Properties of First Row Transition Metal Ions and Salts
Now given in more detail in individual pages covering each Transition Metal
Ti V Cr Mn Fe Co Ni Cu


Titanium(IV) Halides
Formula Colour MP BP Structure
TiF4 white -  284 fluoride bridged 
TiCl4 colourless -24 136.5 -
TiBr4 yellow  38  233.5 hcp I- but essentially monomeric cf. SnI4
TiI4 violet-black 155 377 hcp I- but essentially monomeric cf. SnI4

They can all be prepared by direct reaction of Ti with halogen gas (X2). All are readily hydrolysed.
They are all expected to be diamagnetic.

Titanium(III) halides
Formula Colour MP BP μ (BM) Structure
TiF3 blue 950d - 1.75 -
TiCl3 violet 450d - - BiI3
TiBr3 violet - - - BiI3
TiI3 violet-black - - - -

They can be prepared by reduction of TiX4 with H2.

Vanadium(V) halides
Formula Colour MP BP μ (BM) Structure
VF5 white 19.5 48.3 0 trigonal bipyramid in gas phase

Prepared by reaction of V with F2 in N2 or with BrF3 at 300C.
In the solid state it is an infinite chain polymer with cis-fluoride bridging.

Vanadium(IV) halides
Formula Colour MP BP μ (BM) Structure
VF4 lime-green 100 (a) - 1.68 -
VCl4 red-brown  -25.7 148 1.61 tetrahedral (monomeric)
VBr4 purple  -23d  - - -
(a) sublimes with decomposition at 100 C.

VCl4 is prepared by reaction of V with chlorinating agents such as Cl2, SOCl2, COCl2 etc.
Reaction of VCl4 with HF in CCl3F at -78C gives VF4.

Chromium(III) halides
Formula Colour MP M-X (pm) μ (BM) (b) Structure
CrF3 green 1404 190 - -
CrCl3 red-violet 1152 238 - CrCl3
CrBr3 green-black 1130 257 - BiI3
CrI3 black >500d - - -

(b) all 3.7-4.1 BM.


CrX3 are prepared from Cr with X2, dehydration of CrCl3.6H2O requires SOCl2 at 650C.

Chromium(II) halides
Formula Colour MP μ (BM) Structure
CrF2 green 894 4.3 distorted rutile
CrCl2 white 820-824 5.13 distorted rutile
CrBr2 white 844 - -
CrI2 red-brown 868 - -

Reduction of CrX3 with H2/HX gives CrX2.

Manganese(II) halides
Formula Colour MP BP μ (BM) Structure
MnF2 pale-pink 920 - - rutile
MnCl2 pink 652 1190 5.73 CdCl2
MnBr2 rose 695 - 5.82 -
MnI2 pink 613 - 5.88 CdI2

Prepared from MnCO3 + HX -> MnX2 + CO2 + H2O

Iron(III) halides 
Formula Colour MP Structure
FeF3 green 1000 sublimes -
FeCl3 black 306 sublimes BiI3
FeBr3 dark-red-brown - BiI3

Prepared by reaction of Fe + X2 -> FeX3.
Note that when boiled gives FeBr2.

Iron(II) halides 
Formula Colour MP BP Structure
FeF2 white 1000 1100 rutile
FeCl2 pale yellow-grey 670-674 - CdCl2
FeBr2 yellow-green 684 - CdI2
FeI2 grey red heat - CdI2

Fe +HX at red heat -> FeX2 for X=F,Cl and Br
Fe + I2 -> FeI2

Cobalt(II) halides
Formula Colour MP μ (BM) Structure
CoF2 pink 1200 - rutile
CoCl2 blue 724 5.47 CdCl2
CoBr2 green 678 - CdI2
CoI2 blue-black 515 - CdI2

Co or CoCO3 + HX -> -> CoX2


Nickel(II) halides
Formula Colour MP μ (BM) Structure
NiF2 yellow 1450 2.85 tetragonal rutile
NiCl2 yellow 1001 3.32 CdCl2
NiBr2 yellow 965 3.0 CdCl2
NiI2 Black 780 3.25 CdCl2

