Chemical elements
  Arsenic
      Occurrence
      Ubiquity
      History
    Isotopes
    Energy
    Production
    Application
    Physical Properties
    Chemical Properties
      Aluminium Arsenide
      Antimony Arsenides
      Barium Arsenide
      Bismuth Arsenides
      Cadmium Arsenides
      Calcium Arsenide
      Cerium Arsenide
      Chromium Arsenides
      Cobalt Arsenides
      Copper Arsenides
      Gold Arsenides
      Iridium Arsenide
      Iron Arsenides
      Lead Arsenides
      Lithium Arsenide
      Magnesium Arsenide
      Manganese Arsenides
      Mercury Arsenides
      Molybdenum Arsenide
      Nickel Arsenides
      Niobium Arsenide
      Palladium Di-arsenide
      Platinum Arsenides
      Potassium Arsenides
      Rhodium Arsenide
      Ruthenium Arsenide
      Silver Arsenides
      Sodium Arsenide
      Strontium Arsenide
      Thallium Arsenide
      Tin Arsenides
      Tungsten Arsenide
      Uranium Arsenide
      Zinc Arsenides
      Arsenic Subhydride
      Arsenic Monohydride
      Arsenic Trihydride
      Arsenic Trifluoride
      Arsenic Pentafluoride
      Arsenic Nitrosyl Hexafluoride
      Arsenic Trichloride
      Arsenic Oxychloride
      Arsenic Pentachloride
      Arsenic Tribromide
      Arsenic Oxybromide
      Arsenic Moniodide
      Arsenic Diiodide
      Arsenic Triiodide
      Arsenic Pentiodide
      Arsenic Suboxide
      Arsenious Oxide
      Aluminium Arsenite
      Ammonium Arsenites
      Antimony Arsenite
      Barium Arsenites
      Beryllium Arsenite
      Bismuth Arsenite
      Cadmium Arsenites
      Calcium Arsenites
      Chromic Arsenite
      Cobalt Arsenites
      Copper Arsenites
      Gold Arsenites
      Iron Arsenites
      Lead Arsenites
      Lithium Arsenite
      Magnesium Arsenites
      Manganese Arsenites
      Mercury Arsenites
      Nickel Arsenites
      Palladium Pyroarsenite
      Platinum Arsenites
      Potassium Arsenites
      Arsenites of Rare Earth Metals
      Rubidium Metarsenite
      Silver Arsenites
      Sodium Arsenites
      Strontium Arsenites
      Thallous Orthoarsenite
      Tin Arsenites
      Titanyl Tetrarsenite
      Tungsto-arsenites
      Uranyl Metarsenite
      Zinc Arsenites
      Zirconium Arsenite
      Arsenic Tetroxide
      Arsenic Pentoxide
      Aluminium Arsenates
      Ammonium Arsenates
      Barium Arsenates
      Beryllium Arsenates
      Bismuth Arsenates
      Cadmium Arsenates
      Caesium Arsenate
      Calcium Arsenates
      Chromium Arsenates
      Cobalt Arsenates
      Copper Arsenates
      Hydroxylamine Orthoarsenate
      Iron Arsenates
      Lead Arsenates
      Lithium Arsenates
      Magnesium Arsenates
      Manganese Arsenates
      Mercury Arsenates
      Molybdenum Arsenates
      Nickel Arsenates
      Palladium Arsenate
      Platinic Arsenate
      Potassium Arsenates
      Rare Earth Metals Arsenates
      Rhodium Arsenate
      Rubidium Arsenates
      Silver Arsenates
      Sodium Arsenates
      Strontium Arsenates
      Thallium Arsenates
      Thorium Arsenates
      Tin Arsenates
      Titanyl Arsenate
      Tungsto-arsenic Acids
      Uranium Arsenates
      Vanado-arsenates
      Zinc Arsenates
      Zirconium Arsenates
      Perarsenates
      Arsenic and Sulphur
      Arsenic Subsulphide
      Tetrarsenic Trisulphide
      Arsenic Disulphide
      Arsenic Trisulphide
      Arsenic Pentasulphide
      Thioarsenates
      Ammonium Thioarsenates
      Antimony Thioarsenate
      Barium Thioarsenates
      Beryllium Thioarsenate
      Bismuth Thioarsenate
      Cadmium Thioarsenates
      Calcium Thioarsenates
      Cerium Thioarsenates
      Chromium Thioarsenate
      Cobalt Thioarsenate
      Copper Thioarsenates
      Gold Thioarsenates
      Iron Thioarsenates
      Lead Thioarsenates
      Lithium Thioarsenates
      Magnesium Thioarsenates
      Manganese Thioarsenates
      Mercury Thioarsenates
      Molybdenum Thioarsenates
      Nickel Thioarsenates
      Platinic Thioarsenate
      Potassium Thioarsenates
      Silver Thioarsenates
      Sodium Thioarsenates
      Strontium Thioarsenates
      Thallium Orthothioarsenate
      Tin Thioarsenates
      Uranyl Thioarsenate
      Yttrium Thioarsenate
      Zinc Thioarsenates
      Zirconium Thioarsenate
      Trioxythioarsenic Acid
      Dioxydithioarsenic Acid
      Oxytrithioarsenic Acid
      Arsenic Monosulphatotrioxide
      Arsenic Disulphatotrioxide
      Arsenic Trisulphatotrioxide
      Arsenic Tetrasulphatotrioxide
      Arsenic Hexasulphatotrioxide
      Arsenic Octasulphatotrioxide
      Complex salts of Sulphato-compounds of Arsenic
      Arsenic Nitride
      Arsenic Imide
      Arsenic Amide
      Arsenic Phosphides
      Arsenic oxyphosphides
      Arsenic Phosphate
      Arsenic Thiophosphate
      Arsenic Tricarbide
      Arsenic Pentasilicide
      Boron Arsenate
    Detection of Arsenic
    Estimation of Arsenic
    Physiological Properties
    PDB 1b92-1ihu
    PDB 1ii0-1tnd
    PDB 1tql-2hmh
    PDB 2hx2-2xnq
    PDB 2xod-3htw
    PDB 3hzf-3od5
    PDB 3ouu-9nse

Arsenious Oxide, As2O3






Arsenious Oxide (Arsenic Trioxide), As2O3 or As2O6, also known as white arsenic, arsenic, flowers of arsenic and arsenious acid, was known to the Arabian alchemists of the twelfth century. Roger Bacon described it as a white, transparent substance resulting from the sublimation of a mixture of orpiment and iron filings. It is obtained commercially in large quantities by the roasting of arseniferous minerals, a method of production known from early times, the product being called Huttenrauch (furnace smoke) by Basil Valentine. When arsenopyrite (mispickel), FeAsS, is heated in air below red heat, arsenic sulphide vapour is liberated, but at a higher temperature arsenious oxide and sulphur dioxide are evolved, leaving a residue of ferric oxide, sulphide and arsenate. Likewise, lollingite produces arsenious oxide at a bright red heat. But although some arsenic ores are thus treated for arsenic alone, almost the whole of the world's supply is produced as a by-product in the treatment of ores for gold, copper, lead and tin, etc.

In 1914 the annual world production of white arsenic was about 10,000 tons. Germany was the foremost producer, but the production of the United States and Canada was of increasing importance; England, which in the prosperous days of Cornish tin mining had produced even more than Germany, was still an important producer. In recent years the annual production has expanded until, since 1933, it has reached more than 70,000 tons, the lead being taken by Sweden and the United States. The outputs for 1933 and 1934 of the principal countries concerned are shown in the following table. The available figures do not distinguish between crude and refined white arsenic, such division being practically impossible, for what is designated "crude arsenic" in some countries contains as much as 98 per cent. As2O3 and compares favourably with the refined product exported from other sources.

Sweden, which produced virtually no arsenic prior to 1926, has rapidly become the dominant producer, and the smelting works and electrolytic refineries recently erected at Boliden on the Gulf of Bothnia are able to deal with 400,000 tons of ore per annum; they produce refined copper, gold, silver and white arsenic. The ore used is a mixture of chalcopyrite, arsenopyrite and iron pyrites containing only 2 per cent, of copper, small quantities of gold and silver, and some lead, zinc and antimony. The arsenic content averages nearly 11 per cent, and the annual production is now more than 40,000 tons - more than world requirements. This production on such an unprecedented and, at present, unmarketable scale of so toxic a substance presented an interesting problem both from the economic and metallurgical points of view. At first the crude arsenical fume was rendered innocuous by mixing with cement and water and allowing the mixture to harden before dumping, or the condensed product was placed in large concrete cylinders and sunk in deep water in the Gulf of Bothnia. This method of disposal, however, proved too expensive and the product is therefore stored in huge concrete structures.

In the preliminary roast the arsenic is almost completely removed from the ore, less than 1 per cent, remaining in the calcines. The gases are cooled below 150° C., first in large chambers of sheet iron equipped with air-cooled jackets, where about one-third of the arsenious oxide condenses, the rest being caught in Cottrell " treaters " of the plate type. These are chambers packed with vertical collecting plates of corrugated iron; the gas flow is horizontal across the face of the plates and the latter are cleaned by tapping with " knockers," whereupon the condensed dust falls into hoppers. From these, the dust passes in closed cars or by screw-conveyors to a belt-conveyor which runs in an underground tunnel to an elevator which carries the powder to the top of the arsenic storage into which it is discharged through slots in the roof.

In the United States 2 about 90 per cent, of the white arsenic produced is recovered as a by-product in the smelting of copper ores, the remainder being obtained from lead, gold and iron concentrates. No ores are now mined directly for arsenic, the above sources being sufficient to satisfy existing demands. The chief producing states are Montana, where arsenic occurs chiefly as enargite, and Utah, where the deposits contain arsenopyrite and scorodite. The roaster plants yield flue dusts containing about 20 per cent, of As2O3. Lead " bag-house dust " carries from 30 to 40 per cent. As2O3. The dust is smelted in a reverberatory furnace, producing a matte, and the gases are passed through a hot Cottrell " treater " maintained at a temperature which precipitates the dust but allows the arsenical fume to pass on to a cold Cottrell " treater " where it is condensed, yielding a product containing 70 per cent. As2O3. This crude product is resublimed until the desired colour and purity are obtained, and in the final stages reverberatory type furnaces with iron hearths, heated both below and above, are employed and smokeless fuel is used in order to prevent contamination of the product with soot. The fumes are condensed in " kitchens," which are long chambers divided into compartments, and the white arsenic finally obtained contains 99.9 per cent. As2O3. This constitutes refined white arsenic, but crude white arsenic, " black dust" and " treater dust " are also marketable products. Part of the latter is made directly into sodium arsenate or "weed-killer".

