Transition metal chemistry
1. What defines a transition metal?
2. Write electron configurations for the following atoms and ions.
3. Which of the ions above are not considered transition metal ions and why not?
4. Why do so many different oxidation states exist within the transition metals?
5. Suggest a structure for the following complex ions and draw them. Work out the overall charge on the complex ion.
6. a) What type bonding exists between the ligand and the metal? How does it arise?
b) Which of the ligands in the above complexes are monodendate and which are bidentate?
c) Why are transition metals such good catalysts?
d) Explain why many transition metal complexes are intensely coloured. Use the bright blue complex [Cu(II)(NH3)4(H2O)2]2+ as an example showing the arrangement of d electrons in the metal ion.
e) [Cu(II)(NH3)4(H2O)2]2+ ↔ [Cu(II)(H2O)6] 2+ ↔ [Cu(II)Cl4]2-
What do you need to add to push the equilibrium to the left or right in the above reactions?
7. i) Write balanced equations and describe what happens when excess sodium hydroxide is added to solutions of
a) Cr3+ b) Fe2+ c) Fe3+ b) Cu2+
ii) Name an important industrial (catalytic) application for transition metals.
iii) Name an important biological role for transition metals.
1. A transition metal is one which forms an ion with a partially filled d sub-shell.
2. a) Sc 1s22s22p63s23p63d14s2
b) V 1s22s22p63s23p63d34s2
c) Cr 1s22s22p63s23p63d54s1
d) Fe 1s22s22p63s23p63d64s2
e) Cu 1s22s22p63s23p63d104s1
f) Sc3+ 1s22s22p63s23p6
g) Mn2+ 1s22s22p63s23p63d5
h) Ni2+ 1s22s22p63s23p63d8
i) Cu2+ 1s22s22p63s23p63d9
j) Cu+ 1s22s22p63s23p63d10
k) Zn2+ 1s22s22p63s23p63d10
3. Sc3+, Cu+, Zn2+
4. There is only a small difference in the ionisation energy for the removal of sequential d sub-shell electrons; therefore it is possible to remove different amounts of electrons quite easily. In the s-block the barrier of a new shell is reached quickly limiting oxidation states to +1 and +2.
5. a) [Cu(II)(H2O)6] – 2+ , octahedral
b) [Cu(II)(NH3)4(H2O)2] – 2+ , octahedral
c) [Cu(II)Cl4] – 2- , tetrahedral
d) [Fe(III)(C2O4)3] - 3- , octahedral
e) [Ni(II)(NH2CH2CH2NH2)3] - 2+ , octahedral
6. a) A dative or coordinate bond. The ligand donates a lone pair into the empty d-orbitals on the metal ion.
b) H2O, NH3, Cl- monodentate
C2O42-, NH2CH2CH2NH2 bidentate
c) They accept lone pairs from the reactant molecules to form intermediate complexes. These modify electron arrangements in the reactants, making it easier to break bonds (lowers Ea). The variable oxidation states of transition metals also facilitate redox reactions.
d) The partially filled d-orbitals of similar energy levels are split into two different energy levels. 2 orbitals have a higher energy level than the other 3 (if octahedral). The lower energy orbitals will be filled sequentially and at least one of the higher energy pair will be empty or half full. Electrons can be promoted to this higher level from the lower one absorbing a small amount of energy (hf) to do so. This energy happens to correspond to light in the visible region and so some visible light is absorbed and the rest reflected making the substance appear coloured.
Cu2+ 3d9 →↑↓ / ↑
↑↓ / ↑↓ / ↑↓
add water → add Cl- →
Cu(II)(NH3)4(H2O)2]2+ ↔ [Cu(II)(H2O)6] 2+ ↔ [Cu(II)Cl4]2-
← add NH3 ← add water
7. (i) a) Cr3+(aq) + 3OH-(aq) → Cr(OH)3(s)
Cr(OH)3(s) + 3OH-(aq) → Cr(OH)-4(aq)
b) Fe2+(aq) + 2OH-(aq) → Fe(OH)2(s)
c) Fe3+(aq) + 3OH-(aq) → Fe(OH)3(s)
b) Cu2++(aq) + 2OH-(aq) → Cr(OH)2(s)
(ii) Iron – Haber process
Vanadium – Contact process
Nickel – Hydrogenation
Titanium – Ziegler Natta
(iii) Iron – haemoglobin
Cobalt – vitamin B12
Molybenum – nitrogenase
Copper - cytochromes