Transition metals

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The chemical elements with atomic numbers from 21 to 30, 39 to 48, 57 to 80 and 89 to 112 are usually referred to as transition elements . Since these elements are all metals , the term transition metals is also used. This name is based on its position in the periodic table , since there the transition is shown by the successive increase of electrons in the d - atomic orbitals along each period. IUPAC defines transition elements as elements that have an incomplete d-subshell or form ions with an incomplete d- subshell . According to this stricter definition, the elements of the zinc group are not transition elements since they have the d 10 configuration. Traditionally, however, the simpler and less strict definition has been used.

21
Sc 		22
Ti 		23
V 		24
Cr 		25
mn 		26
feet 		27
Co 		28
Ni 		29
Cu 		30
notes
39
Y 		40
Zr 		41
Nb 		42
Mon 		43
Tc 		44
Ru 		45
Rh 		46
Pd 		47
Ag 		48
Cd
57
La 		72
Hf 		73
days 		74
W 		75
Re 		76
Os 		77
Ir 		78
Pt 		79
Au 		80
ed
89
Ac 		104
para 		105
Db 		106
Sg 		107
hours 		108
ms 		109
m 		110
Ds 		111
Rg 		112
cn

Electron configuration

Main group elements , which are in the periodic table before the transition metals ( i.e. elements number 1 to 20), have no electrons in the d orbitals, but only in the s and p orbitals (although it is assumed that the empty d orbitals play a role in the behavior of such elements as silicon , phosphorus and sulfur ).

For the d-block elements from scandium to zinc , the d orbitals are filled in along the period. Except for copper and chromium , all 4th period 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 in before the outer shells. The s orbitals in the d block elements are, however, in a lower energy state than the d subshells. Since atoms strive to adopt the lowest possible energy state, the s shells are filled first. The exceptions for chromium and copper - which only have one electron in their outer orbital - are due to electron repulsion. The division of electrons into s and d orbitals leads to lower energy states for the atoms than placing two electrons in the outer s orbital.

Not all d -block elements are transition metals. Scandium and zinc do not fit into the definition given above. Scandium has one electron in its d lower shell and 2 electrons in its outer s orbital. Since the only scandium ion (Sc 3+ ) has no electrons in the d orbital, it can of course not have a “partially filled” d orbital. The same applies to zinc, since its only ion, Zn 2+ , has a completely filled d orbital.

Chemical and physical properties

Transition elements are generally characterized by high tensile strengths , densities , melting points and boiling points . Like other properties of transition metals, these are also due to the ability of the electrons in the d orbitals to be delocalized within the metal lattice. In metallic materials, the more electrons are shared between the nuclei, the more pronounced these properties are.

Typical properties of transition metals are:

Oxidation states

Oxidation states of the transition metal compounds: Frequent oxidation states are marked with a filled circle, rarer and energetically less favorable states with a ring.

Compared to elements of group II such as calcium , the ions of the transition elements exist in numerous oxidation states. Calcium atoms only give off two electrons, as this gives them a noble gas configuration. They are therefore in the + II oxidation state, whereas a transition element can emit up to eight electrons. If you look at the ionization enthalpies of both groups, you can also see the reason for this. The energy required to remove electrons from calcium is low until one tries to remove electrons below its outer s orbital. Ca 3+ has an enthalpy of ionization that is so high that it does not normally occur. Transition elements such as vanadium, on the other hand, have fairly linearly increasing ionization enthalpies along their s and d orbitals because of the small energy difference between the 3d and 4s orbitals. Transition elements therefore also occur with very high oxidation numbers. In general, those electron configurations that are either fully or half occupied are preferred.

Certain behavior patterns can be seen along a period:

  • The number of oxidation states increases in the 4th period up to manganese and then decreases again. This is due to the stronger attraction of the protons in the nucleus, which makes it difficult for electrons to be given off.
  • The elements in their low oxidation states usually occur as simple ions. In higher oxidation states, they are usually covalently bound to other electronegative elements such as oxygen or fluorine , often as anions .