In Cornwall and Devon the ores used contain arsenopyrite mixed with iron and copper pyrites, tin ore, zinc blende, galena, etc. Before roasting, the ores are separated as far as possible by hand, and tinstone is removed by washing the finely powdered material. The roasting is conducted in a reverberatory furnace having a revolving floor over which a number of scrapers are fixed. The ore, which contains 10 to 30 per cent. As, is introduced through a hopper on to the floor, which revolves once every 12 minutes or so. It is heated to dull redness for about 10 hours and the arsenious oxide, with some sulphur and carbon, collects in the flues and condensing chambers as dark grey arsenic soot. This is refined by recalcination with smokeless fuel and condensed in zigzag chambers (of which there are several types), the product of the first chamber being reground and recalcined. In the other chambers crystalline white arsenic collects and this is ground and passed through leather pipes into casks under conditions which prevent escape. About 60 per cent, of the arsenic in the ores is thus recovered, the remainder being lost in the slimes formed during the early treatment.

Special methods of treatment are frequently necessary in order to purify the crude oxide or flue dust from the less volatile antimony oxide and also from metallic and other constituents. A flue dust containing up to 6 per cent, of antimony, if volatilised in a tube furnace in a slow current of air and condensed in baffled flues maintained at about 350° C., gives in the zone nearest the furnace a vitreous deposit containing a small amount of arsenious oxide and most of the antimonious oxide, and beyond a crystalline deposit of arsenious oxide of high purity. Purification may also be accomplished by heating the dust under pressure at 150° C. with water, dilute sulphuric acid or dilute alkali solutions; on cooling, after concentration, the pure arsenious oxide crystallises, and the deposition may be assisted by stirring the solution in the presence of a large amount of powdered arsenious oxide. Another method for removing antimony oxide consists in converting the arsenic to the trichloride, which is then repeatedly shaken with concentrated hydrochloric acid, in which antimony trichloride is more soluble than the arsenic compound. The latter may then be hydrolysed to the oxide by slowly adding to a considerable quantity of boiling water. Lead and cadmium may be removed by heating with alkali solution to about 200° C. under pressure. The arsenic and lead pass into solution and the latter is precipitated by passing in carbon dioxide. The arsenious oxide is then crystallised from the alkali and any alkali hydroxide removed by hydrolysis.

Arsenious oxide may be precipitated from acid solutions containing arsenic acid by treatment with reducing agents, for example, by passing sulphur dioxide under pressure into the agitated liquid. Precipitation may be assisted by concentration or cooling, or by addition of finely divided arsenious oxide to the well-stirred solution.

The removal of small quantities of arsenic from metals and ores is a commercial problem which may be mentioned at this point. The Harris process of softening lead, used in several refineries, is based on the principle that such impurities as arsenic, antimony and tin may easily be oxidised and in the presence of certain alkali salts can be converted into arsenates, antimonates and stannates. Certain fluxes, such as sodium nitrate, sodium hydroxide, sodium chloride or lead oxide, are added to the molten lead, the presence of an oxidising agent and an alkali salt being essential. The alkaline slag obtained is fused and poured in a thin stream into a saturated solution of sodium chloride containing sufficient sodium hydroxide to give a liquid of density 1.4. The mixture is agitated at 90° C. and the hot alkali arsenate solution separated by decantation from the insoluble residue of sodium stannate and antimonate. On cooling, the solution deposits trisodium ortho-arsenate. Another method of extracting the sodium arsenate from the alkaline slag is to digest the crushed slag with cold water, which dissolves out most of the stannate; the residue is then leached with hot water, which dissolves all the arsenate and the remaining stannate, and this solution, when concentrated and cooled, deposits the arsenate. Arsenical tin may be treated in a similar manner; the molten metal is mixed with sodium hydroxide and chloride and small quantities of sodium nitrate added from time to time. The slag, containing sodium arsenate, is skimmed off and the latter recovered by boiling with water and evaporation of the clear solution.

From iron and manganese ores traces of arsenic may be eliminated by heating at a temperature above 500° C. in an atmosphere containing carbon dioxide mixed with certain reducing gases, such as hydrogen, carbon monoxide or hydrocarbons, of such composition that the ores are reduced to lower oxides, but not allowing the formation of the metals or carbides; the arsenic is expelled as arsenious oxide. To eliminate arsenic from tungsten ore the latter may be roasted, or free sulphur may be added and the mixture heated to above the boiling point of sulphur and sufficiently high to volatilise the major portion of the arsenic present.

Arsenic-bearing ores or materials may be mixed with carbonaceous material and ignited in a current of air under low pressure in order to volatilise the arsenious oxide. Other methods of de-arsenising depend on converting the element into volatile sulphide or, in the case of metals, into some compound, such as calcium arsenide, insoluble in the molten metal; or again, finely ground ores may be agitated with carbon disulphide until the arsenic compounds are dissolved.

Sulphur may be freed from arsenic by treating it in the molten state with compounds, such as lime or sodium sulphide, which react to form compounds insoluble in the sulphur and which may be separated by settling and filtration; or the sulphur may be treated in the vapour or liquid state with chlorine or sulphur dichloride in excess, to form arsenic chloride, which may be removed by distillation and by scrubbing with air or an inert gas under reduced pressure. Other methods consist in treating sulphur vapour with a molten alkali sulphide or polysulphide, or with alkali or alkali carbonate which, with sulphur, forms sulphides; or again, the sulphur may be digested with a solution of ammonia, ammonium sulphide or ammonium carbonate, preferably under pressure at 120° to 130° C.

The principle of flotation has been applied to the extraction of arsenic concentrates from residual sands obtained in the treatment of gold ores. The sands contained about 5.8 per cent, of arsenic in the form of small grains of arsenopyrite, associated with pyrites and some quartz and felspars. The best results, whereby 90 per cent, of the arsenic was recovered, were obtained with the following flotation agents in acidulated medium: (1) thiocarbanilide (in solution in o-toluidine) with turpentine, (2) the same reagent containing xanthate, (3) the same reagent with xanthate and pine oil.


Polymorphism and Physical Properties of Arsenious Oxide

Two distinct crystalline varieties of white arsenic are well known, namely, the octahedral (α-) form and the prismatic (β-) form. The existence of a third crystalline (γ-) form has been suggested by Smits and Beljaars (vide infra). An amorphous form of the oxide also occurs as vitreous white arsenic. The conditions of formation of the α-, β- and amorphous varieties may be demonstrated by heating either the octahedral or the vitreous form in a sealed tube fixed vertically, the temperature being maintained at about 400° C. at the lower end and at about 200° C. near the top. The oxide sublimes and condenses in the cooler upper part of the tube as the octahedral form, in the hot lower part as the vitreous form, and in the middle region as the prismatic form. When the octahedral crystals are heated, some volatilisation occurs at about 100° to 125° C. and, as the temperature rises, they sublime without melting; under the pressure of its own vapour, however, fusion occurs at about 272° C. (26.1 mm.), and if the temperature is maintained at a somewhat higher level, starlike masses of the prismatic form gradually appear. This change of the octahedral to the prismatic variety is extremely slow, several days being required for completion; in the presence of moisture, however, which acts as an accelerator, the change has been observed after heating for a few hours at 100° C. The melting point of the prismatic crystals is approximately 315° C.

The relation between the octahedral and prismatic modifications has not yet been satisfactorily elucidated. The former is the stable form at ordinary temperatures and the latter at higher temperatures; the transition point according to Rushton and Daniels is 250° C. and according to Smits and Beljaars 200° C., but the prismatic form is persistent at much lower temperatures and the change from octahedral to prismatic may be monotropic. Interesting information has been obtained from measurements of the vapour pressure of the oxide.

Smits and Beljaars investigated the vapour pressure between 240° and 380° C. and found the following values. The oxide was prepared by repeated sublimation at 320° C. under reduced pressure, followed by heating at 200° C. for 6 to 8 hours.

Vapour pressure arsenious oxide
Vapour pressure curves of arsenious oxide


From the vapour pressure-temperature curves, fig., the melting point of octahedral arsenious oxide is found to be 272.1° C. By not heating the samples too long, the following points on the metastable vapour pressure curve, M, were obtained:

Temp., °C.278.3293.4309.0318.9326.7
Vap. Press, (mm.).37.568.0111.2122.994.1


As seen in the figure, this curve passes through a maximum and decreases to cut the liquid-vapour line. On prolonged keeping of the octahedral form at a temperature above 258° C. a product was obtained which melted at 289.6° C. and was considered to be a new y-form; this yielded abnormally high vapour pressure values:

Phase:γ-form
Temp., °C.243.1263.8274.2279.3284.6289.8295.1310.0315.7
Vap. Press, (mm.)39.457.065.769.274.277.780.477.775.9


The following values for the prismatic variety were obtained:

Phase:Prismatic Crystals (β).
Temp., ° C.243.1253.4263.8274.2279.3284.6289.8295.1
Vap. Press, (mm.)4.97.611.717.221.424.831.137.2


After partial distillation of the prismatic crystals, a residue was obtained which on heating gave abnormally low vapour pressures (curve L), thus:

Temp., ° C.263.8274.2279.3284.6289.8295.1
Vap. Press, (mm.)2.39.6115.922.429.537.2


but, on cooling, the β curve was followed. The above product, if partly sublimed in a closed vessel, yielded a sublimate having the high vapour pressures of the γ-form. It is therefore concluded that the prismatic form behaves as a mixed crystal phase in internal equilibrium which is disturbed by partial distillation.

The following stable triple points were determined: α-β-vapour, about 200° C. (0.26 mm.); β-liquid-vapour, 312.3° C. (66.1 mm.); and the metastable points, α-liquid-vapour, 272.1° C. (26.1 mm.); α-γ-vapour, 258.4° C. (13.9 mm.); γ-liquid-vapour, 289.6° C. (40.7 mm.).

The following thermal values were also obtained: - Molar heats of sublimation: α-As2O3, 29,833; β, 23,676; γ, 21,130 calories. Molar heats of fusion: α, 15,099; β, 8942; γ, 6396 calories. Molar heat of vaporisation: β, 14,734 calories.

The octahedral form of arsenious oxide crystallises in the cubic system. It is produced whenever the vapour is condensed on a cold surface under conditions of rapid cooling; it also results by the slow transformation of the vitreous modification. It may be obtained by crystallisation from a hot saturated aqueous solution of the latter; the crystallisation may be attended by the emission of flashes of light, easily seen in a darkened room. This is the case when crystallisation takes place from solutions containing hydrochloric acid, or a mixture of hydrochloric and nitric acids, even when the latter is in sufficient quantity to cause complete oxidation of the arsenious oxide. The phenomenon persists after recrystallisation and has been attributed to tribolumineseence, since light is also emitted when the crystals are crushed or well stirred with a metal rod. The emitted rays exhibit no electrical properties and the spectrum is continuous. The phenomenon has been investigated by Bhatnagar and his co-workers, who suggest that during crystallisation a small quantity of the solution forms a dispersed phase in the crystal and, according to the physical conditions existing, microcrystallisation takes place more or less rapidly with the emission of light.

