benfeen@ddsw1.MCS.COM (Ben Feen) (12/23/88)
Article 4142 of sci.electronics: Path: ddsw1!mcdchg!clyde!watmath!watcgl!awpaeth From: awpaeth@watcgl.waterloo.edu (Alan Wm Paeth) Newsgroups: sci.electronics Subject: Caution: electolysis of water Message-ID: <7395@watcgl.waterloo.edu> Date: 19 Dec 88 17:56:53 GMT References: <2479@ddsw1.MCS.COM> <849@inuxm.UUCP> Reply-To: awpaeth@watcgl.waterloo.edu (Alan Wm Paeth) Organization: U. of Waterloo, Ontario Lines: 22 In article <849@inuxm.UUCP> micl23@inuxm.UUCP (W E Miller) writes: >> >> How do I use electricity (from a battery-6V lantern type) to separate >> water into hydrogen and oxygen? Please E-Mail me. P.S. don't flame me, >> this was the only place I could think of posting to.. I used to use a sawed-off Purex bottle with a Copper (tubing) electrode and a Zinc bar, both wired to an old Selenium stack auto battery charger. I suppose graphite electrodes would have been fine. The caution: avoid poisonous Chlorine gas. 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 (+terminal) 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, there can also be some nascent Chlorine left around. The suggested fix: use a Carbonate salt (eg, baking or washing soda). /Alan Paeth Computer Graphics Laboratory University of Waterloo Article 4150 of sci.electronics: Path: ddsw1!mcdchg!chinet!att!rutgers!mailrus!uflorida!haven!mimsy!chris From: chris@mimsy.UUCP (Chris Torek) Newsgroups: sci.electronics Subject: Re: Caution: electolysis of water Message-ID: <15107@mimsy.UUCP> Date: 21 Dec 88 04:52:22 GMT References: <2479@ddsw1.MCS.COM> <849@inuxm.UUCP> <7395@watcgl.waterloo.edu> Organization: U of Maryland, Dept. of Computer Science, Coll. Pk., MD 20742 Lines: 19 In article <7395@watcgl.waterloo.edu> awpaeth@watcgl.waterloo.edu (Alan Wm Paeth) 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 (+terminal) 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, there can also be some nascent Chlorine left >around. The suggested fix: use a Carbonate salt (eg, baking or washing soda). Hard water is good for something :-) In my young and innocent (hah) days (when I was about 12 or 13), I played with electrolysis using the 117 volt line---much faster than batteries. Anyway, I vaguely remember that salt water formed a green scum. I guess now that it must have been choride salts of metallic ions in the tap water.... -- In-Real-Life: Chris Torek, Univ of MD Comp Sci Dept (+1 301 454 7163) Domain: chris@mimsy.umd.edu Path: uunet!mimsy!chris Article 4145 of sci.electronics: Path: ddsw1!mcdchg!clyde!watmath!utgpu!utzoo!henry From: henry@utzoo.uucp (Henry Spencer) Newsgroups: sci.electronics Subject: Re: Caution: electolysis of water Message-ID: <1988Dec20.204917.21249@utzoo.uucp> Date: 20 Dec 88 20:49:17 GMT References: <2479@ddsw1.MCS.COM> <849@inuxm.UUCP> <7395@watcgl.waterloo.edu> Organization: U of Toronto Zoology Lines: 22 In article <7395@watcgl.waterloo.edu> awpaeth@watcgl.waterloo.edu (Alan Wm Paeth) writes: >The caution: avoid poisonous Chlorine gas. > >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 (+terminal) 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. All that being said, I think I'd check the matter out chemically before using electrolyzed oxygen for breathing. -- "God willing, we will return." | Henry Spencer at U of Toronto Zoology -Eugene Cernan, the Moon, 1972 | uunet!attcan!utzoo!henry henry@zoo.toronto.edu Article 4151 of sci.electronics: Path: ddsw1!mcdchg!chinet!att!ihlpf!rvs From: rvs@ihlpf.ATT.COM (VonSchwedler) Newsgroups: sci.electronics Subject: Re: Caution: electolysis of water Message-ID: <7026@ihlpf.ATT.COM> Date: 21 Dec 88 19:04:52 GMT References: <2479@ddsw1.MCS.COM> <849@inuxm.UUCP> <7395@watcgl.waterloo.edu> <1988Dec20.204917.21249@utzoo.uucp> Reply-To: rvs@ihlpf.UUCP (55229-VonSchwedler,R.) Organization: AT&T Bell Laboratories - Naperville, Illinois Lines: 36 In article <1988Dec20.204917.21249@utzoo.uucp> henry@utzoo.uucp (Henry Spencer) writes: > >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 one that will more likely be produced (between Clorine and Oxygen), is the one that is more easily reduced. That happens to be oxygen. I am not completely sure about this, but let me just blab. When the circuit is completed, the electrons from the battery try to get over to the positive terminal from the negative terminal (through the solution). When they travel through the wire, they sorta stand around in the solution on the Anode. With all these electrons standing around, the Hydrogen atoms (being positively charged) see this as a very nice place to go. They will even leave their oxygen partners to go check it out. Once they get there, they take an electron off of the anode and are free (there is no reason to hang around with the oxygen atoms because they no longer need to barrow their electrons). At the other end, the oxygen atoms (still in water molecules) are near the positive terminal (the Cathode) which is positive. The oxygen atoms see this as a nice place to be and even leave their partners. When they come into contact with the cathode, the electrons they have that are holding the hydrogen atoms to them are striped away, which frees the hydrogen atoms, and the oxygen atoms are now free to come out