Transition metals

• The transition elements are the d-block elements, which have partially filled or empty d orbitals.
• Iron(II) sulfate is an example of a transition metal compound that exhibits colouration.
• Transition metal compounds can be coloured due to electronic transitions between different energy levels within their atoms.
• Copper(II) chloride is another example of a transition metal compound with a characteristic greenish blue coloration.
• Transition metals can form coloured complexes due to their ability to accept electrons into vacant d orbitals.
• Colourless compounds may become coloured when they react with other substances.
• Copper(II) sulfate turns blue when it reacts with ammonia gas.
• Cobalt(III) nitrate hexahydrate has a deep blue coloration.
• Electronic transitions occur when electrons move from one orbital to another, resulting in changes in the electron configuration and absorption/emission of light at specific wavelengths.
• A transition metal is an element that possesses a partially filled d sub-shell as an atom or in its stable ions.

  Although zinc is part of the d-block, it is not a transition element; it does not have a partially filled d sub-shell as an atom or in its stable ion, Zn2+. Copper is a transition metal because it has partially filled d orbitals in most of its compounds (copper(I) compounds have a full d orbital).
• In a transition metal compound, there will be a central metal ion surrounded by ligands which donate pairs of electrons to the metal ion.
• Ligand - A molecule or ion that binds to a central metal ion through coordinate covalent bonds.
• The number of unpaired electrons on the metal ion determines whether the complex ion is paramagnetic or diamagnetic.
• Although transition metals all have similar physical properties like atomic radius, high densities and high melting and boiling points. They also have special chemical properties like :

  
  

  • variable oxidation states
  • complex ion formation
  • formation of coloured ions
  • catalytic activity.
• There are two exceptions to this — chromium and copper. Both only have 1 electron in the 4s orbital, with chromium having a half-filled d subshell and copper a complete d subshell.

  24 Cr   [Ar]3d54s1

  29 Cu   [Ar]3d104s1

  When transition metal atoms form ions, the 4s electrons are lost first. When you write or draw the electronic configuration of any transition metal ion, it always has an empty 4s orbital.
• Transition metals can have different oxidation states. This is because the energies of the 4s and 3d orbitals are very similar, so the energy required to remove any of these electrons (ionisation energy) is similar. The actual oxidation state depends on many different factors.

  The lower oxidation states are found as simple ions, e.g. Co2+ or Fe3+. The higher oxidation states only exist when the metals are covalently bonded to very electronegative elements, such as oxygen, e.g. CrO42- or MnO4-.
• common oxidation states of the following metals:

  • chromium +3 and +6
  • manganese +2, +4 and +7
  • iron +2 and +3
  • cobalt +2 and +3
  • copper +1 and +2
• A transition metal complex consists of a central metal ion surrounded by a number of molecules or ions called ligands. The ligands are attached to the central ion by coordinate bonds. The ligand supplies the two electrons, which is what makes this a coordinate bond.
• Transition metal ions have many orbitals available for bonding, many of which are empty. Ligands form coordinate bonds when an orbital, from the ligand, containing a lone pair of electrons overlaps with empty orbitals on the transition metal ion.
• Ligands are classified by the number of coordinate bonds that they can form in complexes.
• Ligands that have one atom which can bond to the metal ion are called monodentate. Examples are water, H2Ö:, ammonia, :NH3, chloride ion, Cl-, and cyanide ion, CN-.
• Ligands that have two atoms which can bond to the metal ion are called bidentate. An example is 1,2-diaminoethane, H2NCH2CH2NH2 (often shortened to ‘en’).
• Typically, complexes are either octahedral with six ligands arranged around the central metal ion or tetrahedral with four ligands.

  Both shapes can be seen in the same transition metal ion, but with different ligands. The shape of complexes is dependent on the metal and its oxidation state, as well as the ligands themselves.

  Octahedral complexes are more common than tetrahedral complexes.

  Examples of octahedral complexes with their colours are:

  [Fe(H2O)6]2+ pale green

  [Fe(H2O)6]3+ yellow

  [Cr(H2O)6]3+ dark green

  [Co(H2O)6]2+ pink

  [Cu(H2O)6]2+ blue

  [Cu(NH3)4(H2O)2]2+ royal blue
• The coordination number is equal to the number of coordinate bonds formed around the central metal ion.

  Common coordination numbers are 6 and 4. For small ligands like NH3 or H2O, the coordination number is 6. However, for larger ligands like Cl-, only 4 can fit around the central ion.