When exposed to filtered ultraviolet light, the pure oxide does not exhibit any characteristic fluorescence suitable for its identification.

This cubic modification is the stable form of the oxide at ordinary temperatures and occurs in Nature as arsenite or arsenolite, usually accompanying ores of lead, iron, cobalt, nickel, silver, etc. It is a product of the decomposition of arsenical ores. It is frequently found also as a greyish crust on native arsenic, its presence being due to superficial oxidation. The crystals are isomorphous with senarmontite, a cubic variety of antimony trioxide, and a study of the crystal structure, based on powder photographs, shows that the space-lattice in both cases is of the diamond form. The molecules preserve their identity in the crystal and possess the 24-fold symmetry of the regular tetrahedron; the four arsenic atoms are associated with the four corners of the tetrahedron and the oxygen atoms with the six edges. This is in accordance with the experimental evidence which suggests that the molecule of the oxide agrees with the formula As2O6. The unit cube contains eight molecules and the side a = 11.0457 ± 0.0002 A.; the shortest distance between the arsenic and oxygen atoms is 2.01 A., and the calculated density is 3.877. The As-As separation is 3.20 ± 0.05 A.

When either arsenolite or senarmontite is sublimed on to mica, it is deposited in octahedra, respectively isotropic and birefringent, oriented so that similar dimensions of the crystal meshes coincide, for example: As2O6 13.54, Sb4O6 13.64, mica 13.66 A. Such orientation appears only to occur with minerals of ionic structure and when both substances concerned have heteropolar linkings, so that the phenomenon is said to provide evidence of this type of linking. The actual density of the cubic crystals, according to Baxter and Hawkins, is 3.874 at 0° C., 3.865 at 25° C. and 3.851 at 50° C., but lower values have usually been obtained; the density of the natural product usually varies between 3.70 and 3.72. The hardness is 1.5. The compressibility, β, under pressures up to 9000 atm. has been determined with the following results:

At 30° C. β (i.e. - ΔV/V0) = 92.49×10-7p – 272.4×10-12p2
At 75° C. β (i.e. - ΔV/V0) = 92.88×10-7p – 250.5×10-12p2

This very high compressibility is to be expected from the structure of the crystal which, as seen above, is molecular rather than ionic and moreover contains large open spaces between the atoms. The refractive indices at 17° C. are 1.755 for sodium light and 1.748 for lithium light. Klocke observed that for yellow light the sublimed crystals exhibit double refraction, but this could not be confirmed by Brauns.

The crystals on heating volatilise without melting, but when heated under pressure liquefy with initial formation of the vitreous form. The specific heat is 0.1279, and the molar heat over the range 3° to 41° C. is 28.83. The coefficient of cubical expansion is 0.00011 from 0° to 25° C., and 0.00012 from 25° to 50° C.

The heat of formation is as follows:

2As (cry st.) + 3O = As2O3(Octahedral) + 154,670 cal.

From E.M.F. measurements at 25° and 45° C. of the cell

As (metallic) | As2O3(octahedral) + HClO4(0.22 - 0.94M) | H2

Schuhmann derived the free energy and heat content of arsenious oxide. The electrode, which consisted of arsenic deposited on platinum, was immersed in a mixture of perchloric acid and arsenious oxide and the E.M.F. was found to be, at 25° C., -0.2340 volt and, at 45° C., -0.2250 volt. The free energy of formation at 25° C. of As2O3 (octahedral) from metallic arsenic and oxygen was computed to be -137,300 calories, and the heat content, derived from the measurements at the two temperatures, was found to be -153,800 calories, in fair agreement with Thomsen's value, -154,700 calories, obtained by an indirect method. Experiments with an adiabatic calorimeter, in which heat changes as small as 1.5×10-5 calories per gram-hour could be registered, revealed no continuous heat evolution from arsenious oxide.

The heat capacity of the oxide has been investigated at temperatures from about 57° to 296° Abs., and the entropy at 25° C., in gram-calories per degree, is calculated to be 25.6.

The crystals dissolve slowly in cold water, more readily in boiling water, and the solution is feebly acidic.

The prismatic crystals, which belong to the monoclinic system, occur naturally as claudetite, rhomb arsenite or arsenophyllite. They sometimes occur in thin plates, and there may be penetration twins. The artificially produced crystals were first observed by Wohler, who discovered them in the arsenical sublimate of a furnace in which cobalt ores were roasted for the manufacture of cobalt blue. The prismatic oxide is frequently formed in this way during the roasting of arseniferous minerals if the condensation takes place at temperatures above 250° C., and it has been found to occur in mines in which arsenopyrite has been decomposed by combustion. The oxide may also be obtained in this form from solution if crystallisation takes place at a high temperature or if a hot solution is cooled rapidly; thus it may be obtained from a boiling saturated solution of the oxide in aqueous alkali, or from a hot solution in arsenic acid; also by the addition of ammonia to a boiling saturated ammoniacal solution of the oxide and rapidly cooling. Prismatic crystals have also been separated from the solution resulting from the action of nitric acid on silver arsenite.

The crystallographic elements are given as a:b:c = 0.4040:1:0.3445 and β = 86°3'. The cleavage on the (010) face is perfect. The optical axial angle 2H = 66°14' for red light and 65° 21' for yellow light. The crystals exhibit strong double refraction. The optical character is negative. The density of claudetite is 3.85 to 4.15 and the hardness 2.5.

Vitreous Arsenious Oxide
Vitreous Arsenious Oxide from the Furnaces at Wiluna (Australia)
The vitreous or amorphous form of arsenious oxide, " white arsenic glass," may be prepared by heating ordinary white arsenic, preferably under slight pressure but not at a temperature sufficiently high to volatilise too large an amount, and condensing the fume at a temperature just above the point of fusion, say 350° to 400° C. It is sometimes formed even below 315° C., the melting point of the monoclinic crystals. The operation is generally performed in a cast-iron bell-covered pan; the vitreous arsenic collects as a layer in the bell and by continually adding arsenious oxide to the pan the process is continued until the layer is about one inch in thickness. When freshly formed, it is transparent, but it gradually becomes opaque owing to transformation to the octahedral form and in this state it resembles porcelain. The arsenic glass may be kept in the transparent condition by excluding air, or by confining it in thoroughly dried air, hydrogen or carbon dioxide. When placed in boiling water octahedral crystals rapidly form at the surface of the arsenic glass. The latter often appears to retain its transparency under cold water, in which it is more soluble than the octahedral form; according to Winkler, however, transformation to the octahedral takes place thus - on being immersed in water the vitreous arsenic is dissolved at the surface and the layer of solution so formed deposits crystals of the less soluble variety; dissolution of the vitreous and deposition of the octahedral is repeated towards the interior until the transformation is complete. The change is retarded by the presence of alcohol. If a trace of iodine is added to a piece of vitreous arsenic while undergoing transformation, the latter is coloured more intensely than the octahedral form and the progress of the change may thus be observed. The heat of transformation of vitreous arsenic to the octahedral form is 2400 calories, and the heat of transformation of the monoclinic to the vitreous is 1200 calories.

Unlike the octahedral form, vitreous white arsenic on heating melts before volatilisation begins. The density of the glass has been variously given as 3.70 to 3.88; Winkler found the density under water to be 3.7165 at 12.5° C. but under petroleum 4.6815. The glass is brittle and its hardness is comparable with that of Iceland spar.

Colloidal arsenious oxide may be obtained in a highly dispersed condition by the vaporisation of arsenic in the electric arc and oxidation of the fume in a current of air. The size of the particles thus obtainable corresponds with the upper limit of the colloidal state (100 μμ).

In smoke prepared by volatilisation of arsenious oxide, the particles, which consist almost entirely of octahedral crystals, show only a slight degree of aggregation. In aerosols prepared by rapidly cooling the vapour of the oxide the number of particles per unit volume decreases rapidly during the first hour, especially in concentrated sols containing 150 to 500 mg. per cubic metre. After a preliminary ageing period the variation of mean particle weight with concentration is linear.

Dispersed in aqueous medium, arsenious oxide forms a negatively charged colloid; the magnitude of the charge decreases on increasing the concentration of hydrogen ion, but there is no reversal. The coagulating effect of positive ions increases in the order Li, Na, K, Mg, Ca, Ba, Al. If an alkali chloride is first added to the negative sol, the ionic antagonism increases with the series Ca, Ba, Al. If a salt of the metal of higher valency is first added, then a slight ionic antagonism is shown with the alkali salt, but it disappears with time.

Arsenious oxide is not appreciably volatile at ordinary temperatures, but vapour is emitted at 100° C. Vapour density determinations indicate that at lower temperatures the molecules are mainly As2O6 but, as the temperature rises, dissociation occurs, which is appreciable at 850° C. and practically complete at 1800° C., simple As2O3 molecules being formed. The following values for the vapour density (Theory: As2O6 = 13.76) were obtained by Biltz:

Temp., °C.518769851105912561450158417321800
Vap. Density13.9213.6213.1512.7612.369.418.817.326.93


The vapour is odourless.

The molecular weight, determined ebullioscopically by dissolution of the octahedral crystals in nitrobenzene, agrees according to Biltz with the formula As2O6. Determinations in water, which yields a slightly acidic solution, indicate that the solute molecules contain only one atom of arsenic, apparently existing as the very weak acid H3NaO3 or HAsO2 (vide infra). When arsenious oxide is reduced with zinc dust in the presence of carbon disulphide, the product consists in part of the yellow modification of arsenic; Erdmann therefore suggested the following related formulae:

and

The X-ray spectrum has been investigated by Robinson, who also examined the secondary and tertiary radiations emitted by the oxide under the influence of molybdenum K rays as primary X-radiation. Whiddington estimated the frequencies of high-speed electrons ejected from the oxide by impinging X-rays.

Arsenious oxide is not a very polar substance chemically, and it would be expected to have a rather small molecular moment. That this is so was shown by Clark, who found the electrostatic moment to be 1.3×10-19 electrostatic units. This value was not affected by varying the strength of the electric field, so that it appears to be a permanent characteristic of the molecule. The moment is of the right order to conform with Debye's theory of permanent dipoles. There was no evidence of any definite orientation of the molecules save parallel or antiparallel to the field.

In the solid state arsenious oxide is a very poor conductor of electricity.

The rate of dissolution of arsenious oxide in water is extremely slow and solubility data have been very discordant, probably due in part to insufficient time having been given for saturation, but also owing to the difference in solubility of the crystalline and amorphous varieties and to the tendency of the latter to pass into the octahedral form. The following figures are reliable, however, having been obtained by constantly agitating mixtures of the pure octahedral oxide and water in a thermostat, saturation being approached from above; periods of 10 to 14 days were necessary for the attainment of equilibrium and the arsenic was determined iodometrically.

