What is low and high oxidation state

The chemical elements with atomic numbers from 21 to 30, 39 to 48, 57 to 80 and 89 to 112 are commonly called Transition elements designated. Since these elements are all metals, so is the term Transition metals used. This name is based on its position in the periodic table, as there is the transition through the successive increase of electrons in the d-Shows atomic orbital along each period. Transition elements are chemically called Elements that form at least one ion with a partially filled d-shell, Are defined.

Examples of transition metals    

group   3 (III B) 4 (IV B) 5 (V B) 6 (VI B) 7 (VII B) 8 (VIII B) 9 (VIII B) 10 (VIII B) 11 (I B) 12 (II B)
4th period   Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30
5th period   Y 39 Zr 40 Nb 41 Mon 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 CD 48
6th period   La 57 -

Lu 71

Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80
7th period   Ac 89 -

Lr 103

Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 Ds 110 Rg 111 UUb 112

Electron configuration

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

Both d block-Elements from scandium to zinc, are the d-Orbitals filled in along the period. Except for copper and chrome, everyone has d-Block element two electrons in its outer s-Orbital, even elements with incomplete 3d orbitals. This is unusual: lower orbitals are usually filled in before the outer shells. The sOrbitals in the d-Block elements are in a lower energy state than the d- lower shells. Since atoms strive to take the lowest possible energy state, the s-Shells filled first. The exceptions for chromium and copper - which only have one electron in their outer orbital - are due to electron repulsion. Splitting the electrons up s- and d-Orbitals leads to lower energy states for the atoms than two electrons in the outer s-Orbital to place.

Not all d-Block elements are transition metals. Scandium and zinc do not fit into the definition given above. Scandium has an electron in its d- Lower shell and 2 electrons in the outer s-Orbital. Since the only scandium ion (Sc3+) no electrons in d-Orbital, it can of course not be a "partially filled" d-Have orbital. The same applies to zinc, since its only ion, Zn2+, a completely filled one d-Has orbital.

Chemical properties

Transition elements are generally characterized by high tensile strengths, densities, melting points and boiling points. Just like other properties of transition metals, these too are dependent on the ability of the electrons d-Orbitals are attributed to being delocalized within the metal lattice. In metallic materials, these properties are more pronounced the more electrons are divided between the nuclei.

There are four typical properties of transition metals:

Oxidation states

Compared to elements of group II such as calcium, the ions of the transition elements exist in numerous oxidation states. Calcium ions usually lose no more than two electrons, whereas transition elements can donate up to nine. 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 get electrons below its outer two s-Remove orbitals. Approx3+ has an enthalpy of ionization so high that it does not normally occur. Transition elements like vanadium, on the other hand, have fairly linearly increasing ionization enthalpies along them because of the small energy difference between the 3d and 4s orbitals s- and d-Orbitals. Transition elements therefore also occur with very high oxidation numbers.

Certain behavior patterns can be seen along a period:

  • The number of oxidation states increases 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 6 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 only recently confirmed by the representation of the oxidation state + IV for mercury as HgF4 approved.

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 often used for the catalytic linkage of C-C bonds (Suzuki, Heck, Stille etc.). Rhodium (Rh), iridium (Ir) and ruthenium (Ru) are z. 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, i.e. have two phosphorus atoms that "bind" to the transition metal like tongs.

Colored connections

When the frequency of electromagnetic radiation changes, we perceive different colors. 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 in the same element - MnO4 (Mn in oxidation state +7) is purple compound, Mn2+ but is pale pink.

Complex formation can play an important role in coloring. The ligands have a great influence on the 3d shell. They partially attract the 3d 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 3d orbitals, some electrons are lifted into the higher state. 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, specifically 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 3sd-Groups.

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


  • Sebastian Riedel, Martin Kaupp, "Revising the highest oxidation states of the 5th elements: The case of iridium (+ VII)" Angew. Chem. 2006, 118, 3791-3794 (Angew. Chem. Int. Ed. 2006, 45, 3708 - 3711)

Category: transition metal