In alchemical thought, the Philosopher’s Stone is a fantastical artifact which is capable of transmuting base metals into gold. The new Philosopher’s Stone is not an artifact, but knowledge—the teachings of twentieth century chemistry as transmitted by the up-timers and their books—and while it can’t change one element into another, it can and will change how the down-timers think about the world they live in.
About three thousand up-timers were thrown into the seventeenth century by the Ring of Fire. Of those, perhaps a score have significant college training in chemistry, and of course there are many more who have recently taken a high school chemistry course.
That said, there are thousands of chemicals which we would like to make. The knowledgeable up-timers can’t do it all themselves. It is essential that they train new chemists from the vast population of down-timers.
Some of those trainees will be youngsters, and others will be experienced alchemists. The down-time alchemists have a lot of practical knowledge which is still of value. They are familiar with the gross chemical and physical properties of many substances, although the purity of the substances in question is debatable. They have carried out some of the basic manipulations of the chemical laboratory, such as melting, dissolving, crystallizing, filtering and distilling chemicals.
The down-time alchemists are going to be getting a crash course in modern chemistry. Some of the alchemists will become wholesale converts to modern chemistry. Others will treat it more as the Aristotelians did the Copernican cosmology; as a convenient fiction.
Modern science, including chemistry, will also be seeping into the general curriculum. Perhaps some of the students will aspire to become chemists. (The man who is sometimes called the Father of Chemistry—Robert Boyle—was four years old when Grantville was hurled into 1631 Thuringia.)
Chemical Resources in Grantville
In Lord Kalvan of Otherwhen, Corporal Calvin Morrison becomes the eponymous Lord Kalvan because he happens to know the recipe for gunpowder, a combustible mixture of charcoal, saltpeter and sulfur. This is proof that every time traveler should know some chemistry!
The time travelers of the 1632verse know quite a bit of chemistry, actually. There are six up-timers with a bachelor’s degree in chemistry: Allan Dailey (b. 1964), Greg Ferrara (1970), Thomas “Tom Stoner” Stone (1950s)(also has M.A. Pharmacy and doctoral course work), Alexandra (Lilburn) Selluci (1943), Walter Miller (1927-1636), and Dominic Genucci (1977 graduate course work). It is a safe bet that they have kept their chemistry textbooks from college. Each probably also has an edition of the “CRC” and perhaps additional chemistry books.
Christie (Kemp) Penzey has a degree in geology. She teaches chemistry, and is the “technical adviser” for the Kubiak experiments on recreating baking powder (Offord, “The Doctor Gribbleflotz Chronicles, Part 1:Calling Dr. Phil”, Grantville Gazette 10).
Nine more up-timers have degrees in pharmacy, and Jerry Trainer’s degree is in chemical engineering.
Several more up-timers do not have a college degree, but are getting advanced training in chemistry: Amy Kubiak, Tonya Daoud, Tyler Beckworth, Sam Reed, Mark Dalton Higgins, Lewis Philip Bartolli, Mary Lou (Cantrell) Snell and Kerry Burdette Douglas are laboratory technicians.
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 & 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.
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
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.
In reference to fulminate production, Christie Penzey tells Mike, “we can’t just call up our friendly chemicals supplier and ask for a few hundred gallons of pure nitric acid. We have to triple distill everything, even the water we use.” Offord and Boatright, “The Dr. Gribbleflotz Chronicles, Part 2: Dr. Phil’s Amazing Essence Of Fire Tablets” (Grantville Gazette 7)
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 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 the 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 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:
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 much 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.
Thermodynamics/Gibbs Free Energy. There are reference books in Grantville (e.g., the CRC Handbook of Chemistry and Physics) which have tables of thermodynamic values for various elements, cations, anions and solids. You can use these tables to predict whether a reaction involving those entities can occur spontaneously.
Rate. Loosely speaking, the equilibrium is the endpoint of a chemical reaction, and rate is how quickly it gets there. For a reaction to be commercially feasible, it must not only have an equilibrium favoring the products, it must have a high enough reaction rate. Unfortunately, the prediction of reaction rate is difficult and at the very least requires a knowledge of the exact reaction mechanism. Reaction rates increase with concentration (more chance for the reactants to collide) and temperature. Reactions of ions in solution tend to be fast. Other reactions are slower, as some (but not all) of the bonds holding the reactants together will need to be broken.
Planning. In general, synthetic strategies depend on either displacing one metal with another which is higher in the electromotive series, or on causing two soluble salts to react to form an insoluble product, a gas, or water. (See appendix table 1-2.)
Electrochemistry studies the use of spontaneous chemical reactions to create an electric current (as in a battery) or the use of an applied electrical voltage to force a chemical reaction to occur (as in an electrolytic cell).
If the electromotive potential of a reaction is less than zero, then the reaction won’t occur spontaneously. But you can still make it happen by applying electricity. The voltage has to be high enough to counteract the negative potential of the reaction, and the current will determine how much product is produced. The reaction will not be 100% efficient, so you will have to use more current than what is theoretically required.
An electrolytic cell has an electrolyte and two electrodes (cathode and anode). The electrolyte may be a solution or a molten salt; the key point is that it contains mobile ions. An ion is an atom or molecule which has lost one or more electrons giving it a positive charge (cation), or gained one or more electrons, yielding a negative charge (anion). The voltage drives the movement of cations toward the cathode, where they are reduced, and of anions toward the anode, where they are oxidized.