Solubility of As2O., in water

Temperature, ° C.Grams As2O3 per 100 grams H2O.
01.21
151.66
252.05
39.82.93
48.23.43
624.45
755.62
98.58.18


From these results Anderson and Story deduced the following equation for the solubility at θ° C.:

S = 1.21 + 0.021θ + 0.000505θ2

Schnellbach and Rosin, after 131 days' agitation of the oxide with water, the tube being revolved end over end, found the solubility at 25° C. to be 2.03 g. in 100 g. H2O. For the more soluble vitreous form Winkler determined the solubility in 100 c.c. of water to be 37 g. at the ordinary temperature and 11.86 g. at the boiling point. Small octahedral crystals were deposited in the former case within 12 hours and the solubility gradually diminished until, after 3 or 4 weeks, it approximated to that of the octahedral form. For the monoclinic crystals Claudet found the solubility in 100 parts of water to be 1.75 parts at the ordinary temperature and 2.75 parts at 100° C. The rate of dissolution of arsenious oxide is accelerated by the presence of acid or alkali.

The density and refractive index of aqueous solutions increase linearly with concentration. The following values have been determined:

Concentration, g. per. l.1.7963.2125.0607.18410.1312.8514.368
D (25° C.)1.00141.00251.00391.00571.00801.01021.0113
n (22° C.).1.333091.333261.333401.333761.334171.334501.33469


Increase in temperature causes a slight decrease in refractive index. The heat of dissolution of the octahedral form is -7550 calories at 18° C. The velocity of crystallisation from supersaturated solutions corresponds with -dc/dt = kc, where c is the concentration; the temperature coefficient for the interval 0 to 25° C. is zero.

Arsenious oxide is soluble in a number of organic liquids. Thus, 100 parts of absolute alcohol dissolve at 15° C. 0.025 part, and at boiling point 3.402 parts of the octahedral form. The solubility is increased by the addition of water. The vitreous modification dissolves to the extent of 1.060 parts per 100 at 15° C. and addition of water decreases the solubility. Esters of arsenious acid may be obtained by heating, with stirring, a mixture of the oxide and an alcohol in the presence of a hydrocarbon such as benzene, toluene or xylene.

At 25° C. 100 grams of glycerol very slowly dissolve 20.8 grams of the oxide. In ethyl malonate the solubility in 100 g. is 0.058 g. at 15° C. and 0.061 g. at 100° C. Arsenious oxide is volatile in ethyl malonate vapour, 0.09 g. having been observed to be carried over during the distillation of 100 g. of the ester. The oxide dissolves in warm ethylene glycol, but no definite chemical compound is obtainable from the solution. The vitreous form dissolves slightly in ether, carbon disulphide, fatty oils and turpentine.

The absorption spectrum of 0.1N aqueous solutions of arsenious oxide differs from that of aqueous solutions of alkali arsenites. This is characteristic of weak acids, the un-ionised molecules of which appear to be capable of absorbing more light than ionised molecules; there is little or no difference in the absorption spectrum of a strong acid and its salts.

The aqueous solution of arsenious oxide is colourless and sufficiently acidic to cause a slight reddening of litmus, the pH of the saturated solution being approximately 5. The solution is a poor conductor of electricity, there being only slight ionisation, and the purity of the oxide may conveniently be determined by measuring the conductivity of a saturated solution. The most common impurity is arsenic pentoxide, which is indicated by the reaction of the solution to methyl orange or methyl red. The oxide, even in the most dilute solutions which can be examined cryoscopically, is believed to contain some associated molecules and in addition to be converted almost completely into the weak acid HAsO2 or H3NaO3. The following values for the molecular weight in aqueous solution have been obtained: by the boiling point method 92.5, and by the freezing point method 99.17. This supports the view that at 0° C. the trioxide is in the hydrated form, probably as metarsenious acid, HAsO2. With increasing concentration, however, association increases to a limiting value corresponding with As2O3, and evaporation of the aqueous solution yields only crystals of the oxide itself. Titration with standard alkali appears to indicate that the solute behaves as a monobasic acid, the salt produced being NaH2NaO3, and although the electrical conductivity increases on dilution, the increase is accounted for by hydrolysis and is not due to further ionisation of the acid.

Orthoarsenious acid, H3NaO3, corresponding to phosphorous acid, has not been isolated, although alkali salts of the type M3NaO3 are known; even these in solution appear to behave as salts of a monobasic acid. Walden suggested that the acid in solution was dimetarsenious acid, H2As2O4 or HO.OAs:AsO.OH, the solute molecule appears to contain only one arsenic atom.

The solution obtained by neutralisation of an aqueous solution of arsenious oxide with sodium hydroxide exhibits the same electrical conductivity and freezing point depression as an aqueous solution of sodium metarsenite, NaAsO2, of the same concentration. Conductivity measurements also suggest that the potassium salt produced by neutralisation must be of composition KAsO2, since the difference between the limiting equivalent conductivity, corrected for hydrolysis, of a neutral aqueous mixture of 1 mol. of KOH with 0.5 mol. of As2O3, and that of a similar mixture of sodium hydroxide with arsenious oxide, is equal to the difference in the ionic mobilities of K+ and Na+. From a mixture of equivalent amounts of potassium hydroxide and arsenious oxide a salt may be crystallised which, according to cryoscopic measurements, appears to exist in aqueous solution as K2As2O4. The conductivity of an aqueous solution of tripotassium arsenite, K3NaO3, is not the same as that of a mixture of aqueous arsenious oxide and potassium hydroxide in the molecular ratio 1As2O3: 6KOH and of corresponding concentration.

Assuming the formula H3NaO3 for the acid, Goldfinger and von Schweinitz calculated the first dissociation constant, k1, from the neutralisation curve to be 2×10-10 to 8×10-10, whereas spectroscopic determination of the ionic concentrations in the presence of varying amounts of alkali gave 1×10-14 to 6×10-14 for the second dissociation constant. The long wave limit of continuous absorption in molecular solution is 2680 A. for H2NaO3 and 2800 A. for HNaO3. Cernatescu and Mayer deduced the dissociation constant of arsenious acid from the hydrolysis constants of the sodium and potassium salts to be 9×10-10. Wood obtained the value 6.3×10-10 from the velocity of saponification of methyl acetate in the presence of sodium metarsenite, but by the electrical conductivity method the mean value found was 26.5×10-10, a discrepancy attributed to the presence of slight impurity and the fact that air was not excluded from the conductivity cell. The following values for the molecular conductivity, μ, and for the ionisation constant, k, were obtained (c = concentration of the solution of metarsenious acid):

c0.195N0.086N0.0542N0.0385N
μ0.05530.06540.07560.0824
k×101038.623.920.116.9


The view that aqueous arsenious oxide behaves as a weak monobasic acid is supported by the observation of Thomsen that the heat of neutralisation of 1 mole of As2O3 with 2 moles of NaOH was 13,780 calories and that the addition of a further 4 moles of NaOH liberated only 1800 calories. The neutralisation curve, whether determined conductometrically or potentiometrically, indicates the replacement of one equivalent of hydrogen only, as also does the curve obtained by plotting the depressions of the freezing point against the composition of the mixtures during neutralisation. These methods, however, are not able to decide between the formulae HAsO2 and H3NaO3.

The acidity of arsenious acid in aqueous solution is increased by the addition of mannitol, sorbitol or α-mannitan, probably owing to the formation and superior ionisation of an acid of the type HAsD2 (D representing the diol residue). When sublimed arsenious oxide is heated with water on a water-bath for 5 hours, the dissolved acid has less than the normal acidity, but by boiling the solution for 7 hours under a reflux condenser it attains its original acidity.

Solubility of Arsenious Oxide
Solubility of Arsenious Oxide in Aqueous Hydrochloric Acid of Varying Concentration (Temperature 15° C., approx.).
Arsenious oxide in solution exhibits a slightly amphoteric character, but its basic nature is extremely feeble. By determining the solubility in various concentrations of hydrochloric acid at 15° C. (approx.) Wood obtained the curve shown in fig. From the fact that the first effect of the addition of hydrochloric acid is to cause a steady fall in the solubility, it is evident that the acidic dissociation constant of arsenious hydroxide is much greater than the basic constant. It is only when the hydrogen ions have reached such a concentration to make the acidic ionisation of the hydroxide impossible that it becomes possible to see the effect which a further increase in the concentration of hydrochloric acid has on the solubility of arsenious oxide by virtue of the basic character of its hydroxide. Beyond the minimum of the curve the increase in solubility can only be accounted for by the assumption that the hydroxide possesses feeble basic properties. The basic dissociation constant was calculated to be of the order of 1×10-14. The ionisation was considered to be

(i) As(OH)3AsO(OH)2- + H+
(ii) As(OH)3As(OH)2+ + OH-

The basic ionisation constant at 25° C. in a solution having an ion concentration of 0.1 equivalent per litre is given as

[AsO+][OH-]/H3NaO3 = 0.15×10-14

Chemical Properties of Arsenious Oxide

When pure hydrogen is passed over heated arsenious oxide reduction to arsenic occurs with consequent loss in weight and formation of water; the reduction becomes appreciable at 185° C. In aqueous solution, and in the presence of acid or alkali, nascent hydrogen causes reduction to arsine, and a similar reduction may be brought about electrolytically, but the amount of arsine liberated depends upon the nature of the cathode, the following being given in order of efficiency: Pb, Zn, Cd, Sn, Ag, C (graphite), Fe, Pt, Al, Au, Co, Ni and Pd; in the case of the first five metals, the reduction proceeds as a unimolecular reaction. A polarographic investigation of the electro-reduction in acid solutions of arsenious oxide, using the dropping mercury cathode, has been made. In 0.1N or N hydrochloric acid, the current-voltage curve exhibits four sudden increases of current and two maxima. The first rise is due to the electro-reduction of arsenious acid to arsenic, probably by the primarily deposited hydrogen. The second rise is very steep and is due to the formation of arsine; it occurs at the more negative potentials, the greater the concentration of arsenious oxide. The third increase, which is followed by a prominent maximum, is probably caused by absorption of positively charged dissociation products of arsenious acid; it is suppressed by the addition of small quantities of dyes, the maximum practically disappearing on addition of 0.001M solutions of methylene blue and fuchsin hydrochloride. The fourth increase is attributed to the evolution of hydrogen from the hydrogen ions of the strong acid. Similar results were obtained with sulphuric and nitric acids at various concentrations, and the form of the curve was unchanged by addition of potassium chloride or calcium chloride. Under the influence of occluded hydrogen from palladium or platinum, the reduction in aqueous solution produces arsenic only.

