Industrial Alchemy, Part 2: Inorganic Chemical Bestiary

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 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.

Table 2-2: Suggested Availability and Canon Appearances of Non-Metals, Covalent Compounds, and Anions for Combination Into Metal Salts

Non-Metals

Elemental Form

Anions

NM Compounds

Pre-RoF

Carbon, Sulfur, Arsenic

chlorides, sulfides, sulfates, nitrates, tungstates?

HCl; SO2, SO3, H2SO4, HNO3; NH3; borax

1631-32

Phosphorus

sulfites, nitrites, carbonates, bicarbonates, cyanides, manganates, permanganates

CO2, H2S, CO

1633

Chlorine**, Bromine, Oxygen*, Nitrogen (1)

bromides, oxides

HBr, chlorosulfonic acid; ammonium compounds; CS2

1634

Hydrogen*, Ozone, Iodine

chlorates*, borates, iodides

SCl2, PCl3, PCl5, H2O2, boric acid, HI, HF*

1635

Silicon? Boron?

Hydrides; CaF2; carbides*, chromate, dichromate

SOCl2; nitrous oxide*

1636

1637-1639

diborane

1640s

Fluorine;

Later

Helium, Argon, Neon

*explicit canon appearance **canon presence implied by appearance of related chemical (1) canon appearance in 1634.

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 polyanions 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.

Table 2-3: Production and Demand for Selected Metals

element (OTL discovery)

immediate source

natural source

modern demand

tonnes/yr*

Alkali Metals

Lithium

1817

chloride?

mineral springs; brine pools;

spodumene (lithium aluminum inosilicate);

petalite (lithium aluminum tectosilicate)

Ore: 40,000

Metal: 7,500

Sodium

1807

chloride, carbonate, hydroxide

common salt (chloride);

natron (carbonate)

chloride:200,000,000

carbonate:30,000,000

metal: 80,000

Potassium

1807

chloride, hydroxide

potash (carbonate);

sylvite (chloride);

carnallite (potassium magnesium chloride)

ore: 50,000,000

Alkaline Metals

Beryllium

1797

chloride, fluoride

bertrandite (hydrated silicate); beryl (aluminum beryllium silicate)

<500

Magnesium

1755

chloride

seawater; magnesite (carbonate); dolomite (calcium magnesium carbonate), carnallite; brucite (hydroxide), talc (silicate), olivine (magnesium iron silicate), kieserite (sulfate), serpentine (complex silicate)

400,000

Calcium

1808

oxide

calcite (carbonate) in limestone

lime:120,000,000

metal: 2000

Transition Metals

Titanium

1791

chloride (Kroll process), dioxide (Cambridge)

rutile (dioxide); ilmenite (titanium iron oxide)

dioxide:4,300,000

metal:90,000

Vanadium

1801

oxide

carnotite (triple oxide of potassium, uranium and vanadium); roscoelite (a mica); vanadinite (lead chloro vanadate)

ore:50,000

metal: 7000

Chromium

1797

chromate

chromite (iron chromate)

ore: 10,000,000

metal: 20,000

Molybdenum

1781

molybenite (sulfide)

ore: 90,000

Tungsten

1788

oxide

wolframite (iron manganese tungstate); scheelite (calcium tungstate); placer tungsten

40,000

Manganese

1774

pyrolusite (oxide)

ore:25,000,000

Iron

BC

hematite (oxide), magnetite (oxide)

new M:500,000,000

recycled:300,000,000

Cobalt 1735

cobaltite (cobalt arsenic sulfide)

17,000

Rhodium 1803

byproduct of platinum, gold, silver, palladium and nickel mining

16

Nickel 1751

sulfide

sulfides; laterites (oxides, silicates)

500,000

Palladium

1803

associated with gold or platinum

300

Platinum

1735

native platinum, cooperite (sulfide), braggite

155

Copper

BC

sulfide, sulfate

native copper; various sulfides (eg bornite), oxides (eg cuprite), carbonates (eg malachite, azurite)

new:

12,000,000

recycled:

2,000,000

Silver

BC

native silver, argentite (sulfide), chlorargyrite (chloride); associated with lead, copper or gold ore

17,000

Gold

BC

native gold (often microscopic grains in rock)

2500

Zinc

~1500

sphalerite (“zincblende”) (sulfide) (chief ore); smithsonite (carbonate);

hemimorphite (silicate), franklinite; Calamine is a mixture of smithsonite and hemimorphite.

>7,000,000

Cadmium

1817

greenockite; zincblende impurity

14,000

Mercury BC

cinnabar (sulfide)

8000

Poor Metals and Metalloids

Aluminum

1825

oxide

bauxite

new: 20,000,000

recycled 20,000,000

Tin

BC

cassiterite (oxide)

metal: >140,000

concentrates:

130,000

Lead

BC

oxide

galena (sulfide) (Chief ore), anglesite (sulfate), cerussite (carbonate)

6,000,000

Boron 1808

various borates, boric acid

borate: 2,000,000

Antimony BC

stibnite (sulfide)

50,000

Arsenic 1250

realgar, orpiment

oxide:50,000

Bismuth 1400

bismuthinite (sulfide)

3000

Germanium 1886

sphalerite impurity

80

*Emsley.

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.”

That said, given the difficulties of finding and processing the ore, I don’t expect lithium to be in play in the 1630s.

Sodium

Sodium chloride is, of course, a classic food additive, and is also a deicing agent and a desiccant/preservative. The chemical industry uses it as a source of sodium (particularly sodium hydroxide and carbonate) and chlorine.

Sodium hydroxide (lye, caustic soda) has been used as a strong base since medieval times (OED). In Gassage’s method (1853), it was made by reaction of quicklime (calcium oxide) with a boiling solution of crude sodium carbonate. The modern industrial method is by the electrochemical chloralkali process. (EA). Sodium hydroxide is used by Lewis Bartolli in summer 1634 as a forensic reagent, to detect iron in gall inks. (Lewis has other weak and strong acids and alkalis at his disposal, but we don’t know which.) Cooper, “Under the Tuscan Son” (Grantville Gazette 9).

Sodium carbonate (soda ash) was derived down-time from seaweed, or salt-tolerant land plants. The plant material was dried and burnt, and then the ash was washed with water, and the resulting crude solution boiled dry to yield a purer ash. Because sodium carbonate was used as a flux in the glass industry, some of the plants so used were called glasswort. Spain exported soda ash (30% sodium carbonate) made from Salsola soda; it was illegal to export the seed.

The modern American source is trona, which is a mixture of sodium carbonate and sodium bicarbonate. Trona itself is similar to ancient Egyptian natron (the inspiration for the “Na” used to symbolize the element sodium).

Sodium carbonate can also be synthesized, and several of the earliest industrial chemical processes were intended to produce it (EB11, Alkali Manufacture). The LeBlanc process (1790) featured a double replacement reaction; sodium chloride reacted with sulfuric acid to make sodium sulfate and hydrochloric acid. The sodium sulfate was then “fluxed” with calcium carbonate and coal.