At the anode and cathode, the products may undergo further reaction to form secondary products. In a two compartment diaphragm or membrane cell, some kind of barrier prevents undesired reactions between anode and cathode species. For example, in the chloralkali process, hydroxide ions are allowed to react with sodium ions in the cathode compartment (making caustic soda), but not with chloride ions in the anode compartment. And recombination of sodium and chloride ions is also inhibited.
In 1633, Dr. Phil built a “wet cell” battery with a dilute sulfuric acid electrolyte and a zinc electrode. Offord, “Dr. Phil Zinkens a Bundle” (Grantville Gazette 7). That story doesn’t reveal the identity of the second electrode, but it would probably be copper, see Boatright, “So You Want to Do Telecommunications in 1633?” (Grantville Gazette 2).
Here, we are more concerned with electrolysis, which is the decomposition of a chemical by electricity. Dr. Gribbleflotz experimented with electrolysis of an unspecified salt in Offord and Boatright, “The Dr. Gribbleflotz Chronicles, Part 2: Dr. Phil’s Amazing Essence Of Fire Tablets” (Grantville Gazette 7)
In the old time line, water was decomposed into hydrogen and oxygen in 1800; sodium and potassium were isolated by electrolysis of their salts in 1807.
The first electrochemical reaction of industrial importance was probably in the purification of platinum. In 1991, the principal electrochemical products were caustic soda, chlorine, aluminum, copper, zinc, chromium, sodium chlorate, caustic potash, magnesium, sodium, manganese dioxide, permanganates, manganese, perchlorates, and titanium. (KirkOthmer9:125). The most common electrolyte was probably sodium chloride.
Electricity is supplied by power plants as high voltage alternating current, but for electrochemical use, this needs to be rectified into direct current and stepped down by transformers to a lower voltage.
What appears to be a single reaction may occur through a series of steps (addition, elimination, substitution and rearrangement), each with its own molecularity (the number of reacting molecules) and own rate law (a mathematical relationship between the rate of the reaction step and the concentration of the reactants). The slowest step determines the rate of the overall reaction.
Catalysts increase (or decrease, so-called negative catalysts) the rate of a chemical reaction without participating in the net reaction. They have no effect on the equilibrium concentrations of the reactants and products.
Johann Dobereiner discovered that the rate of the conversion of alcohol to acetic acid (1816) or acetic aldehyde (1832) could be increased by conducting the reaction in the presence of platinum wire. He created (1823) a lighter in which the hydrogen flame was produced by the action of sulfuric acid on zinc, in the vicinity of a platinum sponge (EA “Dobereiner”; Jentoft). In 1817, Humphrey Davy studied the effect of wires of different metals on the rate of reaction of coal-gas with oxygen. The term “catalysis” was coined by Jons Jakob Berzelius, who used it to explain additional phenomena, including the rapid decomposition of hydrogen peroxide by metals.
EA “Catalyst” says that “many common catalysts are powders of metals or of metallic compounds,” and by way of example mentions that platinum catalyzes the hydrogenation of double bonds. It also indicates that acids can be catalysts; “sulfuric acid catalyzes the isomerization of hydrocarbons.”
EA “Platinum” says that for use as a catalyst, platinum is used in powdery (“platinum black”, from reduction of platinum chloride) or spongy form, and there is reference to its use in production of nitric acid.
Further “data mining” EA will identify other catalysts, which I have tried to logically group below:
metals: palladium, neodymium, samarium, rhenium, lutetium, ruthenium, molybdenum, silver, mercury, nickel, iron, rhodium, a platinum-rhodium alloy (for preparation of hydrocyanic acid from ammonia, methane and air, or preparation of nitric acid or ammonium nitrate), copper, unidentified transition metals,
metal oxides: iron oxide (to catalyze the direct combination of nitrogen and hydrogen in the Haber Process, EA “Ammonia”), manganese dioxide (to speed the thermal decomposition of potassium chlorate to produce oxygen, EA “Chemical Reactions”), platinum dioxide (from fusion of chloroplatinic acid with sodium nitrate), copper oxides, chromium zinc oxide (used in methanol production), scandium oxide,cadmium oxide, lead oxide (litharge),
acids: hydrobromic acid, chromic acid, hydrogen fluoride, hydrochloric acid (for nitrobenzene),
miscellaneous: copper acetate, aluminum chloride, certain organotin compounds, nickel-aluminum sulfide, sodium nitrate (for manufacture of sulfuric acid), sodium ethylate, peroxides, hot alcoholic solution of potassium cyanide, lithium acetate, n-butyllithium, coordination compounds of zirconium, phosphorus pentaflouride, water (!).
EA apparently overlooks the organometallic catalysts, which were rather important in the late twentieth century.
It is important to note that many catalysts are reaction-specific. Hence, there is going to be a lot of educated trial-and-error going on; systematically testing the effect of each of a series of potential catalysts to see if any of them facilitate a reaction of interest.
A good example of this is the screening carried out by Bosch to make the Haber nitrogen fixation process feasible commercially. Haber initially identified osmium and uranium, both of which were quite expensive, as effective catalysts. Bosch set up test reactors, and tested 4,000 different catalysts over five years, finding that an impure iron oxide catalyst was cheap and operable. (McGrayne 66; KirkOthmer5:323).
Just to complicate matters further, modern catalysts aren’t necessarily simple materials. Because the catalytic material is expensive, it is usually advantageous to use it in small amounts, and disperse it on a support material with a high surface area. Gamma-alumina is the most popular support. (KirkOthmer 5:347).
There are also catalytic promoters. These are substances which don’t act as catalysts themselves, but which potentiate the activity of the “real” catalyst. There are both chemical promoters which change the surface chemistry, and textural promoters which alter the physical characteristics. Alkali metals have been used as chemical promoters.