Fluorine reacts vigorously with arsenious oxide to yield a colourless liquid consisting of arsenic trifluoride and oxyfluoride. The oxide becomes incandescent in hydrogen fluoride and if heated with acid fluorides, or with calcium fluoride and sulphuric acid, the trifluoride may be distilled from the mixture. The action of chlorine and hydrogen chloride has previously been mentioned. When chlorine is passed into an aqueous suspension containing 70 to 80 per cent, of the oxide at a temperature of 60° to 70° C., the absorption of chlorine is rapid and exothermic, about 70 per cent, of the arsenious oxide being converted to the pentoxide and the remainder to the trichloride. When solutions of arsenious oxide in hydrochloric acid are boiled, the arsenic volatilises, the amount escaping depending on the concentration of the acid. With solutions containing less than 180 g. HCl per litre the concentration of the arsenic remaining in the undistilled liquid rises, although some arsenic passes over; the ratio, arsenic: acid, becomes practically constant when the solution contains 185 g. HCl per litre (i.e. HCl.10H2O), but with more than this the concentration of arsenic remaining rapidly falls. When such solutions are exposed to the air, slight oxidation occurs.

Oxidation is readily brought about by hypochlorites and by chlorates. In the latter case, in the presence of hydrochloric acid, the reaction is independent of the concentration of the arsenious acid and, according to Kubina, reduction of the chlorate first to a hypothetical intermediate product occurs, probably as follows:

ClO3- + Cl- + 2H+H2ClO3 + Cl

This reaction, which proceeds at a measurable rate, is succeeded by the following rapid reactions:

H2ClO3 + 4Cl- + 4H+ → 5Cl + 3H2O
and
3Cl2 + 3As3O3- + 3H2O → 3AsO43- + 6Cl- + 6H+

The chlorine ion produced does not accelerate the oxidation, as might be expected, owing to the high initial concentration of this ion.

Arsenious oxide and arsenites may be oxidised similarly by bromine or bromic acid. In hydrochloric acid solution the reaction with bromine may be represented by the equation

As2O3 + 2Br2 + 2H2OAs2O5 + 4HBr

and if the concentration of the hydrochloric acid is less than 24 per cent, the reaction proceeds entirely from left to right. Under such conditions arsenites may be titrated accurately with bromine, the end-point of the titration being unaffected by the actual concentration of hydrochloric acid. If the latter exceeds 24 per cent., however, the reverse reaction may take place, the equilibrium conditions depending on the concentrations of arsenate, bromide and hydrochloric acid. The oxidation by means of bromic acid is extremely slow at ordinary temperatures, but is accelerated by the addition of sulphurous or sulphuric acid. At 40° C., in the presence of an excess of H+ ions, and at lower temperatures in the presence of sulphuric acid, the reaction proceeds at a measurable rate. It is autocatalytic and of the second order, according to the equation dx/dt = kax(1 -x); in the presence of 0.1 mol. of sulphuric acid the velocity constant is 9.7 at 30.7° C. The initial production of hydrobromic acid must be due to the interaction of bromic acid with arsenious acid; the latter, however, does not appear to influence the reaction other than by acting as an inductor of the reaction between hydrobromic and bromic acids. The action of the sulphuric acid is proportional to the square of the concentration of H+ ion and the addition of neutral sulphates, which reduce this concentration, retards the reaction. On the addition of hydrogen bromide, the reaction proceeds in accordance with the formula dx/dt = ka(b + x)(1 - x), where b is the concentration of hydrobromic acid and x the initial concentration of bromic acid; the velocity constant, k, remains as before. Hydriodic acid has a similar but much greater effect. Arsenic acid, which is the final product of the reaction, also acts as a positive catalyst, although its effect is about nine times weaker than that of an equivalent quantity of sulphuric acid. The addition of neutral halides also accelerates the reaction, the relative effects of potassium chloride, bromide and iodide being as 1:15:3000.

Vitreous arsenious oxide is coloured brown by iodine vapour, but the octahedral form appears to be unaffected.

In hydrochloric acid solution arsenious acid is oxidised by iodine, but the reaction is reversible owing to the reducing action of hydriodic acid. The kinetics of the reaction were first investigated by Roebuck, who concluded that the balanced reaction could be represented thus -

H3NaO3 + I3- + H2OH3AsO4 + 2H+ + I- (i)

and assumed that the reverse reaction proceeded in two stages:

H3AsO4 + H+ + I- = H3AsO4.HI (ii)
H3AsO4.HI = H3NaO3 + HIO (iii)

From the reaction velocities in the neighbourhood of equilibrium, he determined the equilibrium constant [As3O3-][I3-]/[As4O3-][I-]3[H+]2 to be 1.5×105 and the temperature coefficient between 10° and 0° C. to be 1.41. The equilibrium adheres to the requirements of the law of mass action over a considerable range of concentration, and Roebuck's views are confirmed by recent work. Liebhafsky considers that, over a sufficient concentration range, the equilibrium constant, k1 expressed as [H3AsO4][H+]2[I-]3/[I3-][H3NaO3] is equal to k2/k3 derived from the velocity equations (a) -dI3-/dt = k2[I3-][H3NaO3]/[H+][I-]2 and (b) + dI3-/dt = k3[H3AsO4][H+][I-]. Both the forward and the reverse reactions are concerned with the rate-determining step

HIO + H3NaO3 = H3AsO4 + H+ + I- (iv)

which can be interpreted by assuming that the concentration of hypoiodous acid is governed by the relatively rapid equilibrium

I3- + H2OH+ + 2I- + HIO (v)
which is the sum of the equilibria
I3-I2 + I- (vi)
and
I2 + H2OH+ + I- + HIO (vii)

since I3- is not hydrolysed directly, while I2 and H2O react with moderate speed. Thus the rate of the latter hydrolysis is the limiting rate realised as the concentrations of H+ and I- ions decrease. At 0° C. k2 = 9.4×10-4 and k3 = 6.3×10-3, so that the equilibrium constant k1 = 1.5×10-1.

The addition of neutral salts, such as chlorides and bromides of the alkali and alkaline earth metals, at concentrations from 0.5 to 3N, also nitrates of sodium and potassium, causes the reversal of reaction (vii), and consequently considerably reduces the rate of oxidation of arsenious acid, whilst augmenting slightly that of the reduction of arsenic acid, thus shifting the equilibrium point in the direction of arsenious acid formation. The effect of sulphates is much less than that of corresponding chlorides. The maximum effect is obtained with lithium chloride, that of the other alkali chlorides being in the order K, Na, NH4. Jozefowicz studied the reaction at 25° C. in the presence of excess of hydriodic acid and found the heat of reaction (i) under these conditions to be -1640 calories, while the heats of reactions (vi) and (vii) from left to right were respectively -4300 and -23,200 calories. Washburn and Strachan found the heat of the reaction between arsenious acid and iodine to be 1360 calories, the effect of temperature being represented by lgK = -1.3495 + 0.00372t, and the free energy of the reaction. RTlogeK = 5690 + 5.42T joules for ionic concentrations of about 0.1 equivalent per litre.

Under appropriate conditions either the direct or the reverse reaction represented in equation (i) may proceed to completion, and the reactions are therefore applied to the volumetric estimation of arsenites and arsenates respectively. Thus an arsenite is oxidised quantitatively to arsenate if the hydriodic acid is removed as quickly as it is formed and the solution kept approximately neutral. This is best accomplished by adding sodium bicarbonate to the arsenite solution; sodium hydroxide and carbonate could not be used owing to their reaction with iodine. On the other hand, the reduction of arsenate to arsenite by means of hydriodie acid proceeds to completion in strongly acid solution. The action of the acid is not catalytic, but appears to be similar to that of a neutral electrolyte as mentioned above, causing reversal of reaction (vii). The reducing action of the hydriodic acid is augmented by the presence of potassium iodide. In a solution containing 25 per cent, of the latter salt and 3.6 per cent, of hydrochloric acid the reaction is complete in 5 minutes if the reacting mixture is heated to 100° C. In employing this reaction for the determination of arsenic acid, the liberation of iodine from hydriodic acid by means of atmospheric oxygen should be prevented by the addition of a small quantity of sodium bicarbonate previous to the addition of the potassium iodide. Pure arsenious oxide may be used as a trustworthy standard for iodometric estimations.

Solubility data at 25° C. for solutions of arsenious oxide in aqueous alkali halides have been obtained.

Aqueous arsenious acid is readily oxidised by hypoiodites, iodates and periodates. The oxidation by iodic acid or iodates is an induced reaction which proceeds with rapid initial acceleration owing to the catalytic effect of the I- on produced. The reaction appears to proceed according to the following scheme, the first reaction only occurring at a measurable rate:

(i) IO3- + 2I- + H+HIO + 2IO-
(ii) 2IO- + 2AsO33- → 2AsO43- + 2I-
and
HIO + AsO33-AsO43- + H+ + I-

The induction period of the reaction may be curtailed by (1) the presence of an excess of iodic acid, (2) an increase in the concentrations of the reactants, (3) the addition of a trace of arsenic acid, (4) the addition of a mineral acid and (5) exposure to sunlight. On the other hand, the period may be prolonged by the addition of mercuric chloride or by violent shaking. The proportion of the iodine liberated increases with the arsenious acid concentration and passes through a maximum. The iodine appears on the surface of the solution even though the latter may be covered with benzene (or occasionally it appears at a nucleus on the glass). The reduction of periodate to iodate by means of arsenite is a bimolecular reaction and is of the first order with respect to both components. At 25° C. it proceeds according to the velocity equation

d[IO4-]/dt = 5.5[IO4-][AsO2-]

the units being minutes and gram-molecules. The reaction velocity is independent of the concentration of H+ ion over the range [H+] = 1.3×10-3 to 3.4×10-7.

Arsenious oxide in the solid state is not affected by oxygen under ordinary conditions, but if subjected at high temperature to oxygen under pressure oxidation to the pentoxide results. Thus, if heated at 400° to 480° C. with oxygen at pressures of 130 to 180 atm., oxidation occurs, but is incomplete; the amount oxidised increases with the temperature. According to Razuvaev and Malinovskij, at 200° to 300° C. the optimum pressure for oxidation by air is 60 to 80 atm., and under these conditions the reaction is complete in about 20 minutes; finely divided iron has a weak and copper a strong catalytic effect. In the presence of potassium iodide or activated carbon, a suspension of the oxide in water is oxidised by air or oxygen at 130° to 140° C. and 4 to 5 atm. pressure. The usually accepted reaction is

As2O3 + O2 = As2O5

but Reissaus, from a study of the effects of heat on the oxide and metallic arsenites both in the absence and presence of air or oxygen, concluded that direct oxidation was not involved, but that the change was invariably based on the reaction:

5As2O3 → 3As2O5 + 4As

Arsenious oxide is formed subsequently from the arsenic liberated and then undergoes further decomposition; the arsenic thus acts as an oxygen carrier. By thus heating arsenious oxide under pressure pure arsenic may be prepared.

In neutral, weakly acid or weakly alkaline solutions, arsenious oxide and arsenites are very stable towards gaseous oxygen and such solutions may be kept indefinitely without undergoing change, but in the presence of an excess of alkali, oxidation readily takes place. For this reason standard arsenious solutions containing alkali gradually diminish in titre. Thus a 0.1N solution of arsenious oxide in N sodium hydroxide suffers a daily loss equal to about 0.176 per cent, of As2O3. The oxidation rate is directly proportional to the alkalinity of the solution. Weakly alkaline solutions (pH 7 to 9) have been found to be unchanged after 18 months. Stable solutions are best prepared by dissolving pure arsenious oxide in carbonate-free sodium hydrogen carbonate solution; in such solutions micro-organisms do not develop nor is arsine produced.