The LeBlanc process was eclipsed by the ammonia-soda process (1838), especially as improved by Solvay (1872). Ammonia, water and carbon dioxide react in situ to make ammonium bicarbonate, and that reacts with sodium chloride to make sodium bicarbonate and ammonium chloride. The bicarbonate, when heated, releases the carbonate and carbon dioxide.

Sodium bicarbonate (baking soda) can be synthesized by adding carbon dioxide to sodium carbonate. Another route is by interrupting the Solvay process of producing sodium carbonate. Sodium bicarbonate was made in 1631-32 by Dr. Phil in Offord, “The Doctor Gribbleflotz Chronicles, Part 1: Calling Dr. Phil,” Grantville Gazette 6. Dr. Phil used the Solvay technique (Offord, private communication).

Sodium nitrate is found in large deposits in Chile, and can be converted to potassium nitrate by reacting it with potassium chloride. Glauberite (sodium calcium nitrate) is found in Stassfurt (EB11).

Glauber’s salt is sodium sulfate and was first prepared (as sal mirabilis) by Johann Glauber in 1658. He reacted sodium chloride with sulfuric acid.

Sodium fluoride is made by treating sodium hydroxide or sodium carbonate with HF.

Natural cryolite (sodium aluminum fluoride) is mined in Greenland, in summer 1633, by an expedition sponsored by Louis De Geer. Mackey, “Land of Ice and Sun” (Grantville Gazette 11). The cryolite is of value in making soda (sodium carbonate) by the “cryolite soda” process, as a flux in smelting aluminum from aluminum oxide, and as an aluminum ore in its own right. By early 1636, Dr. Phil’s HDG Enterprizes has made synthetic cryolite in small quantities. Offord, “Doctor Phil’s Family” (Grantville Gazette 15). Most likely, this was by reacting sodium hydroxide, aluminum hydroxide, and HF.

Sodium metal is a johnny-come-lately. It was first made in 1807 by electrolysis of sodium hydroxide. Then, for a period, it was made by igniting charcoal with sodium hydroxide. The current production method is by electrolysis of molten sodium chloride together with calcium chloride or sodium carbonate.

The metal is used as a reducing and dehydrating agent, and in sodium vapor lamps (EA). It can also be used in preparation of the organic chemical reagent sodium borohydride.

I expect sodium metal to be the first of the group 1 or 2 metals to be produced post-RoF. However, problems (see EB11/Sodium) should be expected in attempting to reduce so reactive an element.

Potassium

Potassium nitrate (saltpeter) is formed by bacterial action on human and animal waste. It is one of the three main ingredients of gunpowder.

Potassium chloride‘s principal use is as a source of potassium ions for plant growth. It’s also used to make potassium carbonate. It’s found naturally as sylvinite (potassium and sodium chloride) or carnallite (potassium and magnesium chloride). The latter is found at Stassfurt (Emsley 336).

Potassium chlorate has been used, by 1634, as a primer for the French “Cardinal” musket. Greg Ferrara, the USE’s R&D boss, had told Mike Stearns that “the production process would be way too complicated,” but that was an oversight on his part. Flint, 1634: The Baltic War, Chapter 55.

Potassium hydroxide (caustic potash) is made by down-timers by reacting calcium hydroxide (slaked lime) with potassium carbonate (EB11). The reaction is driven by the precipitation of the insoluble calcium carbonate. The modern production method (I don’t know if it’s known in Grantville) is by electrolysis of potassium chloride solutions.

Potassium carbonate was derived down-time from hardwood trees. The wood was burnt, and the ash washed with water and then boiled dry to yield the impure salt, potash. This could be baked in a kiln to make a purer form, pearlash. A modern method of making it is by electrolysis of potassium chloride in an aqueous solution, yielding potassium hydroxide, which in turn is carbonated.

Potassium bicarbonate is useful in baking and, interestingly, as a fire suppression agent.

Alum is not a single compound but rather a series of related compounds that are all “double” sulfates, that is, sulfates of two different cations, an alkali metal (or ammonium) and a trivalent metal. Potassium aluminum sulfate occurs in nature in the mineral alunite, from which it can be obtained by treatment with sulfuric acid. In the late twentieth century, the two most important alums were potassium aluminum sulfate and ammonium aluminum sulfate All of the alums can be made by mixing and heating solutions of the appropriate single sulfates. Alums are used as mordants, that is, to fix dyes to fabric.

Potassium metal isn’t used much, because its uses are similar to those of the cheaper metal, sodium. (EA).

Other Group 1 Metals

Rubidium, cesium and francium. Ignored.

Group 2 Metals

Beryllium

Beryllium metal was not readily available until 1957, and is somewhat problematic to use because of the toxicity of beryllium dusts. It’s alloyed with copper (EA) and nickel (Emsley 58), improving their conductivity and elasticity. Beryllium oxide, a ceramic, has a high melting point and is (unusually) both an electrical insulator and a good conductor of heat, making it useful in the electronics industry.

The principal ore is beryl (beryllium aluminum silicate), and European sources exist. (EB11/Beryl). Unfortunately, the beryllium content is rather low (Simons 13). The modern production method involves treatment with, successively, sodium fluoride, caustic soda and HCl or HF, and then electrolytic reduction of the chloride or magnesium reduction of the fluoride (EA). The caustic soda treatiment forms beryllium hydroxide, and this may instead be converted into beryllium oxide.

Calcium

Calcium oxide (quick lime) is produced down-time by heating calcium carbonate (from limestone) sufficiently to decompose it into calcium oxide and carbon dioxide. It is much more important than calcium metal (Table 2-3 above; Emsley 87).

Calcium hydroxide (slaked lime) has also been known since ancient times, and was made by reacting calcium oxide with water.

Calcium sulfate is, in hydrated form, the mineral gypsum. It is used in the manufacture of plaster of Paris and Portland cement.

Calcium hypochlorite. In December, 1633, Nicki Jo tells Scaglia and Rubens that Essen is producing this bleach and a disinfectant. Mackey, “Ounces of Prevention” (Grantville Gazette 5). It’s doing so by the processes described in Wagner, A Handbook of Chemical Technology (1872) (Mackey, private communication). Wagner makes calcium hypochlorite by reacting chlorine gas with slaked lime (calcium hydroxide). That reaction is described in EB11/Alkali Manufacture.

Calcium carbide is made by reducing lime with coke in an electric furnace at 2000oC. It is an “acetylene generator”; add water, and it decomposes into acetylene and lime. This reaction explains the flame of the miner’s safety lamp. Carbide lamps are used in January 1635 by the miners in Huston, “Twenty-Eight Men” (Grantville Gazette 10). They have clearly been made since the RoF since they were retrofitted onto the up-time hard hats.

Calcium carbonate is the principal mineral of chalk, limestone and marble, and is also found in shells.

Calcium metal is used as a reducing and drying agent. Its reducing power is such that it can react with water to generate hydrogen. It’s produced by electrolysis of the fused chloride or fluoride, or reduction of lime with aluminum and heat. (EA).