Ni + F2 55 C /slow -> NiF2
Ni + Cl2 EtOH/ 20 C -> NiCl2
Ni + Br2 red heat -> NiBr2
NiCl2 + 2NaI -> NiI2 + 2NaCl


Copper(II) halides
Formula Colour  MP  BP μ (BM)  Structure
CuF2 white 950d - 1.5  
CuCl2 brown 632 993d 1.75 CdCl2
CuBr2 black 498 - 1.3  

Copper(II) halides are moderate oxidising agents due to the Cu(I)/ Cu(II) couple. In water, where the potential is largely that of the aquo-complexes, there is not a great deal of difference between them, but in non-aqueous media, the oxidising (halogenating) power increases in the sequence:  CuF2 << CuCl2 << CuBr2.

Cu + F2 -> CuF2
Cu + Cl2  450 C -> CuCl2
Cu + Br2 -> CuBr2
or from by heating -> CuX2

Copper(I) halides
Formula Colour MP BP Structure
CuCl white 430 1359 -
CuBr white 483 1345 -
CuI white 588 1293 Zinc Blende

Reduction of CuX2 -> CuX except for F which has not been obtained pure.
Note that Cu(II)I2 can not be isolated due reduction to CuI.

Oxides and Aquo Species

Titanium oxides
Formula Colour MP μ (BM) Structure
TiO2 white  1892 diam. rutile - Refractive Index 2.61-2.90 cf. Diamond 2.42


obtained from hydrolysis of TiX4 or Ti(III) salts.

TiO2 reacts with acids and bases.
In Acid:
TiOSO4 formed in H2SO4 (Titanyl sulfate)
In Base:
MTiO3 metatitanates (eg Perovskite, CaTiO3 and ilmenite, FeTiO3)
M2TiO4 orthotitanates.

Peroxides are highly coloured and can be used for colourimetric analysis.
pH <1         [TiO2(OH)(H2O)x]+
pH 1-2        [(O2)Ti-O-Ti(O2)](OH) x2-x; x=1-6

[Ti(H2O)6]3+ -> [Ti(OH)(H2O)5]2+ + [H+] pK=1.4
TiO2+ + 2H+ + e- -> Ti3+ + H2O E=0.1V


Vanadium oxides
Formula  Colour Common name Oxidation State MP V-O distance (pm)
V2O5 brick-red pentoxide V5+ 658 158.5-202
V2O4 blue dioxide V4+ 1637 176-205
V2O3 grey-black sesquioxide V3+ 1967 196-206

V2O5 is the final product of the oxidation of V metal, lower oxides etc.

Aqueous Chemistry very complex:

In alkaline solution,

VO43- + H+ -> HVO42-
2HVO42- -> V2O74- + H2O
HVO42- + H+ -> H2VO4-
3H2VO4- -> V3O93- + 3H2O
4H2VO4- -> V3O124- + 4H2O
In acidic solution,
10V3O93- + 15H+ -> 3HV10O285- + 6H2O
H2VO4- + H+ -> H2VO4
HV10O285- + H+ -> H2V10O284-
H3VO4 + H+ -> VO2+ + 2H2O
H2V10O284- + 14H+ -> 10VO2+ + 8H2O

The crystal structure of this salt was first determined in 1965. The V=O bond length was 159.4 pm, the aquo group trans to this had the longest V-O bond length (228.4pm) and the equatorial bond lengths were in the range 200.5-205.6 pm. Note that SO42- was coordinated in an equatorial position.
The IR stretching frequency for the V=O in vanadyl complexes generally occurs at 985 +/- 50 cm-1.

