Transition metal

The thirty chemical elements 21 through 30, 39 through 48, and 71 through 80, are commonly referred to as the transition metals. This name comes from their position in the periodic table of elements, which represent the successive addition of electrons to the d orbitals of the atoms as one progresses through each of the three periods. Transition elements are chemically defined as any element which forms at least one ion with a partially filled subshell of d electrons.

Main group elements prior to the appearance of the transition group elements in the periodic chart (ie, elements number 1 through 20) have no electrons in d atomic orbitals, but only in their s and p orbitals. (Though the low-lying, but empty d orbitals are thought to play a role in third-period elements such as silicon, phosphorus and sulfur)

                     /  Sc 21    Y 39   Lu 71   Lr 103    3 (IIIB)
                     |  Ti 22   Zr 40   Hf 72  Unq 104    4 (IVB)
                     |   V 23   Nb 41   Ta 73  Unp 105    5 (VB)
                     |  Cr 24   Mo 42    W 74  Unh 106    6 (VIB)
   transition metal /   Mn 25   Tc 43   Re 75  Uns 107    7 (VIIB)
                    \   Fe 26   Ru 44   Os 76  Uno 108    8 (VIIIB)
                     |  Co 27   Rh 45   Ir 77  Une 109    9 (VIIIB)
                     |  Ni 28   Pd 46   Pt 78            10 (VIIIB)
                     |  Cu 29   Ag 47   Au 79            11 (IB)
                      \ Zn 30   Cd 48   Hg 80            12 (IIB)

Electronic configuration

The elements fron Scandium to Zinc fill up their d-orbitals across the period. They all have two electrons in the outer s-orbital (except for copper and chromium which have one). This is unusual: the d-orbitals are usually filled up first before the s shell. It happens that the s-orbitals in d block elements have a lower energy than the d-subshell. The copper and chromium exceptions - which have one electron in their outer shell - do so because of electron repulsion. Sharing the electrons throughout the s and d-orbitals gives lower energy levels than putting two electrons in the outer s-orbital.

Not all d block elements are transition metals. Scandium and zinc don't qualify, due to chemical definition given above. Scandium has one electron in the d-subshell, and 2 electrons in its outer s-orbital. It forms only one ion, Sc³⁺, where there are no electrons in the d-orbital. Zinc also doesn't qualify. Its only ion Zn²⁺ has a full d-orbital.

Chemical properties

Transition elements tend to have a high tensile strength, density and melting and boiling points. This is due to the d-orbital electrons being able to be delocalized? within the metal lattice. The more electrons released then the stronger the metal, as the nucleus' can bind to many electrons at once.

There are four characteristic properties of transition elements:

• They form coloured compounds
• They can have a variety of different oxidation states
• They are good catalysts
• They form complexes

Variable oxidation states

Compared to a group II element such as calcium, the transition elements form many more oxidation states. The ionisation enthalpies of calcium are low until you try to remove electrons below its outer s-orbitals. Thus Ca³⁺ has an ionisation enthalpy so high that it rarely occurs naturally. However a transition element like vanadium has roughly linear increasing ionisation enthalpies, due to the close energy difference between the 3d and 4s orbitals.

There are certain patterns which emerge across period of transition elements:

• The number of oxidation states increases up to Mn, after which they start to fall again. The drop is due to the stronger pull from the protons in the nucleus towards the electrons, making them harder to remove.
• When the elements are in lower oxidation states, they can be found as simple ions. However elements in higher oxidation states usually bond covalently to electronegative compounds such as O or F, often in an anion.

Properties with respect to stability of oxidation states:

• The stability of higher oxidation states of the ions decreases across the period
• Ions in higher oxidation states tend to be oxidising agents, whereas elements in low oxidation states are good reducing agents
• The 2+ ions across the period start as strong reducing agents, and become more stable
• The 3+ ions start stable and become more oxidising across the period

Catalytic activity

Transition metals form good homogenous or heterogenous catalysts, for example iron is the catalyst for the haber process. Nickel or platinum is used in the hydrogenation of alkenes.