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transition metal chemistry of vanadium complexes oxidation states +2 +3 +4 +5 redox chemical reactions physical properties advanced inorganic chemistry of vanadium

Revision notes 3d block Transition Metals Vanadium for A Level Inorganic Chemistry

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Periodic Table - Transition Metals 3d block Vanadium Chemistry - Doc Brown's Chemistry  Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Part 10. Transition Metals 3d–block:  

10.5 Vanadium Chemistry

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All my periodic table (3d-block) advanced level chemistry revision study notes

All my advanced A level inorganic chemistry revision study notes

GCSE Level Notes on Transition Metals (for the basics)

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Vanadium exhibits oxidation states of +2, +3, +4 and +5.

The principal oxidation states of vanadium are described via their redox reactions of vanadium, ligand substitution displacement reactions of vanadium, balanced equations of vanadium chemistry, formula of vanadium complex ions, shapes colours of vanadium complexes, formula of compounds

10.5. Chemistry of Vanadium V, Z=23, 1s22s22p63s23p63d34s2 

data comparison of vanadium with the other members of the 3d–block and transition metals

Z and symbol 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn
property\name scandium titanium vanadium chromium manganese iron cobalt nickel copper zinc
melting point/oC 1541 1668 1910 1857 1246 1538 1495 1455 1083 420
density/gcm–3 2.99 4.54 6.11 7.19 7.33 7.87 8.90 8.90 8.92 7.13
atomic radius/pm 161 145 132 125 124 124 125 125 128 133
M2+ ionic radius/pm na 90 88 84 80 76 74 72 69 74
M3+ ionic radius/pm 81 76 74 69 66 64 63 62 na na
common oxidation states +3 only +2,3,4 +2,3,4,5 +2,3,6 +2,3,4,6,7 +2,3,6 +2,3 +2,+3 +1,2 +2 only
outer electron config.[Ar]... 3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
Elect. pot. M(s)/M2+(aq) na –1.63V –1.18V –0.90V –1.18V –0.44V –0.28V –0.26V +0.34V –0.76V
Elect. pot. M(s)/M3+(aq) –2.03V –1.21V –0.85V –0.74V –0.28V –0.04V +0.40 na na na
Elect. pot. M2+(aq)/M3+(aq) na –0.37V –0.26V –0.42V +1.52V +0.77V +1.87V na na na

Elect. pot. = standard electrode potential data for vanadium (EØ at 298K/25oC, 101kPa/1 atm.)

na = data not applicable to vanadium

Extended data table for VANADIUM

property of vanadium/unit value for V
melting point V/oC 1910
boiling point V/oC 3380
density V/gcm–3 6.11
1st Ionisation Energy V/kJmol–1 650
2nd IE/kJmol–1 1414
3rd IE/kJmol–1 2828
4th IE/kJmol–1 4507
5th IE/kJmol–1 6294
atomic radius V/pm 132
V2+ ionic radius/pm 88
Relative polarising power M2+ ion 2.3
V3+ ionic radius/pm 74
Relative polarising power V3+ ion 4.1
V4+ ionic radius/pm 60
Polarising power V4+ ion 6.7
oxidation states of V, less common/stable +2, +3, +4, +5
simple electron configuration of V 2,8,11,2
outer electrons of V [beyond argon core] [Ar]3d34s2
Electrode potential V(s)/V2+(aq) –1.18V
Electrode potential V(s)/V3+(aq) –0.85V
Electrode potential V2+(aq)/V3+(aq) –0.26V
Electrode potential [VO]2+(aq)/V3+(aq) +0.34
Electronegativity of V 1.63


  • Uses of VANADIUM

    • Vanadium is one of many transition metals alloyed with iron to make specialist steels.

  • Vanadium(V) oxide, V2O5, is used as a heterogeneous catalyst in the 'Contact Process' in the production of sulfur trioxide for the manufacture of sulfuric acid.

    • The catalysing of the conversion of sulfur dioxide into sulfur trioxide is explained via change in oxidation state changes i.e. some classic transition metal chemistry.

    • 2SO2(g) +  O2(g) ===> 2SO3(g) 

    • The mechanism, somewhat simplified, goes via the catalytic cycle ...

