In chemistry, atransition metal (ortransition element) is achemical element in thed-block of theperiodic table (groups 3 to 12), though the elements ofgroup 12 (and less oftengroup 3) are sometimes excluded. Thelanthanide andactinide elements (thef-block) are calledinner transition metals and are sometimes considered to be transition metals as well.
They are lustrous metals with goodelectrical andthermal conductivity. Most (with the exception ofgroup 11 and group 12) are hard and strong, and have high melting and boiling temperatures. They form compounds in any of two or more differentoxidation states and bind to a variety ofligands to formcoordination complexes that are often coloured. They form many usefulalloys and are often employed ascatalysts in elemental form or in compounds such as coordination complexes andoxides. Most are stronglyparamagnetic because of theirunpaired d electrons, as are many of their compounds. All of the elements that areferromagnetic near room temperature are transition metals (iron,cobalt andnickel) or inner transition metals (gadolinium).
English chemistCharles Rugeley Bury (1890–1968) first used the wordtransition in this context in 1921, when he referred to atransition series of elements during the change of an inner layer of electrons (for examplen = 3 in the 4th row of the periodic table) from a stable group of 8 to one of 18, or from 18 to 32.[1][2][3] These elements are now known as the d-block.
The first row of transition metals, in order
Definition and classification
The 2011IUPACPrinciples of Chemical Nomenclature describe a "transition metal" as any element in groups 3 to 12 on theperiodic table.[4] This corresponds exactly to thed-block elements, and many scientists use this definition.[5][6] In actual practice, thef-blocklanthanide andactinide series are called "inner transition metals". The 2005Red Book allows for the group 12 elements to be excluded, but not the 2011Principles.[7]
The IUPACGold Book[8] defines a transition metal as "anelement whose atom has a partially filledd sub-shell, or which can give rise tocations with an incomplete d sub-shell", but this definition is taken from an old edition of theRed Book and is no longer present in the current edition.[7]
In the d-block, the atoms of the elements have between zero and ten d electrons.
Published texts and periodic tables showvariation regarding the heavier members of group 3.[9] The common placement oflanthanum andactinium in these positions is not supported by physical, chemical, and electronicevidence,[10][11][12] which overwhelmingly favour puttinglutetium andlawrencium in those places.[13][14] Some authors prefer to leave the spaces belowyttrium blank as a third option, but there is confusion on whether this format implies that group 3 contains onlyscandium and yttrium, or if it also contains all the lanthanides and actinides;[15][16][17][18][19] additionally, it creates a 15-element-wide f-block, whenquantum mechanics dictates that the f-block should only be 14 elements wide.[15] The form with lutetium and lawrencium in group 3 is supported by a 1988IUPAC report on physical, chemical, and electronic grounds,[20] 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.[15] 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 ofactinium andthorium, which both can use 5f as avalenceorbital but have no 5f occupancy as single atoms);[14][21][22] the majority of investigators considering the problem agree with the updated form with lutetium and lawrencium.[14]
The group 12 elementszinc,cadmium, andmercury are sometimes excluded from the transition metals.[1] This is because they have theelectronic configuration [ ]d10s2, where the d shell is complete,[23] and they still have a complete d shell in all their knownoxidation states. The group 12 elements Zn, Cd and Hg may therefore, under certain criteria, be classed aspost-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 thecrystal field stabilization energy of first-row transition elements, it is convenient to also include the elementscalcium and zinc, as bothCa2+ andZn2+ have a value of zero, against which the value for other transition metal ions may be compared. Another example occurs in theIrving–Williams series of stability constants of complexes. Moreover, Zn, Cd, and Hg can use their d orbitals forbonding 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.[24][25][26]
The 2007 (though disputed and so far not reproduced independently) synthesis ofmercury(IV) fluoride (HgF 4) has been taken by some to reinforce the view that the group 12 elements should be considered transition metals,[27] but some authors still consider this compound to be exceptional.[28]Copernicium is expected to be able to use its d electrons for chemistry as its 6dsubshell is destabilised by strongrelativistic 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, Cn2+ is predicted to be 6d87s2, unlike Hg2+ which is 5d106s0.
Althoughmeitnerium,darmstadtium, androentgenium are within the d-block and are expected to behave as transition metals analogous to their lightercongenersiridium,platinum, andgold, this has not yet been experimentally confirmed. Whethercopernicium behaves more likemercury or has properties more similar to those of thenoble gasradon is not clear. Relative inertness of Cn would come from the relativistically expanded 7s–7p1/2 energy gap, which is already adumbrated in the 6s–6p1/2 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 elementscalcium,strontium, andbarium do not have filled d-orbitals as single atoms, but are known to have d-orbital bonding participation in somecompounds, and for that reason have been called "honorary" transition metals.[29] The same is likely true ofradium.[30]
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.
