Grantville Gazette Volume 24

by Eric Flint

  The high school in Grantville is modeled on North Marion High School (Farmington, WV). It offers a surprisingly wide range of science courses. Grades 9 and 10 receive an integrated science course ("CATS") that is apparently a continuation of a program begun in Grade 7. Eleventh and twelfth graders can take Advanced Environmental Earth Science, Advanced Chemistry, Advanced Placement Chemistry, Advanced Placement Earth Science, Earth and Sky (a college level class), Microbiology and even Forensics ("topics include ballistics, fingerprinting, and the analysis of inorganic and organic compounds"). Lewis Bartolli's knowledge of forensic science (see my stories "Under the Tuscan Son," Grantville Gazette 9 and "Arsenic and Old Italians," Grantville Gazette 22) is based on more than just reading detective stories!

  What we need most is information on descriptive inorganic chemistry, and this subject tends to get short shrift in modern general chemistry and inorganic chemistry courses. Fairmont State presently uses the fourth edition of Brady, Chemistry: The Study of Matter, and I think there is a good chance of finding the third edition (1988) in Grantville. As for more advanced texts, I am sure that there is at least a copy of Cotton and Wilkinson, Advanced Inorganic Chemistry (CW); I used the third edition at MIT. (A JCE review of the sixth edition called it "the most popular inorganic chemistry textbook ever published"). I was pleasantly surprised to discover that the high school has the McGraw-Hill Encyclopedia of Science amp; Technology (4th ed., 1977; 15 vols.).

  As for equipment, as I said in my aluminum article (Gazette 8), the power plant has a "Metallurgist XR," which is a portable X-ray fluorescence spectrophotometer specifically designed for alloy analysis. (Boyes) And, even more surprisingly, the high school has a $300,000 atomic absorption spectrophotometer given to them in October 1997 by LaFarge Corp.

  Prominent Alchemists

  I referred to "industrial alchemy" rather "industrial chemistry" as a gentle reminder that for every up-time chemist, there are hundreds of down-time alchemists.

  We can expect visits (and perhaps citizenship applications) from the prominent alchemists of early seventeenth century Europe.

  Michael Sendivogius (1566-1636) did pioneering research on the composition of air, discovering that it was a mixture of substances, including one (now called oxygen) that supports life. His patrons are the Polish Vasas. Of course, they are more interested in his claim to be able to transmute mercury into gold.

  Cornelius Drebbel (1572-1633) (died in OTL shortly after the RoF, but this could be butterflied) is perhaps best known for his submarine, but he invented a thermostat and the dye known as "color Kufflerianus."

  Arthur Dee (1579-1651)(the physician to Michael I of Russia) wrote Fasciculus chemicus (1630), a compendium of alchemical bon mots.

  Jan Baptist van Helmont (1580-1644) was an early contributor to the development of the law of conservation of mass. He appears as a character in Mackey, "Ounces of Prevention" (Grantville Gazette 5).

  Johann Rudolf Glauber (1604-1670) was the first to produce hydrochloric acid and sodium sulfate. In OTL 1648 he developed a major method of manufacturing sulfuric acid. In NTL, he developed the potassium chlorate-based percussion caps for the French "Cardinal" rifles (1634: The Baltic War, Chapter 27).

  Several other notable alchemists were born before the Ring of Fire, but were young enough when it occurred that they may be "butterflied" into a different line of work: Elias Ashmole (1617-1684), Robert Boyle (1627-1691)(the "Father of Modern Chemistry"), George Starkey (1628-1665) and Hennig Brand (1630-1670).

  Commodity and Specialty Chemicals

  A commodity chemical is one that is produced in great quantity, whereas a specialty chemical has a more limited market.

  Judging from Posthumus' studies of commodity exchange prices in the Netherlands, the inorganic chemical commodities in 1630s Europe included the elements iron, tin, lead, gold, silver, copper, mercury and sulfur; the alloys steel, brass, and spelter (a zinc); the compounds common salt (sodium chloride), copperas (ferrous sulfate), potash (potassium carbonate), white potash (potassium chloride), soda (sodium carbonate), saltpeter (potassium nitrate), alum (potassium aluminum sulfate), and borax (sodium borate); and gunpowder (a mixture of sulfur, saltpeter and charcoal).

