High Pressure Water ("Bergius"). The temperatures are lower (200-300oC) but high pressure is used (150 atmospheres) to keep the water liquid as it reacts with iron (Ellis 513) or carbon (Ellis 527ff). Common salt, iron chloride or hydrochloric acid accelerate the former and thallium salts catalyze the latter. Bergius built a prototype that produced high (99.95%) purity hydrogen. Since the hydrogen is already pressurized it can be put into bottles without the need for a separate compressor. (Teel 64). While initial cost and floor space requirements were expected to be low (Greenwood 188ff)-a 10 gallon capacity generator supposedly can produce 1000 cubic feet/hour (Teel 65)-I don't think these reactions were ever practiced commercially. Bergius (1913) claimed that hydrogen could be produced for just 2 cents/cubic meter (Greenwood says 1s/4.5p to 1s/11p per 1000 cubic feet.)
Electrolysis of Water. Water was first electrolyzed into hydrogen and oxygen in 1800. (Cleveland 128). Hydrogen is produced at the cathode and oxygen at the anode:
Cathode (reduction): 2 H2O + 2e- -› H2 + 2OH-
Anode (oxidation): 2 H2O -›O2 + 4 H+ + 4e-
The net reaction is
2H2O (36 grams) + electricity -› 2H2 (4 grams) + O2 (32 grams) .
The reaction requires an electrolyte, so either base (such as potassium or sodium hydroxide) or acid (such as sulfuric acid) is added to the water.
We will want an electrode material that is resistant to attack by the electrolyte, and minimizes the internal resistance. (Ellis 536; Greenwood 195). Acid electrolytes caused continuing corrosion problems and hence alkaline electrolytes became the norm. (Taylor 106). Taylor (105) recommends the combination of a nickel-plated anode and an iron cathode to minimize overvoltage.
The level of oxygen in the hydrogen compartments shouldn't exceed 5.3%, and of hydrogen in the oxygen ones, 5.5%. (Greenwood 202). It's critically important that the cell be designed to prevent the mixing of the hydrogen produced at the cathode with the oxygen produced at the anode, which can result in an explosion. This is usually done with a diaphragm separating the two, although there are alternatives. (Ellis 561, 581; Greenwood 201). Likewise, the produced gases should be monitored to detect inadvertent mixing. (This can occur, for example, if the polarity is reversed-Ellis 585, or by injury to the diaphragm, blockage in the system, or too great current density-Greenwood 202.) . This can be done by sampling the gas and igniting the sample under controlled conditions; they should burn not detonate. (Taylor 73).
From a portability standard, decomposing water with electricity has the advantage that you don't need to transport reactants. As to the weight of the apparatus itself, a Schukert 600 amp electrolyzer holding 50 liters solution and producing 5 m3/24h weighed 220 kg (Englehardt 86). A 110V, 150 amp, 16.5 kw Schmidt system producing 66 m3/24 h weighed 14,000 kg, while a 1.65 kw plant producing one-tenth that weighed 700 kg. (33). Unfortunately, because access to electricity is required, this is not really a method suitable for launch site production. (Batteries are heavy.)
I have gotten mixed signals on the issue of the space requirements for electrolysis. Ells first says that it requires a "relatively large floor space." (570) and then that "for small plants electrolysis has much in its favor." (595; cp. Teel 132). It may depend on the design; Schuckert demands relatively more room (Greenwood 198) while Schmidt is compact (199). Still, there was a portable Schukert generator car weighing in at 2000 kg, used together with a scrubber car of 1700-2100 kg. (Ardery).
The only apparatus for which I have specifics is the Levin generator; 100 will occupy 31 feet by 4.5 feet, and produce 320 cubic feet/hour at 200 amperes. (580). With normal room height, they can be installed in two tiers.
But an electrolytic cell can be very small. With 12.2 watts of solar-based electricity, a modern homemade cell in a 3.5"x5.5"x1.5" plastic container produced 0.399 milliliters/second hydrogen. (Businelli). So the real question is, what is the required cell volume to achieve the desired production rate?
