MODULE 2

 

7.5.2      Transition elements

 

General characteristics such as incomplete d-shell, metallic nature, variable oxidation states, catalytic action, formation of coloured aqua and other complexes.

 

Transition Metals

The ten elements from Scandium to Zinc form the first transition metal series.  They closely resemble each other and are hard, dense, shiny metals with high melting and boiling points.  They readily form alloys and have other properties in common.  Crossing the period from Sc to Zn there is a small decrease in atomic radius and increase in electronegativity and ionisation energy.  Most of the properties of transition metals are related to their electronic structures.

 

Electronic Structure

Transition elements are characterised by having a partially filled d sub-shell

Sc        [Ar] 3d1 4s2

Ti         [Ar] 3d2 4s2

V          [Ar] 3d3 4s2

*Cr       [Ar] 3d5 4s1

Mn        [Ar] 3d5 4s2

Fe        [Ar] 3d6 4s2

Co        [Ar] 3d7 4s2

Ni         [Ar] 3d8 4s2

*Cu       [Ar] 3d10 4s1

Zn        [Ar] 3d10 4s2

 

*Note that in Cr the arrangement [Ar] 3d5 4s1 with half-filled 3d and 4s sub-shells is more stable than [Ar] 3d4 4s2.

In Cu [Ar] 3d10 4s1 with a completely filled 3d sub-shell and a half-filled 4s sub-shell is more stable than [Ar] 3d9 4s2.

 

Oxidation States

Transition metals form ions which are characterised by having a partially filled d sub-shell.  The common oxidation states are

Sc                                3

Ti                     2          3

V               2       3       4       5

Cr                     2          3                                  6

Mn                    2          3          4                      6          7

Fe                    2          3

Co                    2          3

Ni                     2          3

Cu        1           2

Zn                    2

 

Sc and Zn are not typical transition metals as they only have one oxidation state which does not have a partially filled d sub-shell. (Sc3+ [Ar],  Zn2+  [Ar] 3d10).

Common oxidation states are +2 and +3, with the +2 state more common towards the end.  The higher oxidation states are shown in compounds with electronegative elements like O, Cl or F (e.g. Cr2O72- [+6], MnO4- [+7]).

Variable oxidation state is found because of the small difference in energy between the 3d and 4s sub-shells.  This allows varying numbers of electrons to be used in bonding.  When forming ions transition metals lose electrons from the 4s sub-shell before the 3d.

e.g.       Fe2+      [Ar] 3d5            (NOT [Ar] 3d4 4s2)

            Ni3+      [Ar] 3d7

            Cr3+      [Ar] 3d3

            Mn2+     [Ar] 3d5

 

 

Catalytic Action

The ability of transition metals to exist in various oxidation states makes them important industrial and biological catalysts.

 

V2O5

Contact process for the manufacture of sulphuric acid.

2SO2  +  O2                         2SO3

Fe or Fe2O3

Haber Process for manufacturing ammonia

N2  +  3H2                        2NH3

Pt

Manufacture of nitric acid from ammonia.

4NH3  +  5O2              4NO  +  6H2O

then NO             NO2            HNO3

Ni

Hydrogenation of alkenes

CnH2n  +  H2                       CnH2n+2

 

Trace amounts of transition metals are needed for the catalytic activity of some enzymes e.g.  Cu for cytochrome oxidase (needed for metabolism).

 

 

Complex Ions

These are composed of a central metal atom or ion surrounded by a number of  oppositely charged anions or neutral molecules called ligands.  The ligands donate lone pairs of electrons into the vacant d-orbitals of the transition metal atom or ion forming dative covalent bonds

 

 

Coloured Solutions

Most transition metal compounds are coloured.  In an isolated atom or ion the 3d orbitals have the same energy (degenerate).  In a complex ion the d orbitals are split due to different overlapping with ligands.

 

 


                                                            2 orbitals of higher energy

 

 


                                                    DE   

 

3 orbitals of lower energy

                                   

Transitions between the two levels will absorb energy of frequency n where DE = hn.  For the transition metals this occurs in the visible part of the spectrum, making the ion coloured.

 

 


absorbance




                  l     violet   blue        green   yellow            red

 

The colour of the ion is complementary to the colour absorbed.  i.e. yellow/ green absorbed -ion appears blue/ violet.  The colour of the ion depends mainly on the transition metal but can be affected by different ligands.

 

 

 

Experimental study of oxidation states of vanadium , use of redox potentials to predict suitable reductants for the successive reduction of VO3- to VO2+, V3+, V2+ and suitable oxidants for the successive oxidation of V2+ to V3+, VO2+ and VO2+; the colours associated with these oxidation states of vanadium.

 

[See Practical work on the oxidation states of vanadium.]

 

The range of oxidation states of vanadium can be shown by shaking a solution containing vanadium (V) with zinc and dilute sulphuric acid.  The yellow solution changes gradually through green (a mixture of the original vanadium (V) solution and vanadium (IV) ) to blue, then to green and finally to violet.

