3.4 Chemistry of the d-block transition metals

what is a transition metal

a d-block element with at least one ion with an incomplete d sub shell

characteristics of transition metal elements

shiny metals

high densities

high melting points

form giant metallic structures

why do transition metals exist in compounds in more than one oxidation state

they have the ability to gain/lose electrons as energies of the 4s and 3d orbitals are similar

why do transition metals form coloured compounds 

due to partially filled d orbitals 

substance absorbs some frequencies of light 

why do transition metals form effective catalysts

metals provide a good surface area - reactants are adsorbed onto the surface and bonds weaken, products are desorbed from the surface

partially filled d orbitals allows them to combine with other molecules

variable oxidation states allows the formation of intermediate compounds with lower activation energies

example of a transition metal catalysed reaction

Iron in the Haber process for manufacture of ammonia

complex ion

a metal ion surrounded by ligands which are bonded by coordinate bonds

ligand 

a molecule or ion which donates a pair of electrons to the central metal ion to form a coordinate bond 

coordination number

number of ligands which surround the central metal ion

monodentate ligands

only donate 1 pair of electrons

bidentate ligands

donate two or more pairs of electrons

shape and bond angle of a complex ion with a coordination number of 6

octahedral 

bond angle 90

shape of a complex ion with a coordination number of 4

tetrahedral

ligand substitution

one ligand in a complex ion is replaced by another ligand

Colour of solution containing [Cu(H2O)6] 2+

pale blue

[Cu(H2O)6] 2+ and ammonia colour

royal blue

[Cu(H2O)6] 2+ and ammonia equation

[Cu(H2O)6] 2+ +4NH3 → [Cu(NH3)4(H2O)2 2+ + 4H2O

Cu(H2O)6] 2+ and HCL colour

green/yellow solution

Cu(H2O)6] 2+ and HCL equation

Cu(H2O)6] 2+ + 4Cl- → [CuCl4]2- + 6H2O

[Co(H2O)6] 2+ and HCL colour change

pale pink to dark blue

[Co(H2O)6] 2+ and HCL equation

[Co(H2O)6] 2+ + 4Cl- → [CoCl4] 2- + 6H2O

d-orbital splitting

when ligands approach a transition metal ion they effect the d orbitals of the ion.

the negative charge density of the lone pairs repel the electrons in the orbitals, making the orbitals less stable and changing their energy.

what is required for electrons to move from lower to higher orbitals

energy obtained from light

why do transition metals show different colours

the gap between orbitals will determine the amount of energy required to promote an electron, corresponding to a specific frequency of light.

different ligands will lead to different splitting of d orbitals therefore giving different colours

why do complexes involving Cu+ or Sc3+ have no colour

they have a full d sub shell so electrons cannot move from one level to another

why will metal ion complexes react with an alkali

they’re often acidic due to high positive charge so will readily lose a H+ ion

colour of Fe2+

pale green

colour of Fe3+

yellow brown

colour of Cr3+

dark green

colour of Cu2+ and NaOH

blue precipitate

colour of Co2+ and NaOH

blue precipitate that turns beige

colour of Fe2+ and NaOH

green precipitate that turns brown

colour of Fe3+ and NaOH

red brown precipitate

colour of Cr3+ and NaOH

green grey precipitate that turns deep green