Plainly, several calcium compounds are going to be available even in 1631. Calcium metal can be made electrolytically in Grantville, which has cheap electricity, at least once we have graphite to serve as the anode. The question is when will the demand be sufficient to warrant the startup costs. My guess is that this will be affected by the demand for metals which can be reduced by calcium but not (at least easily) by carbon—e.g., sodium and magnesium. Outside Grantville, production of calcium metal will be dependent on the availability of aluminum.

Magnesium

Magnesium has a variety of uses. Pure magnesium powder is used in pyrotechnics and incendiaries. Magnesium is also necessary to make Grignard reagents, which are important organic chemical intermediates. The bulk metal can also be used to protect less active metals from corrosion. Magnesium-steel alloys have a high strength-to-weight ratio.

There are several ores (see table 2-3), and magnesium can also be found in seawater and salt well brines.

The secret to extracting magnesium from seawater (0.13% magnesium) is to add a calcium salt, which, by a double replacement, causes production of an insoluble magnesium salt. The latter can then be converted into the chloride.

Magnesium metal is preferably made by electrolytic reduction of the fused chloride, and CW215 suggests that it be a mixture of magnesium, calcium, and sodium chlorides (which avoids a separation step). The alternative is chemical reduction of the oxide with carbon or ferrosilicon, at high temperature, but it’s inefficient. EB11 teaches that carboreduction doesn’t work, but CW215 suggests heating with coke at 2000 degrees C followed by rapid quenching.

Epsom salts (crude magnesium sulfate) were discovered by Henry Wicker in 1618 (Emsley 245).

Strontium

Ignored.

Barium

Barium is found principally as barite (barium sulfate) or witherite (barium carbonate). Barite is associated with lead and silver ore veins (EB11/Barytes), and it can be found near Stuttgart (HCA).

It has been known since 1602 that barite phosphoresces if heated. Barium sulfate is undoubtedly known to Grantville’s doctors because of its use in X-ray studies of the digestive system. Lithopore (mixture of barium sulfate and zinc sulfide) is a white pigment. Both the sulfate and the carbonate are used in fluxes. Barium peroxide was used in 1818 to prepare hydrogen peroxide.

Barium proper, in vapor form, is used to remove oxygen, nitrogen and carbon dioxide from vacuum tubes. Manufacturing barium is by reduction of barium oxide (derived from the sulfate or carbonate) with aluminum.

In conclusion, barium is an element of mild interest. Barite itself can be exploited early on, but use of the metal must await the production of aluminum (and vacuum tubes).

Radium

Ignored.

Group 3 Metals

Scandium, yttrium and lutetium. Ignored.

Group 4 Metals

Titanium

Titanium dioxide occurs in nature as the minerals rutile, rookite and anatase. There are hints that Portuguese entrepreneurs are considering the mining of the Keralese beach deposits in Cooper, “Gajam Raanni,” (Grantville Gazette 25).

Titanium was first used as a structural metal in 1952 (NACE 1978); it has the highest strength-weight ratio of the metals (making it attractive for aircraft), and good corrosion resistance. Some of the early methods of extracting titanium include fusion with potassium bisulfate or with potassium carbonate, or by the reaction of titanium fluoride with steam. (EB11/Titanium). However, EA says that it was not possible to mass produce titanium until the development of the Kroll process (1937), which it nonetheless characterizes as “relatively slow and costly.” The requirements for the Kroll process include chlorine (to form titanium tetrachloride), metallic magnesium or sodium, and some kind of inert atmosphere (typically argon or helium, and very definitely not nitrogen).

In view of the discouraging text, not to mention the difficulties of satisfying the prerequisites for the Kroll process, I think that titanium isn’t likely to be exploited, in elemental form, in the 1630s.

Of course, that doesn’t mean we can’t make use of titanium dioxide directly; it’s a fine white pigment.

Zirconium

The principal ore of zirconium is zircon (zirconium silicate), which pokes its toe into seventeenth century international commerce as a Ceylonese gemstone. EB11 gives instructions for the preparation of zirconia (zirconium oxide); you need potassium fluoride, hydrofluoric acid, and ammonia. Zirconia is good for high-temperature ceramics.

We could probably make zirconia in 1636, but we aren’t like to have a need for it for many years later. The metal (which has good corrosion resistance) is unlikely to be of interest in the up-timers’ lifetimes.

Hafnium

Ignored.

Group 5 Metals

Vanadium

Vanadium is the most important metal of this group. The metal is alloyed with steel and titanium; Runkle listed it as an important ingredient for tool steel and stainless steel.

Vanadium was first discovered in the slag from a Swedish iron smelter; the iron ore came from Taberg. The ore descloizite (lead-zinc vanadate) is found at, among other sites, Eisen-Kappel near Klagenfurt in Carinthia, associated with lead ores (EB11). There are other ores which are oxides or sulfides. There are relatively few mines. (Simons 224).

Considering not only what the encyclopedias say about vanadium extraction, but also other sources, I get the impression that the process is on the difficult side. While the encyclopedias mention alternative methods, it seems that the principal large-scale process is reducing the oxide with calcium (see above) in the presence of calcium chloride or iodide at a high temperature, and possibly in an inert gas (argon) atmosphere. (Emsley, 485; Simons, 224; Patnaik 964). EA mistakenly says that vanadium oxide can be reduced with carbon (Emsley).

The only hope I see for “first decade” production of vanadium-steel alloys is if we can make those alloys without first extracting vanadium. Emsley says that “ferrovanadium,” which is what is added to steel, can be made from vanadium oxide by heating with ferrosilicon.

Niobium

Niobium (EB11/Columbium) is used as an anti-corrosive alloying element in steel. It was produced, at the end of the twentieth century, at a rate of 25000 tonnes annually, of which over 85% came from Brazil (Emsley 284). Its principal ore is columbite, a complex oxide. The closest source to the USE is probably Rabenstein, Bavaria (EB11/Columbite), but I don’t know if it’s an economic one.

Tantalum

In 2000, the demand for Tantalum was only around 1000 tonnes annually. At one time, lamp filaments were made out of Tantalum, but such were superseded by Tungsten. Nowadays, it’s important mostly because of its superb corrosion resistance, which is comparable to that of glass (rembar.com), and as a melting point-enhancing alloying element. It is usually found with niobium.

Group 6 Metals

Chromium

Chromium is needed to make stainless (>10% chromium) steel, and other alloys (nichrome, stellite). It can also be used to plate other metals. The principal ore of chromium is chromite (ferrous chromate). There are chromite prospectors in Kemi, Finland by 1633 (1633 Chap. 26) and in Maryland perhaps by 1634. (Mackey, “Trip to Paris,” Grantville Gazette 9). However, as of July 1634, Lolly Aossi is not aware of any chromium having come on the market yet. Runkle, “Sunday Driver” (Grantville Gazette 13).