 

Oxidation state

+5

+4

+3

+2

Colour

Yellow

Blue

Green

Violet

Ion

VO2+

VO2+

V3+

V2+

 

 

The use of transition metals as heterogeneous catalysts, such as nickel in the hydrogenation of C=C double bonds.  The action of transition element ions with variable oxidation states as catalysts.  Illustrated by the catalysis of the decomposition of OCl- ions to O2 by Co2+.

 

The ability of transition elements to vary their oxidation state makes them very useful as heterogeneous catalysts. The catalyst provides a surface on which the reacting molecules can come into intimate contact and also lowers the activation energy barrier to reaction. (See kinetics notes)

e.g.       C=C  +  H2                                 H-C-C-H

 

 

 H2                   C=C

C=C           H2

 H2                                                               H2                                                                                H-C-C-H

                        C=C

            H2

 C=C                 H2                    C=C      H2        C=C      H2                                         H-C-C-H

 


                                                                                                                                               

      absorption                          reaction                                                desorption

 

 

The decomposition of chlorate(I) ions OCl- is catalysed by addition of Co2+ ions

 

OCl-                              ½ O2  +  Cl-

 

In the catalysed reaction the cobalt(II) ions are initially oxidised to Co(III)    , the electrons released produce two OH- ions

 

2Co2+  + OCl-  +  H2O                  2Co3+ + Cl-  + 2OH-                   

 

the cobalt(III) is subsequently reduced to Co(II), producing water and oxygen gas

 

2Co3+ +  2OH-               2Co2+  +  ½ O2  +  H2O

 

 

 

Complexes understood as consisting of a central metal atom or ion surrounded by a number of ligands, defined as anions or molecules possessing lone pairs of electrons.  Coordination number .  Experimental study of formation of complexes of copper(II) and cobalt(II) with Cl- and NH3  from the hexaquacations, viewed as ligand replacement reactions.  The use of stability constants to determine the feasibility of ligand replacements.  Candidates will not be expected to quote names of complexes and only ligand replacement reactions involving monodentate ligands will be expected in this module.

 

Complex Ions

 

A complex consists of a central metal atom or cation surrounded by ligands.  Ligands are anions or neutral molecules possessing a lone pair of electrons.  The formation of a complex involves the donation of a pair of electrons from the ligand to the empty orbitals of the central metal atom or ion i.e. dative covalent bonding.

The metal behaves as a Lewis acid and the ligand as a Lewis base.  The number of ligands surrounding the metal is known as the coordination number.

e.g. [Co(H2O)6]2+

 


                             H2O                        2+

                 

  H2O                        H2O

 


                                  Co

 


     H2O                    H2O

 

 

 H2O

 

 

 

Complex

Name

Coordination No

[Co(H2O)6]2+

hexaaquacobalt (II)

6

[Ni(CO)4]

tetracarbonylnickel (0)

4

[CoCl4]2-

tetrachlorocobaltate (II)

4

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

tetraamminediaquacopper (II)

6

[Fe(CN)6]4-

hexacyanoferrate(II)

6

[Fe(CN)6]3-

 

 

[Al(H2O)6]3+

 

 

[Pt(NH3)4]2+

 

 

[PtCl4]2-

 

 

[CrCl(H2O)5]2+

 

 

                                   

Ligand Replacement Reactions

Transition metal ions are relatively small and have a high charge density.  In aqueous solution they attract water molecules strongly. If a different ligand is added to a solution of a transition metal complex in water, it will compete with the water molecules to complex the metal.  The reaction can be regarded as a stepwise replacement of water ligands.

 

[M(H2O)6]n+     +  L  ó  [M(H2O)5L]n+   +  H2O

[M(H2O)5L]n+   +  L  ó  [M(H2O)4L2]n+  +  H2O

[M(H2O)4L2]n+  +  L  ó  [M(H2O)3L3]n+  +  H2O etc.

 

This replacement is easily observed because of the change in colour of the complex ion.

[Co(H2O)6]2+ (aq)  +   4Cl- (aq)  ó  [CoCl4]2- (aq)  +  6H2O (l)

pink                                                      blue

 

 

Stability Constants

Ligand replacement reactions are equilibrium reactions.  Each step will have a value for the equilibrium constant.  We can write an  equilibrium constant for the overall reaction e.g.

 

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

 

Kc = [Cu(NH3)4(H2O)2]2+ [H2O]4

                                                                                                                                            

            [Cu(H2O)6]2+ [NH3]4

 

As the reaction is carried out in water, its concentration effectively remains constant, and the equilibrium expression can be modified as

 

Kst =   [Cu(NH3)4(H2O)2]2+                      =   1 x 1013 dm12 mol-4

                                                                                               

           [Cu(H2O)6]2+ [NH3]4

 

where Kst is the stability constant for the reaction.

 

The higher the value of Kst the more stable the complex and the more readily it will be formed when the ligand is added.

 

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

 

Kst  =    [CuCl4]2-                                  

                                                                                                            =  1 x 105

            [Cu(H2O)6]2+ [Cl-]4