Arsenic acid

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Arsenic acid
Structural formula
Ball-and-stick model
Identifiers
CAS number 7778-39-4 YesY
ChemSpider 229 YesY
UNII N7CIZ75ZPN YesY
EC number 231-901-9
KEGG C01478 YesY
ChEBI CHEBI:18231 YesY
RTECS number CG0700000
Jmol-3D images Image 1
Properties
Molecular formula H3AsO4
Molar mass 141.94 g/mol
Appearance White translucent crystals,
hygroscopic.
Density 2.5 g/cm3
Melting point 35.5 °C (95.9 °F; 308.6 K)
Boiling point 120 °C (248 °F; 393 K) decomp
Solubility in water 16.7 g/100 mL
Solubility soluble in alcohol
Vapor pressure 55 hPa (50 °C)
Acidity (pKa) 2.19, 6.94, 11.5
Structure
Molecular shape Tetrahedral
Hazards
EU classification Toxic (T)
Dangerous for the environment (N)
R-phrases R23/25, R45, R50/53
S-phrases S53, S45, S60, S61
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
LD50 48 mg/kg (rat, oral)
Related compounds
Other anions Phosphoric acid
Other cations Sodium arsenate
Related compounds Arsenous acid
Arsenic pentoxide
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Arsenic acid is the chemical compound with the formula H3AsO4. More descriptively written as AsO(OH)3, this colorless acid is the arsenic analogue of phosphoric acid. Arsenate and phosphate salts behave very similarly. Arsenic acid as such has not been isolated, but only found in solution where it is largely ionized. Its hemihydrate form (H3AsO4·½H2O) does form stable crystals. Crystalline samples dehydrate with condensation at 100 °C.[1]

Properties[edit]

It is a tetrahedral species of idealized symmetry C3v with As–O bonds lengths ranging from 1.66 to 1.71 Å.[2]

Being a triprotic acid, its acidity is described by three equilibria:

H3AsO4 + H2O is in equilibrium with H2AsO
4
+ H3O+ (K1 = 10−2.19)
H2AsO
4
+ H2O is in equilibrium with HAsO2−
4
+ H3O+ (K2 = 10−6.94)
HAsO2−
4
+ H2O is in equilibrium with AsO3−
4
+ H3O+ (K3 = 10−11.5)

These Ka values are close to those for phosphoric acid. The highly basic arsenate ion (AsO3−
4
) is the product of the third ionization. Unlike phosphoric acid, arsenic acid is an oxidizer, illustrated by its ability to convert iodide to iodine.

Preparation[edit]

Arsenic acid is prepared by treating arsenic trioxide with concentrated nitric acid and dinitrogen trioxide is produced as a by-product.[3]

As2O3 + 2 HNO3 + 2 H2O → 2 H3AsO4 + N2O3

The resulting solution is cooled to give colourless crystals of the hemihydrate H3AsO4·½H2O, although the dihydrate H3AsO4·2H2O is produced when crystallisation occurs at lower temperatures.[3]

Other methods[edit]

Arsenic acid is slowly formed when arsenic pentoxide is dissolved in water, and when meta- or pyroarsenic acid is treated with cold water. Arsenic acid can also be prepared directly from elemental arsenic by moistening it and treating with ozone.

2 As + 3 H2O + 5 O3 → 2 H3AsO4 + 5 O2

Applications[edit]

Commercial applications of arsenic acid are limited by its toxicity. It has found occasional use as a wood preservative, broad-spectrum biocide, a finishing agent for glass and metal, and a reagent in the synthesis of some dyestuffs and organic arsenic compounds. The LD50 in rabbits is 6 mg/kg (0.006 g/kg).[4]

References[edit]

  1. ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5. 
  2. ^ Lee, C.; Harrison, W. T. A. (2007). "Tetraethylammonium dihydrogenarsenate bis(arsenic acid) and 1,4-diazoniabicyclo[2.2.2]octane bis(dihydrogenarsenate) arsenic acid: hydrogen-bonded networks containing dihydrogenarsenate anions and neutral arsenic acid molecules". Acta Crystallographica C 63 (Pt 7): m308–m311. doi:10.1107/S0108270107023967. PMID 17609552. 
  3. ^ a b G. Brauer, ed. (1963). "Arsenic Acid". Handbook of Preparative Inorganic Chemistry 1 (2nd ed.). New York: Academic Press. p. 601. 
  4. ^ Joachimoglu, G. (1915). Biochemische Zeitschrift 70: 144.