Transition Metals

• Ligands 


  Molecules of ions that bond to the central metal ion in a complex
• Dative covalent bond 


  In this type of covalent bond, both the shared electrons originally came from the same atom
• Coordination number 


  The number of nearest neighbours by which an atom or ion is surrounded in a structure
• Complex 


  A complex consists of a central metal ion surrounded by ligands
• Coordination compounds 


  Compounds in which a central metal ion is attached to a group of surrounding molecules or ions by dative covalent bond (also known as coordinate bonds)
• List uses of transition metals

  Piping and wiring, coins, construction, jewellery and catalysts
• Describe transition metals' valency


  Transition metals have variable valency - they can form ions with different charges by losing different numbers of electrons. At least one of these ions must have an incomplete d-subshell
• Which metal forms 3 ions?


  Iron forms three different ions. The Fe3+ ion is the most common however, as it has a half-filled d-subshell (added stability)
• What are the anomalies/exceptions?


  Scandium and Zinc only form one ion (each) so they are not transition metals as they do not form any ions with an incomplete d-subshell, they therefore do not have the typical transition metal properties
• What are the different properties of transition metals?

  Form coloured ions

  Form complexes

  Have variable oxidation states

  Show catalytic activity
• Describe how to work out oxidation number

  1. Oxidation number in a free element is zero (e.g. Mg=0, Cl2=0)
  2. For monatomic ions, oxidation number is equal to the charge
  3. In most compounds, the oxidation number for oxygen is -2 and for hydrogen is +1, the hydride ion is an exception (-1)
  4. Group 1 metals are always +1, Group 2 metals are always +2
  5. In compounds, fluorine and the CN- ion are always -1
  6. In a molecule, the sum of all oxidation numbers is equal to zero
  7. In a polyatomic ion, the sum of all oxidation numbers is equal to the charge on the ion
• How do ligands bond to metals?


  By donating non-bonding electron pairs into unfilled metal orbitals
• Denticity


  How many bonds/electron pairs that ligands form/donate to the central ion
• Describe how to name transition metal complexes

  Ligands are named first (alphabetically) and then the transition metal completes the name

  Ligand part: A prefix is attached to the ligand name to show the number of them (prefixes do not affect the alphabetical order)

  Metal part: The metal name is followed by the oxidation number in Roman numerals in brackets

  If the complex ion is negative, the transition metal name ends in 'ate'

  The name of a negative ligand is changed from 'e' to 'o' e.g. bromide becomes bromido
• What happens to the five 3d orbitals in a compound with a metal ion surrounded by ligands

  An electrostatic field is created which repels electrons in some of the TM's d-orbitals
• What affects the amount of orbital splitting?

  The type of ligand
• The d-block transition metals are metals with an incomplete d subshell in at least one of their ions.
• The filling of the d orbitals follows the Aufbau principle, with the exception of chromium and copper atoms.

  These exceptions are due to the special stability associated with the d subshell being half-filled or completely filled.
• When atoms from the first row of the transition elements form ions, it is the 4s electrons that are lost first rather than the 3d electrons.
• An element is said to be in a particular oxidation state when it has a specific oxidation number.
• The oxidation number can be determined by the following:

  • Uncombined elements have an oxidation number of 0.
  • Ions containing single atoms have an oxidation number that is the same as the charge on the ion.
  • In most of its compounds, oxygen has an oxidation number of −2 and hydrogen has an oxidation number of +1. The hydride ion has a charge of -1.
  • The sum of all the oxidation numbers of all the atoms in a neutral compound must add up to zero.
  • The sum of all the oxidation numbers of all the atoms in a polyatomic ion must be equal to the charge on the ion.
• A transition metal can have different oxidation states in its compounds.
• Compounds of the same transition metal in different oxidation states may have different colours.
• Oxidation can be defined as an increase in oxidation number. Reduction can be considered as a decrease in oxidation number.
• Changes in oxidation number of transition metal ions can be used to determine whether oxidation or reduction has occurred.
• Compounds containing metals in high oxidation states are often oxidising agents, whereas compounds with metals in low oxidation states are often reducing agents.
• Ligands may be negative ions or molecules with non-bonding pairs of electrons that they donate to the central metal atom or ion, forming dative covalent bonds.
• Ligands can be classified as monodentate, bidentate, up to hexadentate.

  It is possible to deduce the ligand classification from a formula or structure of the ligand or complex
• The total number of bonds from the ligands to the central transition metal is known as the coordination number.
• Names and formulae can be written according to IUPAC rules for complexes containing:

  • central metals that obey the normal IUPAC rules - copper (cuprate) and iron (ferrate)
  • ligands, including water (aqua), ammonia (ammine), halogens(e to o), cyanide, hydroxide, and oxalate.
• In a complex of a transition metal, the d orbitals are no longer degenerate.
• Splitting of d orbitals to higher and lower energies occurs when the electrons present in approaching ligands cause the electrons in the orbitals lying along the axes to be repelled.
• Ligands that cause a large difference in energy between subsets of d orbitals are strong field ligands. Weak field ligands cause a small energy difference.
• Ligands can be placed in an order of their ability to split d orbitals. This is called the spectrochemical series.
• Spectrochemical series:

  CN- >NH3 >H2O >OH- > F- >Cl- >Br- > I-
• Colours of many transition metal complexes can be explained in terms of d-d transitions. Light is absorbed when electrons in a lower energy d orbital are promoted to a d orbital of higher energy.
• If light of one colour is absorbed, then the complementary colour will be observed. Electrons transition to higher energy levels when energy corresponding to the ultraviolet or visible regions of the electromagnetic spectrum is absorbed.
• Transition metals and their compounds can act as catalysts.
• Heterogeneous catalysts are in a different state to the reactants.

  Homogeneous catalysts are in the same state as the reactants.
• Heterogeneous catalysis can be explained in terms of the formation of activated complexes and the adsorption of reactive molecules onto active sites. The presence of unpaired d electrons or unfilled d orbitals is thought to allow activated complexes to form. This can provide reaction pathways with lower activation energies compared to the uncatalysed reaction.