View source for Transition metal

{{short description|Series of chemical elements}}
{{pp-semi-indef}}
{{Periodic table (micro)|mark=Sc,Ti,V,Cr,Mn,Fe,Co,Ni,Cu,Zn,Y,Zr,Nb,Mo,Tc,Ru,Rh,Pd,Ag,Cd,Lu,Hf,Ta,W,Re,Os,Ir,Pt,Au,Hg,Lr,Rf,Db,Sg,Bh,Hs,Mt,Ds,Rg,Cn|title=Transition metals in the [[periodic table]]}}
{{Sidebar periodic table|expanded=metalicity}}
In chemistry, a '''transition metal''' (or '''transition element''') is a [[chemical element]] in the [[d-block]] of the [[periodic table]] (groups 3 to 12), though the elements of [[Group 12 element|group 12]] (and less often [[Group 3 element|group 3]]) are sometimes excluded. The [[lanthanide]] and [[actinide]] elements (the [[f-block]]) are called '''inner transition metals''' and are sometimes considered to be transition metals as well.

They are lustrous metals with good [[Electrical conductivity|electrical]] and [[Thermal conductivity and resistivity|thermal]] conductivity. Most (with the exception of [[Group 11 element|group 11]] and group 12) are hard and strong, and have high melting and boiling temperatures. They form compounds in any of two or more different [[oxidation state]]s and bind to a variety of [[ligand]]s to form [[coordination complex]]es that are often coloured. They form many useful [[alloy]]s and are often employed as [[catalyst]]s in elemental form or in compounds such as coordination complexes and [[Surface properties of transition metal oxides|oxides]]. Most are strongly [[paramagnetism|paramagnetic]] because of their [[unpaired electron|unpaired d electrons]], as are many of their compounds. All of the elements that are [[ferromagnetism|ferromagnetic]] near room temperature are transition metals ([[iron]], [[cobalt]] and [[nickel]]) or inner transition metals ([[gadolinium]]).

English chemist [[Charles Rugeley Bury]] (1890–1968) first used the word ''transition'' in this context in 1921, when he referred to a ''transition series of elements'' during the change of an inner layer of electrons (for example ''n''&nbsp;=&nbsp;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 name=Jensen/><ref>{{cite journal |title=Langmuir's theory of the arrangement of electrons in atoms and molecules |journal=J. Am. Chem. Soc. |volume=43 |doi=10.1021/ja01440a023 |pages=1602–1609 |year=1921 |author=Bury, C. R. |issue=7|bibcode=1921JAChS..43.1602B |url=https://zenodo.org/record/1428812 }}</ref><ref>[http://www.encyclopedia.com/doc/1G2-2830905044.html Bury, Charles Rugeley]. Encyclopedia.com Complete dictionary of scientific biography (2008).</ref> These elements are now known as the d-block.
[[File:First row of transition metals.jpg|thumb|The first row of transition metals, in order|300px]]

==Definition and classification==
The 2011 [[International Union of Pure and Applied Chemistry name|IUPAC]] ''Principles of Chemical Nomenclature'' describe a "transition metal" as any element in groups 3 to 12 on the [[periodic table]].<ref>{{cite book |editor-last=Leigh |editor-first=G. J. |date=2011 |title=Principles of Chemical Nomenclature |url=https://iupac.org/wp-content/uploads/2021/12/Principles_Leigh2011-compressed.pdf |location= |publisher=The Royal Society of Chemistry |page=9 |isbn=978-1-84973-007-5}}</ref> This corresponds exactly to the [[d-block]] elements, and many scientists use this definition.<ref>{{cite book |last1 = Petrucci |first1 = Ralph H. |last2 = Harwood |first2 = William S. |last3 = Herring |first3 = F. Geoffrey |date=2002 |title = General chemistry: principles and modern applications |url = https://archive.org/details/generalchemistry00hill |url-access = registration |edition=8th |location=Upper Saddle River, N.J |publisher=Prentice Hall |isbn = 978-0-13-014329-7 |lccn=2001032331 |oclc=46872308 |pages=[https://archive.org/details/generalchemistry00hill/page/341 341–342]}}</ref><ref>Housecroft, C.&nbsp;E. and Sharpe, A.&nbsp;G. (2005) ''Inorganic Chemistry'', 2nd ed, Pearson Prentice-Hall, pp. 20–21.</ref> In actual practice, the [[f-block]] [[lanthanide]] and [[actinide]] series are called "inner transition metals". The 2005 ''[[IUPAC Red Book|Red Book]]'' allows for the group 12 elements to be excluded, but not the 2011 ''Principles''.<ref name=RedBook>{{cite book
 |editor-last1=Connelly
 |editor-first1=N.G.
 |editor-last2=Damhus
 |editor-first2=T.
 |editor-last3=Hartshorn
 |editor-first3=R.M.
 |editor-last4=Hutton
 |editor-first4=A.T.
 |year=2005
 |url=http://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf
 |title=Nomenclature of Inorganic Chemistry
 |publisher=[[Royal Society of Chemistry|RSC]]–[[International Union of Pure and Applied Chemistry|IUPAC]]
 |isbn=0-85404-438-8
}}</ref>

