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Vanadium(V) oxide
CAS number 1314-62-1 Yes check.svgY
PubChem 14814
ChemSpider 14130
EC number 215-239-8
UN number 2862
RTECS number YW2450000
Molecular formula V2O5
Molar mass 181.88 g/mol
Appearance Orange-yellow solid
Density 3.357 g/cm3
Melting point

690 °C (963 K)

Boiling point

1750 °C decomp.

Solubility in water 0.8 g/100 mL (20 °C)
Crystal structure Orthorhombic
Space group Pmmn, No. 59
Lattice constant a = 1151 pm, b = 355.9 pm, c = 437.1 pm
Distorted trigonal bipyramidal (V)
EU Index 023-001-00-8
EU classification Muta. Cat. 3
Repr. Cat. 3
Toxic (T)
Harmful (Xn)
Irritant (Xi)
Dangerous for the environment (N)
R-phrases R20/22, R37, R48/23, R51/53, R63, R68
S-phrases (S1/2), S36/37, S38, S45, S61
NFPA 704
NFPA 704.svg
Flash point Non-flammable
LD50 10 mg/kg
Related compounds
Other anions Vanadium oxytrichloride
Other cations Niobium(V) oxide
Tantalum(V) oxide
Related vanadium oxides Vanadium(II) oxide
Vanadium(III) oxide
Vanadium(IV) oxide
 Yes check.svgY (what is this?)  (verify)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Vanadium(V) oxide (vanadia) is the chemical compound with the formula V2O5. Commonly known as vanadium pentoxide, this orange solid is the most important compound of vanadium. Upon heating it reversibly loses oxygen. Related to this ability, V2O5 catalyses several useful aerobic oxidation reactions, the largest scale of which underpins the production of sulfuric acid from sulfur dioxide. It is a poisonous orange solid which, because of its high oxidation state, is both an amphoteric oxide and an oxidising agent. Unlike most metal oxides, it dissolves slightly in water to give a pale yellow, acidic solution.

The mineral form of this compound, shcherbinaite,[2] is extremely rare, almost always found among fumaroles. A mineral trihydrate, V2O5·3H2O, is also known under the name of navajoite.[3]


Chemical properties


Acid-base reactions

V2O5 is an amphoteric oxide. Thus it reacts with strong non-reducing acids to form solutions containing the pale yellow salts containing dioxovanadium(V) centers:

V2O5 + 2HNO3 → 2VO2(NO3) + H2O

It also reacts with strong alkali to form polyoxovanadates, which have a complex structure that depends on pH.[4] If excess aqueous sodium hydroxide is used, the product is a colourless salt, sodium orthovanadate, Na3VO4. If acid is slowly added to a solution of Na3VO4, the colour gradually deepens through orange to red before brown hydrated V2O5 precipitates around pH 2. These solutions contain mainly the ions HVO42− and V2O74− between pH 9 and pH 13, but below pH 9 more exotic species such as V4O124− and HV10O285− predominate.

Thionyl chloride converts it to VOCl3:

V2O5 + 3SOCl2 → 2VOCl3 + 3SO2

Redox reactions

V2O5 is easily reduced in acidic media to the stable vanadium(IV) species, the blue vanadyl ion (VO(H2O)52+). This conversion illustrates the redox properties of V2O5. For example, hydrochloric acid and hydrobromic acid are oxidised to the corresponding halogen, e.g.,

V2O5 + 6HCl + 7H2O → 2[VO(H2O)5]2+ + 4Cl + Cl2

Solid V2O5 is reduced by oxalic acid, carbon monoxide, and sulfur dioxide to give vanadium(IV) oxide, VO2 as a deep-blue solid. Further reduction using hydrogen or excess CO can lead to complex mixtures of oxides such as V4O7 and V5O9 before black V2O3 is reached. Vanadates or vanadyl(V) compounds in acid solution are reduced by zinc amalgam through the interestingly colorful pathway:

VO2+ VO2+ V3+ V2+
yellow   blue   green   purple

The ions are, of course, all hydrated to varying degrees.


Technical grade V2O5 is produced as a black powder used for the production of vanadium metal and ferrovanadium.[4] A vanadium ore or vanadium-rich residue is treated with sodium carbonate to produce sodium metavanadate, NaVO3. This material is then acidified to pH 2–3 using H2SO4 to yield a precipitate of "red cake" (see above). The red cake is then melted at 690 °C to produce the crude V2O5.