Catalysts can be deactivated as a result of fouling (they are physically masked by deposited material), poisoning (feed impurities which reduce their catalytic activity), and physical change (e.g., sintering). Catalysts may in turn be regenerated.
The modern catalyst for ammonia synthesis is a combination of iron oxide as the catalyst, aluminum and calcium oxide as textural promoters, and potassium as a chemical promoter.
Some catalysts—common acids, finely divided metals (e.g. platinum), and some metal oxides—can be put to work in the 1632verse in fairly short order. Others are rare materials, or of a complex composition or structure, and it will take years, if not decades, of work to duplicate them.
Temperature affects both the rate and the completeness of a reaction. A typical rule of thumb is that for every 10ÌŠC increase in temperature, the reaction rate will double. The effect of the temperature on the completeness of a reaction depends on whether it is endothermic (needs heat) or exothermic (releases heat). Higher temperatures favor endothermic reactions and hinder exothermic ones.
There are other considerations. Too high a temperature can result in side reactions, including decomposition. So, depending on the reaction, you may want to heat things up, keep the temperature from increasing above a certain point, or bring it below room temperature.
If a reaction is temperature sensitive, then you need a good thermometer. For industrial work, you might prefer a thermostat which controls a heating or cooling device. In 1634, the Essen Instrument Company is manufacturing precision mercury thermometers. (Mackey, “Ounces of Prevention,” Grantville Gazette 5). I would expect that simple spirit thermometers are being made, too.
Both heating and cooling processes are slower to start, and stop, when the reaction is on an industrial scale. As the volume increases, the ratio of the heating or cooling surface to the volume decreases.
In the laboratory, if an elevated temperature is needed for a reaction, the chemist will use a gas-burning Bunsen burner. This can reach a temperature close to 900ÌŠC. Up-time, natural gas is used, but Dr. Phil has an alcohol burner in 1633. Offord and Boatright, “Dr. Phil’s Amazing Essence of Fire Tablets,” Grantville Gazette 7).
On the industrial scale, you may be burning some kind of fuel, which heats air or water surrounding the vessel, or passing through tubes in the vessel. Steam distillation falls in this category. Or you may be converting electrical energy into heat energy. Or running two industrial processes alongside each other, one providing heat for the other.
Chemical reactions tend to be more efficient when the reactants are all in the liquid phase. Solids react only at their surfaces, and gases are low in density. If one of the reactants is solid at room temperature, then to put it in the liquid phase, it must be dissolved or melted. And melting requires heat.
In some cases, it is possible to drastically lower the melting point of the substance of interest by adding a second substance, known as a “flux”. Sodium, potassium and lead oxides lower the melting point of glass from 1700° C. to perhaps 900-1200. Aluminum oxide melts at 2054° C., but it can be dissolved in cryolite, which is molten at a little less than 1000° C.
You may also be trying to lower the melting point of the waste material. For example, in smelting copper, you may want to make sure that the silica forms a very liquid slag, that the copper can sink through. So iron oxide is added.
Smelting metals typically requires a reducing agent (e.g. carbon) and heat. For tin or lead oxide, a campfire (600-650°C) is good enough, but copper requires a temperature of 700-800 and forgeable iron, 1100°C.
Combustion processes cannot exceed the “adiabatic combustion temperature,” which, for combustion in air, is about 2000° C for natural gas, 2150 for oil and 2200 for coal. The fuel is the source of carbon and the air is the source of oxygen. The limiting temperature is a function of the heating value of the fuel, the specific heat capacity of the fuel and the air (and the combustion products), the ratio of fuel to air, and the air and fuel inlet temperatures (Wikipedia, “Combustion”). Even higher temperatures are achievable with rocket engine fuels/oxidizers.
The practical combustion temperatures for industrial chemistry are much lower than the theoretic limit. It is difficult to achieve complete combustion if there is insufficient air, heat is lost (radiated out; carried away by exhaust gases), and so forth. To ensure complete combustion, it is customary to use an excess of air, but air dilution then reduces the temperature of combustion.
In 1920, a coal furnace could achieve a temperature of 1600°C without a blast, and 1800°C with one. A gas-fired furnace, with hot air, both the gas and air under pressure, could reach about 2000°C. (Marsh, 46). For higher temperatures, you need to heat by means other than combustion.
An electric arc furnace uses an electric current to heat a conductive material. That could be an ionic compound, or a conductive metal. Perhaps the first industrial use of the electric arc furnace was in the production of calcium carbide by heating lime and coke to 2000°C (1888). Electric arc furnaces came to play an important role in small-scale steelmaking.
Another option for sidestepping the practical combustion temperature limit is to use a solar furnace. Temperatures of 3000°C have been achieved by focusing solar radiation.
The higher the temperature our technology will generate, the more options we have for chemical synthesis.
To chill things down, you can put the vessel in ice, an alcohol bath, dry ice (solid carbon dioxide), or in liquid nitrogen. (for availability of CO2 and nitrogen, see part 2, and Huston, “Refrigeration and the 1632 World” (Grantville Gazette 17)).
Some reactions cannot be conducted in the air, because it would react. If so, the air is replaced with an inert gas, like nitrogen or argon.
Or you may need an atmosphere whose pressure is higher or lower than normal. It is important to compare the number of gas molecules at the beginning and end of the reaction. If that number decreases (as in ammonia synthesis), increasing the pressure will cause the reaction to shift (per Le Chatelier’s Principle) in favor of reducing the pressure, which means in favor of fewer gas molecules, and thus in the forward direction. On the other hand, if the number of gas molecules is increased by the forward reaction, then you want to conduct the reaction under lower-than-normal pressure.