The IUPAC ''[[IUPAC Gold Book|Gold Book]]''<ref>{{GoldBookRef|file=T06456|title=transition element}}</ref> defines a transition metal as "an [[chemical element|element]] whose atom has a partially filled [[Electron shell|d]] sub-shell, or which can give rise to [[cation]]s with an incomplete d sub-shell", but this definition is taken from an old edition of the ''Red Book'' and is no longer present in the current edition.<ref name=RedBook/>

In the d-block, the atoms of the elements have between zero and ten d electrons.
{| class=wikitable style="margin:0.5em 1em 0.5em 0; text-align:center; background:{{element color|d-block}}; border: 1px solid {{element color|table border}}; padding:2px; font-size: 85%;"
|-
|+ style="background-color:{{element color|table title}}" | Transition metals in the d-block
|-
! style="background:{{element color|table header}}" | [[Group (periodic table)|Group]]
! style="background:{{element color|table header}}" | [[Group 3 element|3]]
! style="background:{{element color|table header}}" | [[Group 4 element|4]]
! style="background:{{element color|table header}}" | [[Group 5 element|5]]
! style="background:{{element color|table header}}" | [[Group 6 element|6]]
! style="background:{{element color|table header}}" | [[Group 7 element|7]]
! style="background:{{element color|table header}}" | [[Group 8 element|8]]
! style="background:{{element color|table header}}" | [[Group 9 element|9]]
! style="background:{{element color|table header}}" | [[Group 10 element|10]]
! style="background:{{element color|table header}}" | [[Group 11 element|11]]
! style="background:{{element color|table header}}" | [[Group 12 element|12]]
|-
!style="background:{{element color|table colheader}}"| [[Period 4 element|Period 4]]
| <sub>21</sub>[[Scandium|Sc]]
| <sub>22</sub>[[Titanium|Ti]]
| <sub>23</sub>[[Vanadium|V]]
| <sub>24</sub>[[Chromium|Cr]]
| <sub>25</sub>[[Manganese|Mn]]
| <sub>26</sub>[[Iron|Fe]]
| <sub>27</sub>[[Cobalt|Co]]
| <sub>28</sub>[[Nickel|Ni]]
| <sub>29</sub>[[Copper|Cu]]
| <sub>30</sub>[[Zinc|Zn]]
|-
!style="background:{{element color|table colheader}}"| [[Period 5 element|5]]
| <sub>39</sub>[[Yttrium|Y]]
| <sub>40</sub>[[Zirconium|Zr]]
| <sub>41</sub>[[Niobium|Nb]]
| <sub>42</sub>[[Molybdenum|Mo]]
| <sub>43</sub>[[Technetium|Tc]]
| <sub>44</sub>[[Ruthenium|Ru]]
| <sub>45</sub>[[Rhodium|Rh]]
| <sub>46</sub>[[Palladium|Pd]]
| <sub>47</sub>[[Silver|Ag]]
| <sub>48</sub>[[Cadmium|Cd]]
|-
!style="background:{{element color|table colheader}}"| [[Period 6 element|6]]
| <sub>71</sub>[[Lutetium|Lu]]
| <sub>72</sub>[[Hafnium|Hf]]
| <sub>73</sub>[[Tantalum|Ta]]
| <sub>74</sub>[[Tungsten|W]]
| <sub>75</sub>[[Rhenium|Re]]
| <sub>76</sub>[[Osmium|Os]]
| <sub>77</sub>[[Iridium|Ir]]
| <sub>78</sub>[[Platinum|Pt]]
| <sub>79</sub>[[Gold|Au]]
| <sub>80</sub>[[Mercury (element)|Hg]]
|-
!style="background:{{element color|table colheader}}"| [[Period 7 element|7]]
| <sub>103</sub>[[Lawrencium|Lr]]
| <sub>104</sub>[[Rutherfordium|Rf]]
| <sub>105</sub>[[Dubnium|Db]]
| <sub>106</sub>[[Seaborgium|Sg]]
| <sub>107</sub>[[Bohrium|Bh]]
| <sub>108</sub>[[Hassium|Hs]]
| <sub>109</sub>[[Meitnerium|Mt]]
| <sub>110</sub>[[Darmstadtium|Ds]]
| <sub>111</sub>[[Roentgenium|Rg]]
| <sub>112</sub>[[Copernicium|Cn]]
|}

