Grantville Gazette Volume 25

by editor Paula Goodlett

  Once you have elemental bromine it is easy enough to make hydrobromic acid (HBr) and the various salts. Silver bromide is a photosensitive salt used in early photography. Sodium and potassium bromide were favored in the nineteenth century as anticonvulsants and sedatives. Lithium bromide is used as an absorbent in absorption refrigeration systems. Huston, "Refrigeration and the 1632 World: Opportunities and Challenges" (Grantville Gazette 8).

  In view of the similarities of bromine and chlorine chemistry, I would predict that bromine, HBr and the common metal bromides could be produced as early as late 1633. However, the demand might not be sufficient to move production along that quickly.

  Iodine

  The concentration of iodine in seawater is very low (0.05 ppm). Fortunately, some seaweeds concentrate it—Laminaria is up to 0.45% iodine. Not surprisingly, seaweeds were the first commercial source of iodine. Particulars are given in EA and EB11; chlorine or manganese dioxide is used to oxidize the iodide ion to iodine (a solid). Originally, the big producers were Normandy and Scotland; later Japan became a major player.

  Another source, of more limited distribution, is Chilean saltpeter, which contains sodium and calcium iodate. The iodate is converted to iodide with sodium bisulfite and the iodide to iodine by adding fresh iodate. (EA).

  Finally, iodides can be found in brine wells, although I am not sure whether this is the case in Europe.

  Lewis Bartolli has access to iodine crystals in 1634, although we don't know when they were prepared. He used them in an unsuccessful attempt to develop a latent print on linen. Cooper, "Under the Tuscan Son" (Grantville Gazette 9). Sharon Nichols also has iodine, but not enough for the operation on Ruy Sanchez. Flint and Dennis, 1634: The Galileo Affair, Chapter 43.

  Hydrogen iodide (a gas) is made by direct combination of the elements over a platinum black catalyst (EB11). The iodides can be formed by direct iodination of a metal, or reaction of hydrogen iodide with a metal or its oxide, hydroxide or carbonate (EB11). Alternatively, potassium iodide is used to form iodides of most other metals, by replacement (EA). Tincture of iodine, an antiseptic is an alcoholic solution of potassium iodide and iodine.

  The USE is not likely to be a big producer of iodine compounds because it lacks ready access to cheap natural sources. The demand for iodine doesn't appear likely to be high enough to stimulate early (pre-1636) production. Whether there is commercial production in 1636 is likely to turn on the political situation in both Scotland and France.

  Group 16 Non-Metals (Chalcogens)

  Oxygen

  The method first used (1770s) to obtain oxygen was by heating a heavy metal oxide, e.g., mercuric oxide. It can also be obtained by chemical decomposition of other oxygen-containing compounds, electrolysis of water, or fractional distillation of liquefied air (ca. 1895). The latter two methods are mentioned in EA.

  Dr. Phil knows by mid-1631 about electrolysis of water; otherwise, nothing has been said in canon about oxygen. But we know that historically, oxygen was isolated in 1774, same as chlorine. Since chlorine is canonically available in 1633, and oxygen is at least as useful as chlorine, I propose that oxygen is available then, too. Indeed, an oxygen cylinder is used by Mary Pat in October 1633, but I don't know whether the oxygen was prepared after RoF. Ewing, "An Invisible War" (Grantville Gazette 8).

  Ozone is a molecule consisting of three atoms of oxygen instead of the usual two. It is produced by exposing oxygen to an electric discharge, by reacting sulfuric acid with certain peroxides (see below), or oxygen with certain heated metal oxides. (EB11). It was used at one time as a water sterilant, before it was replaced by chlorine. It can be used as an oxidizing agent, or to cleave certain organic compounds.

  Some metal oxides occur in nature, including the oxides of copper (cuprite), iron (hematite, magnetite), chromium (chromite), tin (cassiterite), manganese (pyrolusite), titanium (rutile, ilmenite).

  Oxides can be made, straightforwardly, by the reaction of oxygen with the appropriate element (e.g., zinc). It may also be possible to make them by reacting the appropriate element (e.g. potassium) with the nitrate of the same element (yielding nitrogen as a byproduct), or the appropriate nitrate (e.g., silver nitrate) with an alkali hydroxide; or by calcining (heating to decomposition) the appropriate carbonate (e.g., of calcium), nitrate, or hydroxide.

  Metal oxides can be reduced to the elemental metal by heating in the presence of carbon or hydrogen. (We'll discuss specific oxides under the heading of the other element.) They can be reacted with hydrogen sulfide, carbonic acid, nitric acid or sulfuric acid to make the metal sulfide, carbonate, nitrate or sulfate.

