Draft:Potassium molybdate
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 Skeletal formula | |
Ball-and-stick model | |
| Names | |
|---|---|
| IUPAC name
Potassium molybdate | |
Other names
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| Identifiers | |
3D model (JSmol) |
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| EC Number |
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PubChem CID |
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| RTECS number |
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| UNII | |
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| Properties | |
| K2MoO4 | |
| Molar mass | 238.14 g·mol−1 |
| Appearance | white, deliquescent powder |
| Odor | odorless |
| Density | 2.91 g/cm3 |
| Melting point | 919 °C (1,686 °F; 1,192 K) |
| Boiling point | decomposes |
| 184.6 g/100mL (at 25 °C) | |
| Solubility | insoluble in ethanol and diethyl ether |
| Structure | |
| Monoclinic (α-form, space group C2/m) | |
| Tetrahedral (for Mo) | |
| Thermochemistry | |
Heat capacity (C) |
144.8 J/(mol·K) |
Std molar entropy (S⦵298) |
188.3 J/(K·mol) |
Std enthalpy of formation (ΔfH⦵298) |
−1496 kJ/mol |
| Hazards | |
| GHS labelling:[1] | |
 | |
| Warning | |
| H315, H319, H335 | |
| P261, P280, P305+P351+P338 | |
| NFPA 704 (fire diamond) | |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose) |
2400–2680 mg/kg (oral, rat)[2] |
| NIOSH (US health exposure limits): | |
PEL (Permissible) |
5 mg/m3 (as Mo, soluble compounds) |
REL (Recommended) |
5 mg/m3 (as Mo) |
IDLH (Immediate danger) |
1000 mg/m3 (as Mo) |
| Related compounds | |
Other anions |
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Other cations |
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Related compounds |
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Potassium molybdate is an inorganic compound with the chemical formula K2MoO4. It is a white, water-soluble, deliquescent salt composed of potassium cations (K+) and tetrahedral molybdate anions (MoO4^−2).
Potassium molybdate is primarily used as an anodic corrosion inhibitor in closed cooling systems, as a micronutrient fertilizer in agriculture, and as a reagent in chemical analysis for detecting phosphates and silicates. Unlike hexavalent chromium compounds, potassium molybdate exhibits low toxicity and is not classified as a carcinogen.
Production
[edit]Potassium molybdate is produced by dissolving molybdenum trioxide in aqueous potassium hydroxide or potassium carbonate:[3]
![{\displaystyle {\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{3}}{}+{}2\,\mathrm {KOH} {}\mathrel {\longrightarrow } {}\mathrm {K} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{4}}{}+{}\mathrm {H} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {O} }}](https://wikimedia.org/api/rest_v1/media/math/render/svg/000bf1af42728feeb4ff976d5b8ed84888c0e1b8)
![{\displaystyle {\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{3}}{}+{}\mathrm {K} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {CO} {\vphantom {A}}_{\smash[{t}]{3}}{}\mathrel {\longrightarrow } {}\mathrm {K} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{4}}{}+{}\mathrm {CO} {\vphantom {A}}_{\smash[{t}]{2}}\uparrow {}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/b7ff6ef1a6350a40b52e1671baeff9f3ce87bfbb)
Crystallization above 60 °C yields the anhydrous salt. Below this temperature, hydrated forms such as the dihydrate (K2MoO4 . 2H2O) precipitate from solution.[4]
The compound can also be prepared by the oxidative roasting of molybdenite (MoS2) with potassium carbonate in air:
![{\displaystyle {2\,\mathrm {MoS} {\vphantom {A}}_{\smash[{t}]{2}}{}+{}6\,\mathrm {K} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {CO} {\vphantom {A}}_{\smash[{t}]{3}}{}+{}9\,\mathrm {O} {\vphantom {A}}_{\smash[{t}]{2}}{}\mathrel {\longrightarrow } {}2\,\mathrm {K} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{4}}{}+{}4\,\mathrm {K} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {SO} {\vphantom {A}}_{\smash[{t}]{4}}{}+{}6\,\mathrm {CO} {\vphantom {A}}_{\smash[{t}]{2}}\uparrow {}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/f8192a88c454d9fd495cf0f610b55892976b0d5d)
Structure and properties
