MAGNESIUM AND ZINC. IONIC EQUILIBRIA 455 



used, both effects are relatively inconspicuous. It is altogether 

 impossible, for example, to reduce the concentration of the hydro- 

 gen-ion given by an active acid like hydrochloric acid, by addition 

 of a salt containing a common ion, like sodium chloride, below 

 the limit at which indicators are affected, for there is no way of 

 introducing the enormous concentration of the other ion which 

 the theory demands. 



With more crude means of observation than indicators afford, 

 effects like this last, however, may sometimes be rendered visible. 

 This is the case with cupric bromide solution, to which potassium 

 bromide is added. The blue of the cupric-ion disappears from 

 view, while much cupric-ion is still present, because the brown 

 color of the molecular cupric bromide covers it up completely. 



The color changes discussed on p. 235 in connection with 

 the reaction FeCl 3 + 3NH 4 .CHS ^ Fe(CNS) 3 + 3NH 4 C1 also 

 illustrate this point. The red color is due entirely to undis- 

 sociated ferric thiocyanate. The reader should re-examine this 

 equilibrium carefully, writing the equation in full ionic form. 

 He will then be in a position to understand why the addition of 

 any ferric salt or of any thiocyanate to a given mixture will deepen 

 the red color, while the addition of any chloride (except ferric 

 chloride) or any ammonium salt (except ammonium thiocyanate) 

 will tend to lighten it. 



Special Case of Saturated Solutions. The commonest as 

 well as the most interesting application of the conceptions devel- 

 oped above is met with in connection with saturated solutions, 

 especially those of relatively insoluble substances. 



The situation in a system consisting of the saturated solution 

 and excess of the solute has been discussed already (read p. 120). 

 In the case of potassium chlorate, for example, we have the follow- 

 ing scheme of equilibria: 



KC10 3 (solid) ?= KC10 3 (dslvd.) * K+ + ClOr. 



