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Question 1

Explain what is meant by the term enthalpy change of hydration (2)

Answer 1

• (enthalpy change when)
  1 mole of gaseous ions react
  OR
  1 mole of hydrated/aqueous ions are formed
• gaseous ions dissolve in water
  OR
  gaseous ions form aqueous/hydrated ions

Question 2

Predict how the enthalpy changes of hydration of F– and Cl– would differ.

Explain your answer (2)

Answer 2

• ∆hydH (F–) more negative/exothermic (than ∆hydH (Cl–))
  AND
  F– has smaller size (than Cl–)
• Comparison of attraction between F- ions and water
  OR
  smaller sized ion linked to greater attraction to H₂O

Question 3

Explain what is meant by the term average bond enthalpy (2)

Answer 3

Breaking of one mole of bonds

In gaseous molecules

Question 4

Halogen Boiling point / °C
Chlorine –35
Bromine 59
Iodine 184

Explain why the halogens show this trend in boiling points (3)

Answer 4

Forces:
London forces increase
OR
induced dipole(–dipole) interactions increase

Reason:
(Number of) electrons increases

Link to energy and particles:
More energy to break intermolecular forces
OR
to break London forces
OR
to break induced dipole(–dipole) interactions

Question 5

Describe and explain how the student should determine the end point of this titration accurately (2)

Answer 5

Add starch (near the end point)

Blue to colourless

Question 6

Element Hydride Boiling point / °C
N NH3 –33
P PH3 –88
As AsH3 –55

Explain why the boiling point of PH3 is lower than the boiling point of NH3 (2)

Answer 6

• NH3 has hydrogen bonding
  OR
  PH3 does not have hydrogen bonding
• Hydrogen bonding is stronger
  OR
  More energy to overcome hydrogen bonding

Question 7

Element Hydride Boiling point / °C
N NH3 –33
P PH3 –88
As AsH3 –55

Explain why the boiling point of PH3 is lower than the boiling point of AsH3 (2)

Answer 7

• AsH3 / As has more electrons (than PH3 / P)
• in AsH3:
  stronger/more induced dipole–dipole interactions
  OR
  stronger/more London forces (than PH3)
  OR
  more energy required to overcome induced dipole–dipole interactions

Question 8

The reactivity of the Group 2 elements Mg–Ba increases down the group.
Explain why

Answer 8

Increasing size:
Atomic radius increases
OR
more shells
OR
more (electron) shielding

Attraction:
Nuclear attraction decreases
OR
(outer) electron(s) experience less attraction

Ionisation energy:
Ionisation energy decreases
OR
less energy needed to remove electron(s)

Question 9

Define disproportionation (1)

Answer 9

A reaction in which the same species is oxidised and reduced

Question 10

Define entropy (1)

Answer 10

Measure of dispersal of energy in a system

Question 11

[H+]

Answer 11

10 ^-pH

Question 12

enthalpy change of neutralisation

Answer 12

(enthalpy change for) the formation of 1 mole H₂O
from reaction of an acid/H+ with an alkali/base/OH–

Question 13

Explain which block in the Periodic Table sodium and magnesium belong to

Answer 13

s-block
AND
highest energy or outer electron is in a s orbital or s sub–shell

Question 14

The lattice enthalpy of sodium oxide is more exothermic than that of potassium oxide.

Explain why (2)

Answer 14

• ionic radius of sodium / Na+ is smaller
• Comparison of attraction of cation and anion
  Na+ has stronger attraction to O2-

Question 15

The first ionisation energy of sodium is more endothermic than that of potassium.

Explain why (2)

Answer 15

• Atomic radius is smaller
  OR
  fewer shells
• nuclear attraction increases
  OR
  (outer) electron(s) experience more attraction

Question 16

Red blood cells contain haemoglobin.

Explain using ligand substitutions:
* how haemoglobin transports oxygen around the body
* why carbon monoxide is toxic

Answer 16

• Coordinate bond mark
  O₂ (coordinately or datively) bonds with Fe2+ / Fe(II) / Fe / Iron
• Ligand substitution mark
  (When required) O₂ is replaced by H₂O OR CO₂
  OR O₂ is replaced by CO
  OR H₂O OR CO₂ is replaced by O₂
• Ligand strength mark
  CO forms strong(er) bonds (than O₂)

Question 17

Explain why the student used 0.00200 mol dm−3 potassium manganate(VII) solution for this titration, rather than the more usual concentration of 0.0200 mol dm−3 used in manganate(VII) titrations (1)

Answer 17

(0.00200 mol dm–3 solution gives) a large titre which leads to a small (percentage) error / uncertainty

Question 18

what does Kp > >1 mean

Answer 18

equilibrium is far to the right

Question 19

what makes an indicator suitable

Answer 19

pH range / colour change matches vertical section

Question 20

Enthalpy change of formation

Answer 20

Enthalpy change when 1 mole of compound is formed from its elements

Question 21

Ligand

Answer 21

• Donates a lone pair of e- to a central metal ion
• forms dative covalent / coordinate bond

Question 22

Transition metal

Answer 22

• forms ions with incomplete d sub-shells
• coloured ions
• multiple oxidation states
• can act as catalysts

Question 23

Enthalpy Change of Solution

Answer 23

the enthalpy change that takes place when one mole of solute is dissolved

Question 24

Enthalpy Change of Atomisation

Answer 24

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

Question 25

Describe and explain the factors that affect the values of lattice enthalpies (3)

Answer 25

• Decrease in (ionic) size
• Increase in (ionic) charge OR charge density
• Greater attraction between ions gives more negative LE

Question 26

Standard electrode potential

Answer 26

The e.m.f. (of a half-cell) compared with a (standard) hydrogen half-cell/(standard) hydrogen electrode