  Metal peroxides are made by reacting the corresponding oxide with more oxygen, or by direct reaction of the metal with oxygen at elevated temperatures.

  Hydrogen peroxide is used as an oxidizing agent, catalyst, bleach and disinfectant. EA suggests three methods of making it, of which the oldest (1818) is reacting barium peroxide with sulfuric acid. EB11 indicates that the barium peroxide may be decomposed with any of several acids. (Barium peroxide presumably made like other metal peroxides.) The other two EA methods are electrolyzing sulfuric acid and then hydrolyzing the product; and autooxidation of 2-ethyl anthraquinone (discovered 1936).

  The hydroxides of the alkali (e.g., potassium, sodium) and alkaline (e.g., calcium, magnesium) metals are strong bases and find much use in synthetic chemistry. Hydroxides may be obtained by reacting the appropriate oxide with water, and thus should be available on the same terms as the metal oxides.

  Sulfur

  Sulfur is readily available in elemental form, usually associated (as "brimstone") with volcanoes, such as those of Sicily. The Frash process (1890s) piped steam into underground sulfur deposits (particularly, those of Texas, Louisiana and Mexico) to melt the sulfur so it could be pumped out economically.

  It can also be obtained by reduction of sulfides and sulfates, possibly as a byproduct of metal smelting.

  Hydrogen sulfide (H2S) is used as a reagent in the production of metal sulfides, and as a source of elemental sulfur. It is the "rotten egg" smell emanating from volcanoes. It was produced by down-time alchemists as a byproduct of the synthesis of liquor hepatis and pulvis solaris. It can be made by direct combination of the elements, by reaction of a metal (especially iron) sulfide with sulfuric acid, or by decomposing antimony sulfide with hydrochloric acid. In the late twentieth century it was a byproduct of desulfurization of petroleum.

  Many metal ores are sulfides, found in hydrothermal deposits. Such deposits may contain sulfides of several different metals. The sulfide ores include cinnabar (mercury), galena (lead), pyrite (iron), stibnite (antimony), sphalerite (zinc), realgar (arsenic), and less well known, pentlandite (nickel), chalcocite (copper), covellite (copper), molybenite, chalcopyrite (iron and copper) and arsenopyrite (iron and arsenic).

  Metal sulfides can be roasted in the presence of oxygen to yield the corresponding oxide, and sulfur dioxide. There are various routes from the oxide to the elemental metal.

  Carbon disulfide (CS2) is used as a solvent for many organic substances, and in production of others, including carbon tetrachloride, viscose rayon and cellophane. It's made by heating coke and sulfur in an electric furnace. (EA)

  Sulfites are prepared by reacting a metal oxide, hydroxide or carbonate with sulfur dioxide (EB11/Sulphur). Thus, sodium sulfite is made by reacting sodium carbonate with sulfur dioxide (EB11/Sodium).

  Sulfuric acid (oil of vitriol) was first made in the early sixteenth century, at Nordhausen, by "dry distillation" (heating which first decomposes the solid into some kind of liquid mixture which is then distilled) of iron or copper sulfate. The metal sulfate decomposes into metal oxide, water and sulfur trioxide. (Derry 268).

  Derry says that sulfuric acid "was of virtually no industrial importance until the seventeenth century." Historically, dry distillation was superseded, by 1651, by Glauber's method. It had already been known in the sixteenth century that one could react sulfur with air (oxygen s
ource) and obtain a gas (sulfur trioxide). And Biringuccio's De la Pirotechnica (1544) took the next step; burning sulfur under a glass bell, in the presence of water, so that the sulfur trioxide combined with the water to make sulfuric acid (Salzberg 129). Glauber's innovation was the use of saltpeter (potassium nitrate) as a catalyst. He burnt a mixture of saltpeter (potassium nitrate) and sulfur in the presence of steam. The result was called, "oil of vitriol made by the bell." (Some authorities believe that the bell process was invented earlier, by Cornelius Drebbel (1572-1633), but the evidence is wanting.) (Kutney, Sulfur, 9).

  In 1744, it was discovered that you could make a very nice blue water-soluble dye (indigo carmine), very cheaply, by reacting indigo (insoluble once exposed to air) with sulfuric acid. That suddenly increased the demand for sulfuric acid. (Caveman Chemistry) The old glass vessels didn't scale up well; Roebuck (1746) replaced the glass vessels with lead-lined ones. Still, the acid was, at best, of 77% purity.

