1902-3.] Freezing-Point Depression in Electrolytic Solutions. 373 
grams, the ice melted in one hour would in the most unfavourable 
case only reduce the concentration by 0*1 per cent, and the 
freezing-point depression by a like amount. Most of the heat was 
apparently conducted by the thermometer, etc., the Dewar vessel 
forming a very efficient protection. 
Dissolved air . — Kaoult showed that water saturated with air 
had a freezing point 0*002° lower than water which was air-free. 
It is therefore of some importance to have the water and solution 
always saturated with air before the freezing point is observed. 
Owing to the method of air-stirring which we employed, the 
liquids whose freezing point we observed were always fully 
charged with air. 
Analysis . — The solutions investigated were either acids or 
chlorides. The acids were estimated by means of N/20 baryta 
solution with phenol phthaleine as indicator, and the chlorides 
by Volhard’s method with N/50 solutions of silver nitrate and 
ammonium thiocyanate. A decinormal solution of hydrochloric 
acid was taken as ultimate standard both for the baryta and 
for the silver nitrate. The baryta solution, however, was also 
checked with pure succinic acid, and the silver nitrate solution 
with pure potassium chloride. 
As an example of the constancy of composition of the different 
portions filtered off from the experimental solutions, and of 
the magnitude of the analytical error, we here give the analysis 
of the solutions for which the thermometric readings were detailed 
on p. 369. The concentrations of the solutions are given in terms 
of normal solutions. 
Acetic acid. Hydrochloric acid. 
First filtrate . . 0*09697 0*05225 
Second filtrate . . 0*09709 0*05213 
Mean, 0*09703 Mean, 0*05219 
Here the divergence of any result from the mean does not 
greatly exceed 0T per cent. At the outside we may take 0*2 
per cent, as the limit of error from sampling and analysis. 
Quantity of ice . — From the results of analysis the quantity 
of ice in contact with the solution at equilibrium could be roughly 
calculated in the following manner. In the above experiment 
with acetic acid, 11*5 c.c. of normal acetic acid solution were 