To change the pressure, you need two things: a pump, and a vessel with walls strong enough to withstand the pressures generated.
Vacuums may be needed to pull out a gaseous product (to drive a chemical reaction), or to lower the boiling points of the compounds in an organic residue (as in vacuum distillation). Vacuum pumps have been scavenged from refrigerators. (Gorg Huff, “Other People’s Money,” Grantville Gazette 3)
Elevated pressure also may be used to keep the reactants in the liquid phase, or to facilitate a gas phase reaction. In the mid-nineteenth century, autoclaves were built which could achieve pressures of 725-1150 psi (14.7 psi is normal atmospheric pressure). A 1901 ammonia synthesis used a 1450 psi autoclave. In the early twentieth century, large-scale continuous feed reactors had been built which could handle 2000-5000 psi. By the 1990s, there were operations using 51,000 psi. (Kirk-Othmer/”High Pressure Technology”).
High pressure vessels are typically thick -walled, and composed of gun steels. During the 50s, the preferred alloy was nickel-chromium-molybdenum, and later an alloy which additionally contained vanadium gained favor.
The down-timers’ only experience with “pressure vessels” is of a rather specialized nature: cannon barrels. These have to resist the internal pressures generated by the explosion. For a given thickness, bronze is better than cast iron, and the down-timers are familiar with the concept of the “built-up” cannon, in which hot hoops or jackets are fit over the barrel and allowed to cool and shrink.
In 1773-91, Woolwich conducted experiments on muskets, reporting a maximum internal pressure of 2,000 atmospheres. (Ingalls). A Civil War era 15-inch Rodman gun, charged with 130 pounds of black powder, will experience 25,000 psi (1700 atmospheres) pressure. (NPS).
While explosives are not exactly a preferred source of pressure (they’re dangerous, and don’t lend themselves to continuous processing), Alfred Noble “packed steel tubes with gunopowder or cordite and heated them until they exploded with tremendous force, briefly attaining pressures of 8,000 atmospheres at more than 5,000°C.” (Hazen 35).
The up-timers include some steam engine enthusiasts, and a locomotive boiler can be considered a high pressure vessel suitable for continuous processing. Canon is a little vague on the issue, but it appears that there is at least one true locomotive on the main line by September 1633 (Flint, 1633, Chapter 33). That locomotive, of course, is generating high pressure steam. I suspect, based on the nineteenth-century locomotive data which the designers will be studying, that it has a steam pressure in the 75-200 psi range. That’s still short of even a nineteenth-century autoclave, but it’s a start.
To some extent, it will be possible to compensate for having weaker alloys by increasing the thickness of the vessel wall. However, that increases the expense of the vessel and, if it’s externally heated or cooled, it impairs heat transfer. In addition, increasing vessel thickness doesn’t address the Achilles’ heel(s) of the system: the openings needed in order to add raw materials, withdraw product and perhaps supply or remove heat.
Solvents are used as a medium in which the reactants can find each other, as catalysts (to help the reactants make or break bonds), and to control the temperature of the reaction. The traditional solvent for inorganic chemical reactions is water.
If cold water doesn’t dissolve a particular salt, you can try hot water, and, if that fails, a dilute or concentrated solution of an acid (hydrochloric, sulfuric, nitric, hydrofluoric, acetic, etc.). If need be, the inorganic chemist may have recourse to pure acids, carbon disulfide, liquid ammonia, liquid sulfur dioxide, alcohol, benzene, chloroform, acetone, ether, and turpentine. CRC provides detailed information on the solubility of inorganic compounds in various solvents.
The choice of solvent can have interesting consequences. Barium chloride is soluble in water, while silver chloride is not. The reverse is true in liquid ammonia. Hence, in water, barium chloride reacts with silver nitrate to form silver chloride and barium nitrate. The reverse reaction is favored in liquid ammonia. (Purcell, 154).
Sometimes, not only do you not want to use water as a solvent, you need to make sure that there isn’t even a trace of water present in the reactor. If so, you will use various dehydrating agents to prepare the reactor and the reactants for use.
While water was the most important solvent in inorganic chemistry, it has a lesser role in organic chemistry. Over twenty different organic compounds are used as solvents, including methanol, ethanol, acetone, acetic anhydride, pyridine, chloroform, diethyl ether, and benzene (Bordwell 201). In winter 1633-34, Henri Beaubriand-Lévesque uses turpentine and ether as solvents for natural rubber. (Offord, “Letters from France,” Grantville Gazette 12).
The “aprotic solvents” (e.g., dimethyl sulfoxide) are especially interesting because they seem to increase the reactivity of the reagents (M&B 492). DMSO can be obtained from the lignin of wood (EA/Dimethyl Sulfoxide).
The internal surface of chemical reaction vessels and piping is potentially subject to corrosion (deterioration as a result of gradual chemical attack by the industrial chemicals which are being handled). If the equipment is outdoors, then one must also worry about corrosion of the outer surface by the environment (including air, water, and microorganisms).
In the modern chemical industry, direct corrosion costs (upgrading materials to inhibit corrosion, inspection for corrosion, cleaning/replacement of corroded elements) are about 8% of total capital expenditures. (Brongers). The most commonly encountered corrosive agents are strong acids and bases, strong oxidizing agents, dehydrating agents, fluorine and chlorine, certain salts, and some organic chemicals (including phenol), and in the long-term even water (especially salt water and steam) and contaminated air can be problematic.
