Transition metal: Difference between revisions
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{{Periodic table (transition metals)}} |
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In chemistry, the term '''transition metal''' (or '''transition element''') has two possible meanings: |
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*The [[IUPAC]] definition<ref>{{GoldBookRef|file=T06456|title=transition element}}</ref> defines a transition metal as "an [[chemical element|element]] whose atom has an incomplete d sub-shell, or which can give rise to [[cation]]s with an incomplete d sub-shell". |
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*Most scientists describe a "transition metal" as any element in the [[d-block]] of the [[periodic table]] (all are metals), which includes groups 2 to 11 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". |
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Jensen<ref name=Jensen/> 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.<ref>{{cite journal|title=Langmuir's theory of the arrangement of electrons in atoms and molecules|journal=J. Amer. Chem. Soc. |volume=43|doi=10.1021/ja01440a023|pages=1602–1609|year=1921|author=C. R. Bury|issue=7}}</ref> |
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==Classification== |
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In the ''d''-block the atoms of the elements have between 1 and 10 ''d'' electrons. |
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{| border="1" align="center" cellpadding="4" cellspacing="0" style="margin: 1em 1em 1em 1em; background: #ffe; border: 1px #aaa solid; border-collapse: collapse; font-size: 95%;" |
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|- |
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![[Group (periodic table)|Group]] |
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!style="background-color:red" |[[Group 3 element|3]] |
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![[Group 4 element|4]] |
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![[Group 5 element|5]] |
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![[Group 6 element|6]] |
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![[Group 7 element|7]] |
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![[Group 8 element|8]] |
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![[Group 9 element|9]] |
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![[Group 10 element|10]] |
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![[Group 11 element|11]] |
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![[Group 12 element|12]] |
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|- |
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![[Period 4 element|Period 4]] |
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|[[Scandium|Sc]] 21 |
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|[[Titanium|Ti]] 22 |
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|[[Vanadium|V]] 23 |
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|[[Chromium|Cr]] 24 |
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|[[Manganese|Mn]] 25 |
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|[[Iron|Fe]] 26 |
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|[[Cobalt|Co]] 27 |
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|[[Nickel|Ni]] 28 |
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|[[Copper|Cu]] 29 |
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|[[Zinc|Zn]] 30 |
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|- |
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![[Period 5 element|Period 5]] |
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|[[Yttrium|Y]] 39 |
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|[[Zirconium|Zr]] 40 |
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|[[Niobium|Nb]] 41 |
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|[[Molybdenum|Mo]] 42 |
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|[[Technetium|Tc]] 43 |
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|[[Ruthenium|Ru]] 44 |
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|[[Rhodium|Rh]] 45 |
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|[[Palladium|Pd]] 46 |
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|[[Silver|Ag]] 47 |
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|[[Cadmium|Cd]] 48 |
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|- |
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![[Period 6 element|Period 6]] |
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|style="background-color:red" |[[Lanthanides|*]] 57–71 |
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|[[Hafnium|Hf]] 72 |
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|[[Tantalum|Ta]] 73 |
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|[[Tungsten|W]] 74 |
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|[[Rhenium|Re]] 75 |
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|[[Osmium|Os]] 76 |
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|[[Iridium|Ir]] 77 |
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|[[Platinum|Pt]] 78 |
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|[[Gold|Au]] 79 |
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|[[Mercury (element)|Hg]] 80 |
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|- |
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![[Period 7 element|Period 7]] |
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|style="background-color:red" |[[Actinides|**]] 89–103 |
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|[[Rutherfordium|Rf]] 104 |
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|[[Dubnium|Db]] 105 |
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|[[Seaborgium|Sg]] 106 |
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|[[Bohrium|Bh]] 107 |
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|[[Hassium|Hs]] 108 |
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|[[Meitnerium|Mt]] 109 |
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|[[Darmstadtium|Ds]] 110 |
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|[[Roentgenium|Rg]] 111 |
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|[[Copernicium|Cn]] 112 |
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|} |