Redox properties of oxovanadium ions:

VO2+ + 2H+ + e- -> VO2+ + H2O E=1.0v

VO2+ + 2H+ + e- -> V3+   + H2O E=0.34V


Chromium oxides 
Formula  Colour Oxidation State MP
CrO3 deep red Cr6+ 197d
Cr3O8 - intermediate  -
Cr2O5 - - -
Cr5O12 etc - - -
CrO2 brown-black Cr4+ 300d
Cr2O3 green Cr3+ 2437

Dichromate and chromate equilibria is pH dependent:

HCrO4- -> CrO42- + H+ K=10-5.9
H2CrO4 -> HCrO4- + H+ K=10+0.26

Cr2O72- + H2O -> 2HCrO4- K=10-2.2
HCr2O7- -> Cr2O72- + H+ K=10+0.85
pH > 8 CrO42- yellow
2-6 HCrO4- & Cr2O72- orange-red
< 1 H2Cr2O7

[Cr(H2O)6]3+ -> [Cr(H2O)5(OH)]2+ -> [(H2O)4Cr Cr(H2O)4]4+ pK=4 etc.


Manganese oxides
Formula  Colour Oxidation State MP
Mn2O7 green oil Mn7+ 5.9
MnO2 black Mn4+ 535d
Mn2O3 black Mn3+ 1080d
Mn3O4 - Haussmanite black Mn2/3+ 1705
MnO grey-green Mn2+ 1650


Mn3O4 is prepared from the other oxides by heating in air. MnO is prepared from the other oxides by heating with H2 at temperatures below 1200 C

Redox properties of KMnO4.

  strong base
  MnO4- + e-      →  MnO42-      E=0.56V (RAPID)
  MnO42-  + 2H2O  + e-  →  MnO2  + 4OH-  E=0.60V (SLOW)
  moderate base
  MnO4- + 2H2O  + 3e- →  MnO2  + 4OH-    E=0.59V
  dil. H2SO4
  MnO4- + 8H2O  + 5e- →  Mn2+  + 4H2O    E=1.51V


Iron oxides
Formula  Colour Oxidation State MP Structure / comments
Fe2O3 red brown Fe3+ 1560d α-form Haematite, 
β-form used in cassettes
Fe3O4 black Fe2+/3+ 1538d magnetite/lodestone
FeO black Fe2+ 1380 pyrophoric

α-Fe2O3 is obtained by heating alkaline solutions of Fe(III) and dehydrating the solid formed.
  FeO,Fe3O4, γ-Fe2O3 ccp
  α-Fe2O3   hcp
The Fe(III) ion is strongly acidic:
[Fe(H2O)6]3+    + H2O   -> [Fe(H2O)5(OH)]2+ + H3O+   K=10-3.05
[Fe(OH)(H2O)5]2+ +  H2O -> [Fe(OH)2(H2O)4]+ + H3O+   K=10-3.26
  2Fe(H2O)63+ + 2H2O  ->  [Fe2(OH)2(H2O)8]4++ 2H3O+   K=10-2.91
The Fe2+ ion is barely acidic:
  Fe(H2O)62+  + H2O ->  [Fe(OH)(H2O)5]+ + H3O+    K=10-9.5
The Redox chemistry of Iron is pH dependent:
  Fe(H2O)63+  + e-  ->  Fe(H2O)62+        E=0.771V

  E=E-RT/nF  Ln[Fe2+]/[Fe3+]
  at precipitation
  [Fe2+].[OH-]2   ~ 10-14
  [Fe3+].[OH-]3   ~ 10-36

or for OH- =1M then [Fe2+]/[Fe3+] = 1022

  E =0.771 -0.05916 log10(1022)
    =0.771 -1.301
thus in base the value of E is reversed and the susceptibility of Fe2+ to oxidation increased. In base it is a good reducing agent and will reduce Cu(II) to Cu(0) etc. Note the implications for rust treatment.