  The most important improvement, which permitted complete purification of the acid, was the "contact process," invented in 1831 but forgotten until the 1870s. In essence, sulfur trioxide (a waste gas is reacted with oxygen in the presence of a heated platinum wire catalyst. The "contact process" will probably become dominant as soon as the platinum catalyst becomes available.

  The "chamber" and "contact" processes are described in both EA/Sulfuric Acid and, in more detail, in EB11/Sulphuric Acid.

  The large-scale production of sulfuric acid is an early target of Grantville R&D. On Rebecca's talk show, Greg Ferrara explains "the critical importance of sulfuric acid to practically all industrial chemical processes." (Flint, 1632, Chapter 43). A conversation between Amy Kubiak and Lori Fleming in May 1632 implies that sulfuric acid is readily available (although given her subsequent reference to a "flame thrower," she may have been joking). Mackey, "The Prepared Mind" (Grantville Gazette 10). Discussing the synthesis of chloramphenicol with Rubens, Von Helmont comments that he needs "very pure" sulfuric acid, which is "quite difficult" (but he didn't say impossible) to obtain. Mackey, "Ounces of Prevention" (Grantville Gazette 5). In February 1634, Dr. Phil has about fifteen hogsheads of 90% pure sulfuric acid in hand, made from sphalerite. Offord, "Dr. Phil Zinkens A Bundle" (Grantville Gazette 7).

  By fall 1633, Grantville has sulfanilamide, so its chemists must previously have made chlorosulfonic (chlorosulfuric) acid. CW456 says it's made by reaction of sulfur trioxide with dry hydrochloric acid. (This reaction is supposed to be carried out in sulfuric acid.) It's also possible to chlorinate sulfuric acid with phosphorus pentachloride—(Wikipedia/Chlorosulfonic Acid.)

  Sulfates are typically made by reacting an elemental metal, or a metal hydroxide or oxide, with sulfuric acid. It is also possible to oxidize a metal sulfide or sulfite, or to add sulfur trioxide to a metal oxide. In some cases, a metal sulfate can be reacted with a different metal to yield a sulfate of the second metal (e.g., copper sulfate + zinc -> zinc sulfate).

  Sulfur dioxide is known to the down-timers. Sulfur dioxide is formed when sulfur is burnt in air, and it can also be released when a metal sulfide is roasted. Sulfur trioxide was first made (at least by 1675) by distillation of green vitriol (copper sulfate) but can also be obtained by the catalyzed union of sulfur dioxide with oxygen. Both are useful in the preparation of sulfuric acid, and the trioxide may also used, with hydrogen chloride, to make chlorosulfonic acid.

  Elemental sulfur, and the sulfur compounds known to the down-timers, should be coming into Grantville by late 1631. Additional sulfides and sulfates will become available as new sulfide ores are mined, and by chemical conversion of elemental metals, or their oxides, hydroxides or carbonates.

  Group 15 Non-Metals (Pnictogens)

  Nitrogen

  Nitrogen is used in the production of ammonia, and as an inert atmosphere and (in liquid form) a coolant for chemical reactions. Typically nitrogen is obtained, directly or indirectly, from air (which is over 70% nitrogen). First of all, active metals can be burnt with air to form nitrides, and the nitrides subsequently decomposed to release nitrogen. Secondly, air can be passed over heated coke, thus converting the oxygen to carbon dioxide, and the latter absorbed into water. Or you can instead burn phosphorus in air, or pass air over heated copper. The purest form of nitrogen is made by liquefying and fractionating air. Nitrogen was first obtained in 1772, by removing oxygen from air. Ammonia, ammonium nitrite, or ammonium nitrate can also be used as sources of nitrogen.

  There is reference in Offord, "Silencing the Sirens' Song" (Grantville Gazette 23) to an experimental facility, operating in July 1634, for using electricity to extract nitrogen out of the air. It is manufacturing nitric acid.

  Nitrates (NO3-) are fairly common minerals, and metal nitrates, because of their solubility, are useful in the preparation of other metal salts. Potassium and sodium nitrate are both naturally occurring.

  In 1631-32, one of the new chemical firms has someone making the rounds, asking for chicken manure for a nitrate farm. DeMarce, et al., "The Brillo Letters" (1634: The Ram Rebellion). Nitrates are excellent fertilizers. In 1632, this is known to the English. Turner, "Hobson's Choice" (Grantville Gazette 3). By 1634, this information has disseminated at least as far as Russia. Huff and Goodlett, "Butterflies in the Kremlin, Part Five: The Dog and Pony Show" (Grantville Gazette 13).