The general electronic configuration of the d-block atoms is [noble gas](n − 1)d0–10ns0–2np0–1. Here "[noble gas]" is the electronic configuration of the lastnoble gas preceding the atom in question, andn is the highestprincipal quantum number of an occupied orbital in that atom. For example, Ti (Z = 22) is in period 4 so thatn = 4, the first 18 electrons have the same configuration of Ar at the end of period 3, and the overall configuration is [Ar]3d24s2. The period 6 and 7 transition metals also add core (n − 2)f14 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 highZ), but they can contribute to the chemical bonding in transition metal compounds.
TheMadelung rule predicts that the inner d orbital is filled after thevalence-shell s orbital. The typicalelectronic structure of transition metal atoms is then written as [noble gas]ns2(n − 1)dm. 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 exceptpalladium (Pd), with no electron in that s sub shell in its ground state. The s subshell in the valence shell is represented as thens 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 anns2(n − 1)d1 configuration, except forlawrencium (Lr): its 7s27p1 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 = 20) of group 2 with the configuration [Ar]4s2, orscandium (Sc), the first element of group 3 with atomic numberZ = 21 and configuration [Ar]4s23d1, 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:[31]
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 configuration
3d14s2
3d24s2
3d34s2
3d54s1
3d54s2
3d64s2
3d74s2
3d84s2
3d104s1
3d104s2
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 configuration
4d15s2
4d25s2
4d45s1
4d55s1
4d55s2
4d75s1
4d85s1
4d105s0
4d105s1
4d105s2
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 configuration
5d16s2
5d26s2
5d36s2
5d46s2
5d56s2
5d66s2
5d76s2
5d96s1
5d106s1
5d106s2
Fourth (6d) d-block Series (Lr–Cn) (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 configuration
7s27p1
6d27s2
6d37s2
6d47s2
6d57s2
6d67s2
6d77s2
6d87s2
6d97s2
6d107s2
A careful look at the electronic configuration of the elements reveals that there are certain exceptions to theMadelung rule. For Cr as an example the rule predicts the configuration 3d44s2, but the observed atomic spectra show that the realground state is 3d54s1. To explain such exceptions, it is necessary to consider the effects of increasingnuclear charge on the orbital energies, as well as the electron–electron interactions including bothCoulomb repulsion andexchange energy.[31] 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.[32]
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 (ns andnp) 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[33]
the formation of manyparamagnetic 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 ofligands, allowing for a wide variety of transition metal complexes.[34]
Coloured compounds
From left to right, aqueous solutions of:Co(NO 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 predominantlyligandorbital 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 ofchromate,dichromate andpermanganate ions is due to LMCT transitions. Another example is thatmercuric 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.
In general charge transfer transitions result in more intense colours than d–d transitions.
d–d transitions. An electron jumps from oned 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 usingcrystal 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 onTanabe–Sugano diagrams.
Incentrosymmetric complexes, such as octahedral complexes, d–d transitions are forbidden by theLaporte rule and only occur because ofvibronic coupling in which amolecular 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. Themolar absorptivity (ε) of bands caused by d–d transitions are relatively low, roughly in the range 5–500 M−1cm−1 (whereM = mol dm−3).[35] Some d–d transitions arespin forbidden. An example occurs in octahedral, high-spin complexes ofmanganese(II),which has a d5 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. Thespectrum of[Mn(H 2O) 6]2+ shows a maximum molar absorptivity of about 0.04 M−1cm−1 in thevisible spectrum.
Oxidation states
A characteristic of transition metals is that they exhibit two or moreoxidation states, usually differing by one. For example, compounds ofvanadium are known in all oxidation states between −1, such as[V(CO) 6]− , and +5, such asVO3− 4.
Oxidation states of the transition metals. The solid dots show common oxidation states, and the hollow dots show possible but unlikely states.
Main-group elements 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 ofgallium 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 afree radical and generally be destroyed rapidly, but some stable radicals of Ga(II) are known.[36] Gallium also has a formal oxidation state of +2 in dimeric compounds, such as[Ga 2Cl 6]2− , which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom.[37] 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 fromtitanium (+4) up tomanganese (+7), but decreases in the later elements. In the second row, the maximum occurs withruthenium (+8), and in the third row, the maximum occurs withiridium (+9). In compounds such as[MnO 4]− andOsO 4, the elements achieve a stable configuration bycovalent bonding.
The lowest oxidation states are exhibited inmetal carbonyl complexes such asCr(CO) 6 (oxidation state zero) and[Fe(CO) 4]2− (oxidation state −2) in which the18-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.
Transition metal compounds areparamagnetic when they have one or more unpaired d electrons.[38] In octahedral complexes with between four and seven d electrons bothhigh spin andlow spin states are possible. Tetrahedral transition metal complexes such as[FeCl 4]2− arehigh 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 arediamagnetic. 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 alloyalnico 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 heterogeneouscatalytic activity. This activity is ascribed to their ability to adopt multiple oxidation states and to form complexes.Vanadium(V) oxide (in thecontact process), finely dividediron (in theHaber process), andnickel (incatalytic 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 transition metal ions can change their oxidation states, they become more effective ascatalysts.
