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Industrial Alchemy, Part 2: Inorganic Chemical Bestiary

Written by Iver P. Cooper

 

Within a few weeks of the Ring of Fire (RoF), Greg Ferrara tells the "Emergency Committee" that "Sulfuric acid is about as basic for modern industry as steel." The 1911 Encyclopedia Britannica (EB11) and the modern Encyclopedia Americana (EA) agree that sulfuric acid is the most important of all chemicals. But that, of course, doesn't mean that it is the only chemical that the up-timers need more of. If there are a dozen they want at the end of 1632, I guarantee that they will be begging for hundreds by the end of 1634.

Elements, Ions and Compounds

The non-metals, discussed in section I below, are carbon; the pnictogens ("pn" as in phosphorus and nitrogen), the chalcogens (oxygen, sulfur, selenium), the halogens (fluorine, chlorine, etc.), and the noble gases (helium, etc.). Hydrogen is sui generis, the proverbial "sore thumb" of the Periodic Table, but I will treat it as a non-metal.

The non-metallic elements, by themselves, can form molecules (e.g., the two atom molecules of nitrogen, oxygen and chlorine), covalent compounds (e.g., carbon dioxide), and many important anions (e.g., chloride, carbonate, sulfate). Many anions are salts of acids having the form HX, and the X (the anion part) always contains at least one non-metal atom and sometimes is entirely composed of non-metallic elements. Many metal salts are of the form MX, where M is one or more atoms of the same metal, and X is one or more copies of the same anion, each one or more atoms.

In section I, I will identify which non-metallic elements, and compounds and ions composed just of those elements, were known prior to the RoF, which weren't known to the down-timers but occur in nature, and which will first be synthesized after RoF. I will also discuss how these elements and compounds are made and used, and make suggestions as to when they may be first available in the 1632 universe.

The metals and their salts are discussed in section II below, which is organized first by the column (1-16) of the periodic table which the metal falls into, and then by the metal itself.

The metals are sometimes classified as

—the group Ia (column 1) or alkali metals (notably lithium, sodium, potassium)

—the group IIa (2) or alkaline earth metals (notably beryllium, magnesium, and calcium)

—the transition metals (3-12) (notably iron, nickel, platinum, copper, silver, gold, zinc, mercury)

—the inner transition metals (which I will be ignoring)

—the "poor" (lower melting) metals (13-16) (notably aluminum, gallium, tin, lead and bismuth)

There are also metalloids, intermediate in behavior between metals and nonmetals. These are boron, silicon, germanium, arsenic, antimony and tellurium. Note that I have chosen to discuss boron and silicon with the non-metals, and arsenic and antimony with the metals.

I. Non-Metallic Elements and Compounds

Table 2-1 looks at the non-metals from a modern OTL perspective:

 

Table 2-1: Non-Metals: Sources and Demand
Element	Source	Modern Demand tonnes/yr*
Hydrogen	natural gas (methane), release from hydrogen compounds (hydroxides, bicarbonates, acids), water	30,000,000
Oxygen	air	100,000,000
Nitrogen	air	45,000,000
Argon	air	750,000
Helium	natural gas	>100,000,000 m3
Chlorine	brine	salt:160,000,000
Bromine	brine	300,000
Iodine	brine, seaweed	13,000
Fluorine	fluorspar (calcium fluoride)	fluorspar: 120,000,000 HF: 400,000 Element:15,000
Sulfur	hydrogen sulfide in natural gas; elemental sulfur deposits; sulfide minerals	50,000,000
Phosphorus	phosphate rock; bone	ore:153,000,000 acid:50,000,000 element: <1,000,000
Silicon	sand, talc, mica	96% pure:4,000,000 98% pure:500,000

* Emsley.

Hydrogen

Hydrogen, discovered in 1766, is used in the manufacture of ammonia and methanol, and in hydrogenation of unsaturated organic compounds. It also had direct uses; in the early twentieth century, as a buoyancy gas, and in the late twentieth century, as a rocket fuel and welding gas (part of the oxyhydrogen torch).