Published texts and periodic tables show [[Group 3 element#Composition|variation regarding the heavier members of group 3]].<ref name="PTSS">{{cite book | last=Scerri | first=Eric R. | title=The Periodic Table: Its Story and Its Significance | publication-place=New York, NY | date=2020 | isbn=978-0-19-091436-3 | oclc=1096234740}}</ref> The common placement of [[lanthanum]] and [[actinium]] in these positions is not supported by physical, chemical, and electronic [[Scientific evidence|evidence]],<ref>{{cite book
 |author=[[Lev Landau|L. D. Landau]], [[Evgeny Lifshitz|E. M. Lifshitz]]
 |year=1958
 |title=Quantum Mechanics: Non-Relativistic Theory
 |edition=1st |volume=3
 |publisher=[[Pergamon Press]]
 |pages=256–7
}}</ref><ref name="Wittig">{{cite book |last=Wittig |first=Jörg |editor=H. J. Queisser |date=1973 |title=Festkörper Probleme: Plenary Lectures of the Divisions Semiconductor Physics, Surface Physics, Low Temperature Physics, High Polymers, Thermodynamics and Statistical Mechanics, of the German Physical Society, Münster, March 19–24, 1973 |chapter=The pressure variable in solid state physics: What about 4f-band superconductors? |series=Advances in Solid State Physics |volume=13 |location=Berlin, Heidelberg |publisher=Springer |pages=375–396 |isbn=978-3-528-08019-8 |doi=10.1007/BFb0108579}}</ref><ref>{{cite book |last=Matthias |first=B. T. |date=1969 |editor-last=Wallace |editor-first=P. R. |title=Superconductivity |publisher=Gordon and Breach |pages=225–294 <!--precise quote calling it a mistake is on pp. 247–9--> |chapter=Systematics of Super Conductivity |isbn=9780677138107 |volume=1}}</ref> which overwhelmingly favour putting [[lutetium]] and [[lawrencium]] in those places.<ref name="Jensen1982">{{cite journal |title=The Positions of Lanthanum (Actinium) and Lutetium (Lawrencium) in the Periodic Table |author=William B. Jensen |journal=J. Chem. Educ. |year=1982 |volume=59 |issue = 8|pages=634–636 |doi=10.1021/ed059p634|bibcode=1982JChEd..59..634J }}</ref><ref name="Jensen2015">{{cite journal |last1=Jensen |first1=William B. |date=2015 |title=The positions of lanthanum (actinium) and lutetium (lawrencium) in the periodic table: an update |url=https://link.springer.com/article/10.1007/s10698-015-9216-1 |journal=Foundations of Chemistry |volume=17 |issue= |pages=23–31 |doi=10.1007/s10698-015-9216-1 |s2cid=98624395 |access-date=28 January 2021 |archive-date=30 January 2021 |archive-url=https://web.archive.org/web/20210130011116/https://link.springer.com/article/10.1007/s10698-015-9216-1 |url-status=live |url-access=subscription }}</ref> Some authors prefer to leave the spaces below [[yttrium]] blank as a third option, but there is confusion on whether this format implies that group 3 contains only [[scandium]] and yttrium, or if it also contains all the lanthanides and actinides;<ref name=2021IUPAC/><ref name="Thyssen">{{cite book|last1=Thyssen|first1=P.|last2=Binnemans|first2=K.|editor1-last=Gschneidner|editor1-first= K. A. Jr.|editor2-last=Bünzli|editor2-first=J-C.G|editor3-last=Vecharsky|editor3-first=Bünzli|date=2011|chapter=Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis|title=Handbook on the Physics and Chemistry of Rare Earths|publisher=Elsevier|location=Amsterdam|volume=41|pages=1–94|isbn=978-0-444-53590-0|doi=10.1016/B978-0-444-53590-0.00001-7}}</ref><ref name="JWP">{{cite journal |author=Barber, Robert C. |author2=Karol, Paul J |author3=Nakahara, Hiromichi |author4=Vardaci, Emanuele |author5=Vogt, Erich W. |title=Discovery of the elements with atomic numbers greater than or equal to 113 (IUPAC Technical Report) |doi=10.1351/PAC-REP-10-05-01 |journal=Pure Appl. Chem. |date=2011 |volume=83 |issue=7 |page=1485|doi-access=free }}</ref><ref name="Karol">{{cite journal |last1=Karol |first1=Paul J. |last2=Barber |first2=Robert C. |last3=Sherrill |first3=Bradley M. |last4=Vardaci |first4=Emanuele |last5=Yamazaki |first5=Toshimitsu |date=22 December 2015 |title=Discovery of the elements with atomic numbers Z = 113, 115 and 117 (IUPAC Technical Report) |journal=Pure Appl. Chem. |volume=88 |issue=1–2 |pages=139–153 |doi=10.1515/pac-2015-0502|doi-access=free }}</ref><ref>{{cite journal |last1=Pyykkö |first1=Pekka |date=2019 |title=An essay on periodic tables |url=http://www.chem.helsinki.fi/~pyykko/pekka/No330b.pdf |journal=Pure and Applied Chemistry |volume=91 |issue=12 |pages=1959–1967 |doi=10.1515/pac-2019-0801 |s2cid=203944816 |access-date=27 November 2022}}</ref> additionally, it creates a 15-element-wide f-block, when [[quantum mechanics]] dictates that the f-block should only be 14 elements wide.<ref name=2021IUPAC/> The form with lutetium and lawrencium in group 3 is supported by a 1988 [[IUPAC]] report on physical, chemical, and electronic grounds,<ref name="Fluck">{{cite journal |last1=Fluck |first1=E. |year=1988 |title=New Notations in the Periodic Table |journal=[[Pure and Applied Chemistry|Pure Appl. Chem.]] |volume=60 |pages=431–436|doi=10.1351/pac198860030431 |url=https://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |access-date=24 March 2012 |issue=3 |s2cid=96704008 |url-status=live |archive-url=https://web.archive.org/web/20120325152951/https://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |archive-date=25 March 2012}}</ref> and again by a 2021 IUPAC preliminary report as it is the only form that allows simultaneous (1) preservation of the sequence of increasing atomic numbers, (2) a 14-element-wide f-block, and (3) avoidance of the split in the d-block.<ref name="2021IUPAC">{{cite journal |last1=Scerri |first1=Eric |date=18 January 2021 |title=Provisional Report on Discussions on Group 3 of the Periodic Table |url=https://iupac.org/wp-content/uploads/2021/04/ChemInt_Jan2021_PP.pdf |journal=Chemistry International |volume=43 |issue=1 |pages=31–34 |doi=10.1515/ci-2021-0115 |s2cid=231694898 |access-date=9 April 2021 |archive-date=13 April 2021 |archive-url=https://web.archive.org/web/20210413150110/https://iupac.org/wp-content/uploads/2021/04/ChemInt_Jan2021_PP.pdf |url-status=live }}</ref> Argumentation can still be found in the contemporary literature purporting to defend the form with lanthanum and actinium in group 3, but many authors consider it to be logically inconsistent (a particular point of contention being the differing treatment of [[actinium]] and [[thorium]], which both can use 5f as a [[Valence (chemistry)|valence]] [[Orbital hybridisation|orbital]] but have no 5f occupancy as single atoms);<ref name=Jensen2015/><ref name=Scerri2009>{{cite journal |last1=Scerri |first1=Eric |date=2009 |title=Which Elements Belong in Group 3? |url=https://pubs.acs.org/doi/pdf/10.1021/ed086p1188 |journal=Journal of Chemical Education |volume=86 |issue=10 |pages=1188 |doi=10.1021/ed086p1188 |bibcode=2009JChEd..86.1188S |access-date=1 January 2023|url-access=subscription }}</ref><ref name=Chemey>{{cite journal |last1=Chemey |first1=Alexander T. |last2=Albrecht-Schmitt |first2=Thomas E. |date=2019 |title=Evolution of the periodic table through the synthesis of new elements |url= |journal=Radiochimica Acta |volume=107 |issue=9–11 |pages=771–801 |doi=10.1515/ract-2018-3082 |s2cid=104470619 |access-date=}}</ref> the majority of investigators considering the problem agree with the updated form with lutetium and lawrencium.<ref name=Jensen2015/>

