6 7s 2 rn5f 7 7s 2 rn5f 7 6d 1 7s 2 rn5f 9 7s 2 rn5f

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6 7s 2 [Rn]5f 7 7s 2 [Rn]5f 7 6d 1 7s 2 [Rn]5f 9 7s 2 [Rn]5f 10 7s 2 [Rn]5f 11 7s 2 [Rn]5f 12 7s 2 [Rn]5f 13 7s 2 [Rn]5f 14 7s 2
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Chemistry of the transition metals I Electronegativities range from 1.1 (La, Ac) to 2.4 (W, Au) I Many TM compounds are essentially covalent I Often act as Lewis acids I Rough rule: TMs form ionic compounds in lower oxidation states, covalent compounds in higher oxidation states
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I Common and less common oxidation states: 1 Ti V Cr Mn Fe Co Ni Cu Zn [Ar]3d 10 4s 2 +7 +6 +5 +4 +3 +2 +1 [Ar]3d 1 4s 2 [Ar]3d 3 4s 2 [Ar]3d 2 4s 2 [Ar]3d 5 4s 1 [Ar]3d 5 4s 2 [Ar]3d 6 4s 2 [Ar]3d 7 4s 2 [Ar]3d 8 4s 2 [Ar]3d 10 4s Sc
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Naming ionic compounds of the transition metals I Because transition metals can have different oxidation states/ionic charges, we need to name this oxidation state in order to specify an ion. This is done by adding the charge as a Roman numeral glued onto the name of an ion. Examples: iron(II) ion for Fe 2+ iron(III) ion for Fe 3+ I Ionic compounds are then named in the usual way, by naming the cation first, and then the anion. Examples: iron(II) chloride FeCl 2 iron(III) oxide for Fe 2 O 3
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Complexes and ligands Coordinate bond: bond made between a metal ion and a Lewis base
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