21.3: Physical Properties - An Overview

Now given in more detail in individual pages covering each Transition Metal

Halides

Titanium

Preparations:

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.

Preparations:

They can be prepared by reduction of TiX₄ with H₂.

Vanadium

Preparations:

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.

^(a) sublimes with decomposition at 100 C.
Preparations:

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

^(b) all 3.7-4.1 BM.

Preparations:

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

Preparations:

Reduction of CrX₃ with H₂/HX gives CrX₂.

Manganese

Preparations:

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

Iron

Preparations:

Prepared by reaction of Fe + X₂ -> FeX₃.
Note that FeBr₃.aq when boiled gives FeBr₂.

Preparations:

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

Cobalt

Preparations:

Co or CoCO₃ + HX -> CoX₂.aq -> CoX₂

Nickel

Preparations:

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

Copper

Preparations:

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 CuX₂.aq by heating -> CuX₂

Preparations:

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

Preparations:

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 Colored and can be used for Colorimetric analysis.
pH <1 [TiO₂(OH)(H₂O)x]+
pH 1-2 [(O₂)Ti-O-Ti(O₂)](OH) _x^2-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

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

Aqueous Chemistry very complex:

In alkaline solution,

VO₄^3- + H⁺ -> HVO₄^2-
2HVO₄^2- -> V₂O₇^4- + H₂O
HVO₄^2- + H⁺ -> H₂VO₄^-
3H₂VO₄^- -> V₃O₉^3- + 3H₂O
4H₂VO₄^- -> V₃O₁₂^4- + 4H₂O

In acidic solution,

10V₃O₉^3- + 15H⁺ -> 3HV₁₀O₂₈^5- + 6H₂O
H₂VO₄^- + H⁺ -> H₂VO₄
HV₁₀O₂₈^5- + H⁺ -> H₂V₁₀O₂₈^4-
H₃VO₄ + H⁺ -> VO₂⁺ + 2H₂O
H₂V₁₀O₂₈^4- + 14H⁺ -> 10VO₂⁺ + 8H₂O

VO(H₂O)₄SO₄
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 SO₄^2- 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

Dichromate and chromate equilibria is pH dependent:

HCrO₄^- -> CrO₄^2- + H⁺ K=10^-5.9
H₂CrO₄ -> HCrO₄^- + H⁺ K=10^+0.26

Cr₂O₇^2- + H₂O -> 2HCrO₄^- K=10^-2.2
HCr₂O₇^- -> Cr₂O₇^2- + H⁺ K=10^+0.85

CrO₃

pH > 8 CrO₄^2- yellow
2-6 HCrO₄^- & Cr₂O₇^2- orange-red
< 1 H₂Cr₂O₇

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

Manganese

Preparations:

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

Preparations:
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

olation

  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
    =-0.530v

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

Preparations:

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

thermal decomposition of Ni(OH)₂, NiCO₃, or NiNO₃ gives NiO.
[Ni(H₂O)₆]²⁺ not acidic

Copper

[Cu(H₂O)₆]²⁺ not acidic

Preparations:

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 Color 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

Titanium

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            →      TiO2.aq         +       4HCl
                2EtOH           →      TiCl2(OEt)2     +       2HCl

TiCl₃ has less Lewis acid strength but can form adducts also;

                3pyr            →      TiCl3pyr3

Vanadium

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.

Chromium

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.

Manganese

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.

Iron

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 Colors:

        Fe(phen)33+     +       e-      →      Fe(phen)32+                     E=1.12V
        blue                                            red

The ligand is 1,10 phenanthroline and the indicator is called ferroin.

Cobalt

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:

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.

Nickel

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.

• (a) substituted acacs react with Ni²⁺ to give green dihydrates (6 Coord) by heating the waters are removed to give tetrahedral species. The unsubstituted acac complex, Ni(acac)₂ normally exists as a trimer.
  Lifschitz salts containing substituted ethylenediamines can be isolated as either 4 or 6 coordinate species depending on the presence of coordinated solvent.
• (b) Ni(acac)₂ is only found to be monomeric at temperatures around 200 C in non-coordinating solvents such as n-decane. 6-coordinate monomeric species are formed at room temperature in solvents such as pyridine but in the solid state Ni(acac)₂ is a trimer, where each Ni atom is 6-coordinate. Note that Co(acac)₂ actually exists as a tetramer.
• (c) Complexes of the type NiL₂X₂ where L are phosphines can give rise to either tetrahedral or square planar complexes. It has been found that:

        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 .

Copper

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 Color 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.