  Changes in Demand

  The arrival of Grantville will change the chemical marketplace. Some chemicals will be demanded because of their value as end-products, others, for use as starting materials or reagents.

  The principal chemicals in the first decade after the RoF will not necessarily be those that are prominent nowadays. In particular, those inorganic chemicals whose principal utility is in making organic chemicals may be disdained until the necessary organic raw materials are isolated in reasonable quantities.

  That said, it is worth using late-twentieth century compilations as a starting point. The top inorganic chemicals in the late-twentieth century are listed in Table 1-1. Most, if not all, of those compounds are going to be important in the first decade after RoF, too. (I am a bit uncertain about titanium dioxide, since titanium ores have never been mined by the down-timers.)

  Inorganic Chemicals in Canon

  The following inorganic chemicals are known to be in canon. The years given are those of their "canon appearance"; they may in fact have been made earlier (unless canon actually says "this is a first"). The chemicals marked with* were actually known to down-timers before RoF. Further details appear in later parts of this article.

  1631-32: sulfuric acid*, nitric acid*, sodium bicarbonate,

  1633: zinc sulfide* (sphalerite), natural cryolite (sodium aluminum fluoride), mercury fulminate, ammonia*, calcium hypochlorite, ammonium nitrate, [and by implication, chlorine, chlorosulfonic acid, hydrochloric acid*, calcium hydroxide]

  1634: sodium hydroxide*, chromium ore (chromite), potassium chlorate, boric acid, borax*, hydrogen, graphite*

  1635: calcium carbide

  1636: synthetic cryolite, hydrogen fluoride

  Infrastructure Problems

  A twentieth-century chemist can buy, off the shelf, pure chemicals, borosilicate laboratory glassware, and accurate measuring equipment (thermometers, pH meters, analytical balances, etc.) The life of the "industrial alchemist" is going to be more difficult.

  Mackey, "Ounces of Prevention" (Grantville Gazette 5) illustrates this by reference to the ring nitration step in chloramphenicol synthesis, which must be performed at near-freezing temperatures. Von Helmont complains he needs very pure sulfuric and nitric acids, and that the Essen Instrument Company has a six month backlog of orders for precision mercury thermometers.

  There are further requirements for industrial-scale production. You need stainless steel, rather than glass, to handle large quantities of reactants, especially if they are being handled under high temperatures and pressures. (Flint, 1633, Chapter 26). Chromium is a key ingredient in stainless steel, and that is why Josh Modi goes to Paris in August 1633… to persuade Richelieu to permit an expedition to Maryland to mine chromite. (Modi's patron, De Geer, had apparently learned of that part of the terms of the then-secret Treaty of Ostend). Mackey, "The Essen Chronicles, Part Three: A Trip to Paris," Grantville Gazette 9.

  The reactions which I expect will be the most difficult to duplicate early in the new time line are those which require special conditions (high or low pressure, unusual catalysts, or even high or low temperatures) or which are very finicky in their requirements for pure solvents and reagents. Unfortunately, modern industrial chemistry, especially organic chemistry, places great reliance on exotic catalysts.

  Qualitative Analysis

  Qualitative analysis answers the question, "Is it present?" There is a reasonable chance that at least one up-time chemist took a qualitative analysis course and has the textbook for it. If so, then it will be possible to determine the presence or absence of many common ions (electrically charged chemicals). Even without it, there is quite a bit of useful information in the encyclopedias, general chemistry textbooks, and the CRC.

  Dry Analysis. In th
e flame test, the sample solution is dried on a wooden splint, or a platinum or nichrome wire, and waved through an "invisible" flame. The heat excites electrons in metal ions. The electrons eventually release energy, and for some ions, this happens in steps which correspond to one of the colors of visible light. For example, sodium is blue; boron is green, and calcium is red. Note that different ions can produce the same flame color, so this test is far from definitive.

  In the borax bead test, a bead of borax, held on a platinum wire, is dipped in the sample, and then heated in the lower, reduction zone of the flame, and allowed to cool. You then heat it in the upper, oxidation zone, and let it cool. You observe its colors, hot and cold, and oxidized and reduced. The combination is indicative of which metal is present.

