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Transition_metal.
The Organometallic Chemistry of the Transition Metals, 3rd Edition by Robert H. Crabtree
Transition Metals in Organic Synthesis: A Practical Approach (The Practical Approach in Chemistry Series) by Susan Gibson
Electronic Structure and Properties of Transition Metal Compounds: Introduction to the Theory by Isaac B. Bersuker
Transition Metal Oxides : Structure, Properties, and Synthesis of Ceramic Oxides by C. N. R. Rao
Transition Metal Hydrides by A. Dedieu
Late Transition Metal Polymerization Catalysis by Bernhard Rieger
Localized to Itinerant Electronic Transition in Perovskite Oxides (Structure and Bonding, 98) by John B. Goodenough
Application of Transition Metal Catalysts in Organic Synthesis (Springer Desktop Editions in Chemistry) by L. Brandsma
Fundamentals of Molecular Catalysis by Hideo Kurosawa
Photochemistry and Photophysics of Metal Complexes (Modern Inorganic Chemistry) by D. M. Roundhill
Transition Metals in the Synthesis of Complex Organic Molecules by Louis S. Hegedus
d- and f- Block Chemistry by Chris J. Jones
Transition Metal Reagents and Catalysts : Innovations in Organic Synthesis by Jiro Tsuji
Lecture Notes on Electron Correlation and Magnetism (Series in Modern Condensed Matter Physics, Vol. 5) by Patrik Fazekas
Organometallics 2: Complexes With Transition Metal-Carbon Pi-Bonds (Oxford Chemistry Primers, No 13) by Manfred Bochmann
Transition metal
Group
Period 4
Period 5
Period 6
Period 7
3 (III B)
Sc 21
Y 39
Lu 71
Lr 103
4 (IV B)
Ti 22
Zr 40
Hf 72
Rf 104
5 (V B)
V 23
Nb 41
Ta 73
Db 105
6 (VI B)
Cr 24
Mo 42
W 74
Sg 106
7 (VII B)
Mn 25
Tc 43
Re 75
Bh 107
8 (VIII B)
Fe 26
Ru 44
Os 76
Hs 108
9 (VIII B)
Co 27
Rh 45
Ir 77
Mt 109
10 (VIII B)
Ni 28
Pd 46
Pt 78
11 (I B)
Cu 29
Ag 47
Au 79
12 (II B)
Zn 30
Cd 48
Hg 80
A transition metal is any of the thirty chemical elements 21 through 30, 39 through 48, and 71 through 80. This name comes from their position in the periodic table of elements, which represent the successive addition of electrons to the d atomic orbitals of the atoms as one progresses through each of the three periods.
Transition elements are chemically defined as elements which form at least one ion with a partially filled subshell of d electrons.
Table of contents showTocToggle("show","hide")
1 Electronic configuration
2 Chemical properties
3 Variable oxidation states
4 Catalytic activity
5 Colored compounds
Electronic configuration
Main group elements prior to the appearance of the transition group elements in the periodic chart (ie, elements number 1 through 20) have no electrons in d orbitals, but only in the s and p orbitals.
(Though the low-lying, but empty d orbitals are thought to play a role in their d period elements such as silicon, phosphorus and sulfur)
From Scandium to Zinc, d block elements fill up their d orbitals across the period. With the exception of copper and chromium, all d block elements have two electrons in their outer s orbital, even elements with incomplete 3d orbitals.
This is unusual: lower orbitals are usually filled up before outer shells. It happens that the s orbitals in d block elements are at lower energy states than the d subshells. As atoms always strive to be in states of lowest energy, s shells are filled up first. The copper and chromium exceptions - which have one electron in their outer orbital - do so because of electron repulsion. Sharing the electrons throughout the s and d orbitals gives lower energy states to the atoms than putting two electrons in the outer s orbital.
Not all d block elements are transition metals.
Scandium and zinc don't qualify, due to the chemical definition given above. Scandium has one electron in its d subshell, and 2 electrons in its outer s orbital.
As scandium's only ion (Sc3+) has no electrons in its d orbital it is clear that it doesn't have a 'partially filled d orbital'. Similarly, zinc is not applicable because its only ion, Zn2+, has a full d orbital.
Chemical properties
Transition elements tend to have high tensile strength, density and melting and boiling points. As with many properties of transition metals, this is due to d orbital electrons' ability to delocalise within the metal lattice.
In metallic substances, the more electrons shared between nuclei, the stronger the metal.
There are four common characteristic properties of transition elements:
Variable oxidation states
Compared to Group II elements such as calcium, transition elements form ions with a wide variety of oxidation states. Calcium ions typically don't lose more than two electrons, whereas transition metals can lose up to nine.
The reason for this can be obtained by studying the ionisation enthalpies of both groups. The energies required to remove electrons from calcium are low until you try to remove electrons from below its outer two s orbitals.
In fact Ca3+ has an ionisation enthalpy so high that it rarely occurs naturally.
However a transition element like vanadium has roughly linear increasing ionisation enthalpies throughout its s and d orbitals, due to the close energy difference between the 3d and 4s orbitals.
Transition metal ions are therefore commonly found in very high states.
Certain patterns can be seen to emerge across the period of transition elements:
- The number of oxidation states of each ion increases up to Mn, after which they start to drop. This drop is due to the stronger pull from the protons in the nucleus towards the electrons, making them harder to remove.
- When the elements are in lower oxidation states, they can be found as simple ions. However elements in higher oxidation states are usually bonded covalently to electronegative compounds such as O or F, often in an anion.
Properties with respect to the stability of oxidation states:
- Higher oxidation state ions become less stable across the period.
- Ions in higher oxidation states tend to make good oxidising agents, whereas elements in low oxidation states become reducing agents.
- The 2+ ions across the period start as strong reducing agents, and become more stable.
- The 3+ ions start stable and become more oxidising across the period.
Catalytic activity
Transition metals form good homogeneous or heterogeneous catalysts, for example iron is the catalyst for the Haber process.
Nickel or platinum is used in the hydrogenation of alkenes.
Colored compounds
We observe color as varying frequencies of electromagnetic radiation in the visible region of the electromagnetic spectrum. Different colors result from the changed composition of light after it has been reflected, transmitted or absorbed after hitting a substance.
Because of their structure, transition metals form many different colored ions and complexes.
Color even varies between the different ions of a single element - MnO4- (Mn in oxidation state 7+) is a purple compound, whereas Mn2+ is pale-pink.
Complex formation can play a part in determining color in a transition compound.
This is because of the effect that ligands have on the 3d subshell. Ligands pull on some of the 3d electrons and split them in to higher and lower (in terms of energy) groups.
Electromagnetic radiation is only absorbed if its frequency is proportional to the difference in energies between two energy states present in an atom (through the formula e=hf.)
When light hits an atom which has had its 3d orbitals split, some of the electrons become promoted to the higher group. Compared to an un-complexed ion, different frequencies can be absorbed, hence different colors are observed.
The color of a complex depends on:
- The nature of the metal ion, specifically the number of electrons in the d orbitals
- The arrangement of the ligands around the metal ion (for example geometric isomers can display different colors)
- The nature of the ligands surrounding the metal ion. The stronger the ligands then the greater the energy difference between the split high and low 3d groups.
The complex formed by the d block element zinc (though not strictly a transition element) is colorless, because the 3d orbitals are full - no electrons are able to move up to the higher group.
The above article is adapted from from Wikipedia All Wikipedia article text is available under the terms of the GNU Free Documentation License
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