Transition metal

The term transition metal (sometimes also called a transition element) has two possible meanings:

• The IUPAC definition^[1] states that a transition metal is "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell." Group 12 elements are not transition metals in this definition.
• Some authors describe a "transition metal" as any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table. All elements in the d-block are metals. In actual practice, the f-block is also included in the form of the lanthanide and actinide series.

Jensen^[2] has reviewed the history of the terms transition element (or metal) and d-block. The word transition was first used to describe the elements now known as the d-block by the English chemist Charles Bury in 1921, who referred to a transition series of elements during the change of an inner layer of electrons (for example n=3 in the 4th row of the periodic table) from a stable group of 8 to one of 18, or from 18 to 32.^[3]

Classification

In the d-block the atoms of the elements have between 1 and 10 d electrons.

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	* 57–71	Hf 72	Ta 73	W 74	Re 75	Os 76	Ir 77	Pt 78	Au 79	Hg 80
Period 7	** 89–103	Rf 104	Db 105	Sg 106	Bh 107	Hs 108	Mt 109	Ds 110	Rg 111	Cn 112

With a few minor exceptions, the electronic structure of transition metal atoms can be written as [ ]ns²(n-1)d^m, where the inner d orbital has more energy than the valence-shell s orbital. In divalent and trivalent ions of the transition metals, the situation is reversed such that the s electrons have higher energy. Consequently, an ion such as Fe²⁺ has no s electrons: it has the electronic configuration [Ar]3d⁶ as compared with the configuration of the atom, [Ar]4s²3d⁶.

The elements of groups 4–11 are now generally recognized as transition metals, as are Sc and Y in Group 3. For the elements La-Lu and Ac-Lr and also for Group 12, different sets of definitions are used by different authors.

1. Many chemistry textbooks and printed periodic tables classify La and Ac as Group 3 elements and transition metals, since their atomic ground-state configurations are s²d¹ like Sc and Y. The elements Ce-Lu are considered as the “lanthanide” series (or “lanthanoid” according to IUPAC) and Th-Lr as the “actinide” series.^[4][5] The two series together are classified as f-block elements, or (in older sources) as “inner transition elements”.
2. Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides.^[6][7][8] This classification is based on similaritites in chemical behaviour, and defines 15 elements in each of the two series even though they correspond to the filling of an f subshell which can only contain 14 electrons.
3. A third classification defines the f-block elements as La-Yb and Ac-No, while placing Lu and Lr in Group 3.^[2] This is based on the aufbau principle (or Madelung rule) for filling electron subshells, in which 4f is filled before 5d (and 5f before 6d), so that the f subshell is actually full at Yb (and No) while Lu (and Lr) has an [ ]s²f¹⁴d¹ configuration. However La and Ac are exceptions to the Aufbau principle with electron configuration [ ]s²d¹ (not [ ]s²f¹ as the aufbau principle predicts) so it is not clear from atomic electron configurations whether La or Lu (Ac or Lr) should be considered a transition metal.

Zinc, cadmium, and mercury are sometimes not classified as transition metals^[2] as they have the electronic configuration [ ]d¹⁰s², with no incomplete d shell.^[9] In the oxidation state +2 the ions have the electronic configuration [ ] d¹⁰. While these elements can exist in the +1 oxidation state, as in the diatomic ion Hg2+
2, there are no unpaired electrons because of the formation of a covalent bond between the two atoms of the dimer. The group 12 elements Zn, Cd and Hg may be classed as post-transition metals in this case. However, it is often convenient to include these elements in a discussion of the transition elements. For example, when discussing the crystal field stabilization energy of first-row transition elements, it is convenient to also include the elements calcium and zinc, as both Ca²⁺ and Zn²⁺ have a value of zero against which the value for other transition metal ions may be compared. Another example occurs in the Irving-Williams series of stability constants of complexes.

However the recent synthesis of mercury(IV) fluoride (HgF₄) has led to debate as to whether mercury should now always be considered a transition metal.^[10][11]

Characteristic properties

There are a number of properties shared by the transition elements that are not found in other elements, which results from the partially filled d shell. These include

• the formation of compounds whose colour is due to d – d electronic transitions
• the formation of compounds in many oxidation states, due to the relatively low reactivity of unpaired d electrons.^[12]
• the formation of many paramagnetic compounds due to the presence of unpaired d electrons. A few compounds of main group elements are also paramagnetic (e.g. nitric oxide, oxygen)

Coloured compounds

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

Colour in transition-series metal compounds is generally due to electronic transitions of two principal types.