      • (i) SO2 + V2O5 ===> SO3 + V2O4, then (ii) V2O4 + 1/2O2 ===> V2O5 

      • Adding (i) + (ii) gives (iii) SO2(g)1/2O2(g) ===> 2SO3(g)  (the summing up of the catalytic cycle)

      • This is an example of heterogeneous catalysis of a vanadium compound to reduce the activation energy by providing an alternative pathway via the active sites on the surface of the catalyst.

      • The vanadium changes oxidation state from +5 to +4 and back to +5 in the catalytic cycle, a classic combination of two characteristics of transition metals – variable oxidation state and catalytic properties.

      • This is an example of heterogeneous catalysis – reactants (g) and catalyst (s) in different phases.

The Chemistry of VANADIUM

Pd s block d blocks (3d block vanadium) and f blocks of metallic elements p block elements
Gp1 Gp2 Gp3/13 Gp4/14


2 3Li 4Be Part of the modern Periodic Table of Elements: ZSymbol, z = atomic or proton number

Sc to Zn are now considered the head-top elements of groups 3 to 12

3d block of metallic elements: Scandium to Zinc focus on vanadium

5B 6C
3 11Na 12Mg 13Al 14Si
4 19K 20Ca 21Sc







 [Ar] 3d34s2



[Ar] 3d54s1



   [Ar]   3d54s2



[Ar] 3d64s2



[Ar] 3d74s2



[Ar] 3d84s2



[Ar] 3d104s1



[Ar] 3d104s2


31Ga 32Ge
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn
6 55Cs 56Ba 57,58-71 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 81Tl 82Pb
7 87Fr 88Ra 89,90-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn 113Nh 114Fl

Summary of oxidation states of the 3d block metals (least important) Ti to Cu are true transition metals

Sc Ti V Cr Mn Fe Co Ni Cu Zn
  (+2) +2 (3d3) (+2) +2 +2 +2 +2 +2 +2
+3 +3 +3  (3d2) +3 (+3) +3 +3 (+3) (+3)  
  +4 +4  (3d1)   +4     (+4)    
    +5  (3d0)              
      +6 (+6) (+6)        
3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
The outer electron configurations beyond [Ar] and the (ground state of the simple ion)

Note that when 3d block elements form ions, the 4s electrons are 'lost' first.

The oxidation states and electron configuration of vanadium in the context of the 3d block of elements

electrode potential chart diagram of vanadium ions vanadium(V) ion VO22+ Vanadium(IV) ion VO2+ vanadium(III) ion V3+ vanadium(II) ion V2+

The electrode potential chart highlights the values for various oxidation states of vanadium.

The electrode potentials involving chromium ions correspond to hydrated complex ions where the ligands are water, oxide or hydroxide.

As you can see from the chart, changing either the ligand or the oxidation state, will also change the electrode potential for that half-reaction involving a vanadium ion.

Vanadium(II) compounds are readily oxidised to vanadium(III) and vanadium(IV) compounds.

The hexaaquavanadium(II) ion is a strong reducing agent.

The variety of vanadium's oxidation states

  • Vanadium shows a 'classic' display of variable oxidation states of varying colours when a solution of e.g. ammonium vanadate(V), is reduced by a zinc/dilute sulfuric acid mixture.

    • You go from the vanadium(V) vanadate(V) ion ==> vanadium(IV) oxovanadate(IV) ion ==> vanadium(III) ion ==> vanadium(II) ion

    • Acidification changes the vanadate(V) ion into the pale yellow oxo–cation VO2+ (oxovanadium(V) ion)

    • VO43–(aq) + 4H+(aq) rev VO2+(aq) + 2H2O(l) [an acid–base reaction, NOT a redox change]

      • Note: Highly charged cations >3+ rarely exist as the simple 'hydrated' tetra or hexa–aqua ion.

      • The theoretical polarising power of the 'central metal ion' is so strong that they form oxocations (see above) or oxyanions e.g.

      • orange dichromate(VI) Cr2O72–, yellow chromate(VI) CrO42–, purple manganate(VII) MnO4 etc.

      • For transition metals they may be coloured even if electronically the theoretical 'central metal ion'  has a noble gas structure e.g. [Ar] in its maximum oxidation state like V(V), Cr(VI) and Mn(VII).