  Some silver complexes have a coordination number of 2, as they only form 2 coordinate bonds, This gives a complex of a linear shape.

  [NH3--Ag--NH3]+
• The familiar colour of many Cu2+ aqueous solutions is due to the [Cu(H2O)6]2+ ion. One lone pair of electrons on the oxygen atom from each water molecule is used to form a coordinate bond to the metal ion.
• Examples of tetrahedral complexes with their colours are:

  [CoCl4]2- blue

  [CuCl4]2- yellow-green.

  The [CuCl4]2- complex is formed when the Cu2+ ions react with concentrated hydrochloric acid, which displaces the water molecules.
• the overall charge on the complex depends on the charge on the central ion and the total charge due to the ligands.
• A ligand exchange reaction is one in which one ligand in a complex ion is replaced by a different one.
• If ammonia solution is added to a solution containing [Cu(H2O)6]2+, ammonia molecules replace four of the water molecules to form a new complex, [Cu(NH3)4(H2O)2]2+.

  This is an equilibrium process.If more ammonia is added, the equilibrium shifts to the right, forming more [Cu(NH3)4(H2O)2]2+ which gives a royal blue colour. If more water is added, the equilibrium shifts to the left, forming more [Cu(H2O)6]2+.

  Because this complex contains two different ligands, there could be two different arrangements, but usually, the two water molecules are opposite each other.
• Ligand exchange can also lead to a change in the geometry of the complex ion.

  If concentrated hydrochloric acid is added to a solution containing [Cu(H2O)6]2+, the six water molecules are replaced by four chloride ions. This is a reversible reaction.

  
  

  Colour change seen - Blue to yellow green
• Concentrated hydrochloric acid is used because it provides a very high concentration of chloride ions. The high chloride ion concentration shifts the equilibrium to the right giving a yellow-green colour. Adding water shifts the equilibrium to the left and the solution returns to blue.

  A similar reaction occurs with [Co(H2O)6]2+ and chloride ions, but the colours are different. (Pink - Blue)

  If a large amount of chloride is used, for example by adding concentrated hydrochloric acid, the equilibrium shifts from the pink octahedral complex to the blue tetrahedral complex.
• Transition metal ions are only coloured when they form complexes.

  Without ligands, all d orbitals in a transition metal ion have the same energy (they are degenerate). When ligands approach the metal ion, they cause the energy of three of the d orbitals to become different to the other two. This splits the d orbitals into two sets: three of lower energy and two of higher energy.
• An electron in a 3d orbital can move from a lower energy set to a higher energy set if it can gain sufficient energy. When visible light is passed through a solution of this ion, some of the energy is absorbed which promotes an electron to a higher 3d orbital. 
• Only one frequency (colour) of light is absorbed, which corresponds to the energy gap between the orbitals (ΔE = hf). The colour you see is made up of the light frequencies that are not absorbed, i.e. those that are reflected. When light of a particular colour is absorbed, its complementary colour is reflected.
• Compounds containing the complex [Cu(H2O)6]2+ are typically pale blue as they absorb light in the red region of the spectrum.
• The energy difference, ΔE, between the split 3d orbitals, and as such the colour, depends on the oxidation state of the metal and the nature and number of ligands. Different ligands cause different splitting of orbitals, so different frequencies are absorbed (giving a different ΔE) and different colours are produced.
• However, not all complexes are coloured, e.g. Cu+, Sc3+. Copper(I) complexes have an electronic configuration with a full d sub-shell, while Sc3+ ions have an empty d sub-shell, meaning that electrons cannot move from lower to higher orbitals. Therefore, copper(I) and scandium(III) complexes appear colourless.
• colours of aqueous solutions of compounds containing the following ions:

  Cr3+ - Green

  Cr2O72- : Orange

  Fe2+ -Pale green

  Co2+ -Pink

  CrO42- : Yellow

  MnO4- : Purple

  Fe3+ - Red-brown

  Cu2+ - Pale blue
• Transition metals and their compounds are used in industry as catalysts for a large range of chemical processes. Many industrial processes would be uneconomical without them; therefore, they are essential to our economy. Here are some examples of their use:

  • iron in the Haber process
  • nickel to make margarine from the hydrogenation of vegetable oils
  • vanadium(V) oxide in the contact process
  • manganese(IV) oxide in the catalytic decomposition of hydrogen peroxide.
• Homogeneous catalysts are in the same physical state as the reactants. They work by using their variable oxidation states to oxidise/reduce a reactant, making it more reactive. The transition metal can then be converted back to its original oxidation state by reaction with another molecule.