In Huff and Goodlett, "Butterflies in the Kremlin, Part 3: Boris, Natasha . . . But Where's Bullwinkle" (Grantville Gazette 10), set in September 1633, the Russians are experimenting with their third hot air balloon, but they are anxious to move on to hydrogen. By June-July 1634, according to their "Butterflies in the Kremlin, Part 6: The Polish Incident or the Wet Firecracker War" (Grantville Gazette 15), a hydrogen-filled dirigible is flitting about.

In contrast, in September 1635, Marlon Pridmore is flying a hot air blimp in the Grantville area. Kevin and Karen Evans, "Sailing Upwind" (Grantville Gazette 13). Of course, the USE has planes, and therefore less incentive to experiment with dirigibles.

The simplest method of obtaining hydrogen gas is by reacting a metal with a source of hydrogen. Thus, zinc or iron will react with dilute sulfuric acid, and sodium even with cold water. It is also possible to obtain hydrogen by electrolysis of water (which also yields oxygen).

Given the ready availability of zinc or iron, and sulfuric acid, there is no reason someone couldn't have made hydrogen as early as 1631 (Paracelsus supposedly made it in the sixteenth century). And in Grantville, with cheap electricity, the electrolysis route is feasible. Indeed, Tasha Kubiak gives Dr. Phil instructions for "bubbling off hydrogen and oxygen" in July 1631 (Offord, "Dr. Phil's Amazing Lightning Crystal," Grantville Gazette 6).

The problem isn't generating the hydrogen, it's hanging onto it once you have it. Clearly, by 1634, the Russians are doing both, in dirigible-sized quantities.

The classical concept of an acid is as a substance which, in water, dissociates to produce one or more hydrogen cations, and an anion characteristic of the acid. These acids will have formulae like HX (e.g., hydrochloric acid or nitric acid), H2X (e.g., sulfuric acid), or even H3X (orthophosphoric acid). The hydrogen cations behave much like the alkali metal cations. The first three strong acids known to the alchemists—hydrochloric, nitric and sulfuric acids—were all used in assaying, hence the term "acid test." (Salzberg 87).

Hydrogen forms ionic or interstitial hydrides with metals, and covalent hydrides with non-metals. The ionic hydrides are made by passing hydrogen gas over the warmed metal. (CW184 says to use temperatures of 300-700^◦C, 725^◦C for lithium). Hydrides of interest include sodium, potassium, lithium, calcium, strontium, palladium and titanium hydride, and lithium aluminum hydride. (EB11, EA). They are used variously as sources of hydrogen (they will decompose water), reducing agents, and fuels. Lithium aluminum hydride (preparation, see CW273) and sodium borohydride are among the most popular reducing agents in organic chemistry.

After 1634, the availability of the metal hydrides will be limited by the availability of the metal of interest.

Group 17 Non-Metals (Halogens)

The halogens of interest are fluorine, chlorine, bromine and iodine. They combine with hydrogen to form acids of the form HX, where X is halogen. The halides are salts in which the anion is a halogen atom: fluorine, chlorine, bromine or iodine. There are also related oxyanions including hypochlorites, chlorites, chlorates, perchlorates, bromates, perbromates, iodates and periodates.

Some of the metal halides (e.g., sodium chloride) can be extracted from natural sources, others are made by the reaction of 1) the metal with the halogen directly, or 2) the metal, or its oxide, hydroxide or carbonate, with the appropriate acid. For example, you can treat lithium carbonate with hydrochloric acid to make lithium chloride.

Fluorine

The principal natural source of fluorine is fluorspar (calcium fluoride). Fluorine is also found in cryolite (sodium aluminum fluoride) and fluorapatite (calcium fluorophosphate).

Hydrogen fluoride (known as hydrofluoric acid when dissolved in water) can be made by reacting fluorspar (calcium fluoride) with concentrated sulfuric acid at elevated temperature (first carried out by Scheele in 1771) (EB11/Fluorine). HF is extremely nasty stuff. Unfortunately it's critical to the production of many different fluorinated compounds, inorganic and organic. It's also used to etch glass and clean metals.