An arsenical solution containing the equivalent of 1 per cent, of arsenious oxide appears as an official preparation in the British Pharmacopoeia, 1932, under the name liquor arsenicalis or Fowler's solution. It is prepared by dissolving 10 g. of the oxide in 100 c.c. of a 5 per cent, potassium hydroxide solution, warming as may be necessary, and adding 500 c.c. of distilled water; this solution is neutralised with dilute hydrochloric acid and made up to 1 litre. This preparation is extremely stable both to light and air, and is compatible with both acids and alkalies. Fungoid growths, however, frequently develop, the causes apparently being contamination either of the water used or of the air with which the solution has been in contact, and also a suitable pH value which allows the mould to develop. The addition of preventive agents, e.g. 0.25 per cent, of chloroform, has been recommended. Conductivity and potential measurements suggest that the arsenic is present in solution as As2O3 and not as potassium arsenite; dissolution in the potassium hydroxide results in the formation of some potassium metarsenite, but this regenerates the oxide on neutralisation with hydrochloric acid. This official solution replaces the " Fowler's solution " of previous editions of the Pharmacopoeia, the composition of which was similar to that introduced by Fowler about 1778 under the name of Compound Spirit of Lavender. In this the arsenious oxide was dissolved in aqueous potassium carbonate and compound tincture of lavender added; the liquid was thus reddish in colour and alkaline in reaction. Its disadvantage was its frequent incompatibility in modern dispensing; it also undergoes oxidation on keeping. It was therefore supplemented by a 1 per cent, solution containing hydrochloric acid, liquor arsenici hydrochloricus, and also by a simple 1 per cent, solution, liquor acidi arseniosi (B.P. Codex); the latter, however, deposits octahedral crystals on keeping 2 to 3 weeks.

Matignon and Lecanu observed that a concentrated solution of arsenious oxide in sodium hydroxide under an oxygen pressure of 50 atmospheres at 80° C. was oxidised to the extent of 10.9 per cent, after 5 hours. Under ordinary atmospheric conditions the absorption of oxygen by the arsenite occurs more readily in the presence of a second easily oxidisable substance such as sodium sulphite or ferrous sulphate. In the latter case the amount of oxygen absorbed depends upon the quantity of sodium hydroxide present. Manchot considered that the oxidation of Fe++ to Fe+++ involved the activation of 1 equivalent of oxygen which was used in the conversion of arsenite to arsenate, but Wieland and Franks found that in the most concentrated solution of arsenite obtainable, of which the pH was 6, the activation of only 0.88 equivalent occurred; for solutions of pH 10, corresponding with the metarsenite, NaAsO2, activation corresponded with 0.6 equivalent, and for more strongly alkaline solutions corresponding with Na2HNaO3 it corresponded with not more than 1 equivalent. When alkalinity corresponded with Na3NaO3 activation exceeded 1 equivalent owing to spontaneous oxidation of the arsenite independent of the influence of the ferric salt. This catalytic action of readily oxidisable substances, from which it would appear that one chemical change is able to promote another of the same type, is probably due to the formation of intermediate compounds. The addition of a cerous salt dissolved in concentrated potassium carbonate solution to an excess of aqueous potassium arsenite results in the induced oxidation of the latter, which acts as oxygen acceptor. The cerous solution passes to the quadrivalent eerie state and, according to Baur, 2 molecules of oxygen are fixed on the arsenite for 1 molecule on the cerium. In the presence of glucose, however, the cerous salt exerts purely a catalytic action and under suitable conditions a very small quantity can effect the atmospheric oxidation of an unlimited quantity of the arsenite; this is due to the more profound reducing action of the glucose, which continually converts the eerie salt to the cerous state, in which the cerium is again capable of fixing atmospheric oxygen. These reactions have been investigated from the standpoint of the relative oxidation potentials of the reacting substances.

Another interesting example of induced oxidation is the reaction between chromic acid and a manganous salt in the presence of arsenious acid. The chromic acid is reduced to a chromic salt, while the manganous salt is oxidised to the manganic state and the arsenious acid to arsenic acid, probably in accordance with the equations:

(1) CrVI + AsIII = CrIV + AsV
(2) CrVI + MnII = CrIII + MnIII

The relation of MnII oxidised to AsIII oxidised is known as the " induction factor," and is found to be about 0-5 at the beginning of the reaction, whatever amount of manganous salt may be present; but the value diminishes as the action proceeds, for as the chromate concentration falls, the tervalent manganese is reduced by the tervalent arsenic to an increasing extent. The presence of a trace of potassium iodate produces great irregularity in the induction factor and lowers its initial value to about 0.25.

A further example of an induced reaction consists in the atmospheric oxidation of an ammoniacal solution of arsenite, which is brought about by the addition of cobaltous sulphate, the latter also being oxidised.

The oxidation of arsenites by oxygen in the presence of sodium hydroxide is affected by the presence of other types of catalysts. Thus in the presence of copper sulphate and with less alkali than corresponds with Na3NaO3, the velocity of oxidation is very small, but with an increased amount of sodium hydroxide present, copper hydroxide or oxide is formed and the action is accelerated, and indeed copper oxide itself may be used as catalyst. With an excess of sodium hydroxide and a suitable quantity of copper oxide, normal sodium arsenite may be completely oxidised to arsenate in a few hours. Similarly the presence of an excess of sodium carbonate facilitates oxidation.

On the other hand, an aqueous solution containing an alkali di-hydrogen arsenite and sodium hydrogen carbonate is extremely stable towards oxygen and at ordinary temperature no oxidation can be observed after 4 months, although slight oxidation occurs on heating. Shilov and Pevzner, in studying the atmospheric oxidation of potassium dihydrogen arsenite, found that salts of copper, iron and manganese, free iodine and titanic acid were ineffective as oatalysts, whilst chromates, molybdates and cerium salts were only slightly effective. Complete oxidation was possible, however, in the combined presence of oxides of nitrogen and hydriodic acid. Into a column filled with glass beads, aqueous solutions of the arsenite and potassium iodide were introduced dropwise. Into the same column were introduced air, hydrogen chloride and nitrous acid. The addition of these components could be regulated and by proper adjustment of the ratio of the ingredients and the velocity of the gases through the column, complete oxidation to arsenate was attained. In view of the ready oxidation of arsenious acid by iodine, the following series of reactions appears possible:

2HNO2 + 2HI = 2H2O + I2 + 2NO (i)
I2 + KH2NaO3 + H2O = KH2AsO4 + 2HI (ii)
2NO + O2 + H2O = HNO3 + HNO2 (iii)

We have already seen, however, that equation (ii) becomes reversed in acid medium, and the prevention of this reversal here appears to be due to the removal of the hydriodic acid according to equation (i).

Tingle observed that a solution of arsenious oxide in aqueous alcohol, after boiling for 26 hours, was oxidised to arsenic acid, but this was denied by Edgerton. Kessler observed that when sodium arsenite was undergoing oxidation by chromic acid, oxidation by atmospheric oxygen occurred simultaneously.

Arsenious oxide is oxidised to the pentoxide by hydrogen peroxide and by ozone; in alkaline solution the oxidation by ozone is incomplete.

By exposing an aqueous solution of arsenious acid to X-rays oxidation to arsenic acid occurs accompanied by the liberation of an equivalent amount of hydrogen.

When heated with sulphur, arsenious oxide yields the disulphide or trisulphide according to the proportion of sulphur employed; sulphur dioxide is also produced. An aqueous solution of arsenious oxide or of an arsenite is coloured yellow on passing hydrogen sulphide and in the presence of an acid a precipitate of arsenic trisulphide is produced. The addition of arsenious oxide in small quantity (0.2 per cent.) to aqueous sodium carbonate increases the amount of hydrogen sulphide that can be absorbed by the solution, but on regenerating with air the whole of the hydrogen sulphide is not evolved. Arsenious oxide is soluble in aqueous sulphuric acid, the solubility varying somewhat irregularly with the acid concentration and the temperature; the oxide may be recovered either by crystallisation or by distillation of the acid. It also dissolves readily in fuming sulphuric acid and from the solution sulphato-compounds of the type As2O3.nSO3 may be obtained, thus indicating a base-like tendency in arsenious oxide; but the products have not the properties of metallic sulphates. The presence of arsenious oxide retards the oxidation of sodium sulphite.

When dry arsenious oxide is fused with sodium thiosulphate a mixture of the di- and tri-sulphides results. In aqueous solution and in acidified solutions of arsenites the addition of aqueous sodium thiosulphate causes the precipitation of arsenious sulphide after a sharply defined induction period, the duration of which is in inverse proportion to the thiosulphate concentration and practically independent of the concentration of the acid. The nature of the acid, however, influences the time, a longer period being observed with hydrochloric acid than with acetic acid. This reaction may be employed for the experimental demonstration of induction periods and of the relation between concentration and reaction velocity. There is evidence of the formation of intermediate unstable thio-compounds. Thus, from a solution containing arsenious oxide in hydrochloric acid, potassium chloride and sodium thiosulphate, a white compound, potassium arsenothiosulphate, K3As(S2O3)3, may be precipitated by means of alcohol. This compound in aqueous solution deposits arsenious sulphide " after a shorter or longer time according to concentration and temperature," i.e. after an induction period. It is stable when dry but on heating decomposes thus:

2K3As(S2O3)3 = As2S3 + 3K2SO4 + 3SO2 + 3S

In alkaline solutions of arsenites, thioarsenates are formed, which may be crystallised out, while sulphite or bisulphite remains in solution, thus:

Na3NaO3 + Na2S2O3 = Na3NaO3S + Na2SO3
and
Na2HNaO3 + Na2S2O3 = Na3NaO3S + NaHSO3

A small quantity of arsenic is precipitated in each case. Sodium dihydrogen arsenite yields a considerable precipitate of arsenic and also of the red disulphide. The polythionates react similarly to thiosulphates, yielding sulphite, thioarsenate and arsenate. The per-sulphates also cause oxidation to arsenate. Sodium hydrosulphite added to an aqueous solution of arsenious oxide precipitates brown arsenic; if the solution is strongly acid the precipitate also contains arsenious sulphide and sulphur. Arsenious oxide heated at 120° C. in a sealed tube with sulphur monochloride yields arsenic trichloride, thus:

2As2O3 + 6S2Cl2 = 4AsCl3 + 3SO2 + 9S

Arsenious oxide appears to undergo no change when heated in gaseous ammonia; it is insoluble in liquid ammonia, but dissolves readily in hot aqueous ammonia. Heated with solid ammonium chloride, arsenious chloride and ammonia are produced. With nitrogen iodide the following reaction occurs:

3As2O3 + 2N2H3I3 + 6H2O = 3As2O5 + 4NH3 + 6HI

In alkaline solution, oxidation may be brought about by hydroxylamine or nitric oxide; thus with the former the following reactions occur concurrently:

Na3NaO3 + NH2OH = Na3AsO4 + NH3 3NH2OH = NH3 + N2 + 3H2O

With nitric oxide the reaction appears to be termolecular -

Na3NaO3 + 2NO = Na3AsO4 + N2O

the velocity increasing with increasing concentration of alkali hydroxide. A similar change occurs when an excess of sodium arsenite is added to a fresh solution of nitric oxide in a strongly alkaline solution of sodium sulphite. Nitrous acid acts extremely slowly on a solution of the oxide in aqueous sulphuric acid; the reaction speed is at a maximum when the sulphuric acid is of density 1.39 to 1.47.