[edit]At room temperature, potassium molybdate crystallizes in the monoclinic crystal system (space group C2/m). The structure consists of discrete, isolated tetrahedral [MoO4]^−2 anions coordinated to potassium ions (K+) in irregular polyhedra with coordination numbers ranging between 8 and 10.[5]
Unlike potassium chromate, which exhibits an intense yellow color due to low-energy ligand-to-metal charge transfer (LMCT) bands in the visible spectrum, potassium molybdate is colorless because its corresponding charge-transfer absorption bands occur entirely in the ultraviolet region.[4]
Upon heating above 460 °C, anhydrous potassium molybdate undergoes a reversible solid-state phase transition to a hexagonal polymorph isostructural with high-temperature potassium sulfate (α−K2SO4).[6]
Reactions
[edit]Condensation to isopolymolybdates
[edit]In neutral or basic solutions, potassium molybdate exists as discrete tetrahedral [MoO4]^−2 ions. When acidified below pH 6, the molybdate ions condense into edge- and corner-sharing octahedra, forming isopolymolybdate species such as the heptamolybdate (paramolybdate) anion:[7]
![{\displaystyle {7\,[\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{4}}]{\vphantom {A}}^{2-}{}+{}8\,\mathrm {H} {\vphantom {A}}^{+}{}\mathrel {\longrightleftharpoons } {}[\mathrm {Mo} {\vphantom {A}}_{\smash[{t}]{7}}\mathrm {O} {\vphantom {A}}_{\smash[{t}]{24}}]{\vphantom {A}}^{6-}{}+{}4\,\mathrm {H} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {O} }}](https://wikimedia.org/api/rest_v1/media/math/render/svg/65649515af40ea2f5f33dd133596bfb8487fe269)
Further acidification below pH 2 favors the formation of octamolybdate anions:
![{\displaystyle {[\mathrm {Mo} {\vphantom {A}}_{\smash[{t}]{7}}\mathrm {O} {\vphantom {A}}_{\smash[{t}]{24}}]{\vphantom {A}}^{6-}{}+{}[\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{4}}]{\vphantom {A}}^{2-}{}+{}4\,\mathrm {H} {\vphantom {A}}^{+}{}\mathrel {\longrightleftharpoons } {}[\mathrm {Mo} {\vphantom {A}}_{\smash[{t}]{8}}\mathrm {O} {\vphantom {A}}_{\smash[{t}]{26}}]{\vphantom {A}}^{4-}{}+{}2\,\mathrm {H} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {O} }}](https://wikimedia.org/api/rest_v1/media/math/render/svg/950bfdf1e97a065dda57e0cbc1e3f68a6b72e43a)
Redox behavior and molybdenum blue
[edit]Potassium molybdate is resistant to reduction compared to chromates. Reduction of acidified aqueous solutions with mild reducing agents such as ascorbic acid, tin(II) chloride, or sodium sulfite yields mixed-valence molybdenum(V,VI) oxide clusters referred to as molybdenum blue. These species exhibit intense absorption in the near-infrared and red visible spectrum due to intervalence charge transfer.[4]
Heteropoly complex formation
[edit]Acidified potassium molybdate reacts with oxyanions of phosphorus, silicon, or arsenic to form Keggin-type heteropoly compounds. In the presence of orthophosphate, it yields the yellow 12-molybdophosphate anion:[7]
![{\displaystyle {\mathrm {HPO} {\vphantom {A}}_{\smash[{t}]{4}}{\vphantom {A}}^{2-}{}+{}12\,[\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{4}}]{\vphantom {A}}^{2-}{}+{}23\,\mathrm {H} {\vphantom {A}}^{+}{}\mathrel {\longrightarrow } {}[\mathrm {PMo} {\vphantom {A}}_{\smash[{t}]{12}}\mathrm {O} {\vphantom {A}}_{\smash[{t}]{40}}]{\vphantom {A}}^{3-}{}+{}12\,\mathrm {H} {\vphantom {A}}_{\smash[{t}]{2}}\mathrm {O} }}](https://wikimedia.org/api/rest_v1/media/math/render/svg/333606e3a56b0b68fb19e6a64dc0114d07a6a2cc)
This reaction is the basis for the photometric determination of phosphorus in analytical chemistry.
Applications
[edit]Corrosion inhibition
[edit]Potassium molybdate acts as an anodic corrosion inhibitor in closed-loop cooling and heating water systems. It functions by promoting the passivation of ferrous metals, forming a protective surface film of iron(III) molybdate (Fe2(MoO4)3):[8]
![{\displaystyle {2\,\mathrm {Fe} {\vphantom {A}}^{3+}{}+{}3\,[\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{4}}]{\vphantom {A}}^{2-}{}\mathrel {\longrightarrow } {}\mathrm {Fe} {\vphantom {A}}_{\smash[{t}]{2}}(\mathrm {MoO} {\vphantom {A}}_{\smash[{t}]{4}}){\vphantom {A}}_{\smash[{t}]{3}}\downarrow {}}}](https://wikimedia.org/api/rest_v1/media/math/render/svg/421d1b2e951f74d055c51cfcc785335ff750dfa6)
Because of its lower toxicity, molybdate is commonly used as a replacement for chromate-based corrosion inhibitors.