      • These oxyanions are called charge transfer complexes and the theory is beyond pre–university chemistry.

    • Three successive reduction steps then follow to eventually give V2+ ions, shown as half–cell equations:

    • (i) V(V, +5) ==> V(IV, +4): VO2+(aq) + 2H+(aq) + e rev VO2+(aq) + H2O(l)

      • EØhalf–cell potential = +1.00V, pale yellow to the blue oxovanadium(IV) ion

    • diagram of the octahedral shape of the aqueous green hexaaquavanadium(III) ion V3+(aq) [V(H2O)6]3+(ii) V(IV, +4) ==> V(III, +3): VO2+(aq) + 2H+(aq) + e rev V3+(aq) + H2O(l)

      • EØhalf–cell potential = +0.34V, blue to the green vanadium(III) ion

      • Here the vanadium(III) ion, V3+, is actually the green hexaaquavanadium(III) ion,

        • Electron configuration of V3+ is [Ar]3d2

        • e.g. in the ion [V(H2O)6]3+

      • Both V(IV) and V(III) species are slowly oxidised by dissolved oxygen back to the V(V) compound in acid solution.

        • (see electrode potential comments later).

    • diagram of the octahedral shape of the aqueous purple-violet coloured hexaaquavanadium(II) ion V2+(aq) [V(H2O)6]3+(iii) V(III, +3) ==> V(II, +2): V3+(aq) + e rev V2+(aq)

      • EØhalf–cell potential = –0.26V, green to the purple–violet vanadium(II) ion.

      • V2+(aq) is powerful reducing agent and is unstable in the presence of air.

      • Any dissolved oxygen will oxidise V2+(aq) back to the vanadium(III) cation.

      • V2+ is actually the purple–violet hexaaquavanadium(II) ion, [V(H2O)6]2+

    • Note

      1. The standard electrode potential EØZn(s)/Zn2+(aq) is –0.76V, so the reducing power of zinc is sufficient to effect any of the three vanadium oxidation state reduction changes described above.

      2. The reduction occurs on the surface of the zinc metal i.e. the site of electron transfer and you can write the above reductions as fully balanced complete redox equations ...

        • (i) 2VO2+(aq) +  4H+(aq) + Zn(s) ===> 2VO2+(aq) + 2H2O(l) + Zn2+(aq)

          • EØreaction = EØreduction – EØoxidation = +1.00 – (–0.76) = +1.76V

          • The half–cell reaction of the reduction will have the most +ve EØpotential.

        • (ii) 2VO2+(aq) +  4H+(aq) + Zn(s) ===> 2V3+(aq) + 2H2O(l) + Zn2+(aq)

          • EØreaction = +0.34 –(–0.76) = +1.10V

        • (iii) 2V3+(aq) +  Zn(s ===> 2V2+(aq) + Zn2+(aq)

          • EØreaction = –0.26 – (–0.76) = +0.50V

          • BUT the vanadium(II) cation is unstable in the presence of dissolve oxygen in air.

          • 1/2O2(g) +  2H+(aq) + 2e   H2O(l) has a standard electrode potential of +1.23V,

          • so, for the vanadium(II) oxidation reaction ...

          • 1/2O2(g) +  2H+(aq) +  2V2+(aq) ===> 2V3+(aq) + H2O(l)

          • EØreaction = EØreduction – EØoxidation = +1.23 – (–0.26) = +1.49V

          • hence the if left standing open to air, the violet V2+(aq) solution will gradually change to a green V3+(aq) solution and in turn V3+(aq) will revert back to VO2+(aq) in the presence of air because of oxidation by dissolve oxygen unless protected by an inert atmosphere. (see Redox Electrode Potential Chart, V2+/V3+ and V3+/VO2+ potentials are less positive (below) that for O2/H2O/H+ potentials).

      3. You will see hydrogen formed simultaneously from the unavoidable metal–acid reaction.

        • Zn(s) +  2H+(aq) ===> Zn2+(aq) + H2(g)

  • Does vanadium chemistry show an example of disproportionation?

    • This is just a little academic exercise using standard electrode potential data.

    • A disproportionation reaction is where a species in one oxidation state spontaneously and simultaneously changes into two species of different oxidation states – one higher and one lower in oxidation number.