of solution. As someone has mentioned, twice as much hydrogen is produced at the anode then oxygen is produced at the cathode, but the current remains constant. Two electrons for every H20 molecule. I believe that it takes two culombs of current for one mole of water. This is the philosophy behind this, but I may have a few things reversed. Also, you have to figure out what the electron potential of hydrogen and oxygen is because it may (I doubt it) require more than six volts. RvS Article 4152 of sci.electronics: Path: ddsw1!mcdchg!chinet!att!ucbvax!ucsd!rutgers!dayton!jad From: jad@dayton.UUCP (John A. Deters) Newsgroups: sci.electronics Subject: electrolysis of water Keywords: electrolysis Message-ID: <6327@dayton.UUCP> Date: 21 Dec 88 17:15:43 GMT References: <2479@ddsw1.MCS.COM> <849@inuxm.UUCP> <7395@watcgl.waterloo.edu> Reply-To: jad@dayton.UUCP (John A. Deters) Organization: Dayton-Hudson Dept. Store Co. Lines: 20 In article <7395@watcgl.waterloo.edu> awpaeth@watcgl.waterloo.edu (Alan Wm Paeth) writes: >In article <849@inuxm.UUCP> micl23@inuxm.UUCP (W E Miller) writes: >>> >>> How do I use electricity (from a battery-6V lantern type) to separate >>> water into hydrogen and oxygen?... >I used to use a sawed-off Purex bottle with a Copper (tubing) electrode and a >Zinc bar, both wired to an old Selenium stack auto battery charger. I suppose >graphite electrodes would have been fine. > ... description and caution deleted ... >The suggested fix: use a Carbonate salt (eg, baking or washing soda). > /Alan Paeth I found an even more interesting solution to use than plain water: Well-used photographic fixer. Photographic fixer works by removing the light-sensitive silver from the paper, so the silver compound remains in the fixer bath. If you use fixer, you will discover that one of your terminals (the anode, I think) will turn silver colored. You can actually peel off the flakes of silver that form on the terminal. (Well, when you're in high school, these things are pretty neat.) -john Article 4156 of sci.electronics: Path: ddsw1!mcdchg!chinet!att!rutgers!sunybcs!kitty!larry From: larry@kitty.UUCP (Larry Lippman) Newsgroups: sci.electronics Subject: Re: Caution: electolysis of water Summary: Decomposition products of electrodes... Message-ID: <2852@kitty.UUCP> Date: 21 Dec 88 20:26:41 GMT References: <2479@ddsw1.MCS.COM> <849@inuxm.UUCP> <7395@watcgl.waterloo.edu> <15107@mimsy.UUCP> Organization: Recognition Research Corp., Clarence, NY Lines: 32 In article <15107@mimsy.UUCP>, chris@mimsy.UUCP (Chris Torek) writes: > In my young and innocent (hah) days (when I was about 12 or 13), I > played with electrolysis using the 117 volt line---much faster than > batteries. Anyway, I vaguely remember that salt water formed a green > scum. I guess now that it must have been choride salts of metallic > ions in the tap water.... What you saw was probably cupric chloride resulting from decomposition of the copper wires which you (probably) stuck in the solution as "electrodes". If the water was particularly hard, some cupric carbonate may have also formed; cupric carbonate is insoluble in water and would definitely precipitate out. When I was a kid in high school, I went "big time" in electrolysis experiments: I melted table salt in a crucible using my father's Prestolite (acetylene) outfit, and used carbon electrodes salvaged from batteries to produce chlorine and metallic sodium. The sodium production was rather inefficient, but I got enough to prove the point :-) N.B.: I don't recommend this as a home experiment; also, a propane torch will most likely have insufficient heat output to adequately melt and fuse the salt. I always wanted to take the next step and use potassium fluoride to make fluorine, but I chickened out after finally managing to acquire the potassium fluoride. My one and only direct experience with pure fluorine when I was a grad student convinced me that I had made the right decision several years earlier. :-) <> Larry Lippman @ Recognition Research Corp., Clarence, New York <> UUCP: {allegra|ames|boulder|decvax|rutgers|watmath}!sunybcs!kitty!larry <> VOICE: 716/688-1231 {att|hplabs|mtune|utzoo|uunet}!/ <> FAX: 716/741-9635 {G1,G2,G3 modes} "Have you hugged your cat today?" Article 4157 of sci.electronics: Path: ddsw1!mcdchg!clyde!att!kitty!larry From: larry@kitty.UUCP (Larry Lippman) Newsgroups: sci.electronics Subject: Re: Caution: electolysis of water Summary: Introduction to Chlor-Alkalai Processes... Message-ID: <2855@kitty.UUCP> Date: 22 Dec 88 06:13:06 GMT References: <2479@ddsw1.MCS.COM> <849@inuxm.UUCP> <7395@watcgl.waterloo.edu> <1988Dec20.204917.21249@utzoo.uucp> Organization: Recognition Research Corp., Clarence, NY Lines: 61 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 <> UUCP: {allegra|ames|boulder|decvax|rutgers|watmath}!sunybcs!kitty!larry <> VOICE: 716/688-1231 {att|hplabs|mtune|utzoo|uunet}!/ <> FAX: 716/741-9635 {G1,G2,G3 modes} "Have you hugged your cat today?" -- _ /| When you link up to a VAX or UNIX without a "shell", you're linking \'o.O' with every other VAX or UNIX that the admin has been with! =(___)= Stoplights timed for 35 MPH are also timed for 1,244,740 MPH!!!! U Don't respond to watmath!looking!funny.Flame to benfeen@ddsw1.MCS.com