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)
From Scandium to Zinc, d block elements 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 is not applicable because its only ion, Zn2+, has a full d-orbital.
Transition elements tend to have a high tensile strength, density and melting and boiling points. This is due to d-orbital electrons' ability to delocalise within the metal lattice. The more electrons released then the stronger the metal, as the more electrons the nucleus' can bind to, the stronger the bond between the ions.
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. In fact 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, meaning it is found in many different oxidation states.
There are certain patterns which emerge across period of transition elements:
- The number of oxidation states of each ion 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
Colour results from the composition of light after it has been reflected, transmitted or absorbed in the visible region of electromagnetic radiation. A key characteristic of transition metals is the many different coloured ions and complexes that they form. This varies from ion to ion - MnO4- (Mn in oxidation state 7+) is a purple compound, whereas Mn2+ is pale-pink.
Complex formation plays a large part in determining colour in a transition compound. This is because of the effect that ligands have on the 3d subshell. Ligands split the 3d in to two - a higher, and a lower group - by pulling the higher energy group towards them. Electromagnetic radiation is only absorbed if its frequency is proportional to the difference in energies when an electron is promoted from the lower to the higher group (through the formula e=hf.) This means that differing complex properties determine different colours.
The colour if a complex depends on:
- The nature of the metal ion, specificly the number of electrons in the d-orbitals
- The arrangement of the ligands around the metal ion (for example geometric isomers can display different colours)
- The nature of the ligands surrounding the metal ion. The stronger the ligands then the greater the energy difference between the split high and low 3d-groups.
The complex formed by the d block element zinc (though not strictly a transition element) is colourless, because the 3d orbitals are full up - no electrons are able to move up to a higher group.