  Nitric acid (aqua fortis, spirit of nitre, HNO3) is nearly as important as sulfuric acid, and, like it, was made by down-time alchemists. It was made by heating potassium or sodium nitrate with concentrated sulfuric acid (EB11), and was used by the down-timers to dissolve silver and thereby separate it from gold (Derry 268). In 1632-33, the up-timers were producing nitric acid in only limited quantities because they insisted on use of stainless steel reactors and the stainless steel then had to be recycled. However, I would predict that if the "stainless steel bottleneck" isn't solved by 1634-35 the down-timers will simply ignore it and make nitric acid in glass-lined reactors (see "Corrosion Control" in Part 1).

  Nitric acid can be used to make metal nitrates. The acid is also used to add nitro (NO2) groups to organic compounds. Guncotton, for example, is nitrocellulose.

  Nitrites (NO2-) can sometimes be made simply be heating the corresponding nitrate. EB11 recommends making sodium nitrite by heating the nitrate with lead, or with sulfur and sodium hydroxide.

  Nitrogen can react with oxygen to form various oxides. Nitrous oxide(N2O) is made by heating ammonium nitrate (this has to be done gingerly, to avoid an explosion) and is used as an anesthetic. We know from canon that it's being produced by September, 1635 by Dr. Phil's chemical works. Offord, "The Creamed Madonna" (Grantville Gazette 19). Given that ammonium nitrate is available at least by December 1633, and there is demand for anesthetics, I would have expected it to be in production in 1634. (It can't be available before December 1632 since the dentist is still out of anesthetic. Flint, 1632, Chapter 39; Wentworth, "Here Comes Santa Claus", Ring of Fire). But there are so many compounds, and so few chemists. . . .

  Poor Dr. Phil. What he actually wants is to duplicate the effects of VIAGRA® sildenafil. Sildenafil inhibits an enzyme which recycles a metabolite which in turn is released as a result of the action of nitric oxide (NO). Dr. Phil figured that if he couldn't make sildenafil, the next best thing to do was to distribute a tonic containing pressurized nitrogen oxide. As Carl pointed out, his first mistake was to use the wrong nitrogen oxide (nitrous, not nitric). But I also have grave doubts that even nitric oxide, if orally delivered, will have any effect on ED.

  Once he realizes his first mistake, he will find that the encyclopedias say how to make nitric oxide; react nitric acid with ferrous sulfate in sulfuric acid solution. Or combine ammonia with atmospheric oxygen under the benevolent attention of a platinum catalyst. (EA).

  Ammonia (NH3) is primarily used in the manufacture of fertilizers, but also finds application as a refrigerant, and in inorganic and organic chemical synthesis. The compounds synthesized using ammonia include nitric acid, nylon
, dyes, pharmaceuticals and explosives.

  Ammonia was made by down-timers in several ways. First, by treating the distillate of animal horns with hydrochloric acid, and was therefore called spirit of hartshorn. A second route was by reacting ammonium chloride with alkali (hydroxide). Finally, the down-timers knew that it could be extracted from urine, as was done in 1631-32 by Dr. Philip Gribbleflotz of Jena for the Kubiaks. Offord, "The Doctor Gribbleflotz Chronicles, Part 1: Calling Dr. Phil," Grantville Gazette 6. The down-timers used ammonia in the manufacture of alum, and of a lichen-derived dye (archil).

  In the nineteenth century, ammonia was one of the byproducts of coal pyrolysis. But by the early twentieth century, it became possible to make ammonia by direct combination of nitrogen and hydrogen (the Haber process) . . . which, in turn, meant you didn't need access to nitrate deposits or even coal, since nitrogen and hydrogen can be found in air and water, respectively.

  The late twentieth century embodiments of the Haber process use pressures of 200-900 atmospheres and temperatures of 400-650°C. At 300 atmospheres and 500°C, the nitrogen, hydrogen and ammonia will reach an equilibrium in which the mixture in the reactor is 26.5% ammonia. (EA)

  A detailed analysis of the effect of both pressure and temperature on the equilibrium percentage of ammonia appears in EB11/Nitrogen Fixation. As would be predicted based on Le Chatelier's Principle, increasing the pressure increases the yield, whereas increasing the temperature reduces it. So, you logically ask, why not stay at room temperature, or even cool things down? The problem is that the reaction is very slow at room temperature. For a decent production rate, you need elevated temperatures.