  The sample may also be placed on a piece of charcoal, and a blowpipe used to control the flame.

  Wet Analysis. The principal qualitative analysis methods exploit differences in reactivity and solubility. The method described in EB11/Chemistry) divides the metals into six groups; further reactions are needed to identify a particular ion within a group.

  See also the EB11 entries for the tests specific to individual elements. Maria Vorst alludes to the cobalt nitrate test for aluminum in Cooper, "Stretching Out, Part 3: Maria's Mission" (Grantville Gazette 14), and Lewis Bartolli to the turmeric test for boric acid in Cooper, "Under the Tuscan Son" (Grantville Gazette 9).

  Quantitative Analysis

  Quantitative analysis answers the question "how much?" As might be expected, these techniques are more exacting than those of qualitative analysis.

  Gravimetric analysis involves converting all of the chemical of interest (and only that chemical) to a precipitate and then weighing it.

  Volumetric analysis requires adding, drop by drop ("titration"), a known volume of a standard solution of an analytical reagent that reacts with (and only with), the chemical of interest, until a "signal" evidences that all of the target chemical has reacted. The "signal" can be a color change achieved by an "indicator" chemical, or a change in the electrical characteristics of the solution.

  The concentration of a compound in pure solution can be determined by measuring the degree to which it rotates the plane of polarization of linearly polarized light of a particular wavelength passing through the solution. You need to know the specific rotation of the compound (how much it rotates the plane over a unit path length) and the path length through the solution.

  Spectroscopic analysis involves causing the chemical to emit or absorb light of various wavelengths (visible, infrared or ultraviolet) and measuring the emission or absorption.

  Polarography requires measuring the change in the current through an electrochemical cell (see Electrochemistry) containing the solution of interest, as the voltage is varied.

  Natural Sources of Inorganic Chemicals

  Why synthesize a compound if you can isolate it from nature? Many useful compounds occur as minerals in rocks. Minerals are mostly ionic compounds (made of positively and negatively charged ions), and are often classified on the basis of the component anion. The most common classes of minerals are, in descending order of abundance: silicates carbonates/nitrates/borates sulfates/chromates halides oxides/hydroxides sulfides phosphates/arsenates/vanadates/antimonates/molybdates/tungstates native elements organic minerals.

  ("Minerals," Wikipedia).

  Oxides and hydroxides can be found pretty much anywhere, in rocks which would have been exposed to weathering. Silicates are also widely distributed. Sulfides are usually found in volcanic regions, in so-called hydrothermal deposits. Halides, carbonates, sulfates, nitrates and borates are more likely to be in desert regions, as they are formed in water and precipitated as the water evaporates. Phosphates are derived from the skeletons of marine life, and thus are found in former seabeds.

  Other important sources of inorganic compounds are seawater, subterranean brines, natural gas (the main source of helium), air and plants.

  Chemical Reactions 101

  There are only so many chemicals which can be found in nature; the rest must be synthesized. The ideal process is the one-step reaction. However, it may be desirable to take a more circuitous path in order to use a more available, cheaper or less dangerous starting material, or to produce a byproduct which is easier to dispose of or even salable in its own right. Other considerations are minimizing the need for special equipment (e.g., high pressure reactors), reducing energy requirements, and increasing production rate.

  A reaction may seem good on its face but be impractical because the reactants are too expensive to obtain. For example, aluminum will react with iron oxide to produce aluminum oxide and pure iron, but the cost of the aluminum is greater than the value of the iron. (Kotz

  934).

  Planning a chemical synthesis requires thinking about the chemical formula of the product and choosing reactants which provide the necessary building blocks by one or more of the basic forms of reaction. Stoichiometry allows us to express the reaction in quantitative form. Le Chatelier's Principle is used to qualitatively predict the effect of a change in concentration, pressure or temperature on the equilibrium state (ultimate degree of completion) of a reaction. Equilibrium constants, electromotive potentials and Gibbs free energy data are used to make more quantitative predictions as to the completeness of a reaction.

  Basic Forms of Reactions. Combination reactions (A+B-›AB)are most often used to unite elements to make binary compounds (those with just two elements), especially oxides, hydrides, sulfides, nitrides, phosphides and halides. This tends to be most practical when the elements can be cheaply obtained.

