Eric Flint: Grantville Gazette Volume 24

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

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.

Catalysts

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.