Some Chemical Considerations on the Distribution of Iron in the Sea

Abstract

Three purely chemical topics which have aIisen in the course of the investigation describedin the preceding paper (Cooper, 1948)will be discussed; first, a theoreticaltreatment of the formation and stabilityof ferric phosphate in sea water; secondly,an explanation as to how 2:2'-dipyIidyl may react with ferIic iron at high dilutions in natural waters to give the red ferrous complex; and thirdly, a short account of an expeIiment which showed ferrifiuoIide to be unstablein seawater.. FORMATION AND STABILITY OF FERRIC PHOSPHATE IN SEA WATERTwo seIies of experiments in this laboratory by Harvey (1937) and by the writer (unpublished) on the precipitation of iron from sea water containing phosphate and hydroxyl ions (pH 8'1) do not agree. Sometimesmuch of the phosphate is precipitated, sometimesnone. Much closer control is necessary but it has not been clear what form this should take. The followingtheoretical examination, imperfect though it is, may help in the design of such contro1.Pugh & du Toit (1936)examinedthe ionic exchange or hydrolysis of ferric phosphate and silicate in relatively concentrated systems. Solutions of ferric chloIide were titrated with solutions of phosphate or silicate containing varying amounts of sodium hydroxide. A ferIic phosphate with a higher P04/Fe203 mol-ratio than 0'71 could not be prepared since compounds richer in phosphate hydrolysed until this ratio was reached. There is an optimum value of pH above which the ferIic phosphate hydrolyses further with increasing hydroxyl-ion concentration. The same is true of ferric silicate. At pH 8'0 and above, in presence of hydroxyl, silicateand phosphate in solution, ferric hydroxide would be the stable solid phase. The isoelectric point of ferIic hydroxide prepared from its chloIide is said to be pH 7'1 so that at pH 8'1 ferric hydroxide should be regarded rather as the very sparingly soluble ferrous acid, H3Fe03, which may be expected to dissociate H3Fe03~H+ + H2FeO~. Since FeOH++ is the dominant ferric cation at the isoelectric point, for electroneutrality of the micelle, the concentration of its anion, H2FeO~, should be twice that of its cation, FeOH++. This (Cooper, 1937b, equation 12) will be about 10-}2.54 at 18°C. The solubility product of ferrous acid, (H+) (H2FeO~)will therefore be of the order