Grantville Gazette Volume 25

by editor Paula Goodlett

  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, given 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 is made by reacting nitric acid with ammonia.

  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 the ordinary form. The white allotrope is the most reactive, and the black the least. Red phosphorus is what is used in modern matches, the white allotrope being much more poisonous. Phosphorus can be combined with the more electronegative elements like oxygen and halogens to make various covalent or ionic compounds. It also appears as the core atom of the phosphate anion.

  The down-timers were on the brink of making phosphorus (it was extracted from urine in 1669) so it's not surprising that as of the ride to Grantville at the time of the Croat Raid, Harry Lefferts had already been told by Greg that (white) phosphorus bombs were doable. Flint, 1632 (chapter 59). In 1634, Paddy lights a "phosphorus stick." Robison, "O for a Muse of Fire" (Grantville Gazette 11).

  The main commercial source of phosphorus is phosphate rock, which consists primarily of phosphate minerals, especially phosphorite (calcium phosphate). Phosphates are important as fertilizers, and there is reference to this in Turner's "Hobson's Choice" (Grantville Gazette 3).

  Phosphoric acid was first made by treating bone ash (calcium phosphate) with sulfuric acid; phosphate rock is now used instead, resulting in acid of 60-70% purity. A higher grade acid is made by burning phosphorus in an electric furnace, then (the part not explained by EA) reacting the resulting phosphorus pentoxide with carbon to form carbon monoxide and gaseous phosphorus. For ultra pure acid, you boil red phosphorus with nitric acid (the reaction is driven by the escape of gaseous nitrogen oxide).

  Phosphorus is not usually an end-product. To make white phosphorus, you may heat phosphoric acid to decompose it into hydrogen, carbon monoxide and phosphorus. Or reduce calcium phosphate in phosphate rock, using coke (carbon), sand (silica) and a high temperature. The silica reacts with the calcium salt to form calcium silicate and phosphorus, and the latter reacts with the carbon. The red phosphorus is made by calcining (heating without air) the white form. And the black form is obtained by heating the white allotrope under high pressure.

  The phosphoric acid can be used in the manufacture of non-naturally occurring phosphates.

  Manufacture of both phosphoric acid, and the various forms of phosphorus, seems to be of a difficulty on par with making ammonia from urine. So it could have been done as early as 1631-1632, but thanks to Paddy, we can be sure it was achieved by 1634.

  Group 14 Non-Metals

  Carbon

  Carbon occurs in nature as diamond, graphite and various coals. These are all known to the down-timers.

  Diamond-making requires pressures and temperature which are outside the realm of 1630s possibility.

  The telephone people in Grantville want graphite, because of its electrical properties. By April 1634, the USE embassy to Venice has ordered a supply of "good English graphite." Flint and Dennis, 1634: The Galileo Affair, Chapter 29.

  Carbon monoxide was first made in 1776, by heating zinc oxide with carbon. Other metal oxides can be used similarly. Or you can heat a carbonate with a reducing agent (zinc or iron), or pass carbon dioxide over carbon, or burn carbonaceous material with a limited air supply (EB11). Carbon monoxide is used to make the "synthesis gas" of organic chemistry.

  Carbon dioxide gas is used as a reagent and as a carbonating agent. It's frozen to form dry ice, a refrigerant. That requires first liquefying it, which, at room temperature, requires a pressure of about 56 atmospheres. (EA).

  It was identified as a distinct gas by von Helmont, an alchemist alive at the time of the RoF. (He appears in one of Kim's stories.) He called it gas sylvestre, and noted that it was produced by fermentation of sugar into alcohol, and complete combustion of coal.

  While carbon dioxide can be produced by combustion (reaction of oxygen with carbon), an important industrial derivation is by fermentation of sugar to alcohol and carbon dioxide. It also can be made by decomposing carbonates with heat or mineral acids. (EB11); it's a co-product, with lime, of the decomposition of calcium carbonate.

  Carbides are compounds made by combining carbon with a more electropositive element, such as most metals. EB11 states that carbides can be made by (1) "direct union" of the metal with carbon at high temperature, (2) reduction of an oxide with carbon, again at high temperature, (3) reduction of carbonates with magnesium (a very powerful reducing agent) in the presence of carbon, or (4) reaction of a metal with acetylene.

  Once silicon and boron are available, we can try to make their carbides, which are extremely hard (8-9 Mohs scale). CW293 says that silicon carbide (carborundum) and boron carbide are made by reducing the corresponding oxides with carbon in an electric furnace.

  Carbonates (CO3-2) and bicarbonates (HCO3-) are salts of carbonic acid (H2CO3). Potassium, sodium, magnesium, manganese and copper carbonates are usually isolated from natural sources. Metal carbonates also can be prepared by passing carbon dioxide through a solution of the appropriate hydroxide.