• charge transfer transitions. An electron may jump from a predominantly ligand orbital to a predominantly metal orbital, giving rise to a ligand-to-metal charge-transfer (LMCT) transition. These can most easily occur when the metal is in a high oxidation state. For example, the colour of chromate, dichromate and permanganate ions is due to LMCT transitions. Another example is that mercuric iodide, HgI₂, is red because of a LMCT transition.

A metal-to ligand charge transfer (MLCT) transition will be most likely when the metal is in a low oxidation state and the ligand is easily reduced.

• d-d transitions. An electron jumps from one d-orbital to another. In complexes of the transition metals the d orbitals do not all have the same energy. The pattern of splitting of the d orbitals can be calculated using crystal field theory. The extent of the splitting depends on the particular metal, its oxidation state and the nature of the ligands. The actual energy levels are shown on Tanabe-Sugano diagrams.

In centrosymmetric complexes, such as octahedral complexes, d-d transitions are forbidden by the Laporte rule and only occur because of vibronic coupling in which a molecular vibration occurs together with a d-d transition. Tetrahedral complexes have somewhat more intense colour because mixing d and p orbitals is possible when there is no centre of symmetry, so transitions are not pure d-d transitions. The molar absorptivity (ε) of bands caused by d-d transitions are relatively low, roughly in the range 5-500 M⁻¹cm⁻¹ (where M = mol dm⁻³).^[13] Some d-d transitions are spin forbidden. An example occurs in octahedral, high-spin complexes of manganese(II), which has a d⁵ configuration in which all five electron has parallel spins; the colour of such complexes is much weaker than in complexes with spin-allowed transitions. In fact many compounds of manganese(II) appear almost colourless. The spectrum of [Mn(H₂O)₆]²⁺ shows a maximum molar absorptivity of about 0.04 M⁻¹cm⁻¹ in the visible spectrum.

Oxidation states

A characteristic of transition metals is that they exhibit two or more oxidation states, usually differing by one. For example, compounds of vanadium are known in all oxidation states between −1, such as [V(CO)₆]⁻, and +5, such as VO3−
4.

Main group elements in groups 13 to 17 also exhibit multiple oxidation states. The "common" oxidation states of these elements typically differ by two. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No compound of Ga(II) is known: any such compound would have an unpaired electron and would behave as a free radical and be destroyed rapidly. The only compounds in which gallium has a formal oxidation state of +2 are dimeric compounds, such as [Ga₂Cl₆]²⁻, which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom.^[14] Thus the main difference in oxidation states, between transition elements and other elements is that oxidation states are known in which there is a single atom of the element and one or more unpaired electrons.

The maximum oxidation state in the first row transition metals is equal to the number of valence electrons from titanium (+4) up to manganese (+7), but decreases in the later elements. In the second and third rows the maximum occurs with ruthenium and osmium (+8). In compounds such as [MnO₄]⁻ and OsO₄ the elements achieve a stable octet by forming four covalent bonds.

The lowest oxidation states are exhibited in such compounds as Cr(CO)₆ (oxidation state zero) and [Fe(CO)₄]²⁻ (oxidation state −2) in which the 18-electron rule is obeyed. These complexes are also covalent.

Ionic compounds are mostly formed with oxidation states +2 and +3. In aqueous solution the ions are hydrated by (usually) six water molecules arranged octahedrally.

Magnetism

Main article: magnetochemistry

Transition metal compounds are paramagnetic when they have one or more unpaired d electrons.^[15] In octahedral complexes with between four and seven d electrons both high spin and low spin states are possible. Tetrahedral transition metal complexes such as [FeCl₄]²⁻ are high spin because the crystal field splitting is small so that the energy to be gained by virtue of the electrons being in lower energy orbitals is always less than the energy needed to pair up the spins. Some compounds are diamagnetic. These include octahedral, low-spin, d⁶ and square-planar d⁸ complexes. In these cases, crystal field splitting is such that all the electrons are paired up.

Ferromagnetism occurs when individual atoms are paramagnetic and the spin vectors are aligned parallel to each other in a crystalline material. Metallic iron and the alloy alnico are examples of ferromagnetic materials involving transition metals. Anti-ferromagnetism is another example of a magnetic property arising from a particular alignment of individual spins in the solid state.