Of course, the chromium has to be extracted. It can be recovered by several methods, one of which successively requires soda ash, coke and aluminum, and the other, direct reduction with carbon or silicon in an electric arc furnace (EA), and there are additional variations disclosed by EB11. My guess is that the electric arc furnace process would be favored in Grantville, which has plentiful cheap electricity. Essen has also obtained cryolite (see “Aluminum”) and therefore might favor the first method.

EA/Steel says that virtually all stainless are at least 11.5% chromium, and that AISI 302, with excellent corrosion resistance, is 0.15% carbon, 18% chromium, 8% nickel.

Chromic oxide is used as a green pigment, and it’s an intermediate in the soda ash-coke process of producing chromium. Sodium chromate, sodium and ammonium dichromate, and chromic acid are strong oxidizing agents. Lead chromate (the mineral crocoisite) is the pignment “chrome yellow.” Chromic acid is obtained by dissolving chromic oxide in water, and EB11/Bichromates and chromates explains how to make many of its salts.

The earliest I imagine chromium metal could be available is 1635, but 1636 is perhaps more likely. Once we have chromic oxide, we can make the chromates and dichromates.

Molybdenum

Molybdenum is used as a catalyst and an alloying element (it hardens and toughens steel). The main ore is molybdenite (molybdenum sulfide), which looks quite a bit like graphite, and like it was used in “lead” pencils (Sarkar, 504). Chemists didn’t distinguish molybdenum from graphite until 1779. It’s therefore conceivable that a pre-RoF reference to a “drawing lead” is actually to molybdenite, but I think that unlikely, as the principal European deposits are in Norway and Norway was not very developed in the early seventeenth century.

Absent that fortuitous breakthrough, exploitation of molybdenum will require prospecting, by individuals armed with good descriptions of the mineral, and sent to the right vicinity. We know to look for molybdenite in a few European locales (notably, at Slangsvold near Raade in Norway, EB11/molybdenite), but if those fail, our best bet is probably once European civilization reaches Colorado (EA). I have no idea how many years that will take.

Curiously, there are fourteenth century samurai swords which are rich in molybdenum, no doubt as a result of the use of native molybdenite at some point in the forging process.(Emsley 266).

Tungsten

Tungsten is probably best known to the average up-timer as the filament of the electric light bulb. It’s also used, less obviously, to make alloys with iron and aluminum, and in high temperature processing. Josh Modi told magnate De Geer that “tungsten would allow you to make a steel close to what was called ‘hi-speed tool steel.'” Mackey, “The Essen Chronicles, Part 3: Trip to Paris” (Grantville Gazette 9). The cemented carbide tools which Larry Wild broke (Cresswell and Washburn, “When the Chips are Down,” Ring of Fire) are sintered composites of tungsten carbide in a cobalt matrix.

Tungsten’s principal ores are scheelite (calcium tungstate) and wolframite (iron-manganese tungstate), which can accumulate in placer deposits. Josh also declared that, according to the encyclopedias, tungsten can be obtained from the tailings of tin mines. Mackey, “Essen Steel, Part 1: Crucibellus” (Grantville Gazette 7). EB11/Wolframite confirms that wolframite is “commonly associated with tin ores.” However, this should not inspire a false sense of confidence. “This element is far less widely distributed than tin . . . it is probable that on the whole tungsten is not more than one-hundredth as abundant in Cornwall as tin . . . there may be tin without tungsten, but not tungsten without tin.” (Collins 333-4). The principal European source is actually Portugal (EA).

Wolframite does occur with cassiterite in the Erzgebirge (Mineral Industry of the British Empire and Foreign Countries, 24, 1921) and indeed it’s believed to have gotten its name from German tin miners (they thought the wolframite devoured tin) (Emsley 470) So we will be able to produce tungsten . . . just in limited quantities.

To complicate matters further, the standard methods of extracting tungsten are reduction of tungstic acid with aluminum or heating tungsten oxide with carbon in an electric furnace. So you either need aluminum, or lots of electricity.

Group 7 Metals

Manganese

The principal ore of manganese is pyrolusite (manganese dioxide), and it has been used since ancient times to decolorize soda lime glass (which otherwise has a greenish cast) or to instead give it a purple tint (Emsley 250-1). It can be (and probably already is) mined at Ilmenau and elsewhere in Thuringia (EB11/pyrolusite). Another ore is rhodochrosite (manganese carbonate), which is associated with silver.

The alloy ferromanganese (introduced 1839) is used in steelmaking to remove oxygen and sulfur impurities; indeed, it was a key component (1856) in the perfected Bessemer process. Manganese also serves as an alloying element. EA says, somewhat cryptically, “Ferromanganese is produced from manganese ore in blast furnaces in somewhat the same manner as pig iron.” I take this to mean that a mixture of hematite (iron oxide) and pyrolusite (manganese oxide) is reduced by the combination of heat and carbon.

The metal can be obtained from the oxide by reduction with aluminum, or by heating the carbonate with carbon. (EB11). Another option, which I am not sure is documented in Grantville, is to electrolyze magnesium sulfate (Emsley 251). The sulfate, in turn, can be derived from manganese dioxide and concentrated sulfuric acid (EB11).

Manganese dioxide is also useful as an oxygen source in a common dry cell (EA) and in the production of chlorine gas by reaction with HCl (Emsley 251).

The manganates are made by fusion of manganese dioxide with the metal hydroxide in presence of an oxidizing agent. In turn, adding carbon dioxide or chlorine to a manganate should yield the permanganate. Potassium permanganate is an important oxidizing agent. Historically, potassium manganate and permanganate were both prepared by Glauber in 1659.

It seems to me that the down-timers had everything they needed to make ferromanganese and potassium permanganate, all they lacked was knowledge of the underlying chemistry (so they would know what to do with the pyrolusite) and of the uses for these materials (so they would have the motivation).

I figure that until the end of the Baltic War, there will be reluctance to fiddle around with new alloys; the emphasis will be on producing as much basic steel as possible. However, by late 1634, there may well be some ferromanganese production. And I think that the alchemists will be experimenting with potassium permanganate even earlier, perhaps in 1633. After all, there is a connection between strong oxidizing agents and things that go BOOM!

Technetium, Rhenium

Ignored.

Platinum Group” Metals

These are ruthenium, rhodium, paladium, osmium, iridium and platinum, which are group 8, 9 and 10 elements which tend to occur together, because of their high density and similar chemical properties, in the same deposits. However, the same is true, to a lesser degree, of iron (group 8), cobalt (9) and nickel (10).

Group 8 Metals

Iron

Iron, of course, is extremely well known to the down-timers, and methods of using it to make steel are discussed in Boatright, “Iron” (Grantville Gazette 3). Iron forms both ferrous (+2 valence) and ferric (+3) compounds; EA gives uses for ferrous sulfate, sulfide, phosphide, chloride, and ferric sulfate, sulfide (pyrite), chloride, hydroxide, oxide (hematite) and bromide. All of these compounds can be made by disclosed reactions (EA, EB11) of iron, or iron compounds, with reagents such as sulfuric acid, sulfur, phosphorus, hydrochloric acid, chlorine, ammonia and bromine. Chlorine is available in NTL 1633, and most of the others even sooner. The only question mark is bromine.