The group 12 elements [[zinc]], [[cadmium]], and [[mercury (element)|mercury]] are sometimes excluded from the 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> This is because they have the [[electronic configuration]] [&nbsp;]d<sup>10</sup>s<sup>2</sup>, where the d shell is complete,<ref>Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999). ''Advanced Inorganic Chemistry'' (6th ed.). New York: Wiley, {{ISBN|978-0-471-19957-1}}.</ref> and they still have a complete d shell in all their known [[oxidation state]]s. The group 12 elements Zn, Cd and Hg may therefore, under certain criteria, be classed as [[post-transition metal]]s 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 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. Moreover, Zn, Cd, and Hg can use their d orbitals for [[Bonding (chemistry)|bonding]] even though they are not known in oxidation states that would formally require breaking open the d-subshell, which sets them apart from the p-block elements.<ref>{{cite journal |last1=Tossell |first1=J.A. |date=1 November 1977 |title=Theoretical studies of valence orbital binding energies in solid zinc sulfide, zinc oxide, and zinc fluoride |journal=Inorganic Chemistry |volume=16 |issue=11 |pages=2944–2949 |doi=10.1021/ic50177a056}}</ref><ref>{{cite journal |last1=Farberovich |first1=O. V. |last2=Kurganskii |first2=S. I. |last3=Domashevskaya |first3=E. P. |date=1980 |title=Problems of the OPW Method. II. Calculation of the Band Structure of ZnS and CdS |url= |journal=Physica Status Solidi B |volume=97 |issue=2 |pages=631–640 |doi=10.1002/pssb.2220970230 |bibcode=1980PSSBR..97..631F |access-date=}}</ref><ref>{{cite journal |last1=Singh |first1=Prabhakar P. |date=1994 |title=Relativistic effects in mercury: Atom, clusters, and bulk |url= |journal=Physical Review B |volume=49 |issue=7 |pages=4954–4958 |doi=10.1103/PhysRevB.49.4954 |pmid=10011429 |bibcode=1994PhRvB..49.4954S |access-date=}}</ref>