Since canon says that HF and synthetic cryolite are available in 1634, see Offord, "Feng Shui for the Soul" (Grantville Gazette 17) and is then used in an effort to produce artificial cryolite which is successfully produced in early 1636, see Offord, "Dr. Phil's Family" (Grantville Gazette 15) it is likely that fluorspar was being mined at least as early as 1633-34, and sodium and potassium fluoride, and perhaps aluminum fluoride, were probably being made in small quantities by late 1636.

Note that in table 2-4, the reference to cryolite for 1633 is to mined cryolite, per Mackey, Land of Ice and Sun.

The most straightforward way of making fluorine itself is probably by electrolysis of anhydrous HF containing dissolved potassium fluoride (EB11). The addition of the potassium fluoride is necessary since HF itself is non-conductive. (CW460).

I don't think it likely that there will be any fluorine gas production before 1640. Fluorine is not only a gas, it's a gas that can cause stainless steel to burn! In the old time line, elemental fluorine was not produced commercially until World War II (when it was needed for the manufacture of uranium hexafluoride).

Chlorine

 

A dilute form of hydrochloric acid (HCl) is already made by down-timers and used as part of aqua regia (the mixture of HCl and HNO3 used to dissolve gold). Concentrated HCl was obtained by Glauber (1648). The first commercial production was by the LeBlanc process (1790), in which sodium chloride is treated with concentrated sulfuric acid, yielding sodium bisulfate (or sulfate) and HCl. (LeBlanc's purpose was to make sodium carbonate.) The brute force method is to combine hydrogen and chlorine and it is used when you must have ultrapure material.

In OTL, chlorine was discovered in 1774. In the nineteenth century, chlorine was produced by oxidizing HCl with a strong oxidizing agent (air, manganese dioxide, potassium dichromate, etc.) A more modern method is electrolysis of sodium chloride solutions, yielding chlorine, sodium hydroxide, and hydrogen. (EA, EB11).

There are several canonical clues that chlorine is available by 1633. When the Grantville delegation to England left in June 1633, they carried DDT with them, and the chlorine atoms of the DDT were almost certainly introduced by reacting an intermediate with chlorine gas. By winter 1633-34, the Essen Chemical Company is producing small quantities of sulfanilamide (apparently in preference to Grantville's preferred antibiotic, chloramphenicol) as well as calcium hypochlorite. See Mackey, "Ounces of Prevention" (Grantville Gazette 5). By 1634, the French have made potassium chlorate (first synthesized 1786 OTL), possibly by reacting chlorine with potassium hydroxide. (cp. EB11/Chlorates).

Also, Dr. Phil makes bleach (Ethereal Essence of Common Salt) in 1633, by electrolysis of a sodium chloride solution. (Offord, "Dr. Phil's Amazing Lightning Crystal," Grantville Gazette 6) Chlorine is produced at the anode and hydrogen and hydroxide at the cathode. The chlorine then reacts with the hydroxide to produce some hypochlorite. If a membrane (such as asbestos) were placed between the anode and cathode, to block the movement of the chlorine, then you can produce chlorine gas.

Some chlorides are available from natural sources. The best known chloride is certainly sodium chloride (common salt), which is produced by mining rock salt, or evaporating brine from wells or seawater. Potassium chloride can be obtained from the ores sylvite (at Stassfurt, Germany) and sylvinite, or from seawater. It is also a byproduct of manufacturing nitric acid from potassium nitrate and hydrochloric acid.

The other metal chlorides can be obtained by reacting the metal, or its hydroxide, oxide or carbonate, with HCl. An alternative, brute force method is to heat the metal in a stream of chlorine gas. (EB11).

The alkali metal chlorides in general are also useful as sources of their metals; the latter can be produced in elemental form by electrolysis of the corresponding molten metal chloride.

The oxyanions of chlorine are hypochlorite, chlorite, chlorate and perchlorate. The hypochlorites are made by combining chlorine with a cold solution of a strong base; if you want the chlorate, use a hot solution. In both cases, you also produce a chloride.

