Transition metal

From Wikipedia, the free encyclopedia

Jump to: navigation, search

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, cadmium and mercury. This corresponds to groups 3 to 12 on the periodic table.

More strictly, IUPAC defines a transition metal as "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell." By this definition, zinc, cadmium, and mercury are excluded from the transition metals, as they have a d¹⁰ configuration. Only a few transient species of these elements that leave ions with a partly filled d subshell have been formed, and mercury(I) only occurs as Hg₂²⁺, which does not strictly form a lone ion with a partly filled subshell, and hence these three elements are inconsistent with the latter definition.^[1] They do form ions with a 2+ oxidation state, but these retain the 4d¹⁰ configuration. Element 112 may also be excluded although its oxidation properties are unlikely to be observed due to its radioactive nature. This definition corresponds to groups 3 to 11 on the periodic table.

The first definition is simple and has traditionally been used. However, many interesting properties of the transition elements as a group are the result of their partly filled d subshells. Periodic trends in the d block (transition metals) are less prevailing than in the rest of the periodic table. Going across a period, the valence doesn't change, so the electron being added to an atom goes to the inner shell, not outer shell, strengthening the shield. ^[2]

The (loosely defined) transition metals are the 40 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.

Group	3 (III B)	4 (IV B)	5 (V B)	6 (VI B)	7 (VII B)	8 (VIII B)	9 (VIII B)	10 (VIII B)	11 (I B)	12 (II B)
Period 4	Sc 21	Ti 22	V 23	Cr 24	Mn 25	Fe 26	Co 27	Ni 28	Cu 29	Zn 30
Period 5	Y 39	Zr 40	Nb 41	Mo 42	Tc 43	Ru 44	Rh 45	Pd 46	Ag 47	Cd 48
Period 6	Lu 71	Hf 72	Ta 73	W 74	Re 75	Os 76	Ir 77	Pt 78	Au 79	Hg 80
Period 7	Lr 103	Rf 104	Db 105	Sg 106	Bh 107	Hs 108	Mt 109	Ds 110	Rg 111	Uub 112

1 Properties
2 Variable oxidation states
3 Catalytic activity
4 Colored compounds
5 See also
6 References

[edit] Properties

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 several common characteristic properties of transition elements:

They often form colored compounds.
They can have a variety of different oxidation states.
At least one of their compounds has an incomplete d-electron subshell.
They are often good catalysts.
They are silvery-blue at room temperature (except copper and gold).
They are solids at room temperature (except mercury).
They form complex ions (aqua ions included).
They are often paramagnetic.

[edit] Variable oxidation states

As opposed to group 1 and group 2 metals, ions of the transition elements may have multiple stable oxidation states, since they can lose d electrons without a high energetic penalty. Manganese, for example has two 4s electrons and five 3d electrons, which can be removed. Loss of all of these electrons leads to a +7 oxidation state. Osmium and ruthenium compounds are commonly found alone in stable +8 oxidation states, which is among the highest for isolable compounds.

This table shows some of the oxidation states found in compounds of the transition-metal elements.
A solid circle represents a common oxidation state, and a ring represents a less common (less energetically favourable) oxidation state.

Certain patterns in oxidation state emerge across the period of transition elements:

The number of oxidation states of each ion increases up to Mn, after which they decrease. Later transition metals have a stronger attraction between protons and electrons (since there are more of each present), which then would require more energy to remove the electrons.
When the elements are in lower oxidation states, they can be found as simple ions. However, transition metals in higher oxidation states are usually bonded covalently to electronegative elements like oxygen or fluorine, forming polyatomic ions such as chromate, vanadate, or permanganate.

Other properties with respect to the stability of oxidation states:

Ions in higher oxidation states tend to make good oxidizing 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.

[edit] Catalytic activity

Transition metals form good homogeneous or heterogeneous catalysts, for example iron is the catalyst for the Haber process. Vanadium(V) oxide is used for the contact process, nickel is used to make margarine and platinum is used to speed up the manufacture of nitric acid. This is because they are able to form numerous oxidation states, and as such, are able to form new compounds during a reaction providing an alternative route with a lower overall activation energy.

[edit] Colored compounds

From left to right, aqueous solutions of: Co(NO₃)₂ (red); K₂Cr₂O₇ (orange); K₂CrO₄ (yellow); NiCl₂ (green); CuSO₄ (blue); KMnO₄ (purple).

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.

Coordination by ligands can play a part in determining color in a transition compound, due to changes in energy of the d orbitals. Ligands remove degeneracy of the orbitals and split them in to higher and lower energy groups. The energy gap between the lower and higher energy orbitals will determine the color of light that is absorbed, as electromagnetic radiation is only absorbed if it has energy corresponding to that gap. When a ligated ion absorbs light, some of the electrons are promoted to a higher energy orbital. Since different frequency light is absorbed, 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.