Some iron compounds are known to down-timers. Green vitriol (vitriol of Mars) is ferrous sulfate heptahydrate. The down-timers use it in the manufacture of iron gall ink, and as a mordant. In the nineteenth century it was used as a developer in the collodion process. Yellow ochre is hydrated ferric oxide. Colcothar is a red iron oxide made by roasting green vitriol.

Ruthenium

Ignored.

Osmium

Osmium metal is associated with other platinum group metals, and was first discovered in “the residue left when crude platinum was dissolved by aqua regia.” (EB11). Nowadays, it’s a byproduct of nickel refining. It’s useful as a catalyst (it was one of the first good catalysts for the Haber process) and in specialized applications requiring extreme hardness (nibs for ultra-expensive fountain pens). Still, the demand for it is minute. (<100 kg 2000: Emsley 295).

The most intriguing osmium compound is the tetroxide, which is used to detect fingerprints. (EA) EB11 says that it forms “when osmium compounds are heated in air, or with aqua regia, or fused with caustic alkali and nitre.”

Even though it was discovered at a surprisingly early date (1803), Osmium is not likely to be seen commercially in our time frame.

Group 9 Metals

Cobalt

Cobalt is a peculiar case. Cobalt compounds have been used since the days of the Pyramids to make blue glasses and ceramics. This usage appears to have been forgotten in Europe after the fall of Rome as, when medieval miners in Saxony and Bohemia encountered cobalt ores (smaltite, cobalt arsenide) in Saxony and Bohemia, they didn’t consider them to be of value (Emsley 116). In this regard, cobalt’s rather like nickel. However, it appears that “smalt” was rediscovered by the Bohemian glass makers in 1540-60 (Gettens, 158).

At some point in the seventeenth century, it was discovered that if a cobalt ore were dissolved in aqua regia, it formed an “invisible ink” which was revealed by heat. (Emsley 119).

Cobalt is used principally as an alloying element, to impart temperature resistance. Cobalt arsenides are associated with nickel, silver and gold, and other cobalt ores with copper. (EA). Cobalt is mostly produced as a byproduct of nickel refining (Emsley 117) and EA warns that extraction is usually “complicated . . . because of the presence of numerous contaminating elements.”

Rhodium

Rhodium, discovered in 1803, is used as an alloying element, in electroplating other metals, as a reflective coating for mirrors, and as a catalyst (in particular, in catalytic converters of cars). It’s a byproduct of platinum mining. (EA; Emsley, 362).

Iridium

Iridium, discovered in 1802, is perhaps the most corrosion-resistant of the metals, and is used in electrical contacts and pen points (EA). While found in a platinum ore (EB11), it’s usually a byproduct of nickel refining (Emsley 202).

Group 10 Metals

Nickel

The European down-timers have encountered nickel, but without realizing what they were dealing with. The nickel ore niccolite (nickel arsenide) is found with cobalt, silver and copper in Saxon mines (EB11/Niccolite). The miners called it kupfernickel (St. Nick’s copper), because they deemed it a demonic imitation of copper. Lolly Aossi tells Father Smithson in July 1634 that nickel has been found in tailings from more than one mine. Runkle, “Sunday Driver” (Grantville Gazette 13).

The up-time interest in nickel is likely to be mostly in the metal itself, which can be used as a catalyst, as a metal plating agent, or in alloys with other metals, such as iron (Stainless steel), copper (Monel) or chromium (Nichrome). “Alnico,” an aluminum-nickel-cobalt alloy, can be used to make very powerful but compact magnets. Magnets made from Alnico were key to the production of early microphones, headphones, and loudspeakers.

Palladium

Palladium is associated with platinum (and nickel) ores, and with certain placer deposits of gold and silver. It’s used mainly as a catalyst, and, in that guise, can be found in the catalytic converters of cars manufactured shortly before the RoF (older converters used platinum).

It will probably not be sought out independently, but those refining the associated metals may keep an eye out for it.

Platinum

EA says that “platinum was known and used by pre-Columbian Indians in South America,” and this is confirmed by archaeological evidence. Even before RoF, a few Europeans were aware of platinum’s existence: “In 1557, an Italian scholar, Julius Scaliger, wrote of a metal from Spanish Central America that could not be made to melt and this must have been platinum [MP 1772°C].” Serious European investigation did not begin until the early eighteenth century. Even then, the authorities initially considered platinum to be detrimental (it could be used to adulterate gold), and banned it for decades. (Emsley 319).

If the up-time texts excite Spanish interest in mining platinum, then they can probably find Indians who know where the Columbian deposits are located. It appears that sometime after October 1633, Antonio (“Catalina”) de Erauso went to Cartagena to begin a search for platinum. Mackey, “Land of Ice and Sun” (Grantville Gazette 11).

There are three commercial forms of platinum: powder (“platinum black”), spongy and compact. Both the powder (made by reduction of platinum chloride) and the spongy form (made from ammonium chlorplatinate) are used as catalysts, whereas the compact platinum is formed into jewelry. Platinum is also used in many applications in which heat or corrosion resistance are important.

Ammonium chlorplatinate is obtained by dissolving platinum ore in aqua regia, and then precipitating the desired salt by adding acidified ammonium chloride. Platinum chloride is recovered if you heat chlorplatinic acid in dry chlorine or dry hydrogen chloride. (EB11).

Platinum dioxide, which is also a catalyst, is made by fusing chlorplatinic acid with sodium nitrate (EA) or caustic soda (EB11).

Group 11 Metals

The elements of this group are all available pre-RoF as both the elemental metals, and in several salts.

Copper

Copper carbonate occurs naturally as malachite and azurite, copper sulfide as chalcocite, chalcopyrite and bornite, and copper oxide as cuprite.

Copper hydroxide, a pigment, was made down-time by reacting sodium hydroxide (lye) with copper sulfate (blue vitriol). It can also be made electrochemically. (Analogous reactions are used to make iron hydroxide.)

The alchemists’ “blue vitriol” (“blue copperas”) is copper sulfate pentahydrate. Its modern use is as a pesticide and analytical reagent.

Copper sulfate and nitrate are made by reacting copper (or copper salts) with sulfuric or nitric acid, respectively. The hydroxide and chloride are also easy to make. (EB11). The chloride (“resin of copper”) was reportedly made by Robert Boyle in 1664 (Levity).

Silver

Silver nitrate, the most important (and least expensive) silver compound even today, was known to the alchemists as “lunar caustic,” “magisterium argenti,” “crystalli dianae,” or “lapis infernalis.” It was first prepared by the eighth century Geber, and came into medical use in the seventeenth century (Sadtler, 402). Beginning in the nineteenth century, it was placed in newborns’ eyes to prevent eye infections (EA), and used to silver mirrors by the Liebig method. It is made by reacting the metal with nitric acid.