Cobalt oxides
Formula  Colour Oxidation State MP Structure / comments
Co2O3   Co3+    
Co3O4 black Co2+/3+ 900-950d normal spinel
CoO olive green Co2+ 1795 NaCl -antiferromag. < 289 K

Co2O3 is formed from oxidation of Co(OH)2.
CoO when heated at 600-700 converts to Co3O4
Co3O4 when heated at 900-950 reconverts back to CoO.

no stable [Co(H2O)6]3+ or [Co(OH)3 exist.
[Co(H2O)6]2+ not acidic


Nickel oxides
Formula  Colour Oxidation State MP Structure / comments
NiO green powder Ni2+ 1955 NaCl
thermal decomposition of Ni(OH)2, NiCO3, or NiNO3 gives NiO.
 [Ni(H2O)6]2+ not acidic


Copper oxides
Formula  Colour Oxidation State MP
CuO black Cu2+ 1026d
Cu2O red Cu+ 1230
[Cu(H2O)6]2+ not acidic


Cu2O is prepared from thermal decomposition of CuCO3, Cu(NO3)2 or Cu(OH)2. The Fehling's test for reducing sugars also gives rise to red Cu2O. It is claimed that 1 mg of dextrose produces sufficient red colour for a positive test.

The Redox chemistry of Copper:

    Cu2+    + e-  →  Cu+     E=0.15V
    Cu+     + e-  →  Cu      E=0.52V
    Cu2+    + 2e- →  Cu      E=0.34V
By consideration of this data, it will be seen that any oxidant strong enough to covert Cu to Cu+ is more than strong enough to convert Cu+ to Cu2+ (0.52 cf 0.14V). It is not expected therefore that any stable Cu+ salts will exist in aqueous solution.
Disproportionation can also occur:
     2Cu+    →  Cu2+    + Cu    E=0.37V or K=106

Representative Coordination Complexes


TiCl4 is a good Lewis acid and forms adducts on reaction with Lewis bases such as;

                2PEt3           →      TiCl4(PEt3)2
                2MeCN           →      TiCl4(MeCN)2
                bipy            →      TiCl4(bipy)
Solvolysis can occur if ionisable protons are present in the ligand;
                2NH3            →      TiCl2(NH2)2     +       2HCl
                4H2O            →         +       4HCl
                2EtOH           →      TiCl2(OEt)2     +       2HCl
TiCl3 has less Lewis acid strength but can form adducts also;
                3pyr            →      TiCl3pyr3


The Vanadyl ion (eg. from VO(H2O)4SO4 retains the V=O bond when forming complexes.
                VO2+    +       2acac           →      VO(acac)2
Vanadyl complexes are often 5 coordinate square pyramidal and are therefore coordinately unsaturated. They can take up another ligand to become octahedral, eg;
                VO(acac)2       +       pyr     →      VO(acac)2pyr
The V=O stretching frequency in the IR can be monitored to see the changes occurring during these reactions. It generally is found at 985 cm-1 but will shift to lower wavenumbers when 6-coordinate, since the bond becomes weaker.


The Chromium(III) ion forms many stable complexes which being inert are capable of exhibiting various types of isomerism. "CrCl3.6H2O" exists as hydrate isomers, including:

                        trans-[Cr(H2O)4Cl2]Cl.2H2O etc
        CrCl3 anhydrous reacts with pyridine only in the presence of Zinc powder. This allows a small amount of Cr(II) to be formed, which is very labile.
                CrCl3           +       pyr/Zn  →      CrCl3pyr3

[Cr2(OAc)4].2H2O is an example of a Cr(II) complex which is reasonably stable in air once isolated. Each Cr(II) ion has 4 d electrons but the complex is found to be diamagnetic which is explained by the formation of a quadruple bond between the two metal ions. The Cr-Cr bond distance in a range of these quadruply bonded species has been found to vary between 195-255 pm.


Octahedral complexes of Mn(III) are expected to show Jahn-Teller distortions. It was of interest therefore to compare the structures of Cr(acac)3 with Mn(acac)3 since the Cr(III) ion is expected to give a regular octahedral shape. In fact the Mn-O bond distances were all found to be equivalent.

An unusual Mn complex is obtained by the reaction of Mn(OAc)2 with KMnO4 in HOAc. This gives [MnO(OAc)6 3H2O] OAc. It is used as an industrial oxidant for the conversion of toluene to phenol.


