soluble organic matter.
The hydroxides yield on ignition an oxide suitable for weighing
(Al_{2}O_{3}, Cr_{2}O_{3}, Fe_{2}O_{3}).]

DETERMINATION OF SULPHUR

PROCEDURE.--Add to the combined filtrates from the ferric hydroxide
about 0.6 gram of anhydrous sodium carbonate; cover the beaker, and
then add dilute hydrochloric acid (sp. gr. 1.12) in moderate excess
and evaporate to dryness on the water bath. Add 10 cc. of concentrated
hydrochloric acid (sp. gr. 1.20) to the residue, and again evaporate
to dryness on the bath. Dissolve the residue in water, filter if not
clear, transfer to a 700 cc. beaker, dilute to about 400 cc., and
cautiously add hydrochloric acid until the solution shows a distinctly
acid reaction (Note 1). Heat the solution to boiling, and add !very
slowly! and with constant stirring, 20 cc. in excess of the calculated
amount of a hot barium chloride solution, containing about 20 grams
BaCl_{2}.2H_{2}O per liter (Notes 2 and 3). Continue the boiling for
about two minutes, allow the precipitate to settle, and decant the
liquid at the end of half an hour (Note 4). Replace the beaker
containing the original filtrate by a clean beaker, wash the
precipitated sulphate by decantation with hot water, and subsequently
upon the filter until it is freed from chlorides, testing the washings
as described in the determination of iron. The filter is then
transferred to a platinum or porcelain crucible and ignited, as
described above, until the weight is constant (Note 5). From the
weight of barium sulphate (BaSO_{4}) obtained, calculate the
percentage of sulphur (S) in the sample.

[Note 1: Barium sulphate is slightly soluble in hydrochloric acid,
even dilute, probably as a result of the reduction in the degree of
dissociation of sulphuric acid in the presence of the H^{+} ions of
the hydrochloric acid, and possibly because of the formation of a
complex anion made up of barium and chlorine; hence only the smallest
excess should be added over the amount required to acidify the
solution.]

[Note 2: The ionic changes involved in the precipitation of barium
sulphate are very simple:

Ba^{++} + SO_{4}^{--} --> [BaSO_{4}]

This case affords one of the best illustrations of the effect of an
excess of a precipitant in decreasing the solubility of a precipitate.
If the conditions are considered which exist at the moment when just
enough of the Ba^{++} ions have been added to correspond to the
SO_{4}^{--} ions in