The 2007 (though disputed and so far not reproduced independently) synthesis of [[mercury(IV) fluoride]] ({{chem|HgF|4}}) has been taken by some to reinforce the view that the group 12 elements should be considered transition metals,<ref>{{cite journal |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 |last1=Wang |first1=Xuefang |last2=Andrews |first2=Lester |last3=Riedel |first3=Sebastian |last4=Kaupp |first4=Martin |bibcode=2007ACIE...46.8371W }}</ref> but some authors still consider this compound to be exceptional.<ref>{{cite journal |title=Is Mercury Now a Transition Element? |author=Jensen, William B.  |journal=J. Chem. Educ. |year=2008 |volume=85 |pages=1182–1183 |doi=10.1021/ed085p1182|bibcode = 2008JChEd..85.1182J |issue=9 |doi-access=free }}</ref> [[Copernicium]] is expected to be able to use its d electrons for chemistry as its 6d [[Electron shell|subshell]] is destabilised by strong [[relativistic quantum chemistry|relativistic effects]] due to its very high atomic number, and as such is expected to have transition-metal-like behaviour and show higher oxidation states than +2 (which are not definitely known for the lighter group 12 elements). Even in bare dications, Cn<sup>2+</sup> is predicted to be 6d<sup>8</sup>7s<sup>2</sup>, unlike Hg<sup>2+</sup> which is 5d<sup>10</sup>6s<sup>0</sup>.

Although [[meitnerium]], [[darmstadtium]], and [[roentgenium]] are within the d-block and are expected to behave as transition metals analogous to their lighter [[Congener (chemistry)|congeners]] [[iridium]], [[platinum]], and [[gold]], this has not yet been experimentally confirmed. Whether [[copernicium]] behaves more like [[mercury (element)|mercury]] or has properties more similar to those of the [[noble gas]] [[radon]] is not clear. Relative inertness of Cn would come from the relativistically expanded 7s–7p<sub>1/2</sub> energy gap, which is already adumbrated in the 6s–6p<sub>1/2</sub> gap for Hg, weakening metallic bonding and causing its well-known low melting and boiling points.

Transition metals with lower or higher group numbers are described as 'earlier' or 'later', respectively.  When described in a two-way classification scheme, early transition metals are on the left side of the d-block from group 3 to group 7. Late transition metals are on the right side of the d-block, from group 8 to 11 (or 12, if they are counted as transition metals).  In an alternative three-way scheme, groups 3, 4, and 5 are classified as early transition metals, 6, 7, and 8 are classified as middle transition metals, and 9, 10, and 11 (and sometimes group 12) are classified as late transition metals. 

The heavy group 2 elements [[calcium]], [[strontium]], and [[barium]] do not have filled d-orbitals as single atoms, but are known to have d-orbital bonding participation in some [[Chemical compound|compounds]], and for that reason have been called "honorary" transition metals.<ref>{{cite journal |last1=Fernández |first1=Israel |last2=Holzmann |first2=Nicole |last3=Frenking |first3=Gernot |date=2020 |title=The Valence Orbitals of the Alkaline-Earth Atoms |journal=Chemistry: A European Journal |volume=26 |issue=62 |pages=14194–14210 |doi=10.1002/chem.202002986 |pmid=32666598 |pmc=7702052 |bibcode=2020ChEuJ..2614194F |s2cid=220529532 |doi-access=free }}</ref> The same is likely true of [[radium]].<ref>{{cite journal |last1=Pyykkö |first1=Pekka |last2=Desclaux |first2=Jean-Paul |date=1979 |title=Relativity and the Periodic System of Elements |journal=Accounts of Chemical Research |volume=12 |issue=8 |pages=276–281 |doi=10.1021/ar50140a002 |pmid= |pmc= |s2cid= |doi-access= }}</ref>

The f-block elements La–Yb and Ac–No have chemical activity of the (n−1)d shell, but importantly also have chemical activity of the (n−2)f shell that is absent in d-block elements. Hence they are often treated separately as inner transition elements.