An important Fe complex which is used in Actinometry since it is photosensitive is K3[Fe(C2 O4)3.3H2O.
It can be prepared from:
Fe(C2O4) in K2C2O4 by reacting with H2O2 in H2C2O4 to give green crystals. It is high spin m =5.9 BM at 300K and has been resolved into its two optical isomers, although they racemise in less than 1 hour.

In light the reaction is:

        K3Fe(C2O4)3.3H2O        →      2Fe(C2O4)       +       2CO2            +       3K2C2O4
        Another important complex is used as a redox indicator since the Fe(II) and Fe(III) complexes are both quite stable and have different colours:
        Fe(phen)33+     +       e-      →      Fe(phen)32+                     E=1.12V
        blue                                            red
The ligand is 1,10 phenanthroline and the indicator is called ferroin.


The Cobalt(III) ion forms many stable complexes, which being inert, are capable of exhibiting various types of isomerism. The preparation and characterisation of many of these complexes dates back to the pioneering work of Werner and his students.
        Coordination theory was developed on the basis of studies of complexes of the type:

Werner Complexes
[Co(NH3)6]Cl3 yellow
[CoCl(NH3)5]Cl2 red
trans-[CoCl2(NH3)4]Cl green
cis-[CoCl2(NH3)4]Cl purple

Another important complex in the history of coordination chemistry is HEXOL. This was the first complex that could be resolved into its optical isomers that did not contain Carbon atoms. Since then, only three or four others have been found.

An interesting complex which takes up O2 from the air reversibly is Cosalen. This has been used as an emergency oxygen carrier in jet aircraft.


The Nickel(II) ion forms many stable complexes. Whilst there are no other important oxidation states to consider, the Ni(II) ion can exist in a wide variety of CN's which complicates its coordination chemistry.
For example, for CN=4 both tetrahedral and square planar complexes can be found,
for CN=5 both square pyramid and trigonal bipyramid complexes are formed.
The phrase "anomalous nickel" has been used to describe this behaviour and the fact that equilibria often exist between these forms.
Some examples include:
        (a) addition of ligands to square planar complexes to give 5 or 6 coordinate species
        (b) monomer/polymer equilibria
        (c) square-planar/ tetrahedron equilibria
        (d) trigonal-bipyramid/ square pyramid equilibria.

        L=P(aryl)3              are tetrahedral
        L=P(alkyl)3             are square planar
L= mixed aryl and alkyl phosphines, both stereochemistries can occur in the same crystalline substance. The energy of activation for conversion of one form to the other has been found to be around 50 kJ mol-1. Similar changes have been observed with variation of the X group:
        Ni(PPh3)2Cl2    green   tetrahedral             μ = 2.83 BM
        Ni(PPh3)2(SCN)2 red     sq. planar              μ = 0.
Ni2+ reacts with CN- to give Ni(CN)2.nH2O (blue-green) which on heating at 180-200 is dehydrated to yield Ni(CN)2. Reaction with excess KCN gives K2Ni(CN)4.H2O (orange crystals) which can be dehydrated at 100C. Addition of strong concentrations of KCN produces red solutions of [Ni(CN)5]3-.

The crystal structure of the double salt prepared by addition of [Cr(en)3]3+ to [Ni(CN)5]3- showed that two types of Ni stereochemistry were present in the crystals in approximately equal proportions;
50% as square pyramid and 50% as trigonal bipyramid .


The Copper(II) ion forms many stable complexes which are invariably described as either 4 coordinate or distorted 6 coordinate species.
Cu(OH)2 reacts with NH3 to give a solution which will dissolve cellulose. This is exploited in the industrial preparation of Rayon. The solutions contain tetrammines and pentammines. With pyridine, only tetramines are formed eg Cu(py)4 SO4.
A useful reagent for the analytical determination of Cu2+ is the sodium salt of N,N-diethyldithiocarbamate. In dilute alcohol solutions, the presence of trace levels of Cu2+ is indicated by a yellow colour which can be measured by a spectrometer and the concentration determined from a Beer's Law plot. The complex is Cu(Et2dtc)2 which can be isolated as a brown solid.

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Created June 2000. Links checked and/or last modified 1st November 2006.