556 CYCLES OF ORGANIC AND INORGANIC SUBSTANCES 



acid-base equilibria have mainly been taken from the compilation 

 of Bjerrum et al. (1958) and from supplementary tables which are 

 being kept up to date in Stockholm. The equilibrium constants 

 for redox reactions are taken from a compilation which is being 

 prepared in Stockholm and which will eventually be published. 

 In these tables, full references are made to the original literature; 

 in this paper, this is done only occasionally. 



The original equilibrium constants are very often those calcu- 

 lated from "zero activities" (based on an infinite dilution as the 

 standard state) in the law of mass action. In general, measured 

 values of pH and pE can be assumed to correspond to zero activi- 

 ties of H+ and e". For other species, the zero activity constant 

 should be corrected with activity coefficients for the medium of 

 ocean water. We shall use, as a very rough estimate, log A = —0.2 

 for univalent ions, and log/o = —0.8 for bivalent ions. 



Some equilibrium constants given in the literature refer to an 

 ionic medium, such as O.SM NaCl; or 0.5, 1, or 3>M NaC104. The 

 corrections from these media to sea water should be considerably 

 smaller, and have been neglected here in comparison with the 

 variation with temperature and pressure. 



All equilibria are more or less dependent on temperature. The 

 equilibrium constants are often determined only for 25°C, whereas 

 5°C is closer to the real average temperature. As seen from some 

 examples in the text, where log K for more than one temperature 

 is given, the difference may be several tenths of a unit. 



The pressure dependence of an equilibrium constant is given by 

 the formula (5 log K/hp)T - -^V/{RT In 10). At 5°C, the change 

 in log K would be approximately —1.9 X 10~^/>AF, with p in 

 atmospheres and AF in milliliters. For example, for the reaction 

 CaCOs (s) ;=± Ca++ + CO3--, AF ~ -50 ml, and so log K would 

 increase by about 0.2 unit at the average pressure (200 atm), and 

 by about 0.4 unit at the a\erage depth of the ocean floor. However, 

 many other equilibria, such as the dissociation equilibria of 

 H2CO3, will also change with pressure with comparable amounts. 

 No attempt has been made to correct for such variations. 



The following treatment of the equilibria may horrify some 

 physical chemists; certainly, it is not as strict as would be required 

 in treating equilibria in a laboratory. However, considering the 



