Transition Metal

In common terminology, transition metals (or transition elements) are chemical elements that lie in groups 3 through 12 of the periodic table, in the standard view of the table. The name transition comes from their position in the table—they form a transition between the main group elements, which occur in groups 1 and 2 on the left side, and groups 13–18 on the right.

Some transition elements occur naturally in their metallic state and have been known since antiquity. Three of these—gold, silver, and copper—have been used extensively in coinage and jewelry. The use of copper in tools was one of the first historical technological advances. Also, iron, in the form of steel, is used in many structures, from automobiles to bridges. Many transition metals are useful as catalysts in industrial and laboratory settings, and many of these elements form brightly colored compounds.

The Transition Metals

Group →	3	4	5	6	7	8	9	10	11	12
Period ↓
4	21 Sc	22 Ti	23 V	24 Cr	25 Mn	26 Fe	27 Co	28 Ni	29 Cu	30 Zn
5	39 Y	40 Zr	41 Nb	42 Mo	43 Tc	44 Ru	45 Rh	46 Pd	47 Ag	48 Cd
6	57 La	72 Hf	73 Ta	74 W	75 Re	76 Os	77 Ir	78 Pt	79 Au	80 Hg
7	89 Ac	104 Rf	105 Db	106 Sg	107 Bh	108 Hs	109 Mt	110 Ds	111 Rg	112 Uub

Periodic table

Placement of the group of transition elements in the periodic table can be observed by examining the color-coded table shown below.

Group →	1	2	3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
Period ↓
1	1 H																	2 He
2	3 Li	4 Be											5 B	6 C	7 N	8 O	9 F	10 Ne
3	11 Na	12 Mg											13 Al	14 Si	15 P	16 S	17 Cl	18 Ar
4	19 K	20 Ca	21 Sc	22 Ti	23 V	24 Cr	25 Mn	26 Fe	27 Co	28 Ni	29 Cu	30 Zn	31 Ga	32 Ge	33 As	34 Se	35 Br	36 Kr
5	37 Rb	38 Sr	39 Y	40 Zr	41 Nb	42 Mo	43 Tc	44 Ru	45 Rh	46 Pd	47 Ag	48 Cd	49 In	50 Sn	51 Sb	52 Te	53 I	54 Xe
6	55 Cs	56 Ba	*	72 Hf	73 Ta	74 W	75 Re	76 Os	77 Ir	78 Pt	79 Au	80 Hg	81 Tl	82 Pb	83 Bi	84 Po	85 At	86 Rn
7	87  Fr	88 Ra	**	104 Rf	105 Db	106 Sg	107 Bh	108 Hs	109 Mt	110 Ds	111 Rg	112 Uub	113 Uut	114 Uuq	115 Uup	116 Uuh	117 Uus	118 Uuo
* Lanthanides				57 La	58 Ce	59 Pr	60 Nd	61 Pm	62 Sm	63 Eu	64 Gd	65 Tb	66 Dy	67 Ho	68 Er	69 Tm	70 Yb	71 Lu
** Actinides				89 Ac	90 Th	91 Pa	92 U	93 Np	94 Pu	95 Am	96 Cm	97 Bk	98 Cf	99 Es	100 Fm	101 Md	102 No	103 Lr

Chemical Series of the Periodic Table

Alkali metals	Alkaline earth metals	Lanthanides	Actinides	Transition metals
Poor metals	Metalloids	Nonmetals	Halogens	Noble gases

State at standard temperature and pressure

• Elements numbered in red are gases.
• Elements numbered in green are liquids.
• Elements numbered in black are solids.

Natural occurrence

• Elements without borders have not been discovered/synthesized yet.
• Elements with dotted borders do not occur naturally (synthetic elements).
• Elements with dashed borders naturally arise from decay of other chemical elements.
• Elements with solid borders are older than the Earth (primordial elements).
  • Note: Although californium (Cf, 98) is not Earth-primordial, it (and its decay products) does occur naturally: its electromagnetic emissions are regularly observed in supernova spectra.

Definitions

The general definition of transition metals as those that lie in groups 3 through 12 of the periodic table, mentioned above, is simple and has been traditionally used. Although this definition is still widely used, the characteristic properties of transition metals arise because of the electron configuration of their atoms, which have partially filled "d orbitals." Based on this perspective, the term transition element has been defined more strictly. The International Union of Pure and Applied Chemistry (IUPAC) defines a transition element as "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell."^[1]

By this definition, zinc, cadmium, and mercury (group 12 elements) are not considered transition metals. This is because the atoms of these elements and their stable ions contain electrons that completely fill the d orbitals. When these elements form ions, they usually lose electrons from only their outermost s subshell, leaving the d subshell intact. In just a few, exceptional cases, they have formed unstable ions in which the d subshell is partly filled.^[2] Element 112 (in group 12) may also be excluded, because its electron configuration is likely to be similar to that of other members of 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 elements in groups 3 to 11.

Properties

There are several common characteristic properties of transition elements:

• Almost all of them are solids at room temperature, with high tensile strength (ability to withstand stress), density, and melting and boiling points. The one exception is mercury, which is a liquid.
• Most of them are silvery-blue at room temperature. The exceptions are copper and gold.
• They form monatomic ions with a 2+ charge, but can form other ions with a different charge. For example, iron can form Fe²⁺ and Fe³⁺ ions. In addition, they often have higher oxidation states in compounds.
• They form complexes known as "coordination compounds," many of which are brightly colored.
• They are often good catalysts. For example, iron is the catalyst for the Haber process, involving the reaction of nitrogen and hydrogen to produce ammonia. Nickel, palladium, or platinum can be used in the hydrogenation of (addition of hydrogen atoms to) alkenes and alkynes. Platinum is the catalyst in the catalytic converters of automobile exhaust systems.