  The cyanide ion is CN-. The acid, hydrogen cyanide (prussic acid), is found in nature (e.g., cherries, apricots, bitter almonds). Ammonium cyanide was made in 1843
by passing ammonia gas over hot coke or charcoal. EB11 says that the most important salt is potassium cyanide, because of its use in the extraction of gold.

  Cyanogen (CNCN) can be prepared by oxidation of cyanide anions by copper (II) cations (EA/Cyanogen). A derivative, cyanogen bromide, is used as a cleavage agent in studies of proteins.

  There are several important ions related to cyanide: cyanate (OCN-). isocyanate (OCN-). and fulminate (CNO-). The cyanates are obtained by oxidizing the corresponding cyanide. The isocyanates are made from cyanates, although this is not very clearly communicated by the encyclopedias. They do provide synthetic routes to several fulminates, but because they are highly sensitive explosives, I am deliberately not discussing how they are made.

  Potassium ferricyanide is used in the ferro-prussate process for making blueprints. Potassium ferrocyanide is used in assays for zinc. Prussian blue was an early synthetic dye (1704).

  Silicon

  Silicon dioxide (silica) is the basic chemical component of glass and sand, and is the mineral of quartz and various gemstones such as amethyst. The metal silicates can be considered to be combinations of a metal oxide with silicon dioxide.

  Silicon is used in alloys and, when ultra-pure, in the semi-conductor industry. Silicon, discovered in 1824, occurs in both amorphous and crystalline forms. Amorphous silicon can be prepared by heating silica with magnesium in the presence of magnesia, and the crystalline form if magnesia is replaced with zinc. There are other synthetic routes, too. (EB11). The modern method is by heating silica with coke in an electric furnace (EA).

  Silicates are the most common type of minerals. Unfortunately, the silicates are not very useful as ores; it is hard to liberate the metal. Hydrofluoric acid will dissolve silicates, however.

  The more useful silicates include phenacite (beryllium ore), zircon (zirconium ore), willemite (zinc ore), petalite (lithium aluminum silicate), thorite (thorium uranium silicate) and asbestos, the fibrous form of the mineral serpentine (hydrous magnesium silicate).

  Metal silicates can be made by reacting molten silica with the metal carbonate. Thus, sodium silicate is made by combining sand and soda ash.

  Group 13 Non-Metals

  Boron

  Boric acid (H3BO3) occurs naturally in the Maremma of Tuscany. Its salts are the borates, and sodium (borax, kernite), magnesium (boracite) and calcium (colemanite) borate, and combinations (ulexite), are all found in nature. At least calcium borate is in down-time long-distance trade under the name "tincal." See Cooper, "Adventures in Prospecting and Mining for Minerals" at http://www.1632.org/1632tech/faqs/ for more information on boron sources.

  In April 1634, Sharon Nichols expresses concern about the availability of borax: "The Turks seem to be the only ones who've got it, and they're not being real friendly so far." Flint and Dennis, 1634: The Galileo Affair, Chapter 29. Tibetan borax (tincal) was sold in Italy pre-RoF. (Admittedly, it passed through Ottoman middlemen). By April 1634, borax is used by the Antonite hospital in Cologne to facilitate the manufacture of penicillin. Mackey, "The Prepared Mind" (Grantville Gazette 10). Also in 1634, Lewis Bartolli travels to Tuscany and makes arrangements for "mining" of the boric acid of the Maremma. Cooper, "Under the Tuscan Son," Grantville Gazette 9 . However, while he arrives in early 1634, production probably doesn't begin until late 1634.

  Suffice it to say that tincal should be available in small quantities in Grantville by 1632, and that large-scale production of boric acid and desired borates by the Tuscans should commence in 1634-35.

  Elemental boron was isolated in 1808 by (1) heating boron trioxide with potassium (a classic single displacement reaction) and (2) from boric acid. (EB11). The modern methods are by reduction with magnesium (followed by washing with alkali, hydrochloric acid and hydrofluoric acid) and by hydrolysis of boric oxide over tungsten. (EA, CW226).

  Diborane (BH3BH3) is used in the production of certain alcohols from alkenes. It can be made in the lab by reacting boron trifluoride (a strong acid, made by reacting calcium fluoride with sulfuric acid, CW233) with a metal hydride, or industrially by a high temperature, aluminum-catalyzed reaction of boron oxide with hydrogen (CW237-9).

  Group 18 Non-Metals (Noble Gases)

  We backtrack now to the noble gases. They are extremely unreactive with other elements.

