Arsine

Arsine is the chemical compound with the formula AsH₃. This flammable, pyrophoric, and highly toxic gas is the simplest compound of arsenic. Aside from its lethality, it finds applications in the semiconductor industry and for the synthesis of organoarsenic compounds.^[1]

At its standard state, arsine is a colorless, denser-than-air gas that is soluble in water (200 mL/L) and in many organic solvents as well. Whereas arsine itself is odorless, owing to its oxidation by air it is possible to smell a slight, garlic-like scent when the compound is present at about 0.5 ppm. This compound is generally regarded as stable, since at room temperature it decomposes only slowly. At temperatures of ca. 230 °C decomposition to arsenic and hydrogen is rapid. Several factors, such as humidity, presence of light and certain catalysts (namely aluminium) facilitate the rate of decomposition.^[1]

AsH₃ is a pyramidal molecule with H–As–H angles of 91.8° and three equivalent As–H bonds, each of 1.519 Å length. The term arsine is commonly used to describe a class of organoarsenic compounds of the formula AsH_3−xR_x, where R = aryl or alkyl. For example, As(C₆H₅)₃, called triphenylarsine, is referred to as "an arsine."

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Discovery

AsH₃ was reported in 1775 by Carl Scheele from the reduction of arsenic(III) oxide with zinc and acid. This reaction is a prelude to the Marsh test, described briefly below.

Synthesis

AsH₃ is generally prepared by the reaction of As³⁺ sources with H⁻ equivalents.^[2][3]

4 AsCl₃ + 3 NaBH₄ → 4 AsH₃ + 3 NaCl + 3 BCl₃

Alternatively, sources of As³⁻ react with protonic reagents to also produce this gas:

Zn₃As₂ + 6 H⁺ → 2 AsH₃ + 3 Zn²⁺

Reactions

The chemical properties of AsH₃ are well developed and can be anticipated based on an average of the behavior of PH₃ and SbH₃.

Thermal decomposition

Typical for a heavy hydride (e.g., SbH₃, H₂Te, SnH₄), AsH₃ is unstable with respect to its elements. In other words, AsH₃ is stable kinetically but not thermodynamically.

2 AsH₃ → 3 H₂ + 2 As

This decomposition reaction is the basis of the Marsh Test described below, which detects the metallic As.

Oxidation

Continuing the analogy to SbH₃, AsH₃ is readily oxidized by O₂ or even air:

2 AsH₃ + 3 O₂ → As₂O₃ + 3 H₂O

Arsine will react violently in presence of strong oxidizing agents, such as potassium permanganate, sodium hypochlorite or nitric acid.^[1]

Precursor to metallic derivatives

AsH₃ is used as a precursor to metal complexes of "naked" (or "nearly naked") As. Illustrative is the dimanganese species [(C₅H₅)Mn(CO)₂]₂AsH, wherein the Mn₂AsH core is planar.^[4]

Gutzeit test

A characteristic test for arsenic involves the reaction of AsH₃ with Ag⁺, called the Gutzeit test for arsenic.^[5] Although this test has become obsolete in analytical chemistry, the underlying reactions further illustrate the affinity of AsH₃ for "soft" metal cations. In the Gutzeit test, AsH₃ is generated by reduction of aqueous arsenic compounds, typically arsenites, with Zn in the presence of H₂SO₄. The evolved gaseous AsH₃ is then exposed to AgNO₃ either as powder or as a solution. With "solid" AgNO₃, AsH₃ reacts to produce yellow Ag₄AsNO₃, whereas AsH₃ reacts with a "solution" of AgNO₃ to give black Ag₃As.

Acid-base reactions

The acidic properties of the As–H bond are often exploited. Thus, AsH₃ can be deprotonated:

AsH₃ + NaNH₂ → NaAsH₂ + NH₃

Upon reaction with the aluminium trialkyls, AsH₃ gives the trimeric [R₂AlAsH₂]₃, where R = (CH₃)₃C.^[6] This reaction is relevant to the mechanism by which GaAs forms from AsH₃ (see below).

AsH₃ is generally considered non-basic, but it can be protonated by "super acids" to give isolable salts of the tetrahedral species [AsH₄]⁺.^[7]

Reaction with halogen compounds

Reactions of arsine with the halogens (fluorine and chlorine) or some of their compounds, such as nitrogen trichloride, are extremely dangerous and can result in explosions.^[1]

Catenation

In contrast to the behavior of PH₃, AsH₃ does not form stable chains, although H₂As–AsH₂ and even H₂As–As(H)–AsH₂ have been detected. The diarsine is unstable above −100 °C.

