Chapter 1 Chemical Energetics

Question 1

Enthalpy change (ΔH) refers to

Answer 1

the amount of heat energy transferred during a chemical reaction, at a constant pressure

Question 2

Enthalpy change of atomisation

Answer 2

The standard enthalpy change of atomisation (ΔH_at^ꝋ) is the enthalpy change when 1 mole of gaseous atoms is formed from its element under standard conditions

Question 3

Standard conditions in this syllabus are ?

Answer 3

• a temperature of 298 K and a pressure of 101 kPa

Question 4

The ΔH_at^ꝋ is always endothermic/exothermic?

Answer 4

endothermic as energy is always required to break any bonds between the atoms in the element, to break the element into its gaseous atoms

• Since this is always an endothermic process, the enthalpy change will always have a positive value

Question 5

The lattice energy (ΔH_latt^ꝋ) is the

Answer 5

enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions (under standard conditions)

Question 6

The ΔH_latt^ꝋ is always exothermic/endothermic

Answer 6

as when ions are combined to form an ionic solid lattice there is an extremely large release of energy

• Since this is always an exothermic process, the enthalpy change will always have a negative value
• Because of the huge release in energy when the gaseous ions combine, the value will be a very large negative value

Question 7

The large negative value of ΔH_latt^ꝋ suggests what?

Answer 7

that the ionic compound is much more stable than its gaseous ions

• This is due to the strong electrostatic forces of attraction
• Since there are no electrostatic forces of attraction between the ions in the gas phase, the gaseous ions are less stable than the ions in the ionic lattice
• The more exothermic the value is, the stronger the ionic bonds within the lattice are

Question 8

The ΔH_latt^ꝋ of an ionic compound

Answer 8

• cannot be determined directly by one single experiment
• Multiple experimental values and an energy cycle are used to find the ΔH_latt^ꝋ of ionic compounds

Question 9

The lattice energy (ΔH_latt^ꝋ) of an ionic compound can be written as an equation

Answer 9

• For example, magnesium chloride is an ionic compound formed from magnesium (Mg²⁺) and chloride (Cl^-) ions
• Since the lattice energy is the enthalpy change when 1 mole of magnesium chloride is formed from gaseous magnesium and chloride ions, the equation for this process is:

Mg²⁺(g) + 2Cl^-(g) → MgCl₂(s)

Question 10

The first electron affinity (EA₁) is the

Answer 10

• enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms, to form 1 mole of gaseous ions each with a single negative charge under standard conditions

X(g) + e^- → X^-(g)

Question 11

EA₁ is usually exothermic/endothermic

Answer 11

exothermic, as energy is released

• Since this is generally an exothermic process, then the value for EA₁ will usually be a negative number

Question 12

The second and third electron affinities are endothermic/exothermic

Answer 12

as energy is absorbed

• This is because the incoming electron is added to an already negative ion
• Energy is required to overcome the repulsive forces between the incoming electron and negative ion
• Since these are endothermic processes, the values will be positive

Question 13

Second & third electron affinity table

Answer 13

Question 14

Factors affecting electron affinity

Answer 14

• Nuclear charge
• Distance
• Shielding

Question 15

Factors affecting electron affinity

• nuclear charge

Answer 15

the greater the nuclear charge, the stronger the attractive forces between an incoming electron and the nucleus

Question 16

Factors affecting electron affinity

• distance

Answer 16

the greater the distance between the nucleus and the outermost shell/orbital where the electron is added, the weaker the force of attraction

Question 17

Factors affecting electron affinity

• Shielding:

Answer 17

the greater the number of shells, the greater the shielding effect and the weaker the force of attraction

Question 18

The value of the electron affinity depends on

Answer 18

how strongly the incoming electron is attracted to the nucleus

• The greater the attractive forces between the electron and nucleus, the more energy is released and therefore the more exothermic (more negative) the EA₁ value will be

Question 19

Electron affinities of non-metals become endothermic/exothermic?