Silver chloride is the alchemists’ lac argenti (milk of silver) or luna cornea, which occurs in nature as the mineral cerargyrite (horn silver) and in that form was described by Oswald Croll in 1608. Since it is insoluble, it can also be obtained by reacting a soluble chloride with a soluble silver salt (such as the nitrate). Silver bromide and iodide are obtained analogously. These three insoluble silver halides darken on exposure to light, which explains why they are useful in photography. Silver iodide has also been used to seed clouds.

Fulminating silver (silver nitride) is an explosive, and was also known to the alchemists. EB11 says that it can be set off by the touch of a feather.

Silver carbonate and silver chromate are used in organic synthesis. Silver carbonate is made readily from sodium carbonate and silver nitrate. The bottleneck in producing silver chromate will be obtaining a chromate salt; potassium will do.

Silver salts have been used to sterilize drinking water. (Emsley 396).

Gold

Gold is ductile, a good electrical conductor, and resistant to corrosion. Gold coatings are occasionally used in industrial chemical equipment to protect surfaces from corrosive fluids. But because of gold’s unreactiveness, preparation of gold salts isn’t easy.

Gold chloride may be made by dissolving gold in aqua regia, a mixture of nitric and hydrochloric acids. Gold chloride is the secret ingredient of the ruby glass of Bohemia, which would have been invented by Johann Kunkel (1630-1703) but for the RoF.

If gold chloride is reacted with stannous chloride, you obtain Purple of Cassius, which is a mixture of colloidal gold (tiny gold particles) and tin oxide (EA). It was first made by Andrea Cassius in 1685, and it was used to impart a ruby color to glass, or purple to porcelain. (Emsley 167).

Another interesting compound is gold hydrazide (“fulminating gold”). It has been called the world’s first high explosive, and it was probably made inadvertently by alchemists when they fiddled around with gold chloride (EB11). We know that both Robert Hooke and Johann Glauber experimented with it during the seventeenth century. (Lateral Science)

Since gold is readily available, but expensive, the commercial appearance of gold compounds is going to be “demand-driven.”

Group 12 Metals

Zinc

Brass, a zinc-copper alloy, has been used since antiquity, but the ancients thought of it as simply being a form of copper. Elemental zinc was produced in medieval India (hence the name “Malabar Lead” or “Indian Lead”), and zinc smelting technology was transmitted to China by the sixteenth century. In 1597, Libavius received a sample of Indian zinc, but he took it to be a “peculiar kind of tin.” I strongly suspect that the “Japanese zinc” mentioned in 1634: The Galileo Affair (Chapter 33) is actually Chinese.

After the RoF, zinc is in high demand. Chad Jenkins complains in May 1632 that zinc is not available for galvanization at a reasonable price. Rittgers, “Von Grantville” (Grantville Gazette 7). In April-July 1633, the recycling crew carefully strips zinc off any unusable galvanized steel, and “later date American pennies” (mostly zinc) have been pulled out of circulation. Schillawski and Rigby, “Recycling” (Grantville Gazette 6). Nonetheless, in 1634, Lewis Bartolli has arsenic-free rods of zinc metal. Cooper, “Arsenic and Old Italians” (Grantville Gazette 22).

Nowadays, the principal ore of zinc is sphalerite (zinc sulfide), which is found in the Harz. However, prior to the RoF, it was calamine (a zinc carbonate, with some zinc silicate) which was most likely to be used by Europeans to make brass. Indeed, Beckmann (78) asserts that by the mid-sixteenth century, there was isolated use of “furnace calamine” (a calcined zinc) from the Rammelsberg mine in alloying, under the guidance of Erasmus Ebener. Clark, “The Secret Book of Zink” (Grantville Gazette 2) explains how zinc can be extracted from calamine and used to galvanize iron, or to make zinc oxide or zinc chloride.

Nonetheless, Dr. Phil ordered a large quantity of sphalerite in December 1633. He received about five tons worth, in fact, and figured out how to recover, not only the zinc, but various potentially salable byproducts, notably sulfur and sulfuric acid. Offord, “Dr. Phil Zinkens A Bundle” (Grantville Gazette 7). Nonetheless, in April 1634, Sharon Nichols orders two hundred tons “Japanese zinc” from her Venetian connections, and expects delivery by June 1635.

In the twentieth century, the single most important use of metallic zinc was in galvanizing steel, but it’s still used to make brass (and other alloys). Because of its reactivity, it’s useful as an anode in certain batteries, and it can be used as an oxidizing agent in chemistry.

Zinc oxide was produced in thirteenth century Persia (Emsley 502). It’s a normal intermediate in the reduction of zinc ore to zinc, but it is preferably produced by oxidizing zinc vapors (EB11/Zinc). It’s a possible substitute for titanium dioxide as a white pigment.

The alchemists’ “white vitriol” is zinc sulfate heptahydrate, and it appears to have been known by the late sixteenth century (Beckmann, 81). It’s produced by reacting zinc with sulfuric acid. It’s now used in making rayon, and in zinc plating.

Zinc chloride is mentioned as a wart remedy by Clark, “The Secret Book of Zink” (Grantville Gazette 2). It’s used to preserve and fireproof wood, and for other purposes (EA). It can be made by reaction of the metal with chlorine gas or with HCl (EB11).

The main impediment to exploitation of zinc and its standard salts is one of communication; the up-timers have to accurately convey what it is that they are seeking. Zinc and its compounds should be available in limited quantities in 1633, but it will probably be a struggle to keep up with demand until 1635 or so.

Cadmium

The situation of cadmium is a peculiar one, in that the down-timers may have encountered one of its compounds, without knowing that it contained a new element, and it’s uncertain whether the up-timers will be able to enlighten them.

The principal cadmium ore is greenockite (cadmium sulfide). Emsley (77) asserts that it was mined in Classical Greece and used as a yellow pigment. This is plausible—it is found, in association with calamine, at Laurion—but it has been challenged by other authorities (Eastaugh, 176).

Cadmium sulfide entered the historical record in the early nineteenth century, when Stromeyer, the inspector of pharmacies, investigated a complaint by Hannover druggists that the zinc oxide they made by heating calamine sometimes was yellow rather than white. (Emsley 76). The calamine trade is centuries old, and I can’t help but wonder whether this yellow adulterant is known to the seventeenth-century apothecaries. If so, they may question the up-time chemists about it, and Stromeyer’s discovery may be anticipated by several centuries.

In any event, the up-timers certainly know that the element cadmium exists, and that it’s associated with flue dust from zinc ore processing.

Cadmium plating has been used to protect steel from salt water, and cadmium is used in Nicad batteries. Cadmium sulfide is used as a pigment, semiconductor, and photoelectric cell component.

That said, I don’t expect cadmium to come into the chemical marketplace in the NTL 1630s.

Mercury

Mercury, as a pure element, is known to the down-timers. It is obtained from cinnabar (mercuric sulfide). The mineral itself can be powdered to produce the pigment vermillion.