    • Examples: disproportionation in manganese(VI) chemistry and disproportionation in copper(I) chemistry

    • Question: In terms of aqueous ions, is the disproportionation of vanadium(III) into vanadium(II) and vanadium (IV) feasible?

      • (i) VO2+(aq) + 2H+(aq) + 2e rev V3+(aq) + H2O(l)   (EØVO2+/V3+ = +0.34V)

      • (ii) V3+(aq) + e rev V2+(aq)   (EØV3+/V2+ = –0.26V)

      • The disproportionation equation would be (iii) 2V3+(aq) + H2O(l) rev V2+(aq) + VO2+(aq) + 2H+(aq)

      • For equation (iii), (ii) will be the reduction half–cell equation and (i) reversed will be the oxidation half–cell reaction.

      • EØreaction = EØreduction – EØoxidation = = EØV3+/V2+ – EØVO2+/V3+ = (–0.26) – (+0.34) = –0.60V

      • showing the disproportionation is thermodynamically NOT feasible i.e. EØreaction is less than zero.

      • In fact what can actually happen is if you mix salt solutions of vanadium(IV) and vanadium(II) on an equimolar basis, you end up with a solution of vanadium(III) salts, a sort of 'anti–disproportionation' reaction!

  • Summary of some complexes–compounds & oxidation states of vanadium compared to other 3d–block elements

keywords redox reactions ligand substitution displacement balanced redox equations formula of vanadium hexaaqua complex ions complexes ligand exchange reactions redox reactions ligands colours oxidation states: all the colour changes when a vanadium(V) salt is reduced in aqueous solution, vanadium ions V2+ V(+2) V(II) V3+ V(+3) V(III) V4+ V(+4) V(IV) V5+ V(+5) (V) SO2 + V2O5 ==> SO3 + V2O4 + 1/2 O2 ==> V2O5 VO43– + 4H+ VO2+ + 2H2O V(V, +5) ==> V(IV, +4): VO2+ + 2H+ + e– VO2+ + H2O V(IV, +4) ==> V(III, +3): VO2+ + 2H+ + e– V3+ + H2O [V(H2O)6]3+ V(III, +3) ==> V(II, +2): V3+ + e– V2+ VO3+/VO2+ (+1.00V), VO2+/V3+ (+0.34V) and V3+/V2+ (–0.26V) 2 VO2+ + 4H+ + Zn ==> 2 VO2+ + 2H2O + Zn2+ 2VO2+ + 4H+ + Zn ==> 2V3+ + 2H2O + Zn2+ 2V3+ + Zn(s ==> 2V2+ + Zn2+ 1/2O2 + 2H+ + 2V2+ ==> 2V3+ + H2O V2+/V3+ and V3+/VO2+ potentials VO2+ + 2H+ + 2e– V3+ + H2OEØVO2+/V3+ = +0.34V) (ii) V3+ + e– V2+ (EØ V3+/V2+ = –0.26V) EØ V3+/V2+ – EØ VO2+/V3+ oxidation states of vanadium, redox reactions of vanadium, ligand substitution displacement reactions of vanadium, balanced equations of vanadium chemistry, formula of vanadium complex ions, shapes colours of vanadium complexes how to work out redox reactions of vanadium using electrode potentials and half-reactions to test for feasibility of a vanadium redox reaction


GCSE Level Notes on Transition Metals (for the basics)

The chemistry of Scandium * Titanium * Vanadium * Chromium * Manganese

The chemistry of Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

Introduction 3d–block Transition Metals * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations Some pages have a matching sub-index

Advanced Level Inorganic Chemistry Periodic Table Index: Part 1 Periodic Table history Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr AND important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots All 11 Parts have their own sub-indexes near the top of the pages

Group numbering and the modern periodic table

The original group numbers of the periodic table ran from group 1 alkali metals to group 0 noble gases. To account for the d block elements and their 'vertical' similarities, in the modern periodic table, groups 3 to group 0 are numbered 13 to 18. So, the p block elements are referred to as groups 13 to group 18 at a higher academic level, though the group 3 to 0 notation is still used, but usually at a lower academic level. The 3d block elements (Sc to Zn) are now considered the head (top) elements of groups 3 to 12.

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