In the finely divided state, arsenious oxide can be oxidised with oxygen under pressure in the presence of nitric acid, which acts as oxygen carrier. The oxidising action of the nitric acid usually ceases when reduction to nitric oxide has occurred, but nitric acid is regenerated by oxygen under pressure, thus:

4NO + 3O2 + 2H2O = 4HNO3

The oxidation of the arsenious oxide is practically quantitative at 90° C. with nitric acid of 40 to 60 per cent, concentration and the oxygen under 20 atm. pressure; the rate of the oxidation and of re-formation of nitric acid increases with rise in temperature and with increase in oxygen pressure. Increase of acid concentration accelerates oxidation, but retards regeneration. Nitric acid alone also brings about oxidation, oxides of nitrogen and nitrous acid being formed as intermediate products. It is suggested that the nitrous acid produces nitrogen trioxide, which determines the rate of oxidation:

2HNO2 = N2O3 + H2O
N2O3 + H3NaO3 = H3AsO4 + 2NO

The speed of the reaction is thus proportional to the concentration of arsenious acid and to the square of the concentration of nitrous acid present. The effect of the presence of mercuric salts on the velocity of this oxidation is peculiar; at a concentration of 7.7×10-6 mol. per litre the reaction is completely inhibited, but with diminishing concentration the inhibiting effect becomes less marked until, at 7.7×10-9 mol. per litre, there is a definite positive catalytic effect which becomes more marked at the extremely low concentration of 7.7×10-11 mol. per litre. The retarding effect of the mercuric salt may be overcome by the addition of a small amount of halogen acid to the nitric acid. The oxidation by nitric acid may be facilitated by the addition of a small quantity of iodine or potassium iodide.

Phosphorus converts arsenious oxide to phosphide. When an aqueous solution is heated with phosphorus to 200° C., a precipitate of arsenic and arsenic phosphide is formed. A mixture of the dry oxide with phosphorus trichloride heated to 120° C. in a sealed tube reacts as follows:

6PCl3 + 5As2O3 = 3P2O5 + 6AsCl3 + 4As

The addition of phosphorus trichloride to an aqueous solution precipitates brown amorphous arsenic; the reaction does not occur if the trichloride is first dissolved in water, nor when phosphorous acid is used. The reaction, which accords with the equation

3PCl3 + As2O3 + 9H2O = 2As + 3H3PO4 + 9HCl

is very delicate and is able to detect the presence of 0.000075 g. As per c.c. Phosphorus tribromide and triiodide react similarly, but more slowly; in the latter case the precipitate is contaminated with red phosphorus. The reaction proceeds also when an aqueous solution of the tribromide or triiodide is used. Phosphorus oxychloride yields arsenic trichloride at 250° C., thus:

As2O3 + 2POCl3 = 2AsCl3 + P2O5

Phosphorus pentachloride also reacts to yield arsenic trichloride. Arsenious oxide in aqueous solution is reduced by hypophosphorous acid, especially on boiling, when phosphine is liberated and brown arsenic precipitated. The reduction is readily brought about by calcium hypophosphite dissolved in 10 parts of hydrochloric acid (dens. 1.126), this salt being preferable in use to the sodium salt and providing an extremely sensitive reagent, although in the presence of slight traces of arsenic the brown colour may appear only after the lapse of 20 to 30 minutes. Arsenious oxide dissolves in arsenic trichloride to form an oxychloride.

Carbon brings about reduction of arsenious oxide at a temperature below red heat, while in carbon monoxide reduction begins at 60° C. Silicon tetrachloride heated for 30 hours at 270° to 280° C. with the oxide yields arsenic trichloride, whilst silicochloroform when heated with the oxide in the presence of aqueous sodium hydroxide or sodium hydrogen carbonate causes reduction to arsenic, thus:

As2O3 + 3SiHCl3 + 9NaOH = 3Si(OH)4 + 9NaCl + 2As

Reduction to arsenic also results on heating with boron nitride.

When heated to redness with an alkali metal or with zinc or aluminium, reduction of the oxide readily occurs. Reaction with various metallic oxides, hydroxides and carbonates results in the formation of arsenites and arsenates. Certain hydroxides, however, particularly those of iron, aluminium, chromium, magnesium, manganese and zinc, have the power of removing arsenious oxide from solution, although no chemical combination appears to take place. The action was first observed by Bunsen and Berthold in connection with ferric hydroxide; they suggested the use of the latter as an antidote for arsenical poisoning, the removal of arsenic being attributed to the formation of a hydrolysed ferric arsenite. The formation of such a compound could not be demonstrated, however, and Biltz showed that the phenomenon was one of reversible adsorption. He observed that the amount of arsenious oxide adsorbed from solution diminished with the ageing of the ferric hydroxide, and suggested that at equilibrium the distribution of the oxide between the hydroxide and water could be represented by the usual adsorption isotherm, C1 = kC21/n, where C1 is the concentration of arsenious oxide in the ferric hydroxide, C2 the concentration of arsenious oxide in solution, and k and n are constants. Later workers have shown that this relation is not quite true, and that for samples of the adsorbent prepared under different conditions and of different ages, whilst there is little variation in the value of n, k varies between wide limits. Thus in an investigation by Boswell and Dickson it was found that 1/n varied only from 0.183 to 0.284, but that k varied from 33.3 to 200. Sen observed that the adsorptive power of a sample of ferric hydroxide after ageing four months was diminished by 50 per cent. If the hydroxide is prepared by the addition of ammonium hydroxide to a solution of a ferric salt the adsorptive capacity is greater than if prepared by addition of a ferric salt to ammonium hydroxide. Also the adsorptive power decreases the higher the temperature at which the hydroxide is produced, and with specimens prepared at 75° and 100° C. the amount of adsorption deviates greatly from the relation expressed above.

Both arsenious acid and sodium arsenite are strongly adsorbed by ferric hydroxide. The amount adsorbed from a given volume of solution increases with the initial concentration, and for a given concentration the adsorption increases with the volume. The amount of adsorption depends also on the quantity of ferric hydroxide employed, but in a decreasing ratio. The curvature of the adsorption isotherms depends therefore on these factors. Adsorption takes place very rapidly, but the time taken for the limiting value to be reached appears to vary with conditions, from one to twelve hours being recorded. Temperature has little influence.

Sodium hydroxide is adsorbed appreciably by ferric hydroxide and the presence of the alkali in solution diminishes the amount of arsenic adsorbed. On the other hand, salts such as potassium or ammonium chloride have no effect on the adsorption. The ferric hydroxide, especially when fresh, is liable to become peptised by arsenious acid in certain dilutions. Apparently peptisation results when a definite quantity of arsenious acid has been adsorbed by each particle. Excess of arsenious acid, however, causes flocculation.

Colloidal saccharated iron is sometimes used in place of ferric hydroxide as an antidote in arsenical poisoning, but its adsorptive capacity depends on the alkalinity of the medium. Thus a commercial preparation containing 0.75 per cent, of sodium hydroxide was found to adsorb 12.57 per cent, of arsenious oxide (reckoned on the amount of iron present); addition of alkali increased the adsorption until, with 1.28 per cent, of sodium hydroxide present, there was a maximum adsorption of 27 per cent. The addition of acid correspondingly diminished the adsorption. A gel of ferric magnesium hydroxide, if prepared without boiling, also adsorbs arsenic from sodium arsenite solutions.

Ferric hydroxide is far from being the best adsorbent for arsenious oxide. According to Boutaric and Perreau, the most active hydroxides are those of zinc and manganese, while the following are in order of decreasing adsorptive activity: cadmium, chromium, iron, aluminium and magnesium. The last two adsorb arsenic from Fowler's arsenite solution but not from sodium cacodylate. Aluminium hydroxide behaves similarly to ferric hydroxide; the presence of potassium chloride has no effect upon the equilibrium, which is reversible, nor does sucrose or alcohol affect the adsorption. If the aluminium hydroxide gel is boiled with water for 15 minutes, the adsorptive power at the ordinary temperature is greatly reduced. Adsorption of arsenious oxide by chromium hydroxide is much greater than by ferric or aluminium hydroxide and the action differs from that of the two latter gels in that the equilibrium is not reversible and, in fact, no true equilibrium can be attained. Moreover, arsenious acid is adsorbed to a greater extent than is sodium arsenite, whereas in the case of aluminium hydroxide no difference is observed. Zirconium hydroxide exhibits considerable adsorptive power, as also does suitably prepared wood charcoal.

After an extensive study of the adsorption of arsenious oxide by metallic hydroxides, Sen concluded that this type of adsorption resembles that of cations by manganese dioxide, and that the chemical affinity between the adsorbent and the substance adsorbed plays an important part, thus differing from adsorption by charcoal. It has been observed that soils having a high absorption capacity for bases also absorb the arsenite ion from solutions of 0.001 to 0.01N concentration. The absorption increases with time, without reaching an end-point, and the process follows the normal adsorption equation C1 = kC21/n. The addition of ferric oxide or calcium carbonate to the soil considerably increases the capacity for absorption, but such salts as calcium sulphate or copper sulphate have no effect.

The reducing power of arsenious oxide has been the subject of much investigation, especially in regard to its reaction with the more important oxidising agents. Arsenious acid and arsenites undergo oxidation in the presence of chromic acid or dichromates. With a mixture of potassium dichromate and sulphuric acid, the oxidation of arsenious acid proceeds at a rate nearly proportional to the first power of the concentration of the dichromate, to the first power of the concentration of the arsenious acid and to the l.4th power of that of the sulphuric acid; the inexactness of the proportionality is probably due to incomplete ionic dissociation of the dichromate and to the influence of the hydrogen ion concentration on the dissociation of the arsenious acid. Between 0° and 10° C. the temperature coefficient is low, the increase in the rate of oxidation over this interval being only 26 per cent. If potassium iodide is added to the mixture of arsenious acid, dichromate and sulphuric acid, an induced reaction occurs in which the arsenious acid acts as the inductor and the iodide as the acceptor of the oxidation. Under these conditions the rate at which the chromic acid is reduced equals the sum of the rates at which it is reduced in solutions of arsenious acid and potassium iodide separately, the former reaction being retarded as much as the latter is accelerated. The temperature coefficients of the single actions and the joint action are equal.