There are a number of important covalent compounds that contain chlorine. These include sulfur dichloride, thionyl chloride (SOCl2), phosphorus trichloride and phosphorus pentachloride. The latter three are standard chlorinating agents in organic chemistry. (M&B 601). EB11 says to synthesize sulfur dichloride by "distilling sulfur in a chlorine gas," phosphorus trichloride by reacting heated red phosphorus with chlorine, and phosphorus pentachloride by further reaction of the trichloride with chlorine. All three should be feasible in 1634.

The availability of thionyl chloride is more uncertain; neither EB11 nor EA clearly state how to make it. However, CW453 describes a route from phosphorus pentachloride and sulfur dioxide. So perhaps we will be making it by 1635.

Bromine

Bromine was originally isolated from seawater (1826), in which it occurs as bromides in concentrations of just 65 ppm (EA). In 1911, the principal commercial source was the salt deposits at Stassfurt, Germany; the salt is a mixture of potassium, sodium, and magnesium bromide (EB11). The commercial "periodic" process required chlorine gas (which oxidizes the bromide ion), either manganese dioxide or potassium chlorate, and sulfuric acid. EA describes procedures (requiring chlorine gas, and either sodium carbonate or sulfur dioxide) for recovering bromine from seawater bromide. Since bromine is a liquid, it is actually easier to handle than chlorine (although bear in mind that its name comes from the Greek word for "stench").

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 source) 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 be 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.

You can use a catalyst, rather than a higher temperature, to increase the rate without loss of yield, but a catalyst isn't a panacea. Even with a catalyst, you need a fairly high temperature. EB11 says, "the formation of ammonia begins at as low a temperature as 360^◦C," but admits that the reaction is still "exceedingly slow." So that's why the temperature is bumped up to 500^◦C. And with high temperatures, you need high pressures to get respectable yields.

Catalysts can also be expensive (the first ones used were osmium and uranium). They tend to deteriorate over time, so, for economic reasons, you need to know how to recover and regenerate them. If your materials aren't pure enough, the catalyst can be poisoned. The modern catalyst consists "primarily of magnetic iron oxide (Fe₃O₄) or iron oxide mixed with the oxides of other metals" (EA/Ammonia), but we don't know the exact physical form (e.g., particle size, porosity, etc.). And the devil is in the details (Wikipedia/Haber Process; Frankenburg).

Increasing pressure is good for both high yield and fast reaction rate, but it takes energy to maintain a high pressure, and very expensive structures to safely contain it (especially at high temperatures). So plant designers typically use more moderate pressures, and compensate for the reduced equilibrium level in two ways.

First, they remove the ammonia as a liquid, taking advantage of the higher boiling points of nitrogen and hydrogen. (They can remove ammonia much faster than the system can come to equilibrium.) Secondly, they recycle the nitrogen and hydrogen gas, giving them further opportunities to react.

Amides. Active metals can react with ammonia to form amides (NH₂^-); sodium amide is used in organic chemical synthesis.

Ammonium. (NH₄⁺) is a cation consisting of a hydrogen ion added to ammonia, and behaves somewhat like a Group 1 metal.

Ammonium hydroxide, a strong base, is made when ammonia is bubbled into water. The alchemists called it "spirits of hartshorn."

Ammonium chloride (sal ammoniac) was known in antiquity, as it forms in volcanic regions. 

Ammonium nitrate, made by reaction of ammonia with nitric acid, is the fertilizer that Mike Stearns discovers, in December 1633, is stored in a shed near the stricken Magdeburg coal gas plant. (Flint, 1634: The Baltic War, Chapter 3). The ammonia could come from the ammoniacal liquor produced by destructive distillation of coal

"Ammonia" (probably ammonium carbonate) smelling salts are used to awaken Magdalena in Huff and Goodlett, "The Monster" (Grantville Gazette 12).

Clearly, nitric acid, ammonia (albeit not by the Haber process!), and several nitrates (ammonium, potassium) are going to be available in 1631-32, whereas the availability of other nitrite and nitrate salts will be "metal-limited." And I am reluctantly forced to assume that nitrous oxide isn't on the market until late 1635, and nitric oxide later still.

Phosphorus

Phosphorus exists in several different elemental forms (allotropes) with different structures: white (yellow), red and black (violet). White phosphorus is ...