A linear trend for the maximum oxidation states was recently predicted for the transition metals of the 6th period. The maximum oxidation levels from lanthanum to osmium increase gradually from + III to + VIII and then decrease again linearly to oxidation level + IV for mercury. This prediction of the maximum oxidation states for the 5d transition metal series was confirmed by the representation of the oxidation state + IV for mercury as HgF 4 .

Properties depending on the oxidation state:

  • Higher oxidation states become less stable along the period.
  • Ions in higher oxidation states are good oxidizing agents, whereas elements in lower oxidation states are reducing agents.
  • The (2+) ions begin at the beginning of the period as strong reducing agents and then become more and more stable.
  • The (3+) ions, on the other hand, start out stable and then become better and better oxidizing agents .

Catalytic activity

Transition metals are good homogeneous or heterogeneous catalysts , e.g. B. iron is the catalyst for the Haber-Bosch process . Nickel and platinum are used for the hydrogenation of alkenes. Palladium (Pd) is well suited for catalyzed CC coupling reactions ( Suzuki , Heck , Stille etc.). Rhodium (Rh), iridium (Ir) and ruthenium (Ru) are e.g. B. used in the asymmetric hydrogenation of prochiral molecules. In most cases, phosphorus compounds are used here as ligands for stereo control. The best known ligands are z. B. BINAP by R. Noyori (Nobel Prize 2001), DIOP by Kagan , JosiPhos / WalPhos, and DuPhos . All the ligands mentioned have in common that they are bidentate and chelating , ie two phosphorus atoms of the ligand bind to the metal at the same time.

Colored connections

From left to right: Co (NO 3 ) 2 dissolved in water (red); K 2 Cr 2 O 7 (orange); K 2 CrO 4 (yellow); NiCl 2 (green); CuSO 4 (blue); KMnO 4 (purple)

When the frequency of electromagnetic radiation changes, different colors can be perceived. They result from the different composition of light after it has been reflected , transmitted or absorbed after contact with a substance - one also speaks of remission . Because of their structure, transition metals form many different colored ions and complexes. The colors differ even for the same element - e.g. B. MnO 4 - (Mn in the +7 oxidation state) is a purple compound, but Mn 2+ is pale pink. Cr (II) compounds are usually blue, Cr (III) compounds are green, while Cr (VI) compounds are yellow to orange. Complex formation can play an important role in coloring. The ligands have a great influence on the d-shell. They partially attract the d electrons and split them into higher and lower (in terms of energy) groups. Electromagnetic radiation is only absorbed if its frequency corresponds to the energy difference between two energy states of the atom (because of the formula E = hν .) When light hits an atom with split d orbitals, some electrons are raised to the higher state ( dd transition ) . Compared to a non-complexed ion, different frequencies can be absorbed and therefore different colors can be observed.

The color of a complex depends on:

  • the type of metal ion, or more precisely the number of electrons in the d orbitals
  • the arrangement of the ligands around the metal ion ( complex isomers can take on different colors)
  • the nature of the ligands around the metal ion. The stronger the ligands, the greater the energy difference between the two split 3 d groups.

The complexes of the d-block element zinc (strictly speaking not a transition element) are colorless, since the 3d orbitals are completely occupied and therefore no electrons can be lifted.

See also

Web links

Individual evidence

  1. ↑ Entry on transition metal . In: IUPAC Compendium of Chemical Terminology (the “Gold Book”) . doi : 10.1351 / goldbook.T06456 Version: 2.3.1.
  2. ↑ Xuefang Wang, Lester Andrews, Sebastian Riedel, Martin Kaupp: Mercury is a Transition Metal: The First Experimental Evidence for HgF4 ; Angew. Chem. 2007, doi : 10.1002 / anie.200703710 , Homepage of Dr. Sebastian Riedel