Science Fair Project Encyclopedia
In chemistry, the term transition metal (sometimes also called a transition element) has two possible meanings:
- It commonly refers to any element in the d-block of the periodic table, including zinc and scandium. This corresponds exactly to periodic table groups 3 to 12 inclusive.
- More strictly, it can refer to those elements which form at least one ion with a partially filled d shell of electrons. This is exactly the d-block with zinc and scandium excluded.
Both definitions have their uses and supporters. The first has the attraction of apparent simplicity and is the traditional usage. However, many interesting properties of the transition elements as a group are the result of their ability to contribute valence electrons from s orbitals before d orbitals, a property which all members of the d-block except zinc and scandium share, so the more restricted definition is in many contexts the more useful. The d orbitals are contributed after the s orbitals because once the d orbital begins to fill its electrons move closer to the nucleus, leaving the s electrons as the outermost.
|Group||Period 4||Period 5||Period 6||Period 7|
|3 (III B)||Sc 21||Y 39||Lu 71||Lr 103|
|4 (IV B)||Ti 22||Zr 40||Hf 72||Rf 104|
|5 (V B)||V 23||Nb 41||Ta 73||Db 105|
|6 (VI B)||Cr 24||Mo 42||W 74||Sg 106|
|7 (VII B)||Mn 25||Tc 43||Re 75||Bh 107|
|8 (VIII B)||Fe 26||Ru 44||Os 76||Hs 108|
|9 (VIII B)||Co 27||Rh 45||Ir 77||Mt 109|
|10 (VIII B)||Ni 28||Pd 46||Pt 78||Ds 110|
|11 (I B)||Cu 29||Ag 47||Au 79||Rg 111|
|12 (II B)||Zn 30||Cd 48||Hg 80||Uub 112|
The (loosely defined) transition metals are the forty chemical elements 21 to 30, 39 to 48, 71 to 80, and 103 to 112. The name transition comes from their position in the periodic table of elements. In each of the four periods in which they occur, these elements represent the successive addition of electrons to the d atomic orbitals of the atoms. In this way, the transition metals represent the transition between group 2 elements and group 13 elements.
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. The 3rd period p block elements have empty d orbitals. In the fourth period from scandium to zinc, d-block elements fill up their d orbitals across the period. With the exception of the copper group and the chromium group, all d-block elements in the ground state have two electrons in their outer s orbital. The electronic configuration of the d-block elements is ns2(n-1)d1-10, where n is the ground state principal quantum number.
The outer s orbitals in the d-block elements are at lower energy states than the d orbitals of the n-1 levels. As atoms always strive to be in states of lowest energy, s orbitals are filled up first. The copper (4s13d10) and chromium (4s13d5) exceptions, which have one electron in their outer orbital, occur because half- and fully-filled orbitals are more stable than any other configurations (this occurs when there are 5 or 10 electrons in the d-orbitals).
Scandium has one electron in its d orbital, 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 does not have a 'partially filled d orbital', and is not a transition metal in the stricter sense. Similarly, zinc is not a transition metal in the stricter sense because its only ion, Zn2+, has a full d orbital, which does not participate in bonding.
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:
- They form colored compounds
- They can have a variety of different oxidation states
- They are good catalysts
- They are silvery-blue at room temperature (except copper and gold)
- They are solids at room temperature (except mercury)
- They form complexes, which is described by crystal field theory.
Variable oxidation states
Compared to Group II elements such as calcium, transition elements form ions with a wide variety of oxidation states. The transition metals show such a range of oxidation states because their partially filled d orbitals can accept or donate electrons in chemical reactions. Calcium ions typically do not lose more than two electrons, whereas transition metals can lose up to nine. The reason for this can be obtained by studying the ionisation enthalpies of both groups. The energies required to remove electrons from calcium are low until you try to remove electrons from below its outer two 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 throughout its s and d orbitals, due to the close energy difference between the 3d and 4s orbitals. Transition metal ions are therefore commonly found in very high states.
Certain patterns can be seen to emerge across the period of transition elements:
- The number of oxidation states of each ion increases up to Mn, after which they start to drop. This 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 are usually bonded covalently to electronegative compounds such as O or F, often as a polyatomic ion such as chromate, vanadate , and permanganate ions.
Properties with respect to the stability of oxidation states:
- Higher oxidation state ions become less stable across the period.
- Ions in higher oxidation states tend to make good oxidising agents, whereas elements in low oxidation states become 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 oxidizing across the period.
We observe color as varying frequencies of electromagnetic radiation in the visible region of the electromagnetic spectrum. Different colors 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 colored ions and complexes. Color 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 color in a transition compound. This is because of the effect that ligands have on the 3d orbital. 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 colors are observed.
The color of a complex depends on:
- the nature of the metal ion, specifically the number of electrons in the d orbitals
- the arrangement of the ligands around the metal ion (for example geometric isomers can display different colors)
- 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 colorless, because the 3d orbitals are full - no electrons are able to move up to the higher group.
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