Transition metal

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 Hg₂^{2+, 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. (Here's a site that has some info: http://www.jce.divched.org/Journal/Issues/2005/Nov/abs1660.html)

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.

NB. Strictly speaking, neither Sc nor Zn are actually Transition Metals as they are unable to form partially complete d-orbital subshells

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 s 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.
• 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 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, 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 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. 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

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

1. ^ Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999). Advanced Inorganic Chemistry (6th ed.). New York: Wiley.

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