Transition metal: Difference between revisions
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*''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 diagram]]s. |
*''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 diagram]]s. |
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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<sup>−1</sup>cm<sup>−1</sup> (where [[Molar concentration|M]] = mol dm<sup>−3</sup>).<ref>{{cite book|last=Orgel|first=L.E.|title=An Introduction to Transition-Metal Chemistry, Ligand field theory|publisher=Methuen|location=London|year=1966|edition=2nd.}}</ref> Some ''d''-''d'' transitions are [[spin forbidden]]. An example occurs in octahedral, high-spin complexes of [[manganese]](II), |
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 butt range 5-500 M<sup>−1</sup>cm<sup>−1</sup> (where [[Molar concentration|M]] = mol dm<sup>−3</sup>).<ref>{{cite book|last=Orgel|first=L.E.|title=An Introduction to Transition-Metal Chemistry, Ligand field theory|publisher=Methuen|location=London|year=1966|edition=2nd.}}</ref> Some ''d''-''d'' transitions are [[spin forbidden]]. An example occurs in octahedral, high-spin complexes of [[manganese]](II), |
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which has a ''d''<sup>5</sup> 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 [[Tanabe-Sugano diagram#Manganese(II) Hexahydrate|spectrum of {{chem|[Mn(H|2|O)|6|]|2+}}]] shows a maximum molar absorptivity of about 0.04 M<sup>−1</sup>cm<sup>−1</sup> in the [[visible spectrum]]. |
which has a ''d''<sup>5</sup> 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 [[Tanabe-Sugano diagram#Manganese(II) Hexahydrate|spectrum of {{chem|[Mn(H|2|O)|6|]|2+}}]] shows a maximum molar absorptivity of about 0.04 M<sup>−1</sup>cm<sup>−1</sup> in the [[visible spectrum]]. |
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In chemistry, the term transition metal (or transition element) has two possible meanings:
- The IUPAC definition[1] 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".
- Most scientists describe a "transition metal" as any element in the d-block of the periodic table (all are metals), which includes groups 3 to 12 on the periodic table. In actual practice, the f-block lanthanide and actinide series are also considered transition metals and are called "inner transition metals".
Jensen[2] reviews 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 |
The typical electronic structure of transition metal atoms can be written as [ ]ns2(n-1)dm, 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 3–12 are now generally recognized as transition metals, although the elements La-Lu and Ac-Lr and Group 12 attract different definitions from different authors.
- 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”.
- Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides.[6][7][8] This classification is based on similarities 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.
- 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 as transition metals. Eric Scerri has proposed placing Lu and Lr in group 3 on the grounds of continuous sequences of atomic numbers in an expanded or long-form periodic table.[9]
Zinc, cadmium, and mercury are sometimes excluded from the transition metals[2] as they have the electronic configuration [ ]d10s2, with no incomplete d shell.[10] In the oxidation state +2 the ions have the electronic configuration [ ] d10. However, these elements can exist in many other oxidation states, including the +1 oxidation state, as in the diatomic ion Hg2+
2. The group 12 elements Zn, Cd and Hg may be classed as post-transition metals in this case, because of the formation of a covalent bond between the two atoms of the dimer. 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.
The recent synthesis of mercury(IV) fluoride (HgF
4) reinforces that these elements should be considered a transition metal,[11] but some authors still consider this compound to be exceptional.[12]
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.[13]
- 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
3)
2 (red); K
2Cr
2O
7 (orange); K
2CrO
4 (yellow); NiCl
2 (turquoise); CuSO
4 (blue); KMnO
4 (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 butt range 5-500 M−1cm−1 (where M = mol dm−3).[14] 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(H
2O)
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 [Ga
2Cl
6]2−
, which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom.[15] 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 [MnO
4]−
and OsO
4 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
Transition metal compounds are paramagnetic when they have one or more unpaired d electrons.[16] 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 [FeCl
4]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 (nanomaterial-based catalysts) 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 lowered). Also because the other people are transition metals, ions can also 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 −38.83 °C (−37.89 °F) and is a liquid at room temperature.
Many transition metals can be bound to a variety of ligands.[17]
See also
- Inner transition element, a name given to any member of the f-block
- Main group element, an element other than a transition metal.
- Ligand field theory a development of crystal field theory taking covalency into account.
- Crystal field theory a model that describes the breaking of degeneracies of electronic orbital states.
- Poor metal or Post-transition metal, a metallic element to the right of the transition metals in the periodic table.
References
- ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "transition element". doi:10.1351/goldbook.T06456
- ^ a b c Jensen, William B. (2003). "The Place of Zinc, Cadmium, and Mercury in the Periodic Table" (PDF). Journal of Chemical Education. 80 (8): 952–961. Bibcode:2003JChEd..80..952J. doi:10.1021/ed080p952.
- ^ Bury, C. R. (1921). "Langmuir's theory of the arrangement of electrons in atoms and molecules". J. Amer. Chem. Soc. 43 (7): 1602–1609. doi:10.1021/ja01440a023.
- ^ Petrucci, R. H. et al. (2002), “General Chemistry”, 8th edn, Prentice-Hall, pp. 49–50, 951
- ^ Miessler, G. L. and Tarr, D. A. (1999) “Inorganic Chemistry”, 2nd edn, Prentice-Hall, p. 16
- ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
- ^ Cotton, F.A. and Wilkinson, G. (1988) “Inorganic Chemistry”, 5th edn, Wiley , pp. 626–7
- ^ Housecroft, C. E. and Sharpe, A. G. (2005) “Inorganic Chemistry”, 2nd edn, Pearson Prentice-Hall, p. 741
- ^ Scerri,E.R. (2011), "A Very Short Introduction to the Periodic Table, Oxford University Press.
- ^ Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999). Advanced Inorganic Chemistry (6th ed.). New York: Wiley, ISBN 0471199575.
- ^ Wang, Xuefang; Andrews, Lester; Riedel, Sebastian; Kaupp, Martin (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.
- ^ Jensen, William B. (2008). "Is Mercury Now a Transition Element?". J. Chem. Educ. 85 (9): 1182–1183. Bibcode:2008JChEd..85.1182J. doi:10.1021/ed085p1182.
- ^ Matsumoto, Paul S (2005). "Trends in Ionization Energy of Transition-Metal Elements". Journal of Chemical Education. 82 (11): 1660. Bibcode:2005JChEd..82.1660M. doi:10.1021/ed082p1660.
- ^ Orgel, L.E. (1966). An Introduction to Transition-Metal Chemistry, Ligand field theory (2nd. ed.). London: Methuen.
- ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. p. 240
- ^ Figgis, B.N. (1960). Lewis, J. and Wilkins, R.G. (ed.). The Magnetochemistry of Complex Compounds. Modern Coordination Chemistry. New York: Wiley Interscience. pp. 400–454.
{{cite book}}: Unknown parameter|coauthors=ignored (|author=suggested) (help)CS1 maint: multiple names: editors list (link) - ^ Hogan, C. Michael (2010). "Heavy metal" in Encyclopedia of Earth. National Council for Science and the Environment. E. Monosson and C. Cleveland (eds.) Washington DC.