Microelectronics applications

AsH₃ is used in the synthesis of semiconducting materials related to microelectronics and solid-state lasers. Related to P, Arsenic is an n-dopant for silicon and germanium.^[1] More importantly, AsH₃ is used to make the semiconductor GaAs by CVD at 700–900 °C:

Ga(CH₃)₃ + AsH₃ → GaAs + 3 CH₄

Chemical warfare applications

Since before WWII AsH₃ was proposed as a possible chemical warfare weapon. The gas is colorless, almost odourless, and 2.5 times more dense than air, as required for a blanketing effect sought in chemical warfare. It is also lethal in concentrations far lower than those required to smell its garlic-like scent. In spite of these characteristics, arsine was never officially used as a weapon, because of its high flammability and its lower efficacy when compared to the non-flammable alternative phosgene. On the other hand, several organic compounds based on arsine, such as lewisite (β-chlorovinyldichloroarsine), adamsite (diphenylaminearsine), Clark I (diphenylchlorarsine) and Clark II, (diphenylcyanoarsine) have been effectively developed for use in chemical warfare.^[8]

Forensic science and the Marsh test

AsH₃ is also well known in forensic science because it is a chemical intermediate in the detection of arsenic poisoning. The old (but extremely sensitive) Marsh test generates AsH₃ in the presence of arsenic.^[3] This procedure, developed around 1836 by James Marsh, is based upon treating a As-containing sample of a victim's body (typically the stomach) with As-free zinc and dilute sulfuric acid: if the sample contains arsenic, gaseous arsine will form. The gas is swept into a glass tube and decomposed by means of heating around 250–300 °C. The presence of As is indicated by formation of a deposit in the heated part of the equipment. The formation of a black mirror deposit in the cool part of the equipment indicates the presence of Sb.

The Marsh test was widely used by the end of the 19th century and the start of the 20th; nowadays more sophisticated techniques such as atomic spectroscopy, inductively coupled plasma and x-ray fluorescence analysis are employed in the forensic field. Though neutron activation analysis was used to detect trace levels of arsenic in the mid 20th century it has fallen out of use in modern forensics.

Toxicology

For the toxicology of other arsenic compounds, see Arsenic, Arsenic trioxide and Arsenicosis.^[7]

The toxicity of arsine is distinct from that of other arsenic compounds. The main route of exposure is by inhalation, although poisoning after skin contact has also been described. Arsine binds to the haemoglobin of red blood cells, causing them to be destroyed by the body.

The first signs of exposure, which can take several hours to become apparent, are headaches, vertigo and nausea, followed by the symptoms of haemolytic anaemia (high levels of unconjugated bilirubin), haemoglobinuria and nephropathy. In severe cases, the damage to the kidneys can be long-lasting.

Exposure to arsine concentrations of 250 ppm is rapidly fatal: concentrations of 25–30 ppm are fatal for 30 min exposure, and concentrations of 10 ppm can be fatal at longer exposure times. Symptoms of poisoning appear after exposure to concentrations of 0.5 ppm. There is little information on the chronic toxicity of arsine, although it is reasonable to assume that, in common with other arsenic compounds, a long-term exposure could lead to arsenicosis.

See also

• Arsenic
• Cacodylic acid
• Cacodyl oxide

Bibliography

1. ^ ^{a b c d e} Institut National de Recherche et de Sécurité (2000). "Fiche toxicologique nº 53: Trihydrure d'arsenic" (PDF). Retrieved on 2006-09-06.
2. ^ Bellama, J. M.; MacDiarmid, A. G. "Synthesis of the Hydrides of Germanium, Phosphorus, Arsenic, and Antimony by the Solid-Phase Reaction of the Corresponding Oxide with Lithium Aluminum Hydride" Inorganic Chemistry, 1968, vol. 7, page 2070–2
3. ^ ^{a b} Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001
4. ^ Herrmann, W. A.; Koumbouris, B.; Schaefer, A.; Zahn, T.; Ziegler, M. L. "Generation and Complex Stabilization of Arsinidene and Diarsine Fragments by Metal-Induced Degradation of Monoarsine" Chemische Berichte (1985), vol. 118, pages 2472–88
5. ^ King, E. J. "Qualitative Analysis and Electrolytic Solutions" Harcourt, Brace, and World; New York (1959)
6. ^ Atwood, D. A.; Cowley, A. H.; Harris, P. R.; Jones, R. A.; Koschmieder, S. U.; Nunn, C. M.; Atwood, J. L.; Bott, S. G. "Cyclic Trimeric Hydroxy, Amido, Phosphido, and Arsenido Derivatives of aluminum and gallium. X-ray Structures of [tert-Bu₂Ga(m-OH)]₃ and [tert-Bu₂Ga(m-NH₂)]₃" Organometallics (1993), vol. 12, pages 24–29
7. ^ ^{a b} R. Minkwitz, R.; Kornath, A.; Sawodny, W.; Härtner, H. "Über die Darstellung der Pnikogenoniumsalze AsH₄⁺SbF₆⁻, AsH₄⁺AsF₆⁻, SbH₄⁺SbF₆⁻" Zeitschrift für anorganische und allgemeine Chemie Vol. 620, pages 753–756.
8. ^ Suchard, Jeffrey R. (March 2006). "CBRNE — Arsenicals, Arsine" (HTML). eMedicine. Retrieved on 2006-09-05.

• Hatlelid K. M. (1996). "Reactions of Arsine with Hemoglobine". Journal of Toxicology and Environmental Health Part A 47 (2): 145–157. doi:10.1080/009841096161852.
• Nielsen H. H. (1952). "The Molecular Structure of Arsine". The Journal of Chemical Physics 20 (12): 1955–1956. doi:10.1063/1.1700347.
• Fowler B. A., Weissberg J. B. (1974). "Arsine poisoning". New England Journal of Medicine 300: 1171–1174.