  Combination reactions are also used to convert oxides to carbonates (by adding carbon dioxide), nitrates (by adding nitrogen oxide), and sulfates (by adding sulfur oxide), or to hydrate (add water) to a compound.

  The simplest and most important decomposition reaction (AB-›A+B) is electrolysis, in which a compound made of several ions is dissociated into its component ions. The various combination reactions can also be reversed.

  Double displacement reactions (AB+CD -› AD + CB) occur between ionic compounds, but are only useful if the reaction is driven forward by the "disappearance" of one of the products; see Le Chatelier's Principle, below.

  A redox reaction is one in which one atom or group gains electrons (reduction) and another loses electrons (oxidation). There are many inorganic compounds which comprise a positively charged metal ion. If the metal ion is reduced to the point that it is electrically neutral, then you have obtained the elemental metal. This is the one of the goals in metallurgy.

  If any of the reactants or products in a combination, decomposition, or single replacement (AB + C -› AC + B, or -› CB + A) reaction is an element then the reaction is a redox reaction. A double replacement reaction is a redox reaction if any of the atoms changes its oxidation state (e.g., iron from +2 to +1).

  Tables of reduction potentials can be used to predict whether a particular redox reaction will occur spontaneously, or needs to be driven by an applied voltage (see "Electrochemistry").

  The most important single replacement reactions are those in which one of the reactants is a free metal or a halogen molecule. The more reactive metal displaces the less reactive one (e.g., copper + silver nitrate -› copper nitrate + silver), the more reactive halogen displaces the less reactive one (e.g., bromine + potassium iodide -› potassium bromide + iodine). The goal may be to make the new salt, to reduce the less reactive metal to elemental form, or both.

  We can determine which metal or halogen is more reactive by inspecting a table of reduction potentials; the list of metals, from most active to least, is called the electromotive series.

  Stoichiometry. Knowing the chemical formulae of the reactants and products, we can "balance" the equation of a chemical reaction, e.g., know that "x" molecules of compound 1 (#1) react with "y" molecules of #2 to make "m" molecules of #3 and "n" molecules of #4. And that in turn means we don't have to guess how mu
ch of compound #1 to add in order to fully react it with #2. And likewise we can calculate the theoretical yield of #3 and #4, given the amounts of #1 and #2 provided.

  Le Chatelier's Principle. If a chemical system is in equilibrium, and a variable (pressure, temperature, concentration of reactant or product) is changed, the equilibrium shifts to resist the change. This has a number of interesting implications:

  1) if the chemical reaction is chosen so that one of the products is

  – insoluble, and thus precipitated out of the solution,

  – a gas, and so escapes the solution then the reaction will be driven forward as the system shifts to try to replace the "lost" products.

  2) In a reaction of ionic compounds, if one of the products (ion combinations) is a compound which is itself a poor electrolyte (a compound which only minimally dissociates into ions, such as water), then its component ions are "depleted" which drives the reaction forward.

  3) the chemist can shift the equilibrium of the reaction forward (toward the products)

  – by adding one of the reactants in excess.

  – if any of the reactants or products are gases (e.g., hydrogen, oxygen, carbon dioxide, ammonia), and there are more molecules of gas on one side of the reaction than the other, the equilibrium can be shifted in one direction or another by a suitable change in pressure (see Pressure Control, below).

  – by a suitable change in temperature (see Temperature Control, below) by "coupling" it to a second reaction-a starting material of which is a product of the first reaction-so the second reaction helps pull the first one forward.

  Chemical Equilibrium. Many chemical reactions are reversible, that is, they can proceed in either the forward or reverse directions. If the forward and reverse reaction rates are equal, an equilibrium can occur, in which the reaction is incomplete, but there is no further propensity toward change in the concentrations of the reactants and the products. The equilibrium relationship can be expressed quantitatively as a concentration-dependent ratio which equals an equilibrium constant. (The equilibrium constant is also dependent on temperature and sometimes also on pressure.) Once the equilibrium constant is determined for one set of concentrations of the particular reactants and products, the equilibrium formula can be used to calculate the changes in the concentration of the product if the concentrations of the reactants is changed.