  Argon and neon are produced, along with oxygen and nitrogen, by liquefaction and fractional distillation of air. Air is 78% nitrogen, 21% oxygen, 0.94% argon, 0.03% carbon dioxide, 0.0012% neon, 0.0004% helium, 0.00005% krypton, and 0.000006% xenon. (EA)

  Argon is used to provide a protective atmosphere; it's used in inert-gas-shielded arc welding and to protect molten metals from oxygenation. (EA). Neon, krypton and xeon are mainly used in lighting.

  Helium is also found in natural gas, but not necessarily all natural gas. The Great Plains (Texas, Kansas, Oklahoma) is one source (EA), and Poland is another. (Emsley 177-9) What about Grantville? The Transactions of the Electrochemical Society (39:47, 1921) reported that an unidentified "part of West Virginia" had natural gas bearing 0.1-1.5% helium. Helium is used in ballooning, in air supply for deep-sea divers, and in cryogenics.

  Other Non-Metals

  I have chosen not to discuss the noble gas Radon; the radioactive halogen Astatine, and the chalcogens Selenium, Tellurium and Polonium.

  Predictions

  Table 2-2 shows when various non-metals, covalent compounds, and anions appeared in canon, or, if they haven't yet made an appearance, when I predict they could have first been available.

  II. Metals and Their Salts

  To make a chemical compound which includes a metallic element, we need a source of that metal, whether that be the metal in elemental form, or a salt of that metal. The salt may occur in nature as a mineral, and those minerals from which the metal can be recovered in an economically feasible way are said to be its ores.

  Once we have one salt of the metal, we can convert it into other salts. By way of example, if you had sodium chloride, it can be used in the production of, e.g., sodium carbonate, bicarbonate, sulfate, silicate, and fluoride, if you had the appropriate reactants. (And of course sodium chloride isn't just a source of sodium, it's a source of chlorine, too.) Sodium carbonate also occurs in nature, and is used in a similar way.

  Alternatively, you can reduce the metal ion in the salt to the elemental metal. Sodium metal is typically prepared by electrolysis of molten sodium chloride.

  The elemental metal, in turn, can be alloyed with other metals, or used in further reactions to make additional salts of the metal of choice. Sodium metal can be used to make any of the previously mentioned sodium salts, as well as sodium oxide, peroxide, superoxide, hydride, phosphide, arsenide, bismuthide, bromide, iodide, sulfide, selenide, and amide. (It can also be used in the reduction of other metals, such as potassium and titanium, incidentally forming a sodium salt in the process.)

  Table 2-3 lists, for selected metallic elements, the immediate commercial source of the element (the substance that is directly reduced to yield the element) and the natural commercial source—the naturally occurring substance, such as a mineral, from which the element is directly or indirectly produced. For example, potash (potassium carbonate) is mined and converted directly or indirectly to potassium hydroxide, and in the final reaction, the potassium hydroxide is electrolyzed to yield potassium metal.

  This article identifies the principal ores of the more interesting metals, but doesn't go into details as to how or where the ores are found. For an introduction to the problems of prospecting for ores not previously of interest to the down-timers, see Runkle, "Mente et Malleo: Practical Mineralogy and Minerals Exploration in 1632" (Grantville Gazette 2).

  Some anions—silicates, carbonates, nitrates, sulfates, chlorides, oxides, hydroxides, sulfides, and phosphates—occur widely enough in nature that the problem in making a salt containing that anion is more likely to be finding the metal to go with it than finding the anion. Of the rarer anions, some are the polyanio
ns which contain a metal or metalloid themselves—chromate, tungstate, molybdenate, arsenate, etc.—and those are conveniently discussed together with the metal(loid) itself. Other rarer anions, such as fluorides and borates, were discussed in section I.

  Group 1 Metals

  Lithium

  The chloride is said to be the most common lithium salt. (EA). In the time frame of this article, the only one likely to be of interest is lithium aluminum hydride (as a reducing agent in organic synthesis). And perhaps lithium carbonate, if we have any manic-depressives we want to treat.

  The standard processing of lithium ores, which unfortunately are silicates, results in production of either (1) first a hydroxide and then a chloride, or (2) first a sulfate, then a carbonate, and finally a chloride (EA). Lithium also occurs in sea and spring water (EB11), and Chilean brines are actually the principal modern source of the element; other salts are crystallized out and then the lithium removed as lithium carbonate by reaction with sodium carbonate. (Emsley 237).

  Metallic lithium can be obtained by heating lithium hydroxide with magnesium (EB11) but the more modern approach is by electrolysis of the chloride (EA). There is some demand for the metal; "Lithium-magnesium alloys have the highest strength-to-weight ratio of all structural materials."