State whether you would expect the value of the theoretical enthalpy of lattice dissociation for silver chloride to be greater than, equal to or less than that for silver bromide. Explain your answer. (3)

Answer 64

* Greater * Chloride ions are smaller than bromide * More strongly attracted to silver ions

Question 65

Explain the meaning of the term perfect ionic model (1)

Answer 65

Ions can be regarded as point charges (or perfect spheres)

Question 66

Suggest why the experimental value is greater than the theoretical value for the enthalpy of lattice dissociation for silver chloride (2)

Answer 66

* AgCl has covalent character * Forces in the lattice are stronger than pure ionic attractions

Question 67

Explain the interaction between water molecules and fluoride ions when the fluoride ions become hydrated (2)

Answer 67

* Water is polar / H on water is δ⁺ * (F^– ions) attract water / on H δ⁺

Question 68

Suggest why electron affinity is an endothermic process for this reaction: Ca²⁺_(g) + e^- + S^-_(g) ⇒ Ca²⁺_(g) + S^2-_(g)

Answer 68

* Negative S^- ion * Repels electron being added

Question 69

Explain why the hydration enthaply of a fluoride ion is more negative than the hydration enthaply of a chloride ion (2)

Answer 69

* Fluor**_ide_** (ions) smaller (than chloride) / have larger charge density * So (negative charge) attracts (δ+ hydrogen on) water more strongly

Question 70

By describing the nature of the attractive forces involved, explain why the value of the enthaply of hydration for a chloride ions is more negative than that for the bromide ion (3)

Answer 70

* Chloride (ions) are smaller (than bromide ions) * So the force of attraction between chloride ions and water is stronger * Chloride ions attract the δ+ on H of water / electron deficient H on water

Question 71

Suggest why a value for the enthaply of solution of magnesium oxide is not found in any data books. (1)

Answer 71

Magnesium oxide reacts with water / forms Mg(OH)₂

Question 72

Suggest why using a Born-Haber cycle produces more accurate results than using mean bond enthalpies (1)

Answer 72

Mean bond enthalpies are from a range of compounds