O - Water, Enthalpies and Dissolving

Question 1

State the intermolecular bonds between water molecules and why

Answer 1

Hydrogen bonds form because water is a polar molecule, so the electronegative O atoms draw electrons away from the H atoms making the H atom partially positively charged.
The partially positive charged hydrogen is attracted to the lone pairs of electrons on the Oxygen atom of an adjacent molecule

Question 2

State what 4 unusual properties of water are due to its ability to form hydrogen bonds

Answer 2

1) High boiling point compared to other Group 6 hydrides
2) High specific heat capacity
3) High enthalpy of vaporisation
4) Ice is less dense than water

Question 3

Define enthalpy of vaporisation

Answer 3

The amount of energy required to change a substance from its standard state to a vapour

Question 4

Explain why ice is less dense than water

Answer 4

As water freezes more hydrogen bonds form, forming a lattice structure with greater space between the molecules. Thus it is less dense

Question 5

Explain why water has a higher boiling point than other group 6 hydrides

Answer 5

Because most group 6 hydrides can’t form hydrogen bonds, since the group 6 elements aren’t as electronegative than Oxygen (even Sulphur) so they can only form ID-ID bonds which aren’t as strong as hydrogen bonds. Therefore less energy is needed to break the intermolecular bonds between other Group 6 hydrides

Question 6

Name the 2 main types of solvent and give an example

Answer 6

Polar (e.g. water)

Non-polar (e.g. hexane)

Question 7

Describe which solvents dissolve which solutes best

Answer 7

Solutes dissolve best in solvents with similar intermolecular forces

Question 8

Why don’t covalent molecules dissolve very well in polar molecules?

Answer 8

Because the hydrogen bonds between water molecules are stronger than the bonds that would form between water and the covalent molecules

Question 9

Why don’t ionic substances dissolve in non-polar solvents?

Answer 9

Because the non-polar molecules in the solvent don’t interact strongly enough with the ions to pull them away from the ionic lattice
This is because the electrostatic forces between the ions are much stronger than any bonds that could be formed between the ions and solvent molecules

Question 10

Name the process of ions being surrounded by:

i) water molecules
ii) polar solvent molecules other than water

Answer 10

i) Hydration

ii) Solvation

Question 11

Define standard lattice enthalpy, ΔHᶱᴸᵃᵗᵗ

Answer 11

The standard lattice enthalpy, ΔHᶱᴸᵃᵗᵗ, is the enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions

Question 12

State what Ꝋ means for enthalpy changes (ΔHᶱ)

Answer 12

Under standard conditions

Question 13

Why is ΔHᶱᴸᵃᵗᵗ always negative?

Answer 13

Because energy is released when an ionic lattice is formed

Question 14

State the significance of the magnitude of ΔHᶱᴸᵃᵗᵗ

Answer 14

The more negative the lattice enthalpy, the stronger the bonding

Question 15

Define enthalpy change of hydration, ΔHᴴᵞᵈ

Answer 15

The enthalpy change of hydration, ΔHᴴᵞᵈ, is the enthalpy change when 1 mole of aqueous ions is formed from gaseous ions
(e.g. Na⁺(g) → Na⁺(aq) )

Question 16

Is ΔHᴴᵞᵈ always positive, negative or both?

Answer 16

It is always negative (exothermic) because energy is given out as the bonds between the solvent and solute form

Question 17

Define enthalpy change of solution, ΔH(solution)

Answer 17

The enthalpy change of solution, ΔH(solution), is the enthalpy change when 1 mole of solute is dissolved in sufficient solvent that no further enthalpy change occurs on further dilution
(e.g. NaCl (s) → NaCl (aq) )
It is the net effect of the -ve lattice enthalpy and the enthalpy change of hydration/solvation)

Question 18

Define first ionisation enthalpy

Answer 18

The enthalpy change when 1 electron is removed from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous +1 ions
(e.g. O (g) → O⁺ (g) + e⁻)

Question 19

Define second ionisation enthalpy

Answer 19

The enthalpy change when 1 electron is removed from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions
(e.g. O⁺ (g) → O²⁺ (g) + e⁻)

Question 20

State and explain the 3 factors that affect the size of the ionisation enthalpies

Answer 20

1) Atomic radius - the further the outer shell electrons are from the positive nucleus, the less they’ll be attracted towards the nucleus. So, the ionisation enthalpy will be lower
2) Nuclear charge - the more protons there are in the nucleus, the more it’ll attract the outer shell electrons - it’ll be harder to remove the electrons, so the ionisation enthalpy will be higher
3) Electron shielding - the inner electron shells shield the outer shell electron from the attractive force of the nucleus. Because more inner shells mean more shielding, the ionisation enthalpy will be lower

Question 21

Describe and explain the trend in first ionisation enthalpies as you go down the group

Answer 21

First ionisation enthalpies decrease down a group because there’s less attraction between the nucleus and outer electrons, because:

• The outer electrons are in shells further from the nucleus
• There’s more shielding from inner shells

Question 22

Describe and explain the trend in first ionisation enthalpies as you go across the period

Answer 22

First ionisation enthalpies generally increase across a period because:

• The number of protons is increasing, causing stronger nuclear attraction
• Since all the outer-shell electrons are roughly the same energy level - there’s generally a little extra shielding effect or extra distance to lessen the attraction from the nucleus