Mercury, although used by down-timers as a treatment for syphilis, is poisonous, which comes as an unpleasant surprise to potion maker Guba Ivashka Kalachnikov. Huff and Goodlett, “Butterflies in the Kremlin, Part Four” (Grantville Gazette 11) and is mentioned by Dr. Abrabanel in “Venus and Mercury” (Lee Grantville Gazette 24).

Heating the metal in air yields mercuric oxide, and then dissolving the oxide in nitric acid provides the nitrate (mercurous nitrate if dilute acid, mercuric if strong).

Mercurous chloride (calomel) occurs in nature, and was used in medicine (as laxative and diuretic) before RoF. It can be produced by heating mercury in chlorine, or reducing mercuric chloride, or reacting mercurous sulfate and sodium chloride.

Mercuric chloride (corrosive sublimate) is another alchemical favorite, and can be obtained by heating mercuric sulfate with sodium chloride, or mercurous sulfate with HCl, or by simply chlorinating mercury or calomel. (EB11). The twelfth century Indian method was “heating mercury, salt, brick dust and alum for 3 days in a closed earthenware pot, and then adding water to dissolve out the corrosive sublimate before crystallizing it” (Emsley 255).

Fulminate of mercury is used to make percussion caps. It is very dangerous stuff. Flint, 1633, Chapter 28. By September 1633 it is available in the USE, manufactured in pounds per week quantities, by workers receiving “hazardous duty” pay. Offord and Boatright, “The Dr. Gribbleflotz Chronicles, Part 2: Dr. Phil’s Amazing Essence Of Fire Tablets” (Grantville Gazette 10) and Offord “A Change of Hart” (Grantville Gazette 25).

In April 1632, Jan de Vries, one of De Geer’s lieutenants, mentions plans to make mercury fulminate by “trial and error.” Mackey, “The Essen Steel Chronicles, Part 2: Louis de Geer” (Grantville Gazette 8). By 1634, the Bernese are experimenting with it. Evans, “Thunder in the Mountains” (Grantville Gazette 12). And the French did the same, until Glauber gave them a better solution (see potassium chlorate). Flint, 1634: The Baltic War, Chapter 27.

With the exception of the dangerous fulminate, mercury and its “standard” salts should be readily available, in limited quantities, even in 1631-32.

Group 13 Metals

Aluminum

Aluminum is a precious metal, post-RoF, at least until we start producing it again. Massey, “Ultralight” (Grantville Gazette 9); Bergstralh, “One Man’s Junk” (Grantville Gazette 4); Schillawski and Rigby, “Recycling” (Grantville Gazette 6). Anneke has a brand-new aluminum slide rule; made from a recycled strip. Carroll, “Stepping Up” (Grantville Gazette 14). Dr. Phil is driving people a little crazy with his quest for aluminum. Offord and Boatright, “Dr. Phil’s Aeolian Transformers” (Grantville Gazette 6); DeMarce “Songs and Ballads” (Grantville Gazette 14); Cooper, “Stretching Out, Part Three: Maria’s Mission” (Grantville Gazette 14).

Cooper, “Aluminum: Will O’ the Wisp?” (Grantville Gazette 8) explains where to find aluminum ores, how to extract alumina, how to refine it to obtain elemental aluminum, and finally how to use it. To mass produce it, we need bauxite as the ore, cryolite as a flux, and lots of cheap electricity. There are older methods which involve use of sodium or potassium as reducing agents. Right now, I am guessing “new” aluminum will be available in 1635. That’s earlier than I said in my article, because I wasn’t expecting cryolite to be mined as early as 1633. In Mackey, “Land of Ice and Sun” (Grantville Gazette 11), de Erauso brings back over 80 tons.

Aluminum sulfate occurs in nature as keramohalite, but is more likely to be obtained by treating an aluminum-rich kaolin or china clay with sulfuric acid (EB11/Aluminium).

Aluminum oxide (alumina) is an intermediate in the processing of bauxite into aluminum It is also the chemical composition of corundum, rubies and sapphires..

This is probably as good a place as any to mention the thermite reaction. Thermite is a mixture of aluminum powder and a metal oxide, usually iron oxide. When heated to the ignition temperature, the aluminum reacts with the iron oxide to form aluminum oxide. This is an extremely exothermic reaction, so it creates intense heat. Erwin O’Keefe demonstrates thermite welding to Dr. Phillip Gribbleflotz in Offord and Boatright, “The Dr. Gribbleflotz Chronicles, Part 2: Dr. Phil’s Amazing Essence Of Fire Tablets”(Grantville Gazette 7), and Dr. Phil’s Candles of the Essence of Light are apparently fabricated from thermite powder. Thermite reactions are more than just a curiosity, or a useful welding technique; they underlie the method used to extract a number of metals from their ores.

Aluminum hydroxide is found in nature as the mineral gibbsite, in the ore bauxite. It is also produced by the Bayer process (1887) from alumina in bauxite. See Cooper, “Aluminum: Will O’ The Wisp?” (Grantville Gazette 8).

Gallium

A byproduct of aluminum refining, hence unlikely to be exploited in the 1630s. It’s used in high temperature thermometers because of its large liquid range. It is also a useful alloy. Gallium arsenide is a semiconductor.

Other Group 13 Metals

Indium, Thallium. Ignored

Group 14 Metals/Metalloids

Germanium

Germanium is associated with silver, copper and zinc ores. (EA). In its heyday, germanium, suitably “doped,” was a major semiconductor material. (You needed to provide it in extremely pure form, of course.) Nowadays, it has been superseded by other semiconductors, and its main use is in glass for infrared devices.

For germanium to be of interest post-RoF, we would have to have advanced to the point of needing semiconductors, but not be able to make silicon in the necessary purity (Wikipedia).

Tin

Tin is obtained by heating cassiterite (tin oxide) with coke. Tin was used in the ancient world; the tin was alloyed with copper to make bronze. There is also tin foil, which was used pre-RoF as a reflective backing for glass mirrors.

Tin is an old commodity, but the up-timers will teach the old dog some new tricks. These will include new alloys (e.g., Babbitt metal), tin plating of steel (the tin can be transferred to the steel surface by the chloride or sulfate, or deposited electrolytically), and the use of molten tin as the “float” in the Pilkington float glass process.

Some tin compounds are already known to the down-timers. Stannous chloride was discovered in 1630, by Cornelius Drebbel, to be useful as a mordant. Satterlund, “Dyes and Mordants” (Grantville Gazette 5). Stannic chloride (spiritus fumans) was made by Libavius in 1605, by reaction with mercuric chloride (Levity).

Other tin compounds should be fairly easy to make by instructions given in EB11, assuming the reagent is available.

Lead

Lead is also a metal of antiquity. Its principal ore is galena (lead sulfide), but cerussite (lead carbonate) and anglesite (lead sulfate) are also of interest. We will want to use lead, which is corrosion resistant, in sulfuric acid processing, and in making storage batteries.

Lead oxide (litharge) is known to the down-timers. But the up-timers will reveal to them that it can be used to make lead-alkali (flint) glass.