One nice thing about electrochemistry is that you can predict performance. A chemist would expect that 96,500 coulombs (ampere-seconds) of electricity would liberate 1 gram of hydrogen and 8 grams of oxygen, those being the "equivalent weights" (ionic weight/ionic charge). So one ampere-hour produces 0.03731 gram-equivalents, which works out as 0.01482 cubic feet of hydrogen at STP (0oC, 760 mm Hg), 0.00741 cubic feet oxygen. At 20oC we do better; 0.01585 of hydrogen and 0.00792 of oxygen. (Taylor 103). The current supplied is typically 200-600 amperes. So, if the current were 400 amps, production would be 5.93 cubic feet hydrogen and 2.96 cubic feet oxygen per hour. (Ellis 536).
Teed (39) considered electrolysis to be suitable only for production of up to 1000 cubic feet hydrogen/hour.
In theory, the required voltage is 1.23, but because of secondary effects (overvoltage) it will probably be found that 1.7 volts are needed for continuous decomposition of water. (Taylor 104; Ellis 536). However, the diaphragm will tend to increase the resistance of the cell, necessitating a voltage of 2-4 volts. (106). As a result, energy efficiencies are in the 50-60% range. (Engelhardt 18, 20, 31).
The first electrolytic oxygen generator constructed for laboratory purposes was that of D'Arsonval (1885), and the first large scale apparatus was that of Latchinoff (1888). (Taylor 108). The first with significant industrial adoption was probably Schmidt's (1899), which produced 99% pure hydrogen and 97% pure oxygen. (Engelhardt 31).
2002McGHEST says that "although comparatively expensive, the process generates hydrogen of very high purity (over 99.9%). However, I think it's a mistake to count out the electrolytic process. Water, of course, is cheap, so the main expenses are those of providing electricity, and separating out the oxygen.
1890-1910 prices for electricity ran around 0.25 cents/kwh for hydroelectric and 1.25 for coal-fired steam plants. (Engelhardt 17). (Ellis 538 assumes 1 cent/kwh, and 569 quotes prices of 3-4 cents in New York City and 0.5 cents in South Chicago.)
Chances are that the recovered oxygen can be sold, thus defraying at least some of the production costs. In fact, the zeppelin hydrogen produced in 1934-38 was a byproduct of the electrolytic production of oxygen. (Dick 193). in 1904, oxygen sold for $1/m3 and hydrogen for $0.3125. (Engelhardt 40).
Bear in mind that "a normal military balloon requires, in order to be filled in twenty-four hours, a plant of about 200 kilowatts." An airship requires a lot more hydrogen than that.
Still, in 1904, the Italian, French and Swiss armies all relied on electrolytic hydrogen. (123).
In 2004, the average cost of electricity in America was 5 cents/kwh and at that price, with 80% electrolysis efficiency and 90% compression efficiency, the power cost for compressed electrolytic hydrogen was $2.70/kg. (Doty).
The higher the temperature, the less electrical energy is needed. If heat is cheaper than electricity, then higher operating temperatures are desirable. (Kirk-Othmer 13:868).
Electrolysis of Alkali. Historically, the first electrolytic hydrogen was a byproduct of the processing of brine to yield sodium hydroxide (caustic soda):
2Na+ + 2Cl- + 2H2O -› 2NaOH + H2 + Cl2
With 100% current efficiency, each ampere-hour would produce 1.322 grams of chlorine, 1.491 grams of sodium hydroxide and 0.0373 grams of hydrogen." (Actual current efficiencies were 90-98%.) The caustic soda may be used to make soap, and the chlorine bleach, and of course there are other uses, too. (Taylor 120). At 15oC, each ton of salt electrolyzed produces 72320 cubic feet hydrogen. Hydrogen purity is 90-97%. (Greenwood 203). Naturally, you want to prevent intermixing of hydrogen and chlorine after production.