Determination of molecular hydrocyanic acid in water and studies of the chemistry and toxicity to fish of metal-cyanide complexes Public Deposited

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  • A reliable, easy, and inexpensive method for determination of molecular hydrocyanic acid (HCN) in solutions of simple and complex metal cyanides is described. The method was used to determine molecular HCN concentrations as low as 0.005 milligram per liter, and can be used for determination of even lower levels. It is a modification of a previously published method. A concentration column of glass beads coated with NaOH is employed, on which HCN displaced by air that has been bubbled through solutions under examination is trapped and concentrated for measurement of cyanide by a conventional analytical method. The apparatus could easily be modified for use in both field and laboratory situations where only limited facilities are available. Time periods required for attainment of equilibria upon dilution of solutions of metal-cyanide complexes, and also when metal salts and free cyanide are combined, were quite variable and ranged from several hours for the silver-cyanide complex to many months for iron-cyanide complexes kept in the dark. In solutions in which CuCN and NaCN were combined so that the molar ratio of CN to Cu was either 2.5 to 1 or 3 to 1, constancy of the HCN concentration usually was not attained even 110 days after preparation. The time to attainment of equilibrium through dissociation of the nickelocyanide complex ions generally was longer than that required for equilibrium to be attained in comparable experiments on complex formation, and it increased as the pH or the total cyanide concentration decreased; it is directly related to the percentage of total cyanide present as HCN at equilibrium. Results obtained at high total cyanide concentrations in nickelocyanide formation experiments were anomalous but verifiable by bioassay with fish. The HCN concentrations were at first unexpectedly low and then increased very slowly to the higher equilibrium levels. Cumulative dissociation constants (K[subscript D]) at 20°C for the Ag(CN)₂⁻, Cu(CN)₂⁻, Ni(CN)₄⁻², Fe(CN)₆⁻⁴, and Fe(CN)₆⁻³ complex ions, calculated from equilibrium levels of HCN, are 1.94 ± 2.82 x 10⁻¹⁹, 3.94 ± 1.75 x 10⁻²⁴, 1.00 ± 0.37 x 10⁻³¹, approximately 10⁻⁴⁷, and 10⁻⁵², respectively. The calculated constants for the tetracyanonickelate (II) and dicyanoargentate (I) complex ions inexplicably varied somewhat, increasing slightly with increase in total cyanide concentration and pH. Those for the tetracyanonickelate (II) and dicyanocuprate (I) complex ions showed close agreement with values recently reported in the literature, whereas the constants for the dicyanoargentate (I) and hexacyanoferrate (II) and (III) complex ions were materially different from presently accepted values. Possible unreliability of presently accepted stepwise constants for the cuprocyanide complex ions also was indicated. The acute toxicity of solutions of the different metal-cyanide complexes was generally found to be a function of the molecular HCN level, which increases with increase of total cyanide concentration and with decrease of pH. In some solutions however, a metal-cyanide complex ion per se was shown to be the major toxic component. The 48-hour median tolerance limits for bluegills of the dicyanoargentate (I) and dicyanocuprate (I) ions at 20°C were found to be approximately 9 and 4 mg/l as CN, respectively. The metallocyanide complex ions studied can be arranged in order of decreasing toxicity as follows: Cu(CN)₂⁻, Ag(CN)₂⁻, Ni(CN)₄⁻², and Fe(CN)₆⁻³ or Fe(CN)₆⁻⁴. A published empirical relationship between pH and 48-hour median tolerance limits of the nickelocyanide complex for a fish, determined without assurance that equilibria had been attained in test solutions, was compared with a calculated, theoretical relationship. Considerable divergence of the empirical and theoretical curves at pH values less than about 7.2 is ascribable mostly to the introduction of fish into test solutions long before equilibria had been attained in the solutions of low pH. Divergence at pH values greater than about 7.8 is attributable largely to moderate toxicity of the Ni(CN)₄⁻² complex ion itself. Slightly alkaline solutions of the silver cyanide complex, Ag(CN)₂⁻, become more toxic to sticklebacks with increase of chlorinity. The high toxicity in saline solutions, as compared with the toxicity in fresh water, is clearly attributable, at least in part or in some instances, to a molecular HCN content of the saline solutions much greater than that of comparable solutions prepared with fresh water. The two ligands CN⁻ and Cl⁻ compete for the silver ion, with which both ligands form complexes, and dissociation of the Ag(CN)₂⁻ ion, with production of HCN, consequently increases as the Cl⁻ ion concentration increases. Additional reasons for the observed increase of toxicity of solutions of the complex with increase of chlorinity can be an observed increase of the toxicity of HCN and a possible, similar increase of the toxicity of the complex ion. Experiments with ¹⁴C-labeled cyanide complexed with nickel showed that the complex does not penetrate readily into the body of a bluegill. The ¹⁴C accumulated in gill tissues much more markedly than it did in the blood and in tissues of internal organs sampled. When bluegills were exposed to solutions of the cyanide complexes of copper (I) and silver (I), considerable amounts of the metals accumulated in the blood and in tissues of internal organs, but little accumulation in gill tissues was observed. These results indicate that the cuprocyanide and silver-cyanide complexes enter the body of a bluegill much more readily than does the much less toxic nickelocyanide complex. The silver cation, however, apparently enters even more readily than does the silver-cyanide complex anion, the silver accumulating most markedly in the gill tissues of bluegills exposed to silver nitrate solutions, but also in their internal organ tissues.
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