The oxidation of arsenious oxide or arsenites by means of potassium permanganate has attracted considerable attention owing to its analytical applications. It was early recognised that the course and speed of the reaction and the nature of the products depended on conditions of temperature, acid concentration and the nature of the ions present. In neutral or alkaline solution the oxidation proceeds slowly, but it is more rapid in acid solution. In alkaline solution the reduction of the permanganate results in the precipitation of manganese dioxide according to the equation

3As2O3 + 4KMnO4 = 3As2O5 + 2K2O + 4MnO2

but an excess of permanganate is necessary for the complete oxidation of the arsenious oxide. Some manganese remains in colloidal solution as manganic hydroxide, and this is precipitated if an electrolyte such as potassium sulphate be present. If the arsenite is in excess, reduction proceeds further, the precipitate consisting of an indefinite mixture of manganous oxide and manganese dioxide; in complete absence of air the product is manganous hydroxide. These precipitates adsorb a considerable quantity of arsenic from the solution. In slightly acid medium a green colloidal solution of manganese dioxide may be obtained; this on keeping deposits as a black precipitate containing arsenic.

In the presence of a mineral acid the complete oxidation of arsenic to the quinquevalent state by means of permanganate is difficult, probably owing to the fact that whereas the permanganate is rapidly reduced to the manganic condition, the complete reduction to the manganous state usually occurs slowly. With sulphuric or nitric acid this is the case even at 100° C., but with hydrochloric acid the reaction proceeds more rapidly to completion. Under the influence of certain catalysts, however, the reaction may be accelerated so that it takes place practically instantaneously in the cold, and accurate estimation of arsenious acid may thus be made by titration with permanganate. Suitable catalysts are potassium iodide, potassium iodate, silver chloride, iodine chloride and osmium tetroxide. In the presence of hydrochloric acid, the concentration of which should be from 0.5 to 2N, the addition of one drop of 0.0025N solution of potassium iodide or iodate to the portion to be titrated is sufficient. In the presence of sulphuric acid, the latter should be at least 0.5N and should contain about 1 g. of sodium chloride per 100 c.c., the same amount of iodide being added. The effect of the catalyst is presumably due to the primary formation of hypoiodous acid by the action of the tervalent manganese on the iodide, the hypoiodous acid immediately oxidising the arsenious acid. The favourable effect of the iodide is hindered by the presence of such compounds as mercuric salts or hydrofluoric, metaphosphoric and tungstic acids, which allow the formation of undissociated or slightly ionised complexes. The end-point of the titration may be determined potentiometrically. If osmium tetroxide is used as catalyst, a few drops of a 0.01M aqueous solution should be added to each portion to be titrated.

Quadrivalent cerium salts may be volumetrically determined by arsenious acid in aqueous sulphuric acid solution using a manganese salt as catalyst, with potassium iodate present as a promoter. Excess of arsenious acid is used and back-titrated with permanganate. If nitric acid is present in place of sulphuric acid, an alkali chloride and a trace of iodine are used to promote the action of the manganese salt. Direct potentiometric titration with arsenious acid may also be employed.

If arsenious oxide is heated in a sealed tube with ferric chloride in aqueous hydrochloric acid (1.5 to 4N), reaction proceeds according to the equation

2FeCl3 + H3NaO3 + H2O ⇔ 2FeCl2 + H3AsO4 + 2HCl

The equilibrium constant K as given by [H3NaO3][FeCl3]2/[H3AsO4] [FeCl2]2[HCl]2 is 0.0354 at 107° C. and 0.117 at 127° C. The forward and reverse reactions are both termolecular and are accelerated by hydrochloric acid. The thermal value of the reaction from left to right is 18,000 calories. Sarma studied the reaction in aqueous solution at 50° C. in the presence of potassium iodide as catalyst. Equilibrium was attained in 10 hours. The reaction is bimolecular in respect of the ferric salt and unimolecular in respect of the arsenious acid. There appear to be two consecutive reactions -

(i) 2Fe+++ + 2I- = 2Fe++ + I2
(ii) I2 + H3NaO3 + H2O = 2HI + H3AsO4

the latter being relatively slow, but increasing with the concentration of the catalyst. The heat evolved in the oxidation of 1 gram-molecule of arsenious acid was found to be 13,600 calories.

Silver nitrate in ammoniacal solution may be completely reduced to silver by aqueous arsenious oxide. The reduction is hindered by the presence of ammonium sulphate, owing to the decrease in concentration of the hydroxyl ions; neutral salts such as sodium sulphate or sodium nitrate have no effect. Similarly, auric chloride may be reduced to gold. At 20° C. an aqueous solution of vitreous arsenious oxide reacts 4 to 5 times as rapidly as an aqueous solution of the octahedral form; the greater rate of dissolution in water of the former variety has been mentioned, but from supersaturated solutions of the two forms there is no appreciable difference in the rates of deposition. The explanation of the inferior reducing power of the crystalline variety may be that there exist "anisotropic molecules " which only slowly lose their anisotropic properties. An ammoniacal solution of arsenious oxide heated with cupric sulphate in a sealed tube at 100° C. causes reduction to the cuprous salt. The reduction is retarded by the presence of ammonium salts, but is not affected by the presence of neutral salts such as potassium chloride or nitrate.

The reduction of arsenious acid by means of stannous chloride, the well-known " Bettendorff's test" depending upon the reaction. In the presence of concentrated hydrochloric acid, a voluminous brown precipitate, consisting mainly of yellow arsenic with small amounts of tin, is formed. Velocity measurements lead to the view that the reaction takes place between arsenious and chloride ions and the complex H2SnCl4. Dilution with water decreases the velocity and ultimately prevents precipitation. The reaction is endothermic.

Arsenious oxide exhibits considerable catalytic activity, which may act either positively or negatively. The effect on many oxidation and reduction reactions has been mentioned above. Other examples are the increase in the rate of dissolution of zinc by dilute acids and the retardation of the dissolution of marble. In the latter case a concentration of 0.005N As2O3 reduces the velocity constant for the dissolution of marble in 0.1N HCl by 12 per cent.

Uses of Arsenious Oxide

The most important application of arsenious oxide is in the manufacture of arsenates, which are used extensively, in the form of poison sprays or dusts, as insecticides to control the various pests which attack fruit and vegetable crops. The most important arsenates in this respect are those of calcium and lead. The former is effective in destroying the boll weevil of the cotton fields, while the latter controls the codling moth, plum curculio, cabbage worm, potato beetle, tobacco hornworm and other pests. Emerald green (copper acetoarsenite), magnesium arsenate and manganese arsenate are also used as insecticides. The refined white arsenic is usually employed in the manufacture of the above products, but the crude arsenical flue dust or "treater" dust is used for the production of weed killers, fungicides and wood preservatives, while some is converted into sodium arsenite for sheep dipping purposes. Railroads may be kept free from grass and weeds by the judicious use of arsenical preparations. Wood preservatives may contain aqueous arsenious oxide, copper arsenite or acetoarsenite, or zinc metarsenite, such being useful for telegraph poles, fences, mine props, etc. In India and in Germany arsenious oxide is used with chromates and chromic acid, which fix the arsenic in the wood. Emerald green is used in preservative paints for such purposes as the painting of the bottoms of ships to prevent weed and barnacle growth. It is also used to a limited extent in the arts for giving vivid green tints. In the form of arsenical soap arsenic is used in taxidermy to prevent insect damage. It has been stated, however, that skins treated with arsenates do not dry-out well, and tend to heat when piled.

Arsenious oxide is used in the glass industry as a decoloriser, opacifier and refining agent. When the oxide is added to either soda-lime-silica or potash-lead oxide-silica glasses, the arsenic is either wholly or mainly retained on melting. Thus Turner found that with the former type of glass and with 10 parts of arsenious oxide to 1000 parts of sand, practically the whole of the arsenic remains in the glass when melted in closed pots at 1400° C., while with 250 parts of the oxide to 1000 of sand the amount retained at 1350° C. was 78 per cent, and at 1400° C. 53 per cent, of the total added, the amounts varying somewhat with charging conditions. During melting, 40 to 70 per cent, of the arsenious oxide, even in the absence of oxidising agents, is converted to arsenic pentoxide or arsenates, the solutions of which in glass at high temperatures are remarkably stable. The arsenic exerts no beneficial effect on the process of melting, the temperature of which is raised by large amounts, and the time of melting is prolonged. The rate of melting may be accelerated by the addition of nitre. In batches rich in silica, addition of more than 2 parts of arsenic to 1000 parts of sand tends to cause a surface scum. The green tint due to iron is definitely reduced by the arsenic, the improvement being continuous with increasing proportions of the oxide. When the latter is added to the extent of 150 to 250 parts per 1000 of sand, an opalescent glass is produced, the degree of opacity depending on the amount of arsenious oxide added. The latter is similarly employed as an opacifying agent in the production of opalescent glazes.

A soda-lime glass rendered colourless by means of arsenic may turn yellow under the influence of sunlight owing to conversion of ferrous to ferric iron. Also a glass containing arsenic tends to darken in colour if reworked; the presence of the element facilitates clarifying in lead glasses, but is objectionable where these must be subjected to subsequent heating. The presence of arsenic in chemical glass is undesirable because of its ready extraction by strong acids and alkalis. Foodstuffs, however, may be safely contained in suitable arsenic- bearing glass without contamination. In 6-hour boiling tests with vinegar and a concentrated sugar solution, a good soda-lime glass containing 0.27 per cent, of arsenious oxide showed no loss of the latter by extraction, and the amount extracted under similar conditions by a 2N-sodium carbonate solution corresponded to only 0.2 per cent, of the total loss of weight.

The temperature of devitrification of glass is lowered by the presence of arsenic. With a series of glasses made from soda-lime batches containing from 0 to 250 parts of arsenious oxide per 1000 parts of sand a minimum devitrification temperature (575° to 600° C.) occurred with a glass made from the batch in which 100 parts of the oxide had been employed. Devitrification of glasses with higher arsenic content is probably hindered by the increased hardness of the glass.

In recent years use of the oxide as a constituent of cement has been advocated, especially in Sweden. Thus, a mixture of Portland cement (60 to 70 per cent.) and white arsenic (40 to 30 per cent.) heated to 200° to 250° C. affords a hydraulic cement of normal setting time and of less solubility than ordinary cement, so that lime liberation is inhibited and the resistance to water improved. Wooden structures exposed to the action of sea water may be protected by spraying with a concrete composed of white arsenic, cement and sand in the proportions 1:3:12. The arsenic makes the mixture elastic and helps the cement to adhere to the wood. There is, however, danger in the too widespread application of arsenic in the directions described above.
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