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In article <1988Dec20.204917.21249@utzoo.uucp>, henry@utzoo.uucp (Henry Spencer) writes:
> >There is a strong temptation to table add salt to decrease the resistivity of
> >the solution. Problem is that with NaCl in solution the anode (+) frees
> >both O-- Oxygen ions as gas AND CHLORINE Cl- ions (a problem related to half-
> >cell potentials). Most of the latter returns to solution yielding a HCL/HOCL
> >concoction resembling bleach...
> 
> My own experience as a kid, deliberately trying for chlorine, is that it's
> virtually impossible to get any noticeable quantity.  Maybe I didn't hit
> the right combination of conditions, but even determined electrolysis of
> saturated salt solutions didn't yield anything much.  (The commercial
> chlorine-by-electrolysis processes use tricks to isolate the reaction
> products from the solution.)  And even a pinch of salt can do wonders for
> conductivity, which is important for electrolysis.

	If you want to produce hydrogen and oxygen from water, a _slight_
amount of sulfuric acid will do wonders for increasing electrolyte
conductivity without creating any significant side reactions.

	However, when any significant amount of salt is added to the
solution (presumably in an effort to increase electrolyte conductivity),
it's a whole new ballgame...

	The electrolysis of brine (i.e., salt solutions) is used to form
chlorine at the anode and hydrogen _and_ sodium hydroxide at the cathode.
However, if the anode and cathode products are allowed to commingle,
then various reverse reactions will proceed almost as fast as the forward
reactions, with the net result of not much chlorine or hydrogen being
evolved.  However, sodium hypochlorite will be formed; this reaction
is used for commercial production of sodium hypochorite under carefully
controlled conditions and with an optimized cell design.  The addition
of calcium hydroxide will result in production of sodium chlorite and
(as a further step) chlorine dioxide.

	Oxygen does NOT form in the electrolysis of brine unless OH-
ions manage to reach the anode (an undesireable situation).

	Commercial chlor-alkalai processes always use electrochemical
cells which separate the anode and cathode products.  The first such
practicable cell to perform this process was called the mercury cell,
and was invented in 1892 in Niagara Falls, NY, and used by the Mathieson
Alkalai Works (now Olin Corp.).  The mercury cell was a two compartment
cell using a graphite anode and a circulating mercury cathode.  The
first compartment produced chlorine, and the second compartment produced
sodium hydroxide solution (which required further purification) and
hydrogen.  The mercury cell was reasonably efficient, and it launched
the U.S. chlor-alkalai industry.  Unfortunately, it had one significant
disadvantage - it used mercury, and has the dubious distinction of
being the primary cause for mercury contamination of aquatic life in the
Great Lakes.

	While mercury cells are still used within the continental U.S.,
most chlor-alkalai production uses a membrane cell in which a semi-
permiable membrane separates the anode from cathode.  The membrane allows
ions to pass, but not reaction products.  There is also another type
of cell called the Dow cell.

<>  Larry Lippman @ Recognition Research Corp., Clarence, New York
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