Answer 19

• more exothermic across a period, with a maximum at Group 17
• There is generally a downwards trend in the size of the electron affinities of the elements in Group 16 and 17
  
  • The electron affinities generally become less exothermic for each successive element going down both Groups, apart from the first member of each Group (oxygen and fluorine respectively)

Question 20

Electron affinity table Group 16 & 17

Answer 20

Question 21

Fluorine is an exception and has a lower/higher EA₁ than chlorine

Answer 21

has a lower EA₁ than chlorine

• Fluorine has a very small atomic radius
• This means that the electron density of fluorine is high
• There is more repulsion between the incoming electron and the electrons that are already present in fluorine
• These repulsive forces reduce the attractive forces between the incoming electron and nucleus
• As a result, the EA₁ of fluorine is less exothermic than expected

Question 22

Going down Group 16 and 17 affinity

Answer 22

• The outermost electrons are held less tightly to the nucleus as they are further away
• The number of electron shells increases causing an increased shielding of the outermost electrons
• It gets more difficult to add an electron to the outer shell
• Less energy is released upon adding an electron to the outer shell
• So generally, the EA₁ becomes less exothermic

Question 23

To calculate the lattice energy (ΔH_latt^ꝋ) of an ionic compound

Answer 23

a Born-Haber cycle is used, which is a special type of energy cycle

Question 24

To find the ΔH_latt^ꝋ of an ionic compound, the following values must be known:

Answer 24

• Enthalpy change of formation (ΔH_f^ꝋ)
• The various enthalpy changes involved when going from elements in their standard states to their gaseous ions (the sum of all of these can be referred to as ΔH₁^ꝋ)

Question 25

The various enthalpy changes involved when going from elements in their standard states to their gaseous ions include:

Answer 25

• Enthalpy change of atomisation of each element
• First ionisation energy of the metal
• Successive ionisation energies of the metal if applicable
• First electron affinity of the non-metal
• Successive electron affinities of the non-metal if applicable

Question 26

The order that these are written in the Born-Haber cycle is important

Answer 26

• First ionisation energy cannot take place before atomisation, because first ionisation energy is the process of turning gaseous atoms into gaseous ions
• Electron affinity cannot take place before first ionisation energy, since the electrons must be lost from the metal first during ionisation, to be present in order for the non-metal to gain them

Question 27

Hess’s law states that

Answer 27

‘the enthalpy change in a chemical reaction is the same regardless of the route taken, as long as the final and initial conditions and states of reactants and products are the same for each route’

Question 28

• Once a Born-Haber cycle has been constructed, it is possible to calculate the lattice energy (ΔH_latt^ꝋ) by applying Hess’s law and rearranging:

Answer 28

ΔH_f^ꝋ = ΔH_at^ꝋ + ΔH_at^ꝋ + IE + EA + ΔH_latt^ꝋ

^{which becomes:} ΔH_f^ꝋ = ΔH₁^ꝋ + ΔH_latt^ꝋ

Question 29

sometimes a value may need to be doubled or halved, depending on the ionic solid involved

Answer 29

• For example, with MgCl₂ the value for the first electron affinity of chlorine would need to be doubled in the calculation, because there are two moles of chlorine atoms
• Therefore, you are adding 2 moles of electrons to 2 moles of chlorine atoms, to form 2 moles of Cl^- ions

Question 30

The two key factors which affect lattice energy, ΔH_latt^ꝋ

Answer 30

are the charge and radius of the ions that make up the crystalline lattice

Question 31

Ionic radius

Answer 31

• The lattice energy becomes less exothermic as the ionic radius of the ions increases
• This is because the charge on the ions is more spread out over the ion when the ions are larger
• The ions are also further apart from each other in the lattice
  
  • The attraction between ions is between the centres of the ions involved, so the bigger the ions the bigger the distance between the centre of the ions
  • therefore, the electrostatic forces of attraction between the oppositely charged ions in the lattice are weaker

Question 32

Ionic charge

Answer 32

• The lattice energy gets more exothermic as the ionic charge of the ions increases
• The greater the ionic charge, the higher the charge density
• This results in stronger electrostatic attraction between the oppositely charged ions in the lattice
• As a result, the lattice energy is more exothermic

Question 33

The lattice energies get less exothermic as the ionic radius of the ions increases

Answer 33

Question 34

Enthalpy change of solution

Answer 34

• The standard enthalpy change of solution (ΔH_sol^ꝋ) is the enthalpy change when 1 mole of an ionic substance dissolves in sufficient water to form a very dilute solution
• The symbol (aq) is used to show that the solid is dissolved in sufficient water
• ΔH_sol^ꝋ can be exothermic (negative) or endothermic (positive)