The single most important thing to remember is that it’s very dangerous to generalize about corrosion resistance. You absolutely must run a model to test the containment system against the expected threat under the exact conditions you expect to encounter in production. And you must monitor what actually happens once production begins.
There are four basic methods of controlling corrosion: using a corrosion-resistant structural material, applying a corrosion-resistant coating or lining to the structural material, adding a soluble corrosion inhibitor to the liquid contents, and, for metals, electrical (“cathodic” or “anodic”) protection.
Metals are the dominant structural materials, because of their strength. Corrosion of metals is usually an electrochemical reaction in which the metal is oxidized, so metals at the bottom of the electromotive series (e.g., gold) are typically more resistant. Certain other metals, notably aluminum and lead, are resistant because the initial contact results in passivation (the formation of a resistant metal oxide layer on the surface). Impurities (notably chloride, fluoride and sulfide ions) can interfere with passivation.
Materials, including non-metals, can also be corroded by non-electrochemical mechanisms, such as oxidation by an oxidizing agent (oxygen, sulfur, chlorine, etc.), or leaching and hydrolytic dissolution of glass.
Non-metals (glasses, rubbers, plastics, ceramics, concrete, graphite, etc.) generally have good corrosion resistance, but they are relatively weak and brittle. Natural rubber is vulnerable to many organic solvents; most synthetic rubbers are better but each has its own weaknesses. (Fontana 260). Plastics are usually more resistant to HCl than metals, but less resistant to sulfuric and nitric acid. Fluorocarbons are the “noble plastics”. (265). Both rubbers and plastics have temperature limitations. Ceramics and glasses are vulnerable to hydrofluoric acid and its salts, and strong alkalis. (274).
Unfortunately, the corrosion resistance of materials is very dependent on just what chemical is doing the attacking. Stainless steel is resistant to sulfuric and nitric acid, but is less suitable for hydrochloric or hydrofluoric acid service. Lead can handle sulfuric acid and cold hydrochloric acid, but is not the metal of choice for containment of nitric acid. Silver can handle boiling hydrofluoric acid, but not sulfuric or nitric acid.
Concentration is also important. For example, lead is readily attacked by both dilute and fuming nitric acid, but is quite resistant when the acid strength is 52-70% (Craig 542). Carbon steel can handle concentrations of sulfuric acid above 90% but is attacked by more dilute acid (Schweitzer, 423). Titanium resists boiling nitric acid below 25% and at 65-90%, but not at 25-65% (424).
A material which works fine at room temperature may be corroded if the temperature is elevated. However, increasing temperature can sometimes reduce corrosion, because it reduces dissolved oxygen. (Fontana 281). Other factors affecting corrosion rate include additional chemical species (including oxygen, water, and various ions), and flow rate (high flow rates usually increase corrosion, but stainless steel has better resistance under flowing than stagnant conditions).
In the 163x universe, we have limited choices when it comes to structural (high strength) materials available in commercial quantities (see table below).
A corrosion-resistant coating or lining can be a metal (lead, zinc, tin, and, eventually, also nickel, chromium, etc.) or metal alloy, or a nonmetal (graphite, quartz, fluorite, glass, ceramic, and eventually rubbers and plastics). The advantage of using a corrosion-resistant coating or lining is that this can be chosen without regard to the structural strength of the material. However, one must instead worry about how to apply the coating economically and uniformly, whether the coating is not too thin (less resistance) or thick (uneconomical, might affect heat transfer), and is properly distributed, and its susceptibility to failure by cracking, peeling, etc. Such failure can be the result of differential thermal expansion of the coating and the underlying substrate. There will of course be labor costs for lining the reactor, and for cleaning and occasionally replacing it.
A metal coating can be applied by electroplating, flame spraying, rolling, hot dipping, and vapor deosition. A glass coating can be formed by fusing powdered glass on the substrate (enamelling). Rubber linings are fastened on.
Soluble inhibitors act by forming a protecting coating on the substrate, retarding hydrogen evolution, or removing dissolved oxygen. Inhibitors are used mostly in closed systems, so they don’t need to be replenished frequently. Inhibitors are often substrate-specific and, in the wrong concentration, can make matters worse rather than better. There is limited identification of inhibitors in the general Grantville encyclopedias: silicates (EA/Detergent), hydrazine (EB; EA/hydrazine), triethanolamine (EB; EA/amines). The high school’s 1977 McGraw-Hill Encyclopedia of Science and Technology may say more; the modern version draws attention to chromates, nitrites and phosphates, and comments that 0.1% palmitic acid protects mild steel from nitric acid.
In the original form (1824) of cathodic protection, the structural metal is electrically connected to a sacrificial anode, made of a more active metal (zinc, magnesium, aluminum). An alternative form of cathodic protection is apply an externally-generated protective current to the metal.
Anodic protection is of much more limited applicability. An applied current causes a protective coating to form on the surface of certain metals. First demonstrated in 1954, it can be used to protect stainless steel from hot sulfuric acid. I am not sure that anyone in Grantville knows about it as its acceptance was very slow. Misapplied, it accelerates corrosion.
Stainless steel is one of the USE’s R&D targets. It’s worth noting that there isn’t just one “stainless steel”; it’s a generic name for high-chromium (11%+) steels. There are over 30 varieties.