In addition to these common characteristics, there are some trends in properties as we go through a period, much like those in the main group elements, but with less dramatic changes. Going across the transition metals of a period, the atomic radius generally tends to decrease, and the first ionization energy (energy required to remove an electron from the neutral atom) increases. Also, as we go across the period, the metals tend to become softer, and mercury is a liquid at room temperature. Group 11 elements (copper, silver, and gold) are particularly unreactive. These "noble" metals can occur naturally in their elemental metallic state, and they are sometimes known as coinage metals as they have been useful for minting coins.

Electronic configuration

Main article: electron configuration

The properties of transition metals arise from their defining characteristic of partially filled d orbitals. They are metals because the d orbital electrons are delocalized within the metal lattice, forming metallic bonds.

Most transition metals have two electrons in their outermost, s subshell. As we consider these elements across a period, the number of d electrons increases by one. Thus, in the fourth period, scandium (Sc, group 3) has the configuration [Ar]4s²3d¹, and the next element Titanium (Ti, group 4) has the configuration [Ar]4s²3d², and so forth. There are, however, some exceptions to this progression. For instance, in the fourth period, copper has the configuration ([Ar]4s¹3d¹⁰) and chromium is ([Ar]4s¹3d⁵). These exceptions occur because the atoms acquire additional stability when their subshells are half-filled or fully filled. Copper has a completely filled d subshell, and chromium has a half-filled d subshell. Similar exceptions are more prevalent in the fifth, sixth, and seventh periods.

When these metals lose electrons to form monatomic ions, they generally lose their s electrons first. Thus, most transition metals form ions with a 2+ charge. Higher oxidation states involve d electrons as well. Monatomic ions with a charge greater than 3+ are rare, and the higher oxidation states of transition metals occur in compounds with highly electronegative elements such as oxygen.

Variable oxidation states

Unlike ions of most main group metals, monatomic ions of the transition metals may have more than one stable charge, and, in compounds, they can have several higher oxidation states. (Oxidation state is a measure of the degree of oxidation of an atom in a compound; it is the electrical charge an atom would have, at least hypothetically, if its bonds to all other atoms in the compound were entirely ionic.)

This variability of oxidation state is because the atoms of transition elements can lose or share d electrons without a high energetic penalty. The atom of manganese, for example, has two 4s electrons and five 3d electrons, which can be removed or shared with other atoms. Loss or sharing 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.

Moving across a period of transition elements, certain patterns in their oxidation states emerge:

• The number of oxidation states of each element increases up to manganese (group 7), after which they decrease. Later transition metals have a stronger attraction between protons and electrons (because there are more of them present), requiring more energy to remove the electrons.
• When these elements are in lower oxidation states, they can be found as simple ions. In their higher oxidation states, these elements are usually bonded covalently to electronegative elements like oxygen or fluorine, forming polyatomic ions such as chromate, vanadate, or permanganate.

Other properties associated with the stability of oxidation states are as follows:

• Ions in higher oxidation states tend to make good oxidizing agents, whereas elements in low oxidation states become reducing agents.
• Going across a period, the 2+ ions start as strong reducing agents and increase in stability.
• Conversely, the 3+ ions start at higher stability and become more oxidizing across the period.

Colored compounds

As noted above, the chemistry of transition metals is characterized by the partially filled d orbitals allowing for multiple oxidation states. Another consequence of their electron configuration is that these elements can form stable complexes, or coordination compounds. In such a complex, the transition metal atom or ion forms weak covalent bonds to other small molecules or ions known as "ligands." In some cases, the oxidation state of the transition metal may be zero or a negative number.

Transition metal compounds are often highly colored and coordination by ligands plays a large part in determining the compound's color. In the absence of ligands, the d orbitals of an atom all have the same energy, but when surrounded by ligands, the energies of the d orbitals change and are no longer equal. This phenomenon is described by the cystal field theory. For many compounds of this type, the resulting difference in energy of the d orbitals is in the energy range of visible light. As a result, they strongly absorb particular wavelengths of visible light and appear vividly colored. Many different colors can be observed, and the color can vary even between different ions of the same element. A striking example is the different ions of vanadium (V): VO₂⁺ is yellow in solution, VO²⁺ is blue, V³⁺(aq) is green and V²⁺(aq) is purple.

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; and
• the nature of the ligands surrounding the metal ion. (The stronger the ligand, the greater the energy difference between the different d orbitals.)

Interestingly, though zinc can form complexes, they are colorless because the 3d orbitals of zinc are completely filled. The full d orbitals prevent the complex from absorbing visible light when the energies of the d orbitals are altered by ligands. As zinc is in group 12, it is not considered a transition metal by the newer IUPAC definition.

See also

• Chemical element
• Metal
• Periodic table

Notes

References

ISBN links support NWE through referral fees

• Cotton, F. Albert, G. Wilkinson, C.A. Murillo, and M. Bochmann. 1999. Advanced Inorganic Chemistry, 6th ed. New York: Wiley. ISBN 0471199575

• Crabtree, Robert H. 2005. The Organometallic Chemistry of the Transition Metals. Hoboken, NJ: Wiley. ISBN 978-0471662563

• Greenwood, N. N., and A. Earnshaw. 1997. Chemistry of the Elements, 2nd ed. Oxford: Butterworth-Heinemann. ISBN 0750633654

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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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