Elemental silver, in the 0 oxidation state, is fairly common as is the +1 monovalent argentous silver ion. Most toxicological studies have been conducted on these two forms of silver. The argentic +2 and rarer +3 oxidation states have not been well studied. The +2 state is generated by oxidizing acidified solutions of +1 ions with persulphate, fluorine or ozone. Once formed, the +2 and +3 states are fairly stable in acidic solutions, especially phosphoric acid; reduction to the +1 state occurs slowly. Tables 3.1 to 3.6 give the physical and chemical properties of silver and some of its salts.
In acidic solutions the +1 state is soluble and mobile but silver precipitates as the solution becomes alkaline, and at a pH between 7.5 and 8.0 silver hydrolyzes as the oxide or a basic salt. The precipitation of silver from alkaline solutions is also dependent on the other ions present. If ammonia or some other complexing agent is present, the silver may stay in solution. Silver will precipitate as a halide if Cl-, Br-, or I- ions are introduced, as Ag2S if H2S or S2- is present, and as a thiosulphate, phosphate, chromate or arsenate in the presence of S2O3-, PO43-, CrO42- or AsO43-, respectively (Anon. 1990, Boyle 1968, MRI 1975, Kharkar et al. 1968, Thompson 1973).
Studies using radiosilver released from nuclear power stations show that in fresh water at around pH 8.0, at least 50% of the silver precipitated, and within 48 hours 80% was adsorbed onto sediments. In marine and estuary conditions where the chloride ion is prevalent, silver is rapidly desorbed, 70% in 24 hours, and remains in solution (Murray and Murray 1973, Fukai and Murray 1973).
In natural waters, silver is in the +1 state as a sulphate, bicarbonate or nitrate. Silver chloride complexes, such as Na(AgCl2), form when NaCl or KCl are present. This is the most likely form of silver in seawater. In areas of high biological activity, such as reef complexes in the ocean, [Ag(S2O3)2]3- is probably only stable at low temperatures for a short time near the surface. It is soluble in alkaline and neutral solutions but decomposes to S, Ag2S and Ag2SO4 in acidic conditions (Taylor 1980).
Complex sulphide or polysulphide ions or hydrosulphides may form in water with high levels of H2S or S2-. Silver may also be dissolved as double complexes of sulphur, arsenic, antimony, tellurium and selenium. Silver may exist in colloidal form in water as an integral part of, or adsorbed onto, various humic complexes as AgCl, Ag2S, Ag2Se and Ag3AsS3, or dissolved as acetates, tartrates and other organic compounds. It may also be adsorbed onto plankton or inside the tissues of micro-organisms (Taylor 1980).
Theoretical equilibrium calculations indicate that Ag+ and AgCl are the dominant forms of silver in aqueous solution; at 1 µg/L, 60% is Ag+. Adding organic complexing agents such as NTA, citrate, glycine and cysteine to the equilibrium calculations has little effect on the outcome. Model calculations for an estuary with a dissolved silver concentration of 0.04 µg/L and an H2S and HS- concentration of 0.01 µg/L, have shown that at the river input end of the estuary AgHS was the most prevalent form of silver followed by Ag+ and AgCl. At the marine end of the estuary, AgCl- was dominant followed by AgHS, AgCl32- and AgCl43- (Morel et al. 1973, Hem 1977, Jenne et al. 1978).
In photographic film development the silver bromide grains dispersed in the emulsion are `activated', by a poorly understood process, and thus become more susceptible to reduction by mild reducing agents in the `developer'. These exposed grains are reduced to black metallic silver. To prevent the unexposed grains from reduction when subsequently exposed to light, they are dissolved in `hypo' or `fixer' which contains S2O32- or thiosulphate ions. This results in the formation of a silver thiosulphate complex, [Ag(S2O3)2]3-, which is washed out of the emulsion leaving behind only the black metallic silver deposits in areas where the light struck the emulsion (Sienko and Plane 1966).
The majority of silver from photoprocessing occurs in an insoluble form. Theoretical calculations of organic and inorganic silver complexes indicate that due to the low solubility of silver sulphide, and the high affinity of silver for sulphide, little free silver, <10 to 12 µg/L, would occur at equilibrium in effluents or surface waters that contained any sulphide (Lockhart 1980).
There is no evidence that silver is naturally transformed to a hazardous biologically available form (such as mercury into methyl mercury). Ionic silver is more toxic to aquatic organisms than silver compounds. Thiosulphate-complexed silver breaks down to silver sulphide which is less toxic than the silver ion.