==Electronic configuration==
{{Main|Electron configuration}}
The general electronic configuration of the d-block atoms is [noble gas](''n''&nbsp;−&nbsp;1)d<sup>0–10</sup>''n''s<sup>0–2</sup>''n''p<sup>0–1</sup>. Here "[noble gas]" is the electronic configuration of the last [[noble gas]] preceding the atom in question, and ''n'' is the highest [[principal quantum number]] of an occupied orbital in that atom. For example, Ti (''Z''&nbsp;=&nbsp;22) is in period 4 so that ''n'' = 4, the first 18 electrons have the same configuration of Ar at the end of period 3, and the overall configuration is [Ar]3d<sup>2</sup>4s<sup>2</sup>. The period 6 and 7 transition metals also add core (''n''&nbsp;−&nbsp;2)f<sup>14</sup> electrons, which are omitted from the tables below. The p orbitals are almost never filled in free atoms (the one exception being lawrencium due to relativistic effects that become important at such high ''Z''), but they can contribute to the chemical bonding in transition metal compounds.

The [[Aufbau principle#Madelung energy ordering rule|Madelung rule]] predicts that the inner d orbital is filled after the [[valence electron|valence-shell]] s orbital. The typical [[electron configuration|electronic structure]] of transition metal atoms is then written as [noble gas]''n''s<sup>2</sup>(''n''&nbsp;−&nbsp;1)d<sup>''m''</sup>. This rule is approximate, but holds for most of the transition metals. Even when it fails for the neutral ground state, it accurately describes a low-lying excited state.

The d subshell is the next-to-last subshell and is denoted as (''n'' − 1)d subshell. The number of s electrons in the outermost s subshell is generally one or two except [[palladium]] (Pd), with no electron in that s sub shell in its ground state. The s subshell in the valence shell is represented as the ''n''s subshell, e.g. 4s. In the periodic table, the transition metals are present in ten groups (3 to 12).

The elements in group 3 have an ''n''s<sup>2</sup>(''n''&nbsp;−&nbsp;1)d<sup>1</sup> configuration, except for [[lawrencium]] (Lr): its 7s<sup>2</sup>7p<sup>1</sup> configuration exceptionally does not fill the 6d orbitals at all. The first transition series is present in the 4th period, and starts after Ca (''Z''&nbsp;=&nbsp;20) of group 2 with the configuration [Ar]4s<sup>2</sup>, or [[scandium]] (Sc), the first element of group 3 with atomic number ''Z''&nbsp;=&nbsp;21 and configuration [Ar]4s<sup>2</sup>3d<sup>1</sup>, depending on the definition used. As we move from left to right, electrons are added to the same d subshell till it is complete. Since the electrons added fill the (''n'' − 1)d orbitals, the properties of the d-block elements are quite different from those of s and p block elements in which the filling occurs either in s or in p orbitals of the valence shell.
The electronic configuration of the individual elements present in all the d-block series are given below:<ref name=Miessler>Miessler, G. L. and Tarr, D. A. (1999) ''Inorganic Chemistry'', 2nd edn, Prentice-Hall, p. 38-39 {{ISBN|978-0-13-841891-5}}</ref>

{| class="wikitable"
|+ First (3d) d-block Series (Sc–Zn)
|-
! Group 
| 3 || 4 || 5 || 6 || 7 || 8 || 9 || 10 || 11 || 12
|-
! Atomic number
| 21 || 22 || 23 || 24 || 25 || 26 || 27 || 28 || 29 || 30
|-
! Element 
| Sc|| Ti || V || Cr|| Mn|| Fe|| Co|| Ni || Cu || Zn
|-
! Electron<br/>configuration 
| 3d<sup>1</sup>4s<sup>2</sup> || 3d<sup>2</sup>4s<sup>2</sup> || 3d<sup>3</sup>4s<sup>2</sup> || 3d<sup>5</sup>4s<sup>1</sup> || 3d<sup>5</sup>4s<sup>2</sup> || 3d<sup>6</sup>4s<sup>2</sup> || 3d<sup>7</sup>4s<sup>2</sup>|| 3d<sup>8</sup>4s<sup>2</sup>|| 3d<sup>10</sup>4s<sup>1</sup> || 3d<sup>10</sup>4s<sup>2</sup>
|}

