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= Transition_metal =
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Introduction
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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 (and less often 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 and thermal
conductivity. Most (with the exception of 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 states and
bind to a variety of ligands to form coordination complexes that are
often coloured. They form many useful alloys and are often employed as
catalysts in elemental form or in compounds such as coordination
complexes and oxides. Most are strongly paramagnetic because of their
unpaired d electrons, as are many of their compounds. All of the
elements that are 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' = 3 in the 4th row of the periodic table)
from a stable group of 8 to one of 18, or from 18 to 32. These
elements are now known as the d-block.
Definition and classification
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The 2011 IUPAC 'Principles of Chemical Nomenclature' describe a
"transition metal" as any element in groups 3 to 12 on the periodic
table. This corresponds exactly to the d-block elements, and many
scientists use this definition. In actual practice, the f-block
lanthanide and actinide series are called "inner transition metals".
The 2005 'Red Book' allows for the group 12 elements to be excluded,
but not the 2011 'Principles'.
The IUPAC 'Gold Book' defines a transition metal as "an element whose
atom has a partially filled d sub-shell, or which can give rise to
cations 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.
In the d-block, the atoms of the elements have between zero and ten d
electrons.
d-block}}; border: 1px solid ; padding:2px; font-size: 85%;"
|table title}}" | Transition metals in the d-block |table header}}" |
Group 3 4 5 6 7 8 9 10 11 12
table colheader}}"| Period 4 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni
29Cu 30Zn
table colheader}}"| 5 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag
48Cd
table colheader}}"| 6 71Lu 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au
80Hg
table colheader}}"| 7 103Lr 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds
111Rg 112Cn
Published texts and periodic tables show variation regarding the
heavier members of group 3. The common placement of lanthanum and
actinium in these positions is not supported by physical, chemical,
and electronic evidence, which overwhelmingly favour putting lutetium
and lawrencium in those places. 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;
additionally, it creates a 15-element-wide f-block, when quantum
mechanics dictates that the f-block should only be 14 elements wide.
The form with lutetium and lawrencium in group 3 is supported by a
1988 IUPAC report on physical, chemical, and electronic grounds, 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. 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
orbital but have no 5f occupancy as single atoms); the majority of
investigators considering the problem agree with the updated form with
lutetium and lawrencium.
The group 12 elements zinc, cadmium, and mercury are sometimes
excluded from the transition metals. This is because they have the
electronic configuration [ ]d10s2, where the d shell is complete, and
they still have a complete d shell in all their known oxidation
states. The group 12 elements Zn, Cd and Hg may therefore, under
certain criteria, be classed as post-transition metals in this case.
However, it is often convenient to include these elements in a
discussion of the transition elements. For example, when discussing
the crystal field stabilization energy of first-row transition
elements, it is convenient to also include the elements calcium and
zinc, as both and 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 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.
The 2007 (though disputed and so far not reproduced independently)
synthesis of mercury(IV) fluoride () has been taken by some to
reinforce the view that the group 12 elements should be considered
transition metals, but some authors still consider this compound to be
exceptional. Copernicium is expected to be able to use its d electrons
for chemistry as its 6d subshell is destabilised by strong
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, Cn2+ is predicted
to be 6d87s2, unlike Hg2+ which is 5d106s0.
Although meitnerium, darmstadtium, and roentgenium are within the
d-block and are expected to behave as transition metals analogous to
their lighter congeners iridium, platinum, and gold, this has not yet
been experimentally confirmed. Whether copernicium behaves more like
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-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 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 compounds, and for that reason have been
called "honorary" transition metals. The same is likely true of
radium.
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
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The general electronic configuration of the d-block atoms is [noble
gas]('n' − 1)d0-10'n's0-2'n'p0-1. 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' = 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]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 high 'Z'), but they
can contribute to the chemical bonding in transition metal compounds.
The Madelung rule predicts that the inner d orbital is filled after
the valence-shell s orbital. The typical electronic structure of
transition metal atoms is then written as [noble gas]'n's2('n' −
1)d'm'. 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's2('n' − 1)d1 configuration, except
for lawrencium (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, or scandium (Sc), the first element of group 3
with atomic number 'Z' = 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:
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 the Madelung rule. For Cr as an
example the rule predicts the configuration 3d44s2, but the observed
atomic spectra show that the real ground state is 3d54s1. 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. 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.
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
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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
* 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
* 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.
===Coloured compounds===
From left to right, aqueous solutions of: (red); (orange);
(yellow); (turquoise); (blue); (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.
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
diagrams.
In centrosymmetric complexes, such as octahedral complexes, d-d
transitions are forbidden by the Laporte rule and only occur because
of vibronic coupling in which a molecular vibration occurs together
with a d-d transition. Tetrahedral complexes have somewhat more
intense colour because mixing d and p orbitals is possible when there
is no centre of symmetry, so transitions are not pure d-d transitions.
The molar absorptivity (ε) of bands caused by d-d transitions are
relatively low, roughly in the range 5-500 M−1cm−1 (where M = mol
dm−3). 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 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 spectrum of shows a maximum molar absorptivity
of about 0.04 M−1cm−1 in the visible spectrum.
Oxidation states
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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 , and
+5, such as .
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 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. Gallium also has a formal
oxidation state of +2 in dimeric compounds, such as , which contain a
Ga-Ga bond formed from the unpaired electron on each Ga atom. 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 and , the
elements achieve a stable configuration by covalent bonding.
The lowest oxidation states are exhibited in metal carbonyl complexes
such as (oxidation state zero) and (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
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Transition metal compounds are paramagnetic when they have one or more
unpaired d electrons. 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 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.
Antiferromagnetism is another example of a magnetic property arising
from a particular alignment of individual spins in the solid state.
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 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
transition metal ions can change their oxidation states, they become
more effective as 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)). 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
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As implied by the name, all transition metals are metals and thus
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, which again tends to differentiate them
from the accepted transition metals. Mercury has a melting point of
−38.83 °C and is a liquid at room temperature.
See also
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*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
*Post-transition metal, a metallic element to the right of the
transition metals in the periodic table
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Original Article:
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