Certain of the up-timers are adamant that stainless steel is critical for industrial development. In November 1632, Josh Modi says “the big holdup is the lack of stainless steel. It’s pretty critical if we want to produce large amounts of nitric acid. Not to mention antibiotics, DDT and sulfa drugs. Right now we are barely at the bucket stage.” Mackey, “The Essen Chronicles, Part Three: A Trip to Paris,” Grantville Gazette 9.
In September 1633, Ferrara explained to Gustav Adolf that an “insufficient supply” of stainless steel had created a “bottleneck” in manufacturing nitric acid, which in turn was needed to make guncotton. (1633, Chap. 28). Later, Ferrara complained that “while Grantville had quite a bit of stainless steel lying around in one form of another, almost all of it was in the form of thin sheet. And they were still a long ways off from being able to make stainless steel from scratch.” So they didn’t have a single heavy stainless-steel pressure tank. (Chap. 34).
I fear that Modi and Ferrara have failed to fully appreciate the lessons of history. Stainless steels weren’t successfully produced until after 1900, and production wasn’t significant for several decades later. That implies that all of the economic growth seen in the nineteenth century was accomplished without the assistance of stainless steel. World production of nitric acid in 1890 was 100,000 tons, and that nitric acid was produced in glass or cast iron retorts. (Thorp 137; Molinari 385, 402).
Glass-lined reactors are still used for some chemical production, including sulfuric acid (Gaverick 184) and pharmaceuticals. Glass-coated steel resists nitric acid up to 70% concentration and 125oC. (186). Modest deviation from normal pressure is possible; a Pfaulder Glassteel® RS-96 reactor can hold 4000 gallons at 125 psi or in full vacuum. The danger with use of glass linings is that they are susceptible to mechanical damage and thermal shock (Schweitzer 641). The glass lining option nonetheless may explain how come, in December 1633, the coal gas plant in Magdeburg has so much ammonium nitrate (made by reacting ammonia with nitric acid). (Flint, 1634: The Baltic War, Chapter 3)
Ultimately, we will recreate stainless steel; formulae are in the encyclopedias. Knowing that chromium imparts corrosion resistance, Josh Modi went to Paris in August 1633 . . . to persuade Richelieu to permit an expedition to Maryland to mine chromite. He was successful. However, it will take several years to assemble the expedition, find the chromite, ship it back to Europe, extract the chromium, and finally make stainless steel.
Since production of stainless steel won’t begin until 1635 or later, chances are that there will be plenty of fledgling industrial chemical plants using reactors made of cast iron, and lined with lead, graphite, glass, ceramic, and perhaps even plastic or rubber. They may have to change the linings frequently. They may have occasional catastrophic failures which could have been avoided with stainless. So be it.
The table below contains my suggestions as to which materials will initially be used to contain particular corrosive agents, and which more resistant materials commonly in modern use for the application need to be recreated before they can be used. I have not researched what would be known in Grantville about these compatibilities, but of course they can be determined empirically.
<70%: lead (1)
>70%: carbon steel
high-silicon cast iron
specialty stainless (Durimet 20)(for pumps and valves)
high-silicon cast iron
high-silicon iron with molybdenum
temperature-limited: most plastics and rubbers
>60%: carbon steel (low silicon)
cupronickel (inferior to Monel)
<60%: butyl rubber, neoprene
rubber (natural and synthetic)
nickel and alloys
Source: Fontana 317ff.
(1) preferably “chemical lead” (0.06% copper) or “hard lead” (4-15% antimony).
(2) Borosilicate glass is better than ordinary soda lime glass if there will be significant temperature changes but isn’t necessarily more corrosion resistant (Clark 35). Ceramics may provide service analogous to that of glass.
*expensive, used only when contamination must be minimized or when the conditions are unusually corrosive.
The weight, and hence the mass, of a chemical is measured in chemistry labs by using an equal-arm balance. This has two pans, one holding the unknown, the other a known weight. EA/Balance says that the key to precision measurements is to use a knife edge as a fulcrum, whereas EB11/Weighting Machines warns that the knife-edges and their bearings must be extremely hard. All else being equal, a long arm balance will be more sensitive than a short arm one. Precautions must be taken vis-a-vis temperature, humidity, vibration (from air currents or through the ground), and other disturbances.
In industry, where the weights involved are much greater, the measurement will probably be with an unequal arm balance (“steelyard”), a spring scale, or a platform scale with multiplying levers. (EA/Weighing Machines).
The volume of a liquid is measured by introducing it into a graduated cylinder of suitable size. The flow rate of a gas can also be measured (suitable examples exist at homes which buy natural gas for heating purposes).
Temperature is measured, of course, by a thermometer. The first thermometers were of the liquid-in-glass type; first water, then alcohol, and finally mercury. The liquid expands as the temperature rises. Sealing the tube was essential to avoiding pressure effects. Mercury is liquid from -39 to 357°C. To measure higher temperatures, a gas-in-tube thermometer can be used. Hydrogen thermometers are used up to 1100°C, and nitrogen to 1550°C.
There are many other principles on which a thermometer can be constructed. The platinum resistance thermometer (1886) has been used to measure temperatures in the -259 to 630°C range.
Gas pressures are measured with a pressure gauge designed to handle a suitable range of pressures. There are hydrostatic gauges (manometers) which observe the movement of a column of mercury in a U-shaped tube, flexible pressure sensors like the 1849 Bourdon tube (coiled tube which expands and causes arm to rotate) or the diaphragm gauge (membrane deforms under differential pressure), and thermal gauges which detect the change in heat conductivity of a gas. A primitive manometer was invented by Torricelli in 1643. In Grantville, we probably have diaphragm barometers in several homes, and the steam buffs have pressure gauges which can work up to probably ten or twenty times atmospheric pressure.
pH is a measure of the acidity or basicity of a solution. You can measure it quantitatively with a pH meter; which is really a voltmeter with a glass electrode sensitive to hydrogen ions. There isn’t any useful information about them in the encyclopedias, but it might be possible to reverse engineer them, and, if one of the chemists took a course in chemical analysis, they would be described there.