An interesting type of catalysis occurs when the products of a reaction catalyse the reaction producing more catalyst (autocatalysis). One example is the reaction ofoxalic acid with acidifiedpotassium permanganate (or manganate (VII)).[39] Once a little Mn2+ has been produced, it can react with MnO4− forming Mn3+. This then reacts with C2O4− ions forming Mn2+ again.
Physical properties
As implied by the name, all transition metals aremetals and thus conductors of electricity.
In general, transition metals possess a highdensity and highmelting points andboiling points. These properties are due tometallic bonding by delocalized d electrons, leading tocohesion 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 −38.83 °C (−37.89 °F) and is a liquid at room temperature.
^Wittig, Jörg (1973). "The pressure variable in solid state physics: What about 4f-band superconductors?". In H. J. Queisser (ed.).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. Advances in Solid State Physics. Vol. 13. Berlin, Heidelberg: Springer. pp. 375–396.doi:10.1007/BFb0108579.ISBN978-3-528-08019-8.
^Matthias, B. T. (1969). "Systematics of Super Conductivity". In Wallace, P. R. (ed.).Superconductivity. Vol. 1. Gordon and Breach. pp. 225–294.ISBN9780677138107.
^William B. Jensen (1982). "The Positions of Lanthanum (Actinium) and Lutetium (Lawrencium) in the Periodic Table".J. Chem. Educ.59 (8):634–636.Bibcode:1982JChEd..59..634J.doi:10.1021/ed059p634.
^Thyssen, P.; Binnemans, K. (2011). "Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis". In Gschneidner, K. A. Jr.; Bünzli, J-C.G; Vecharsky, Bünzli (eds.).Handbook on the Physics and Chemistry of Rare Earths. Vol. 41. Amsterdam: Elsevier. pp. 1–94.doi:10.1016/B978-0-444-53590-0.00001-7.ISBN978-0-444-53590-0.
^Chemey, Alexander T.; Albrecht-Schmitt, Thomas E. (2019). "Evolution of the periodic table through the synthesis of new elements".Radiochimica Acta.107 (9–11):771–801.doi:10.1515/ract-2018-3082.S2CID104470619.
^Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999).Advanced Inorganic Chemistry (6th ed.). New York: Wiley,ISBN978-0-471-19957-1.
^Tossell, J.A. (1 November 1977). "Theoretical studies of valence orbital binding energies in solid zinc sulfide, zinc oxide, and zinc fluoride".Inorganic Chemistry.16 (11):2944–2949.doi:10.1021/ic50177a056.
^Farberovich, O. V.; Kurganskii, S. I.; Domashevskaya, E. P. (1980). "Problems of the OPW Method. II. Calculation of the Band Structure of ZnS and CdS".Physica Status Solidi B.97 (2):631–640.Bibcode:1980PSSBR..97..631F.doi:10.1002/pssb.2220970230.
^Pyykkö, Pekka; Desclaux, Jean-Paul (1979). "Relativity and the Periodic System of Elements".Accounts of Chemical Research.12 (8):276–281.doi:10.1021/ar50140a002.
^abMiessler, G. L. and Tarr, D. A. (1999)Inorganic Chemistry, 2nd edn, Prentice-Hall, p. 38-39ISBN978-0-13-841891-5
^Jørgensen, Christian (1973). "The Loose Connection between Electron Configuration and the Chemical Behavior of the Heavy Elements (Transuranics)".Angewandte Chemie International Edition.12 (1):12–19.doi:10.1002/anie.197300121.
^Hogan, C. Michael (2010)."Heavy metal" inEncyclopedia of Earth. National Council for Science and the Environment. E. Monosson and C. Cleveland (eds.) Washington DC.
^Orgel, L.E. (1966).An Introduction to Transition-Metal Chemistry, Ligand field theory (2nd. ed.). London: Methuen.
^Protchenko, Andrey V.; Dange, Deepak; Harmer, Jeffrey R.; Tang, Christina Y.; Schwarz, Andrew D.; Kelly, Michael J.; Phillips, Nicholas; Tirfoin, Remi; Birjkumar, Krishna Hassomal; Jones, Cameron; Kaltsoyannis, Nikolas; Mountford, Philip; Aldridge, Simon (16 February 2014). "Stable GaX2, InX2 and TlX2 radicals".Nature Chemistry.6 (4):315–319.Bibcode:2014NatCh...6..315P.doi:10.1038/nchem.1870.PMID24651198.
^Figgis, B.N.; Lewis, J. (1960). Lewis, J.; Wilkins, R.G. (eds.).The Magnetochemistry of Complex Compounds. Modern Coordination Chemistry. New York: Wiley Interscience. pp. 400–454.
