Arsenious acid As

The toxic action of white arsenic has been attributed to its inhibitory action on the oxidative processes,9 partly owing to the effect of the change of pH on the enzyme concerned. Small quantities of arsenious acid reduce the power of suitably prepared extracts of animal tissues to oxidise reduced phenolphthalein. The oxidation of tartaric acid at the ordinary temperature and at 37-5° C. is inhibited by arsenious acid, as also is the respiration and fermentation of yeast,10 but the latter... 

Prepare a faintly alkaline solution of arsenious acid as follows Dissolve a minute quantity of arsenious oxide in not more than 2 or 3 drops of hydrochloric acid dilute to 10 cc. and add, without heating, a considerable amount of sodium bicarbonate in excess of what is necessary to neutralize the acid. To this solution add, drop by drop, a solution of iodine, and determine whether any free iodine disappears. Write the equation. So far as the state of oxidation of the arsenic is concerned, the reaction is exactly the reverse of the one preceding. Recall a previous instance in which the direction of a reaction of oxidation and reduction is changed on passing from an acid to an alkaline solution. 

Alkalies. Alkalies completely decompose the molecule. Green and Price have shown that on adding even very dilute cold solutions of sodium hydroxide or carbonate, chlorovinyl arsenious oxide is never obtained, but acetylene and arsenious acid, as follows ... 

Arsenious acid, As(OH)s, behaves as a weak acid with a dissociation constant of about 8X10 at 25°. As is the case with boric acid, the addition of n-mannitol to aqueous solutions of arsenious acid increases the acidity of the solution. The formation constants for complexes between polyhydroxy compounds and the arsenite ion have been found to be considerably smaller than those for the corresponding borate complexes. This is also reflected in the absolute mobilities of such compounds during electrophoresis in arsenite solution. 

The tendency for the lower oxidation state to become more stable than the higher oxidation state as foe atomic number increases within a group in foe periodic table is a general phenomenon. For example, we find that it is more difficult to reduce phosphoric acid to phosphorous acid than to reduce arsenic acid to arsenious acid, as indicated by foe following standard electrode potentials ... 

Arsenic Peroxides. Arsenic peroxides have not been isolated however, elemental arsenic, and a great variety of arsenic compounds, have been found to be effective catalysts ia the epoxidation of olefins by aqueous hydrogen peroxide. Transient peroxoarsenic compounds are beheved to be iavolved ia these systems. Compounds that act as effective epoxidation catalysts iaclude arsenic trioxide, arsenic pentoxide, arsenious acid, arsenic acid, arsenic trichloride, arsenic oxychloride, triphenyl arsiae, phenylarsonic acid, and the arsenates of sodium, ammonium, and bismuth (56). To avoid having to dispose of the toxic residues of these reactions, the arsenic can be immobi1i2ed on a polystyrene resia (57). 

Elemental phosphoms from the electrothermal process is a distilled product of high purity and yields phosphoric acid pure enough for most industrial uses without any further treatment. The main impurity is ca 20—100 ppm arsenic present in the phosphoms as the element and in the phosphoric acid as arsenious acid. To remove the arsenic, the phosphoric acid destined for food, pharmaceutical, and some industrial-grade appHcations is treated with excess hydrogen sulfide, filtered, and blown with air to strip out excess H2S. This treatment generally reduces the arsenic content of the phosphoric acid to less than 0.5 ppm. The small amount of filter cake is disposed of in approved chemical landfills. 

The demand for metallic arsenic is limited and thus arsenic is usually marketed in the form of the trioxide, referred to as white arsenic, arsenious oxide, arsenious acid anhydride, and also by the generally accepted misnomer arsenic. 

Chlorate Analysis. Chlorate ion concentration is determined by reaction with a reducing agent. Ferrous sulfate is preferred for quaHty control (111), but other reagents, such as arsenious acid, stannous chloride, and potassium iodide, have also been used (112). When ferrous sulfate is used, a measured excess of the reagent is added to a strong hydrochloric acid solution of the chlorate for reduction, after which the excess ferrous sulfate is titrated with an oxidant, usually potassium permanganate or potassium dichromate. 

The solubility of AS2O3 in water, and the species present in solution, depend markedly on pH. In pure water at 25°C the solubility is 2.16 g per lOOg this diminishes in dilute HCl to a minimum of 1.56g per lOOg at about 3 m HCl and then increases, presumably due to the formation of chloro-complexes. In neutral or acid solutions the main species is probably pyramidal As(OH)3, arsenious acid , though this compound has never been isolated either from solution or otherwise (cf. carbonic acid, p. 310). The solubility is much greater in basic solutions and spectroscopic evidence points to... 

In iodimetry, quantitative oxidation of reducing agents, such as arsenious acid (H2As03) may be carried out by employing standard solutions of iodine as shown under ... 

By the action of nascent hydrogen upon soluble arsenic compounds, as by the introduction of arsenious acid into on apparatus evolving hydrogen —... 

The rate of a reaction is expressed as the increase in the amount of a particular product, or the decrease in the amount of a reactant, per second at a time t. For reactions in a fluid phase, it is usual to substitute concentrations, or for gases partial pressures, for amounts. (See Section 5.6 for treatment of the kinetics of some solid-state reactions.) Note that the numerical value of the rate may depend on how we choose to define it. For example, in the oxidation of aqueous iodide ion by arsenic acid to give triiodide ion and arsenious acid,... 

It follows that the equilibrium constant K is given by kf/kr. The reverse reaction is inverse second order in iodide, and inverse first-order in H+. This means that the transition state for the reverse reaction contains the elements of arsenious acid and triiodide ion less two iodides and one hydrogen ion, namely, H2As03I. This is the same as that for the forward reaction, except for the elements of one molecule of water, the solvent, the participation of which cannot be determined experimentally. The concept of a common transition state for the forward and reverse reactions is called the principle of microscopic reversibility. 

In a notorious case in Manchester in 1900-1901 6000 people were affected by drinking beer containing 15 p.p.m. arsenic 70 people died as a result. The arsenic originated from sulfuric acid containing 1.4% arsenious acid, produced from As-containing iron pyrites. This sulfuric acid was used in a process to manufacture glucose for the brewers, and the glucose contained several hundred parts per million arsenic. 

Pure arsenic has been prepared by reducing carefully purified ammonium dihydrogen arsenate at 1000° C. in a current of ammonia, the arsenic being finally resublimed in a vacuum.7 The element may also be obtained as an amorphous precipitate by reduction of aqueous arsenious acid, for example, by means of sodium hypophosphite,8 or by the addition of a few drops of phosphorus trichloride.9 The reaction in the latter case probably takes the following course ... 

The oxidation of arsine may be accomplished by means of the halogen oxyacids and their salts,8 although not so readily as with the halogens themselves. Hypochlorites and hypobromites cause complete oxidation to arsenic acid, but side reactions are liable to occur, especially if the gas is present in excess. Chloric add slowly oxidises arsine to arsenious acid a trace of silver nitrate catalyses the reaction. Chlorates are quite inactive. More complete oxidation results with solutions of bromic acid and bromates, iodic acid and iodates, especially in the presence of catalysts. The reactions are of the type represented by the equation8... 

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