If a pH meter isn’t available, then you can estimate pH by using one or more acid-base indicators. Those are chemicals which change color depending on the pH. The oldest indicator, litmus paper, was known to the down-timers. EA/Indicator mentions that a mixture of methyl orange, methyl red, bromothymol blue and phenolphthalein will change color continuously from red to violet as the pH varies from 3 to 10. Several of these indicators are discussed in slightly more detail in EB11/Indicator.
The hazards posed by chemicals are fire, explosion, and irritation, burning or poisoning through inhalation of vapor, or skin or eye contact.
Borosilicate glass or stainless steel vessels, goggles, wash stations, specialized fire extinguishers, and fume hoods are all taken for granted in the late twentieth century laboratory but will be quite new to the down-timers.
For further details on hazard control in industrial processes, see Cooper, “Industrial Safety” (part 1 in Grantville Gazette 17 and part 2 in 18).
Many naturally occurring inorganic chemicals are found together with other chemicals, from which they must be separated.
Chemical processes may also yield a mixture of products. When the reaction is performed, you have to separate the product from whatever else is present. At the very least, there will be solvent. If the reaction didn’t go to completion, then there will be some starting materials still around. If your reactants and solvent weren’t pure, then you have to worry either about the original contaminants or what they might have been converted into.
Some reactions, by their very nature, create more than one product. For example, there are decomposition reactions, which break a large molecule into two or more smaller ones. And there are many reactions in which there is a “change of partners” (compound AB reacts with CD to form AC and BD, where A, B, C and D represent pieces of the reactants).
Separation of the mixture usually depends on the physical properties that differentiate the desired chemical from the others with which it is associated. But suppose that you need to separate X (desired) from Y (undesired), and you can’t do so directly. Well, there are tricks that depend on the different chemical reactivities of X and Y. You could chemically convert X to Z, separate Z from Y, then convert Z back to X. Or convert Y to Z, and separate X from Z or even convert X to Z and Y to W, separate Z and W, then convert Z back to X.
The more common separation processes, and the related physical properties, are:
Distillation/Boiling/Condensation: Boiling point (vapor pressure)
Recrystallization: Solubility of Pure versus Mixed Solutes
Decanting/Filtration: Solubility in a particular solvent, and particle size
Extraction: Difference in solubility between two immiscible liquids
Stripping: Difference in solubility in a liquid and in a gas
Magnetic Separation: Magnetism
The down-timers are familiar with simple boiling (distillation), and have occasionally used a simple form of fractional distillation, but know nothing of vacuum distillation. I will discuss the more advanced techniques in a forthcoming article on the organic chemical industry. Since the down-timers have only the vaguest concept of gases, they are unaware of the elements that can be collected by the liquefaction of air.
Recrystallization was used by Birringucio in the sixteenth century to purify leached saltpeter. (Bohm). In the simplest form of recrystallization, the crude material is dissolved in a minimum quantity of a single solvent, heated enough to bring it all into solution, and then allowed to cool. The principal component crystallizes out first, in a purer form. I am not sure that the down-timers know about multi-solvent crystallization. In any event, modern chemistry increases the number of solvents from which to choose. We are also now more aware of the importance of initiating the crystallization step by providing a seed crystal or creating a seeding surface.
The down-timers also know that some reactions form precipitates, which can then be separated from the remaining liquid by decanting the latter. And they filtered liquids through felt, paper, and porous stones. (Bolton). However, they only practiced gravity filtration, not vacuum filtration, and their filter materials can be improved upon.
The down-timers have prepared extracts, usually with water, of various plant tissues (and the aforementioned leaching is also a form of extraction). However, they haven’t really exploited extraction with organic solvents.
Since the down-timers don’t know of any gas other than air, and use air in chemical processes only as an oxidant, they aren’t aware of the use of a gas to selectively remove a chemical from a liquid.
Density separation by gravity has been used since antiquity. However, centrifugal separation didn’t begin until the nineteenth century (a centrifuge was first used separate cream from milk). A centrifuge artificially achieves a sedimenting force much greater than gravity, and hence can separate materials of different density much faster than gravity can.
The down-timers are barely aware of the existence of magnetism, and they lack powerful magnets. Hence, they haven’t performed magnetic separations, e.g., of ferrous from non-ferrous metals in recycling operations.
In the seventeenth century, there were chemical processes, like dyeing and tanning, which could be called industrial processes. Nonetheless, there was no industrial production of chemicals, with the arguable exception of refining ores to metals.
The first chemical compounds produced in reasonably pure form on a large scale were sulfuric acid (late eighteenth century) and soda ash (early nineteenth century). Hence, the down-time alchemists are not accustomed to operations on an industrial scale.
Nowadays, the scaling up of a chemical process is the work of the chemical engineer. In the nineteenth century, chemists teamed up with mechanical engineers. The emphasis of chemical engineers is on “unit processes”—for example, different types of separation.
There are a variety of process changes that must be made when scaling up from laboratory scale (batch size under a kilogram) to industrial scales (tons of material)(White, 117-18). The most obvious one is that the reaction vessels change from possibly glass-lined to metal, but there are others.