They are also familiar with lead acetate (sugar of lead), which is obtained by reacting lead oxide with vinegar.

Lead nitrate (calx plumb dulcis) was known to Libavius (EB11).

Lead chromate is a useful yellow pigment, known since the early nineteenth century (Eastaugh 99).

The availability of lead compounds is going to be “anion-limited”; we need HCl to make the chloride, and chromate to make chrome yellow.

Group 15 Metals/Metalloids

The elements of this group are all known to down-timers (although not qua elements) and hence can be exploited fairly rapidly.

Bismuth

The down-timers have encountered bismuth (which occurs naturally in elemental form), although it’s often confused with lead, tin, and antimony. It can be mined in Schneeburg, Saxony and Joachimsthal, Bohemia (EB11).

The metal’s most interesting use is perhaps as an alloying element in the low-melting Wood’s metal. (EA). Unidentified bismuth compounds have been used to treat syphilis. Bismuth oxychloride might be used by the “new” cosmetics industry to impart an iridescent look.

Arsenic

Arsenic is known to the down-timers in the form of the naturally occurring realgar (disulfide) and orpiment (trisulfide), and the synthesized “white arsenic” (trioxide). The alchemists have probably made the metal itself, too, by reduction with carbon.

The most common mineral is arsenopyrite (iron sulfide + iron arsenide, and arsenic is usually obtained as a byproduct of copper, gold, silver, lead, nickel or cobalt mining.

Realgar and orpiment were in pre-RoF use as pigments (and white arsenic as the infamous “inheritance powder”).

The reaction of arsenic with concentrated nitric acid produces arsenic acid. Several arsenate salts (copper, calcium, lead) were used in decades past as insecticides.

Arsphenamine, an organo-arsenic compound, was introduced in 1909 as a treatment for syphilis. Its formula appears in the Merck Index.

The most interesting use of elemental arsenic is probably in the hardening of lead shot (EB11/Lead).

Antimony

Antimony is found in nature as white antimony (oxide) and black antimony (trisulfide; stibnite; kohl). The down-timers have isolated the pure element and quite a few antimony compounds. The Triumphal Chariot of Antimony, 1604 (attributed to Basil Valentine but actually written by Johann Tholde) apparently refers to antimony trichloride (“butter of antimony”), antimony nitrate (“fixed antimony”) and antimony oxysulfide (“glass of antimony”). Oswald Croll’s Basilica Chymica (1604) discussed antimony trichloride and antimony oxychloride (“powder of algaroth”). (Emsley II, 201). Beguin’s Elements of Chemistry (1615) describes the reaction of stibnite with mercuric chloride to make antimony trichloride (Salzberg 151). The Chaldeans used lead antimonate, and there is reason to believe that Greek fire included antimony sulfide.

Antimony metal is used to harden lead. Antimony oxide is used as a flame retardant.

Predictions

Table 2-4 suggests a chronology for when the metals are available, whether as cations of salts, or in elemental form. This is a summary (and gross simplification) of the analysis earlier in the article. If you are interested in a particular chemical, read the detailed analysis. In general, the metal will first be available as a cation of a naturally occurring salt, and only later (sometimes much later) as the metal itself. How soon will depend on both the demand for the metal and the ease of extraction.

This table is not canon! If enough effort is devoted, early enough, chemicals can be produced earlier than what I forecast.

Some of the elemental metals (e.g., potassium) are listed relatively late because I don’t expect them to be in great demand. If you want to write a story whose character has an earlier need for potassium, that’s fine with me and the Grantville Gazette Editorial Board.

Production of elemental metals is dependent on access to the ores. For the purpose of this article, I ignore which countries control that access. For control issues, see Cooper, “Mineral Mastery: Discovery and Control of Ore Deposits After the Baltic War” (Grantville Gazette 23).

And I assume a minimum of problems in getting access. Please note that it can take years to get an expedition approved by a government (especially Spain), and without bureaucratic blessing, you can find your mother lode and have it taken away from you immediately. And it can take more years to find the deposit, especially if you have to hack your way through jungle or fend off unfriendly natives.

Production of metals can also be dependent on technology, e.g., an electric furnace or electrochemical cells. I have assumed that these are available, on a laboratory scale, by 1633.

There are also “connections” between chemicals. For example, some elements are found as byproducts of mining for other elements. Others are extracted using another element (e.g., aluminum, sodium) as a reducing agent. That increases the demand for the reducing agent, but means that if there are problems obtaining it, the chronology gets thrown off.

Table 2-4: Suggested Availability of Elemental Metals and Metal Cations

Metals

(elemental)

Metal Cations (Salts)

Pre-RoF

Iron, Platinum (Pre-Columbian), Copper, Silver, Gold, Zinc, Mercury, Tin, Lead, Bismuth, Arsenic, Antimony

Sodium (chloride, carbonate, hydroxide, nitrate), Potassium (nitrate, chloride, hydroxide, carbonate, alum), Calcium (oxide, hydroxide, sulfate, carbonate), Magnesium (sulfate), Barium (sulfate), Manganese (dioxide), Iron (sulfate, oxide), Copper (sulfate), Silver (nitrate, chloride, acetate, nitride), Gold (chloride, hydrazide), Zinc (sulfate, oxide), Mercury (sulfide, chloride),

Tin (chloride), Lead (oxide, nitrate, acetate),

Arsenic (sulfides, oxide), Antimony (oxide, sulfide, etc.), Cobalt (arsenide), Aluminum (alum)

1631-32

additional “common” salts of above metals (e.g., sodium sulfate and bicarbonate); Nickel (arsenide); manganates, permanganates, arsenides, arsenates, antimonates

1633

Nickel

Aluminum (cryolite)

1634

Sodium, Magnesium, Manganese

Aluminum (bauxite), aluminates

1635

Chromium, Platinum, Aluminum

chromates, dichromates

Molybdenum (sulfide)

tungstates

1636

Tungsten, Cobalt, Molybdenum

vanadates

1637-1639

Calcium, Vanadium

Titanium (oxide)

1640s

K, Ti, Zr, Cd, Ga, In, Rh, Ir

Later

Li, Be, Ba, Nb, Ta, Ge

Li, Be

Next Part . . . Organic Chemistry!

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About Iver P. Cooper

Iver P. Cooper, an intellectual property law attorney, lives in Arlington, Virginia with his wife and two children. Two cats and a chinchilla rule the household with iron paws. Iver has received legal writing awards from the American Patent Law Association, the U.S. Trademark Association, and the American Society of Composers, Authors and Publishers, and is the sole author of Biotechnology and the Law, now in its twenty-something edition. He has frequently contributed both fiction and nonfiction to The Grantville Gazette.

 

When not writing (or trying to get an “orange blob” off his chair so he can start writing), he has been known to teach swing dancing and folk dancing, or to compete in local photo club competitions. Iver adds, “I can’t get my wife to read my fiction, but she has no trouble cashing the checks.”

Iver’s story “The Chase” is in Ring of Fire II