{| class="wikitable"
|+ Second (4d) d-block Series (Y–Cd)
|-
! Atomic number 
| 39 || 40 || 41 || 42 || 43 || 44 || 45 || 46 || 47 || 48
|-
! Element 
| Y || Zr || Nb || Mo || Tc || Ru || Rh || Pd || Ag || Cd
|-
! Electron<br/>configuration 
| 4d<sup>1</sup>5s<sup>2</sup> || 4d<sup>2</sup>5s<sup>2</sup> || 4d<sup>4</sup>5s<sup>1</sup> || 4d<sup>5</sup>5s<sup>1</sup> || 4d<sup>5</sup>5s<sup>2</sup> || 4d<sup>7</sup>5s<sup>1</sup> || 4d<sup>8</sup>5s<sup>1</sup> || 4d<sup>10</sup>5s<sup>0</sup> || 4d<sup>10</sup>5s<sup>1</sup> || 4d<sup>10</sup>5s<sup>2</sup>
|}

{| class="wikitable"
|+ Third (5d) d-block Series (Lu–Hg)
|-
! Atomic number 
| 71 || 72 || 73 || 74 || 75 || 76 || 77 || 78 || 79 || 80
|-
! Element 
| Lu || Hf || Ta || W || Re || Os || Ir || Pt || Au || Hg
|-
! Electron<br/>configuration 
| 5d<sup>1</sup>6s<sup>2</sup> || 5d<sup>2</sup>6s<sup>2</sup> || 5d<sup>3</sup>6s<sup>2</sup> || 5d<sup>4</sup>6s<sup>2</sup>|| 5d<sup>5</sup>6s<sup>2</sup> || 5d<sup>6</sup>6s<sup>2</sup>|| 5d<sup>7</sup>6s<sup>2</sup>|| 5d<sup>9</sup>6s<sup>1</sup>|| 5d<sup>10</sup>6s<sup>1</sup>|| 5d<sup>10</sup>6s<sup>2</sup>
|}

{| class="wikitable"
|+ Fourth (6d) d-block Series (Lr–Cn)<br>(Configurations predicted for Mt–Cn)
|-
! Atomic number 
| 103 || 104 || 105 || 106 || 107 || 108 || 109 || 110 || 111 || 112
|-
! Element 
| Lr || Rf || Db || Sg || Bh || Hs || Mt || Ds || Rg || Cn
|-
! Electron<br/>configuration 
| 7s<sup>2</sup>7p<sup>1</sup> || 6d<sup>2</sup>7s<sup>2</sup> || 6d<sup>3</sup>7s<sup>2</sup> || 6d<sup>4</sup>7s<sup>2</sup> || 6d<sup>5</sup>7s<sup>2</sup> || 6d<sup>6</sup>7s<sup>2</sup> || 6d<sup>7</sup>7s<sup>2</sup>|| 6d<sup>8</sup>7s<sup>2</sup> || 6d<sup>9</sup>7s<sup>2</sup> || 6d<sup>10</sup>7s<sup>2</sup>
|}

A careful look at the electronic configuration of the elements reveals that there are certain exceptions to the [[Madelung rule]]. For Cr as an example the rule predicts the configuration 3d<sup>4</sup>4s<sup>2</sup>, but the observed atomic spectra show that the real [[ground state]] is 3d<sup>5</sup>4s<sup>1</sup>. To explain such exceptions, it is necessary to consider the effects of increasing [[nuclear charge]] on the orbital energies, as well as the electron–electron interactions including both [[Coulomb repulsion]] and [[exchange energy]].<ref name=Miessler/> The exceptions are in any case not very relevant for chemistry because the energy difference between them and the expected configuration is always quite low.<ref name="Jorgensen">{{cite journal |last1=Jørgensen |first1=Christian |date=1973 |title=The Loose Connection between Electron Configuration and the Chemical Behavior of the Heavy Elements (Transuranics) |journal=Angewandte Chemie International Edition |volume=12 |issue=1 |pages=12–19 |doi=10.1002/anie.197300121}}</ref>

The (''n'' − 1)d orbitals that are involved in the transition metals are very significant because they influence such properties as magnetic character, variable oxidation states, formation of coloured compounds etc. The valence s and p orbitals (''n''s and ''n''p) have very little contribution in this regard since they hardly change in the moving from left to the right in a transition series.
In transition metals, there are greater horizontal similarities in the properties of the elements in a period in comparison to the periods in which the d orbitals are not involved. This is because in a transition series, the valence shell electronic configuration of the elements do not change. However, there are some group similarities as well.

==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 energy gap between different possible oxidation states<ref>{{cite journal|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>
* 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]])
Most transition metals can be bound to a variety of [[ligands]], allowing for a wide variety of transition metal complexes.<ref>Hogan, C. Michael (2010). [http://www.eoearth.org/article/Heavy_metal?topic=49498 "Heavy metal"] in ''Encyclopedia of Earth''. National Council for Science and the Environment. E. Monosson and C. Cleveland (eds.) Washington DC.</ref>

===Coloured compounds<span class="anchor" id="Colored compounds"></span>===
[[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).]]