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With a few [[Periodic table (electron configurations)|minor exceptions]], the [[electron configuration|electronic structure]] of transition metal atoms can be written as [ ]''ns''<sup>2</sup>(''n-1'')''d''<sup>m</sup>, where the inner ''d'' orbital has more energy than the [[valence electron|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 {{chem|Fe|2+}} has no ''s'' electrons: it has the electronic configuration [Ar]3d<sup>6</sup> as compared with the configuration of the atom, [Ar]4s<sup>2</sup>3d<sup>6</sup>. |
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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. |
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# 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 s<sup>2</sup>d<sup>1</sup> 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.<ref>R. H. Petrucci ''et al.'', “General Chemistry”, 8th edn, Prentice-Hall 2002, pp. 49–50, 951</ref><ref>G. L. Miessler and D. A. Tarr “Inorganic Chemistry”, 2nd edn, Prentice-Hall 1999, p. 16</ref> The two series together are classified as [[f-block]] elements, or (in older sources) as “inner transition elements”. |
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# Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides.<ref>{{Greenwood&Earnshaw}}</ref><ref>F.A.Cotton and G.Wilkinson, “Inorganic Chemistry”, 5th edn, Wiley 1988, pp. 626–7</ref><ref>C. E. Housecroft and A. G. Sharpe, “Inorganic Chemistry”, 2nd edn, Pearson Prentice-Hall 2005, p. 741</ref> 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. |
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# A third classification defines the f-block elements as La-Yb and Ac-No, while placing Lu and Lr in Group 3.<ref name=Jensen/> 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 [ ]s<sup>2</sup>f<sup>14</sup>d<sup>1</sup> configuration. However La and Ac are exceptions to the Aufbau principle with electron configuration [ ]s<sup>2</sup>d<sup>1</sup> (not [ ]s<sup>2</sup>f<sup>1</sup> 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. |
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[[Zinc]], [[cadmium]], and [[mercury (element)|mercury]] are sometimes incorrectly classified as non-transition metals<ref name=Jensen>{{cite journal|journal = Journal of Chemical Education|volume = 80|issue = 8|year = 2003|url = http://www.uv.es/~borrasj/ingenieria_web/temas/tema_1/lecturas_comp/p952.pdf|title = The Place of Zinc, Cadmium, and Mercury in the Periodic Table|first = William B.|last = Jensen|pages=952–961|doi = 10.1021/ed080p952|bibcode = 2003JChEd..80..952J }}</ref> as they have the [[electronic configuration]] [ ]''d''<sup>10</sup>s<sup>2</sup>, with no incomplete ''d'' shell.<ref>Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999). ''Advanced Inorganic Chemistry'' (6th ed.). New York: Wiley.</ref> In the [[oxidation state]] +2 the ions have the electronic configuration [ ] d<sup>10</sup>. However, these elements can exist in many other oxidation states, including the +1 oxidation state, as in the diatomic ion {{chem|Hg|2|2+}}. The group 12 elements Zn, Cd and Hg may be classed as [[post-transition metal]]s 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 theory#Crystal field stabilization energy|crystal field stabilization energy]] of first-row transition elements, it is convenient to also include the elements [[calcium]] and zinc, as both {{chem|Ca|2+}} and {{chem|Zn|2+}} 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. |
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The recent synthesis of [[mercury(IV) fluoride]] ({{chem|HgF|4}}) reinforces that these elements should now always be considered a transition metal.<ref>{{cite journal |author=Xuefang Wang |coauthors=Lester Andrews; Sebastian Riedel; and Martin Kaupp |title=Mercury Is a Transition Metal: The First Experimental Evidence for HgF<sub>4</sub> |journal=Angew. Chem. Int. Ed. |year=2007 |volume=46 |issue=44 |pages=8371–8375 |doi=10.1002/anie.200703710 |pmid=17899620}}</ref><ref>{{cite journal |title=Is Mercury Now a Transition Element? |author=William B. Jensen |journal=J. Chem. Educ. |year=2008 |volume=85 |pages=1182–1183 |doi=10.1021/ed085p1182|bibcode = 2008JChEd..85.1182J |issue=9 }}</ref> |
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==Characteristic properties== |
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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 |
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* the formation of compounds whose colour is due to ''d'' – ''d'' electronic transitions |
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* the formation of compounds in many oxidation states, due to the relatively low reactivity of unpaired ''d'' electrons.<ref>{{cite journal|url=http://www.jce.divched.org/Journal/Issues/2005/Nov/abs1660.html|author=Matsumoto, Paul S|journal=Journal of Chemical Education|year=2005|volume=82|page=1660|title=Trends in Ionization Energy of Transition-Metal Elements|doi=10.1021/ed082p1660 |bibcode = 2005JChEd..82.1660M|issue=11 }}</ref> |
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* 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]]) |
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===Coloured compounds=== |
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[[Image:Coloured-transition-metal-solutions.jpg|thumb|right|250px|From left to right, aqueous solutions of: {{chem|link=cobalt(II) nitrate|Co(NO|3|)|2}} (red); {{chem|link=potassium dichromate|K|2|Cr|2|O|7}} (orange); {{chem|link=potassium chromate|K|2|CrO|4}} (yellow); {{chem|link=nickel(II) chloride|NiCl|2}} (turquoise); {{chem|link=copper(II) sulfate|CuSO|4}} (blue); {{chem|link=potassium permanganate|KMnO|4}} (purple).]] |
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Colour in transition-series metal compounds is generally due to electronic transitions of two principal types. |
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*[[Charge transfer complex|charge transfer]] transitions. An electron may jump from a predominantly [[ligand]] [[Atomic orbital|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]], HgI<sub>2</sub>, is red because of a LMCT transition.<!--As this example shows, charge transfer transitions are not restricted to transition metals.<ref>T.M. Dunn in {{cite book|last=Lewis|first=J.|coauthors=Wilkins,R.G.|title=Modern Coordination Chemistry|publisher=Interscience|location=New York|year=1960|pages= Chapter 4, Section 4, "Charge Transfer Spectra", pp. 268–273}}</ref>--> |