Catalytic Properties

The transition metals and their compounds are known for their homogeneous and heterogeneous catalytic activity. This activity is ascribed to their ability to adopt multiple oxidation states and to form complexes. Vanadium(V) oxide (in the contact process), finely divided iron (in the Haber process), and nickel (in Catalytic hydrogenation) are some of the examples. Catalysts at a solid surface involve the formation of bonds between reactant molecules and atoms of the surface of the catalyst (first row transition metals utilize 3d and 4s electrons for bonding). This has the effect of increasing the concentration of the reactants at the catalyst surface and also weakening of the bonds in the reacting molecules (the activation energy is lowering). Also because the transition metal ions can change their oxidation states, they become more effective as catalysts.

Other properties

As implied by the name, all transition metals are metals and conductors of electricity.

In general transition metals possess a high density and high melting points and boiling points. These properties are due to metallic bonding by delocalized d electrons, leading to cohesion which increases with the number of shared electrons. However the group 12 metals have much lower melting and boiling points since their full d subshells prevent d–d bonding. In fact mercury has a melting point of −39 °C and is a liquid at room temperature.

Many transition metals can be bound to a variety of ligands.^[16]

See also

• Inner transition element, a name given to any member of the f-block
• Ligand field theory a development of crystal field theory taking covalency into account
• Post-transition metal

References

1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "transition element".
2. ^ ^{a b c} Jensen, William B. (2003). "The Place of Zinc, Cadmium, and Mercury in the Periodic Table". Journal of Chemical Education 80 (8): 952–961. Bibcode 2003JChEd..80..952J. doi:10.1021/ed080p952. http://www.uv.es/~borrasj/ingenieria_web/temas/tema_1/lecturas_comp/p952.pdf. 
3. ^ C. R. Bury (1921). "Langmuir's theory of the arrangement of electrons in atoms and molecules". J. Amer. Chem. Soc. 43: 1602–1609. doi:10.1021/ja01440a023. 
4. ^ R. H. Petrucci et al., “General Chemistry”, 8th edn, Prentice-Hall 2002, pp. 49–50, 951
5. ^ G. L. Miessler and D. A. Tarr “Inorganic Chemistry”, 2nd edn, Prentice-Hall 1999, p. 16
6. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0080379419. 
7. ^ F.A.Cotton and G.Wilkinson, “Inorganic Chemistry”, 5th edn, Wiley 1988, pp. 626–7
8. ^ C. E. Housecroft and A. G. Sharpe, “Inorganic Chemistry”, 2nd edn, Pearson Prentice-Hall 2005, p. 741
9. ^ Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999). Advanced Inorganic Chemistry (6th ed.). New York: Wiley.
10. ^ Xuefang Wang; Lester Andrews; Sebastian Riedel; and Martin Kaupp (2007). "Mercury Is a Transition Metal: The First Experimental Evidence for HgF₄.". Angew. Chem. Int. Ed. 46 (44): 8371–8375. doi:10.1002/anie.200703710. PMID 17899620. 
11. ^ William B. Jensen (2008). "Is Mercury Now a Transition Element?". J. Chem. Educ. 85: 1182–1183. Bibcode 2008JChEd..85.1182J. doi:10.1021/ed085p1182. 
12. ^ Matsumoto, Paul S (2005). "Trends in Ionization Energy of Transition-Metal Elements". Journal of Chemical Education 82: 1660. Bibcode 2005JChEd..82.1660M. doi:10.1021/ed082p1660. http://www.jce.divched.org/Journal/Issues/2005/Nov/abs1660.html. 
13. ^ Orgel, L.E. (1966). An Introduction to Transition-Metal Chemistry, Ligand field theory (2nd. ed.). London: Methuen. 
14. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0080379419.  p. 240
15. ^ Figgis, B.N.; Lewis, J. (1960). Lewis, J. and Wilkins, R.G.. ed. The Magnetochemistry of Complex Compounds. Modern Coordination Chemistry. New York: Wiley Interscience. pp. 400–454. 
16. ^ C.Michael Hogan. 2010. Heavy metal. Encyclopedia of Earth. National Council for Science and the Environment. E. Monosson and C. Cleveland (eds.) Washington DC.