Vanadium(V) oxide is also the main product when vanadium metal is heated with excess oxygen, but this product is contaminated with other lower oxides. A more satisfactory laboratory preparation involves the decomposition of ammonium metavanadate at around 200 °C:

2NH4VO3 → V2O5 + 2NH3 + H2O


Ferrovanadium production

In terms of quantity, the major use for vanadium(V) oxide is in the production of ferrovanadium (see above). The oxide is heated with scrap iron and ferrosilicon, with lime added to form a calcium silicate slag. Aluminium may also be used, producing the iron-vanadium alloy along with alumina as a by-product.[4] In 2005 a shortage of V2O5 caused a price rise to around $40/kg, which in turn caused a rise in the price of ferrovanadium.

Sulfuric acid production

Another important use of vanadium(V) oxide is in the manufacture of sulfuric acid, an important industrial chemical with an annual production of 165 million metric tons in 2001, with an approximate value of US$8 billion. Vanadium(V) serves the crucial purpose of catalysing the mildly exothermic oxidation of sulfur dioxide to sulfur trioxide by air in the contact process:

2SO2 + O2 is in equilibrium with 2SO3

The discovery of this simple reaction, for which V2O5 is the most effective catalyst, allowed sulfuric acid to become the cheap commodity chemical it is today. The reaction is performed between 400 and 620 °C; below 400 °C the V2O5 is inactive as a catalyst, and above 620 °C it begins to break down. Since it is known that V2O5 can be reduced to VO2 by SO2, one likely catalytic cycle is as follows:

SO2 + V2O5 → SO3 + 2VO2

followed by

2VO2 +½O2 → V2O5

Paradoxically, it is also used as catalyst in the selective catalytic reduction (SCR) of NOx emissions in some power plants. Due to its effectiveness in converting sulfur dioxide into sulfur trioxide, and thereby sulfuric acid, special care must be taken with the operating temperatures and placement of a power plant's SCR unit when firing sulfur-containing fuels.

Other oxidations

Maleic anhydride is another important industrial material, used for the manufacture of polyester resins and alkyd resins.[5] Vanadium(V) oxide can catalyse its production from a variety of organic starting materials such as n-butane, furfural and benzene, the last of which is the usual commercial method. In a related process, phthalic anhydride, used for making plasticisers for PVC manufacture, may be produced by V2O5 catalysed oxidation of ortho-xylene or naphthalene at 350–400 °C.

Other applications

Due to its high thermal coefficient of resistance, vanadium(V) oxide finds use as a detector material in bolometers and microbolometer arrays for thermal imaging. It also find an application as ethanol sensor in ppm levels (up to 0.1 ppm).

Possible new uses include the preparation of bismuth vanadate ceramics for use in solid oxide fuel cells.[6] Another new application is in vanadium redox batteries, a type of flow battery used for energy storage, including large power facilities such as wind farms.

Biological activity


Despite being highly toxic in humans, vanadium occurs in some organisms, notably the Ascidiacea (sea squirts). These organisms contain the protein vanabins, the role of which is unclear. Vanadate (VO43−), formed by hydrolysis of V2O5 at high pH, appears to inhibit enzymes that process phosphate (PO43−). However the exact mode of action remains elusive.[4]


  1. ^ Weast, Robert C., ed. (1981), CRC Handbook of Chemistry and Physics (62nd ed.), Boca Raton, FL: CRC Press, p. B-162, ISBN 0-8493-0462-8 .
  2. ^ Shcherbinaite,,, retrieved 2009-11-30 .
  3. ^ Navajoite,,, retrieved 2009-11-30 .
  4. ^ a b c d Greenwood, Norman N.; Earnshaw, A. (1984), Chemistry of the Elements, Oxford: Pergamon, pp. 1140, 1144, ISBN 0-08-022057-6 .
  5. ^ Tedder, J. M.; Nechvatal, A.; Tubb, A. H., eds. (1975), Basic Organic Chemistry: Part 5, Industrial Products, Chichester, UK: John Wiley & Sons .
  6. ^ Vaidhyanathan, B.; Balaji, K.; Rao, K. J. (1998), "Microwave-Assisted Solid-State Synthesis of Oxide Ion Conducting Stabilized Bismuth Vanadate Phases", Chem. Mater. 10 (11): 3400–4, doi:10.1021/cm980092f 

Further reading

External links


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