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 atomic orbitals of the atoms as one progresses through each of the three periods. Transition elements are chemically defined as elements which form at least one ion with a partially filled subshell of d electrons.
|Table of contents|
2 Chemical properties
3 Variable oxidation states
4 Catalytic activity
5 Coloured compounds
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 orbitals, but only in the s and p orbitals.
(Though the low-lying, but empty d orbitals are thought to play a role in their d period elements such as silicon, phosphorus and sulfur)
From Scandium to Zinc, d block elements fill up their d orbitals across the period. With the exception of copper and chromium, all d block elements have two electrons in their outer s orbital, even elements with incomplete 3d orbitals.
This is unusual: lower orbitals are usually filled up before outer shells. It happens that the s orbitals in d block elements are at lower energy states than the d subshells. As atoms always strive to be in states of lowest energy, s shells are filled up first. The copper and chromium exceptions - which have one electron in their outer orbital - do so because of electron repulsion. Sharing the electrons throughout the s and d orbitals gives lower energy states to the atoms 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 the chemical definition given above. Scandium has one electron in its d subshell, and 2 electrons in its outer s orbital. As scandium's only ion (Sc3+) has no electrons in its d orbital it is clear that it doesn't have a 'partially filled d orbital'. Similarly, zinc is not applicable because its only ion, Zn2+, has a full d orbital.
Transition elements tend to have high tensile strength, density and melting and boiling points. As with many properties of transition metals, this is due to d orbital electrons' ability to delocalise within the metal lattice.
In metallic substances, the more electrons shared between nuclei, the stronger the metal.
There are four common characteristic properties of transition elements:
Certain patterns can be seen to emerge across the period of transition elements:
We observe colour as varying frequencies of electromagnetic radiation in the visible region of the electromagnetic spectrum. Different colours result from the changed composition of light after it has been reflected, transmitted or absorbed after hitting a substance.
Because of their structure, transition metals form many different coloured ions and complexes.
Colour even varies between the different ions of a single element - MnO4- (Mn in oxidation state 7+) is a purple compound, whereas Mn2+ is pale-pink.
Complex formation can play a part in determining colour in a transition compound. This is because of the effect that ligands have on the 3d subshell. Ligands pull on some of the 3d electrons and split them in to higher and lower (in terms of energy) groups. Electromagnetic radiation is only absorbed if its frequency is proportional to the difference in energies between two energy states present in an atom (through the formula e=hf.) When light hits an atom which has had its 3d orbitals split, some of the electrons become promoted to the higher group. Compared to an un-complexed ion, different frequencies can be absorbed, hence different colours are observed.
The colour of a complex depends on: