Download IB Chemistry II Paper 2 Problem: 3/13 and 4 1. When concentrated

IB Chemistry II Paper 2 Problem: 3/13 and 4
1. When concentrated hydrochloric acid is added to a solution containing hydrated copper(II)
ions, the colour of the solution changes from light blue to green. The equation for the reaction
is:
[Cu(H2O)6]2+(aq) + 4Cl–(aq) → [CuCl4]2–(aq) + 6H2O(l)
(i)
Explain what the square brackets around the copper containing species represent.
(1)
(ii)
Explain why the [Cu(H2O)6]2+ ion is coloured and why the [CuCl4]2– ion has a
different colour.
(2)
2. Explain why copper is considered a transition metal while scandium is not.
(Total 3 marks)
3. The ten elements in the first-row d-block have characteristic properties and many uses.
(i)
State and explain the type of reaction that takes place between Fe3+ and H2O to form
[Fe(H2O)6]3+ in terms of acid-base theories.
(2)
(ii)
Explain why [Fe(H2O)6]3+ is coloured.
(3)
4. Explain why:
(i)
calcium has a higher melting point than potassium.
(2)
(ii)
sodium oxide has a higher melting point than sulfur trioxide.
(3)
5. Describe the bonding in iron and explain the electrical conductivity and malleability of the metal.
(Total 4 marks)
6. The graph below shows the boiling points of the hydrides of group 5. Discuss the variation in the
boiling points.
(Total 4 marks)
1. (i) complex (ion) / the charge is delocalized over all that is contained in
the brackets;
(ii)
1
colour is due to energy being absorbed when electrons are promoted
within the split d orbitals;
the colour observed is the complementary colour to the energy
absorbed / OWTTE;
Accept either answer for first mark.
changing the ligand / coordination number / geometry changes
the amount the d orbitals are split/energy difference between the
d orbitals / OWTTE;
2 max
2. Sc has no d electrons as an ion / Cu has d electrons;
Cu compounds are coloured / Sc compounds are colourless;
Cu has more than one oxidation state / Sc has only one oxidation state;
Cu compounds can act as catalysts / Sc cannot act as catalysts; 3 max
3. (i)
Lewis acid-base (reaction);
H2O: e– pair donor, Fe3+: e– pair acceptor / H2O donates an electron pair to Fe3+;
(ii)
d sub-levels are split into two sets of orbitals (of different energies);
electron transitions between (d) orbitals of different energies /
d-d transition(s);
transmitted (visible) light is complementary colour;
4. (i)
calcium ionic charge is twice/greater than the potassium ionic charge /
calcium has more delocalized electrons than potassium;
greater attraction of delocalized electrons and Ca2+ / less attraction
between the delocalized electrons and K+;
Do not accept calcium ion has a 2+ without comparison to K+.
(ii)
Na2O ionic/(stronger electrostatic) attractions between Na+ and O2–;
SO3 has (weak) intermolecular/van der Waals’/London/dispersion/
dipole-dipole attractions;
intermolecular/van der Waals’/London/dispersion/dipole-dipole forces
are weaker/more easily broken than (strong) ionic bonds / ionic bonds
are stronger/harder to break than intermolecular bond/van der
Waals’/London/dispersion/dipole-dipole forces;
5. metallic (bonding);
positive ions/cations and delocalized/sea of electrons;
electrostatic attraction between the two;
Award [2 max] for description of bonding
2
3
2
3
Conductivity:
electrons delocalized/free to move;
Malleability:
atoms/ions/cations can move without breaking bonds / atoms/ions/
cations can slide past each other;
4
6. boiling points increase going down the group (from PH3 to AsH3 to SbH3);
Mr/number of electrons/molecular size increases down the group;
Accept electron cloud increases down the group for the second marking point.
greater dispersion/London/van der Waals’ forces;
NH3/ammonia has a higher boiling point than expected due to the hydrogen
bonding between the molecules;
Do not accept hydrogen bonding alone.
4