Question 35

Enthalpy change of hydration

Answer 35

The standard enthalpy change of hydration (ΔH_hyd^ꝋ) is the enthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form a very dilute solution

Question 36

When an ionic solid dissolves in water

Answer 36

• positive and negative ions are formed
• Water is a polar molecule with a δ- oxygen (O) atom and δ+ hydrogen (H) atoms which will form ion-dipole attractions with the ions present in the solution
• The oxygen atom in water will be attracted to the positive ions and the hydrogen atoms will be attracted to the negative ions
• Since the ΔH_hyd^ꝋ of KCl is -685 kJ mol^-1, 685 kJ mol^-1 is released in forming these ion-dipole attractions when KCl dissolves in water
  
  • This compensates for the remaining +685 kJ mol^-1 which was needed to break down the KCl lattice

Question 37

Energy cycle involving enthalpy change of solution, lattice energy, and enthalpy change of hydration

Answer 37

• The energy cycle shows that there are two routes to go from the ionic lattice to the hydrated ions in an aqueous solution:
  
  • Route 1: going from ionic solid → ions in aqueous solution (this is the direct route)

ΔH_sol^ꝋ= Enthalpy of solution

• Route 2: going from ionic lattice → gaseous ions → ions in aqueous solution (this is the indirect route)

-ΔH_latt^ꝋ + ΔH_hyd^ꝋ = reverse lattice enthalpy + hydration enthalpies

Question 38

Lattice enthalpy usually means Lattice

Answer 38

• formation* enthalpy, in other words bond forming. If we are breaking the lattice then this is reversing the enthalpy change so a negative sign is added in front of the term (alternatively it is called lattice dissociation enthalpy)
• According to Hess’s law, the enthalpy change for both routes is the same, such that:

ΔH_sol^ꝋ = -ΔH_latt^ꝋ + ΔH_hyd^ꝋ

ΔH_hyd^ꝋ = ΔH_sol^ꝋ + ΔH_latt^ꝋ

• Each ion will have its own enthalpy change of hydration, ΔH_hyd^ꝋ, which will need to be taken into account during calculations
  
  • The total ΔH_hyd^ꝋ is found by adding the ΔH_hyd^ꝋ values of both anions and cations together

Question 39

• The energy cycle involving the enthalpy change of solution (ΔH_sol^ꝋ ), lattice energy (ΔH_latt^ꝋ), and enthalpy change of hydration (ΔH_hyd^ꝋ) can be used to calculate the different enthalpy values

Answer 39

• According to Hess’s law, the enthalpy change of the direct and of the indirect route will be the same, such that:

ΔH_hyd^ꝋ = ΔH_latt^ꝋ + ΔH_sol^ꝋ

• This equation can be rearranged depending on which enthalpy value needs to be calculated
• For example, ΔH_latt^ꝋ can be calculated using:

ΔH_latt^ꝋ = ΔH_hyd^ꝋ - ΔH_sol^ꝋ

• Remember: the total ΔH_hyd^ꝋ is found by adding the ΔH_hyd^ꝋ values of both anions and cations together
• Remember: take into account the number of each ion when completing calculations

Question 40

The standard enthalpy change of hydration (ΔH_hyd^ꝋ) is affected by the amount that

Answer 40

• the ions are attracted to the water molecules
• The factors which affect this attraction are the ionic charge and radius

Question 41

ΔH_hyd^ꝋ becomes more exothermic with (ionic radius)

Answer 41

decreasing ionic radii

• Smaller ions have a greater charge density resulting in stronger ion-dipole attractions between the water molecules and the ions in the solution
• Therefore, more energy is released when they become hydrated and ΔH_hyd^ꝋ becomes more exothermic

Question 42

ΔH_hyd^ꝋ is more exothermic for ions with (Ionic Charge)

Answer 42

larger ionic charges

• Ions with large ionic charges have a greater charge density resulting in stronger ion-dipole attractions between the water molecules and the ions in the solution
• Therefore, more energy is released when they become hydrated and ΔH_hyd^ꝋ becomes more exothermic

Question 43

The enthalpy of hydration is more exothermic for smaller ions and for ions with a greater ionic charge

Answer 43

Question 44

The polar water molecules will form ion-dipole bonds with the ions in solution (a) causing the ions to become hydrated (b)

Answer 44

Question 45

energy cycle

Answer 45

Question 46

energy level simple

Answer 46

Question 47

energy level complicated

Answer 47