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With so-called "strong field ligands" such as cyanide, the five electrons pair up as best they can. Thus ferricyanide ([Fe(CN) 6] 3− has only one unpaired electron. It is low-spin. With so-called "weak field ligands" such as water, the five electrons are unpaired. Thus aquo complex ([Fe(H 2 O) 6] 3+ has only five unpaired electrons. It is ...
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Its 26 electrons are arranged in the configuration [Ar]3d 6 4s 2, of which the 3d and 4s electrons are relatively close in energy, and thus a number of electrons can be ionized. [ 17 ] Iron forms compounds mainly in the oxidation states +2 ( iron(II) , "ferrous") and +3 ( iron(III) , "ferric").
Although Fe(III) chloride can be octahedral or tetrahedral (or both, see structure section), all of these forms have five unpaired electrons, one per d-orbital. The high spin d 5 electronic configuration requires that d-d electronic transitions are spin forbidden , in addition to violating the Laporte rule .
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Complexes with zero to two unpaired electrons are considered low-spin and those with four or five are considered high-spin. [ 12 ] Iron(II) complexes are less stable than iron(III) complexes but the preference for O -donor ligands is less marked, so that for example [Fe(NH 3 ) 6 ] 2+ is known while [Fe(NH 3 ) 6 ] 3+ is not.
Iron(II) is a d 6 center, meaning that the metal has six "valence" electrons in the 3d orbital shell. The number and type of ligands bound to iron(II) determine how these electrons arrange themselves. With the so-called "strong field ligands" such as cyanide, the six electrons pair up.
Therefore, addition or removal of electron has little effect on complex stability. In this case, there is no restriction on the number of d-electrons and complexes with 12–22 electrons are possible. Small Δ oct makes filling e g * possible (>18 e −) and π-donor ligands can make t 2g antibonding (<18 e −).