Transition metal

The elements that divide the main groups of the periodic table in the standard view of the table, that is the elements in groups 3 through 12, are termed transition metals or transition elements. The name transition comes from their position in the table. They divide the periods of the table so are a transition between groups 2 and 13. Some of these elements occur naturally in their metallic state and have been known since antiquity. Three of these, gold, silver, and copper are important economically and have been extensively used in coinage and jewelry. Use of copper in tools was one of the first historical technological advances. Transition metals provide some of the important metallic catalysts used in industry and laboratory settings, and iron in the form of steel is used in many things from cars to bridges.

The Transition Metals
Group	3	4	5	6	7	8	9	10	11	12
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

Definitions

The general definition of tranistion metals as groups 3 through 12 (the d-block elements) decribed in the introduction is simple and has been traditionally used. Though this definition is still widely used the characteristic properties of transition metals arise because these elements have partially filled d orbitals. This has lead to a stricter definition of the term. Therefore, more strictly, IUPAC defines the transition metals as any element with an incomplete d subshell or that may form stable ions with an incomplete d subshell (IUPAC definition: "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 (group 12) are not considered to be transition metals. This is because their d orbitals are completly filled in both their elemental and stable ionic states. When they form ions they loose only electrons in their outermost s subshell leaving the d subshell intact. Only a few unstable transient species of these elements that leave ions with a partly filled d subshell have been formed. Element 112 may also be excluded since its electronic configuration is likely to be similar to other members of the group 12, and its oxidation properties are unlikely to be observed due to its radioactive nature. Thus this stricter definition of transition metals limits the term to groups 3 to 11.

The 40 transition metals

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.

Electronic configuration

Elements with atomic numbers 1 through 20 have only electrons in s and p orbitals, with no filled d orbitals in their ground states.

In the fourth period, elements with atomic numbers 21 to 29 (scandium to copper) have a partially filled d subshell or ions with partly filled d subshell. The outer ns orbitals in the d-block elements are of lower energy than the (n-1)d orbitals. As atoms occur in their lowest energy state, the transition metals tend to have their ns orbitals filled with electrons. Hence, these elements all have two electrons in their outer s orbital, with the exception of copper ([Ar]4s¹3d¹⁰) and chromium ([Ar]4s¹3d⁵). These exceptions occur because half- and fully-filled subshells impart unusual stability to the atoms. Similar exceptions are more prevalent in the fifth, sixth and seventh period.

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 coloured compounds.
• They can have a variety of different oxidation states.
• 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 complexes.
• They are often paramagnetic.

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 isolated 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 transistion 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 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.

Catalytic activity

Transition metals form good homogeneous or heterogeneous catalysts, for example iron is the catalyst for the Haber process. Nickel or platinum is used in the hydrogenation of alkenes.

Colored compounds

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.

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.

See also

• inner transition element, a name given to any member of the f-block
• bioinorganic chemistry
• crystal field theory describes the magnetic and optical properties of complexes

Reference

Periodic tables

Standard table | Vertical table | Table with names | Names and atomic masses (large) | Names and atomic masses (small) | Names and atomic masses (text only) | Inline F-block | Elements to 218 | Electron configurations | Metals and non metals | Table by blocks | List of elements by name

Groups:   1 -  2 -  3 -  4 -  5 -  6 -  7 -  8 -  9 - 10 - 11 - 12 - 13 - 14 - 15 - 16 - 17 - 18

Periods:  1  -  2  -  3  -  4  -  5  -  6  -  7  -  8

Series:   Alkalis  -  Alkaline earths  -  Lanthanides  -  Actinides  -  Transition metals  -  Poor metals  -  Metalloids  -  Nonmetals  -  Halogens  -  Noble gases

Blocks:  s-block  -  p-block  -  d-block  -  f-block  -  g-block

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