Analytical chemistry
[edit]Potassium molybdate is used in spectrophotometric assays for quantifying inorganic phosphate and silica. In the Fiske–Subbarow method, phosphate reacts with molybdate in acid to form phosphomolybdic acid, which is reduced to molybdenum blue and measured at 660 nm.[9] Similar reactions are utilized for the colorimetric quantification of dissolved silicates in water treatment facilities.
Agriculture
[edit]In agriculture, potassium molybdate is applied as a water-soluble fertilizer to supply micronutrient molybdenum to crops. Molybdenum is required by plants as part of the molybdenum cofactor (Moco) in enzymes such as nitrate reductase, and by symbiotic rhizobia in root nodules for the FeMoco center of nitrogenase, which catalyzes nitrogen fixation.[10]
Catalysis
[edit]The compound is used as a soluble molybdenum precursor in the manufacture of heterogeneous mixed-metal oxide catalysts. These catalysts are employed industrially in the gas-phase selective oxidation and ammoxidation of propene to acrolein and acrylonitrile.[11]
Safety
[edit]Potassium molybdate exhibits low acute toxicity compared to chromate salts. The oral median lethal dose (LD50) in rats is reported between 2400 and 2680 mg/kg.[1] It is not listed as a carcinogen by the IARC, the NTP, or OSHA.
Dust from the compound may cause mild, mechanical irritation to the eyes, skin, and respiratory system. In animals, particularly ruminants, high dietary intake of molybdates can induce secondary copper deficiency (molybdenosis) by forming insoluble copper thiomolybdates in the digestive tract.[12]
References
[edit]- 1 2 "Potassium molybdate Safety Data Sheet, Product 251216". Sigma-Aldrich. 2024-02-15. Retrieved 2026-09-22.
- ↑ "Potassium molybdate(VI)". National Center for Biotechnology Information. Retrieved 2026-09-22.
- ↑ Brauer, G. (1965). Handbook of Preparative Inorganic Chemistry. Vol. 2 (2nd ed.). New York: Academic Press. pp. 1412–1413.
- 1 2 3 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 1011. ISBN 978-0-08-037941-8.
- ↑ Gatehouse, B. M.; Leverett, P. (1969). "Crystal Structure of Potassium Molybdate, K2MoO4". Journal of the Chemical Society A: Inorganic, Physical, Theoretical. 1969: 849–854. doi:10.1039/J19690000849.
- ↑ van den Berg, A. J.; Tuinstra, F. (1978). "The Space Group and Structure of High-Temperature Potassium Molybdate". Acta Crystallographica Section B. 34 (10): 3177–3181. doi:10.1107/S056774087801041X.
- 1 2 Pope, Michael Thor (1983). Heteropoly and Isopoly Oxometalates. Berlin, Heidelberg: Springer-Verlag. pp. 15–27. doi:10.1007/978-3-642-87225-9. ISBN 978-3-540-11889-3.
- ↑ Vukasovich, M. S.; Farr, J. P. G. (1986). "Molybdate in Corrosion Inhibition—A Review". Materials Performance. 25 (5): 9–18.
- ↑ Fiske, Cyrus H.; Subbarow, Yellapragada (1925). "The Colorimetric Determination of Phosphorus". Journal of Biological Chemistry. 66 (2): 375–400. doi:10.1016/S0021-9258(18)84756-1.
- ↑ Marschner, Horst (2012). Marschner's Mineral Nutrition of Higher Plants (3rd ed.). Academic Press. pp. 251–257. ISBN 978-0-12-384905-2.
- ↑ Grasselli, R. K.; Burrington, J. D. (1981). "Selective Oxidation and Ammoxidation Catalysis". Advances in Catalysis. Vol. 30. Academic Press. pp. 133–163. doi:10.1016/S0360-0564(08)60317-0.
- ↑ Underwood, E. J. (1977). Trace Elements in Human and Animal Nutrition (4th ed.). Academic Press. pp. 109–131. ISBN 978-0-12-709065-8.
External links
[edit]
- Potassium molybdate at PubChem
- NIOSH Pocket Guide to Chemical Hazards: Molybdenum (soluble compounds, as Mo)