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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. |
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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. |
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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), |
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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]]. |
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===Oxidation states=== |
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A characteristic of transition metals is that they exhibit two or more [[oxidation state]]s, usually differing by one. For example, compounds of [[vanadium]] are known in all oxidation states between −1, such as {{chem|[V(CO)|6|]|-}}, and +5, such as {{chem|VO|4|3-}}. |
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[[Main group element]]s 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 {{chem|[Ga|2|Cl|6|]|2-}}, which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom.<ref>{{Greenwood&Earnshaw}} p. 240</ref> 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. |
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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 {{chem|[MnO|4|]|-}} and {{chem|OsO|4}} the elements achieve a stable [[octet rule|octet]] by forming four [[covalent]] bonds. |
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The lowest oxidation states are exhibited in such compounds as {{chem|Cr(CO)|6}} (oxidation state zero) and {{chem|[Fe(CO)|4|]|2-}} (oxidation state −2) in which the [[18-electron rule]] is obeyed. These complexes are also covalent. |
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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. |
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<!-- [[Image:Transition metal oxidation states 3.png|center|frame|This table shows some of the oxidation states found in compounds of the transition-metal elements.<br> A solid circle represents a common oxidation state, and a ring represents a less common oxidation state.]] --> |
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===Magnetism=== |
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{{main|magnetochemistry}} |
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Transition metal compounds are [[paramagnetic]] when they have one or more unpaired ''d'' electrons.<ref>{{cite book|last=Figgis|first=B.N.|coauthors=Lewis, J.|title=The Magnetochemistry of Complex Compounds|editor=Lewis, J. and Wilkins, R.G.|publisher=Wiley Interscience|location=New York|year=1960|series=Modern Coordination Chemistry|pages=400–454}}</ref> 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 {{chem|[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, ''d''<sup>6</sup> and square-planar ''d<sup>8</sup> complexes''. In these cases, [[crystal field]] splitting is such that all the electrons are paired up. |
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[[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. [[Antiferromagnetic interaction|Anti-ferromagnetism]] is another example of a magnetic property arising from a particular alignment of individual spins in the solid state. |
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=== Catalytic properties === |
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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 [[Hydrogenation|catalytic hydrogenation]]) are some of the examples. Catalysts at a solid surface ([[Nanomaterial-based_catalyst|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 transition metal ions can change their oxidation states, they become more effective as [[Catalysis|catalysts]]. |
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===Other properties=== |
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As implied by the name, all transition metals are [[metal]]s and conductors of electricity. |
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In general, transition metals possess a high [[density]] and high [[melting point]]s and [[boiling point]]s. These properties are due to [[metallic bond]]ing by delocalized d electrons, leading to [[Cohesion (chemistry)|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. |
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Many transition metals can be bound to a variety of [[ligands]].<ref>C.Michael Hogan. 2010. [http://www.eoearth.org/article/Heavy_metal?topic=49498 ''Heavy metal''. Encyclopedia of Earth. National Council for Science and the Environment.] E. Monosson and C. Cleveland (eds.) Washington DC.</ref> |
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==See also== |
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*[[Inner transition element]], a name given to any member of the [[f-block]] |
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*[[Main group element]], an element other than a transition metal. |
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*[[Ligand field theory]] a development of crystal field theory taking covalency into account. |
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*[[Crystal field theory]] a model that describes the breaking of [[Degenerate energy levels|degeneracies]] of electronic orbital states. |
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*[[Post-transition metal]], a metallic element to the right of the transition metals in the periodic table. |
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==References== |
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{{reflist|35em}} |
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{{PeriodicTablesFooter}} |
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{{Compact periodic table}} |
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[[Category:Periodic table]] |
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[[Category:Transition metals|*]] |
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[[Category:Chemical element groups]] |
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