Colour in transition-series metal compounds is generally due to electronic transitions of two principal types.
*[[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 ion|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>{{cite book|author=Dunn, T.M.|editor=Lewis, J. and Wilkins, R.G.|title=Modern Coordination Chemistry|publisher=Wiley Interscience|location=New York|year=1960|pages= Chapter 4, Section 4, "Charge Transfer Spectra", pp. 268–273}}</ref>-->

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.

In general charge transfer transitions result in more intense colours than d–d transitions.
*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.

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),
which has a d<sup>5</sup> configuration in which all five electrons have parallel spins; the colour of such complexes is much weaker than in complexes with spin-allowed transitions. 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]].

===Oxidation states===
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-}}.

[[File:Transition metal oxidation states.svg|frame|center|Oxidation states of the transition metals. The solid dots show common oxidation states, and the hollow dots show possible but unlikely states.]]

[[Main-group element]]s in groups 13 to 18 also exhibit multiple oxidation states. The "common" oxidation states of these elements typically differ by two instead of one. For example, compounds of [[gallium]] in oxidation states +1 and +3 exist in which there is a single gallium atom. Compounds of Ga(II) would have an unpaired electron and would behave as a [[free radical]] and generally be destroyed rapidly, but some stable radicals of Ga(II) are known.<ref>{{cite journal|last1=Protchenko|first1=Andrey V.|last2=Dange|first2=Deepak|last3=Harmer|first3=Jeffrey R.|last4=Tang|first4=Christina Y.|last5=Schwarz|first5=Andrew D.|last6=Kelly|first6=Michael J.|last7=Phillips|first7=Nicholas|last8=Tirfoin|first8=Remi|last9=Birjkumar|first9=Krishna Hassomal|last10=Jones|first10=Cameron|last11=Kaltsoyannis|first11=Nikolas|last12=Mountford|first12=Philip|last13=Aldridge|first13=Simon|title=Stable GaX<sub>2</sub>, InX<sub>2</sub> and TlX<sub>2</sub> radicals|journal=Nature Chemistry|date=16 February 2014|volume=6|issue=4|pages=315–319|doi=10.1038/nchem.1870|pmid=24651198|bibcode = 2014NatCh...6..315P }}</ref> Gallium also has a formal oxidation state of +2 in 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.

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 row, the maximum occurs with [[ruthenium]] (+8), and in the third row, the maximum occurs with [[iridium]] (+9). In compounds such as {{chem|[MnO|4|]|-}} and {{chem|OsO|4}}, the elements achieve a stable configuration by [[covalent bonding]].

The lowest oxidation states are exhibited in [[metal carbonyl]] complexes such 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.

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.
<!-- [[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.]] -->

===Magnetism===
{{main|Magnetochemistry}}
Transition metal compounds are [[paramagnetic]] when they have one or more unpaired d electrons.<ref>{{cite book|last1=Figgis|first1=B.N.|last2=Lewis|first2=J.|title=The Magnetochemistry of Complex Compounds|editor=Lewis, J.|editor2=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.

[[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. [[Antiferromagnetism]] 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 [[Hydrogenation|catalytic hydrogenation]]) are some of the examples. Catalysts at a solid surface ([[nanomaterial-based catalyst]]s) 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]].

An interesting type of catalysis occurs when the products of a reaction catalyse the reaction producing more catalyst ([[autocatalysis]]). One example is the reaction of [[oxalic acid]] with acidified [[potassium permanganate]] (or manganate (VII)).<ref>{{cite journal | title = Revising the Mechanism of the Permanganate/Oxalate Reaction |vauthors=Kovacs KA, Grof P, Burai L, Riedel M | journal = J. Phys. Chem. A | doi = 10.1021/jp047061u | year = 2004 | volume = 108 | pages = 11026–11031 | issue = 50| bibcode = 2004JPCA..10811026K |url=https://infoscience.epfl.ch/handle/20.500.14299/2553 }}</ref> Once a little Mn<sup>2+</sup> has been produced, it can react with MnO<sub>4</sub><sup>−</sup> forming Mn<sup>3+</sup>. This then reacts with C<sub>2</sub>O<sub>4</sub><sup>−</sup> ions forming Mn<sup>2+</sup> again.

===Physical properties===
As implied by the name, all transition metals are [[metal]]s and thus conductors of electricity.

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, which again tends to differentiate them from the accepted transition metals. Mercury has a melting point of {{convert|−38.83|°C|F}} and is a liquid at room temperature.

==See also==
{{Scholia}}
*[[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 [[Degenerate energy levels|degeneracies]] of electronic orbital states
*[[Post-transition metal]], a metallic element to the right of the transition metals in the periodic table

==References==
{{reflist|35em}}

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[[Category:Periodic table]]
[[Category:Transition metals| ]]