Process development is the redesign of a laboratory process to work on the industrial scale. This development work is done on a “pilot plant” scale, intermediate between the laboratory and industrial scales.
The raw material samples that are run through the pilot plant process are only those that are available, if accepted for production use, in commercial quantities. The idea is to avoid using raw materials that will require synthesis, or extensive purification.
Solvents are chosen, whenever possible, so that they don’t present severe fire, explosion or toxicity hazards, and so they are recoverable, in reasonable yield (e.g., at least 85%) for reuse.
Since recovery is incomplete, it is a good idea to find ways of minimizing the amount of solvent needed in the first place.
If expensive liquids are involved in the process, whether as solvents or reactants, mockup studies can be performed. That is, an inexpensive fluid with the right physical properties is used as a surrogate to test flow through the system. (Euzen 16).
Many physical processes are size sensitive because of surface/volume ratio considerations. Heating, cooling or filtering material may take minutes on the lab scale but hours on the industrial scale. Extraction of solute from one liquid to another is also on the slow side. The elongated time scale can cause a variety of problems.
There is a general preference for a short time cycle from beginning to end of the production process, but this can cause other problems. For example, a short time cycle may be achievable only if the temperature is allowed to rise rapidly. A temperature rise that is acceptable on the lab scale may result in a fire or explosion when large quantities are involved. The rate of addition of reactants may need to be reduced to compensate.
Significant byproducts of the reaction need to be identified. If you can obtain samples of these byproducts, you can add them to the product and see how the properties change. In this way, you can determine the tolerance limits to be enforced by quality control personnel on the industrial scale.
Many chemical reactions do not yield a single product, even in theory. Others would do so if the reactants were pure, but the required purity may not be obtainable in the early post-RoF period. Separation processes are chosen so that yield is high; crystallization, if necessary, is preferably the last step, because yields are 90% at best.
Ideally, the byproducts are useful in their own right, and recoverable for sale. For example, Spanish pyrites (iron disulfide) were not only used to make sulfuric acid, they usually contained 3-4% copper, which could be profitably extracted from the cinders. (EB11 “Sulphuric Acid”).
The good news is that there are economies of scale. Euzen (9) says, “the capital investment normally required for the transformation of the raw material into a given product varies by the power of 0.7 with the capacity of the unit.”
Batch versus continuous. In a batch process, the raw materials are loaded into the reactor, the reaction is carried out to completion, the products are removed, and the reactor is cleaned out, ready to repeat the cycle. In a continuous process, the reactor is (almost) never shut down. As product is pulled out, new raw material is added.
Continuous processes are typically very efficient; they are amenable to production of extremely large volumes at a very low operating cost. In part, that low operating cost is attributable to the relative ease with which a continuous process can be automated.
However, there are a few catches. First, continuous processes typically use equipment specially designed for the process in question. If the demand for the product drops, you have equipment which is going to waste. If there is an emergency demand for a different product, you need to set up a separate (batch) reactor to deal with it.
Second, continuous processes must be much more closely monitored. You need real time, or near real time, surveillance of the levels of all the raw materials and products so that, if you’re running a little low on one reactant, you can toss more in. And if the product mix isn’t correct, you can try to figure out why, and fix the problem.
Third, and this is related to the first two points, continuous process plants tend to have high start up costs.
Fourth, you are at the mercy of your suppliers (and the transportation infrastructure). If you run out of on one of the reactants because a delivery isn’t made, or because the material delivered isn’t up to spec, then you may have to shut down the process. Idle equipment “burns” money, it doesn’t make money. And with some continuous processes, it is difficult and expensive to “restart.” You can alleviate these problems by keeping a large reserve of the raw materials, but even when that is practical (some materials don’t store well) it is expensive.
This means that we aren’t going to see much in the way of continuous processing during the first decade after RoF.
In parts 2 and 3 we will analyze the prospects for the production of specific elements, molecules and compounds.
Table 1-1: Top Inorganic Chemicals
Sulfuric Acid and Derivatives
manufacture of sulfates, hydrochloric acid and phosphoric acid; acid catalyst,
rust removal, acidification of foods, phosphate (including fertilizer) manufacture, soft drinks
mordant, water purification, concrete additive
Calcium Oxide (Lime)*
steel and cement manufacture
Sodium Carbonate (Soda)*
glass flux; pH adjustment, electrolyte, water softener
Sodium Silicate (Water Glass)
cement, egg preservative, timber preservative, porosity-reducer in concrete, fire protection
ammonia production, petroleum recovery, perishables protection
desulfurization of steel; manufacture of etheylene oxide; welding, rocket fuel oxidizer, oxygen therapy
pressurized gas, fire control, welding, solvent (as liquid), refrigerant (as solid), reagent
Sodium Chloride Derivatives
production of chlorine, chloride, and sodium compounds
Sodium Hydroxide (Caustic Soda)*
strong base in soap, paper, detergent, synthetic fiber manufacture
disinfecting water, bleaching paper, production of vinyl chloride plastics and chlorinated organics
regeneration of ion exchangers, pickling steel, pH control, production of chlorides and chlorinated organics, including PVC
raw material for making nitric acid, ammonium sulfate, chloramine; refrigerant; fertilizer (as water solution); fuel
manufacture of nitrates; oxidizing agent
fertilizer, oxidizing agent (in explosives)
fertilizer, preparation of ammonium salts, protein precipitant
white pigment, photocatalyst
Potassium Carbonate (Potash)*
soap, glass production; drying agent; fire suppressant
pigment, tire filler
(Source: Chenier, Survey of Industrial Chemistry, Table 2.1. Uses from Wikipedia.)