Phosphorous acid (or phosphonic acid) is the compound described by the formulaH3PO3. This acid is diprotic (readily ionizes two protons), not triprotic as might be suggested by this formula. Phosphorous acid is an intermediate in the preparation of other phosphorus compounds. Organic derivatives of phosphorous acid, compounds with the formula RPO3H2, are called phosphonic acids.
Nomenclature and tautomerism
Solid HP(O)(OH)2 has tetrahedral geometry about the central phosphorus atom, with a P−H bond of 132 pm, one P=O double bond of 148 pm and two longer P−OH single bonds of 154 pm. In common with other phosphorus oxides with P−H bonds (e.g.hypophosphorous acid and dialkyl phosphites),[2] it exists in equilibrium with an extremely minor tautomerP(OH)3. (In contrast, arsenous acid's major tautomer is the trihydroxy form.) IUPAC recommends that the trihydroxy form P(OH)3 be called phosphorous acid, and the dihydroxy form HP(O)(OH)2 phosphonic acid.[3] Only the reduced phosphorus compounds are spelled with an "-ous" ending.
Phosphorous acid has a pKa in the range 1.26–1.3.[6][7]
HP(O)(OH)2 → HP(O)2(OH)− + H+ pKa = 1.3
It is a diprotic acid, the hydrogenphosphite ion, HP(O)2(OH)− is a weak acid:
HP(O)2(OH)− → HPO2−3 + H+ pKa = 6.7
The conjugate base HP(O)2(OH)− is called hydrogen phosphite, and the second conjugate base, HPO2−3, is the phosphite ion.[8] (Note that the IUPAC recommendations are hydrogen phosphonate and phosphonate respectively).
The hydrogen atom bonded directly to the phosphorus atom is not readily ionizable. Chemistry examinations often test students' appreciation of the fact that not all three hydrogen atoms are acidic under aqueous conditions, in contrast with H3PO4.
This reaction is used for laboratory-scale preparations of PH3.
Phosphorous acid slowly oxidizes in air to phosphoric acid.[5]
Both phosphorous acid and its deprotonated forms are good reducing agents, although not necessarily quick to react. They are oxidized to phosphoric acid or its salts. It reduces solutions of noble metal cations to the metals. When phosphorous acid is treated with a cold solution of mercuric chloride, a white precipitate of mercurous chloride forms:
H3PO3 + 2 HgCl2 + H2O → Hg2Cl2 + H3PO4 + 2 HCl
Mercurous chloride is reduced further by phosphorous acid to mercury on heating or on standing:
Upon treatment with metals of d6 configuration, phosphorous acid is known to coordinate as the otherwise rare P(OH)3 tautomer. Examples include Mo(CO)5(P(OH)3) and [Ru(NH3)4(H2O)(P(OH)3)]2+.[10][11]
The most important use of phosphorous acid (phosphonic acid) is the production of basic lead phosphite, which is a stabilizer in PVC and related chlorinated polymers.[5]
It is used in the production of basic lead phosphonate PVC stabilizer, aminomethylene phosphonic acid and hydroxyethane diphosphonic acid.
It is also used as a strong reducing agent and in the production of synthetic fibres, organophosphorus pesticides, and the highly efficient water treatment agent ATMP.
Ferrous materials, including steel, may be somewhat protected by promoting oxidation ("rust") and then converting the oxidation to a metalophosphate by using phosphoric acid and further protected by surface coating. (See: Passivation (chemistry)).
Organic derivatives
The IUPAC (mostly organic) name is phosphonic acid. This nomenclature is commonly reserved for substituted derivatives, that is, organic group bonded to phosphorus, not simply an ester. For example, (CH3)PO(OH)2 is "methylphosphonic acid", which may of course form "methylphosphonate" esters.
^Janesko, Benjamin G.; Fisher, Henry C.; Bridle, Mark J.; Montchamp, Jean-Luc (2015-09-29). "P(═O)H to P–OH Tautomerism: A Theoretical and Experimental Study". The Journal of Organic Chemistry. 80 (20). American Chemical Society (ACS): 10025–10032. doi:10.1021/acs.joc.5b01618. ISSN0022-3263. PMID26372089.
^Larson, John W.; Pippin, Margaret (1989). "Thermodynamics of ionization of hypophosphorous and phosphorous acids. Substituent effects on second row oxy acids". Polyhedron. 8 (4): 527–530. doi:10.1016/S0277-5387(00)80751-2.
^CRC Handbook of Chemistry and Physics (87th ed.). pp. 8–42.
^Novosad, Josef (1994). Encyclopedia of Inorganic Chemistry. John Wiley and Sons. ISBN0-471-93620-0.
^ abXi, Chanjuan; Liu, Yuzhou; Lai, Chunbo; Zhou, Lishan (2004). "Synthesis of molybdenum complex with novel P(OH)3 ligand based on the one-pot reaction of Mo(CO)6 with HP(O)(OEt)2 and water". Inorganic Chemistry Communications. 7 (11): 1202–1204. doi:10.1016/j.inoche.2004.09.012.
^Sernaglia, R. L.; Franco, D. W. (2005). "The ruthenium(II) center and the phosphite-phosphonate tautomeric equilibrium". Inorg. Chem. 28 (18): 3485–3489. doi:10.1021/ic00317a018.
^Alexander, K. A.; Bryan, S. A.; Dickson, M. K.; Hedden, D.; Roundhill (2007). "Potassium Tetrakis[Dihydrogen Diphosphito(2–)]Diplatinate(II)". Inorganic Syntheses. Vol. 24. pp. 211–213. doi:10.1002/9780470132555.ch61. ISBN9780470132555.
Further reading
Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN0-12-352651-5.
Corbridge., D. E. C. (1995). Phosphorus: An Outline of its Chemistry, Biochemistry, and Technology (5th ed.). Amsterdam: Elsevier. ISBN0-444-89307-5.
Lee, J.D. (3 January 2008). Concise Inorganic Chemistry. Oxford University Press. ISBN978-81-265-1554-7.
