Our word salt is cognate to Latin sal. The greek word is 'als or "hals" in Roman letters. Many place names contain these words, such as Salisbury, Salzburg, Halle, Hallstatt, Reichenhall and so forth, showing a connection with salt. The ending -wich in English, as in Droitwich, Nantwich and Greenwich, indicates a place where salt was prepared. The book by Kurlansky in the References explains very well the historical importance of salt. This article considers the more technical aspects of salt.
Salt has a characteristic taste that is distinguished by receptors on the tongue. Similar ionic solutions also may taste salty, but not with the clean salt taste of salt itself. I do not know why the organism finds it important to be able to detect salt, except that salt is an essential nutrient. It has recently been recognized that athletes, particularly long-distance runners, often drink too much water and dilute the sodium ion concentration in the blood to dangerous levels, a condition called hyponatremia. In the past, salt pills were recommended to replace the sodium lost in sweat, but this was later deemed useless. Hyponatremia is a greater threat than dehydration. Carnivorous animals derive sufficient salt from their diet, but herbivorous animals may not. Grazing animals in the wild search out salt licks, and salt must be supplied to domesticated cattle. Perhaps the salt sense indicates that humans did not evolve as meat-eaters.
The most important use of salt until modern methods of food preservation (freezing, canning, drying) were introduced was in the preservation of food, as in salt fish, sausages, pickles and cheese. Salt immediately prevents bacterial rotting, often until lactic acid fermentation can take over to preserve the food more permanently. Salt fish, salt pork and salt beef appeared commonly in the diet. The taste for these items has survived into modern times, particularly since the availability of refrigeration can reduce the amount of salt used, rendering them more palatable. Ham, bacon and corned beef are examples. In general, salt was rarer than the food, and its supply determined the scale of the industry.
The greatest modern use of salt is for spreading on roads to combat ice, usually in combination with sand. A large amount is used in the chemical industry, to produce sodium carbonate, sodium bicarbonate and sodium hydroxide in tonnage amouints, all of wide utility in industry. Chlorine compounds, for bleaching and other uses, are also supplied by salt. Culinary salt is only a minor use by tonnage, but is the most common form in which salt is known to the public. This salt is pure, white, and comprised of uniform small cuboid crystals. This salt is a relatively modern invention. Indeed, impure salt and salt of less uniformity generally command a premium, as in "sea salt." These salts may have bitter overtones that are prized. Culinary salt contains a small amount of finely divided calcium silicate or sodium silicoaluminate that prevents the grains from adhering in a humid atmosphere. Caking of salt is due to hygroscopic impurities, such as magnesium chloride. Often, a small amount of potassium iodide is added to supply iodide in inland regions where it may be deficient in the environment, thereby preventing goiter. Simple goiter wouuld otherwise be prevalent in the Great Lakes and western interior mountain regions of North America. For some superstitious reason, iodized salt is not kosher, though (kosher) sea salt contains iodide. Iodide is present in such a small amount that it is only detectable by sensitive chemical tests, and has no systemic effects other than beneficial absorption by the thyroid gland. An essential metabolism-regulating hormome, thyroxine, cannot be formed without iodine. Fluoride could also be added to fight tooth decay, but this is not currently done.
Although salt is a common and inexpensive sodium compound, the rarer sodium carbonate is more useful. The first process for manufacturing sodium carbonate from salt was the LeBlanc process, originating in 1787. This process used solid-state reactions and evaporations. First, the salt was treated with sulphuric acid, and the hydrogen chloride was driven off by heat, leaving sodium sulphate. The sulphate was mixed intimately with coke and heated. The coke extracted the oxygen as carbon monoxide, leaving sodium sulphide. Finally, the sulphide was roasted with limestone to produce sodium carbonate and calcium sulphide. On treatment with water, the calcium sulphide decomposes and the sodium carbonate can be separated. The Solvay process, introduced in 1864, dissolves ammonia in brine, then bubbles carbon dioxide through it, making ammonium carbonate. Sodium bicarbonate, which is less soluble, precipitates because of the excess of sodium ion. The ammonia is recycled. The bicarbonate is heated to make anhydrous sodium carbonate, which may be redissolved in water and crystallized as the decahydrate, which is known as washing soda. Treated with additional carbon dioxide, a pure sodium bicarbonate is formed that is called baking soda. While the LeBlanc process produces noxious gases, the Solvay process does not and is much more environmentally acceptable. Finally, carbonate and bicarbonate can also be made by treating the sodium hydroxide solution obtained by the electrolysis of brine with carbon dioxide (obtained by calcining limestone). This is the leading current process.
Salt is a solid crystal formed from alternating Na+ and Cl- ions, held together by the electrical attraction between these very stable assemblies, whose charges exactly compensate. The Na+ ion is nothing like metallic sodium, and the Cl- ion nothing like chlorine gas. If it were not for their electrical charges, they would probably form heavy, nearly inert, gases. They do not absorb visible light, so the NaCl crystal is clear and transparent. When the salt is divided into small grains, multiple scattering returns incident light without change, and the powder appears pure white. This is characteristic not only of salt, but of any other colorless crystals, so that white powders are very common. This was a severe embarrassment to the traditional views of matter, which emphasized inherent qualities, so that "white powder" would really be a chemical description. However, one white powder (sodium bicarbonate) might extinguish a fire on which it is thrown, while another (potassium nitrate) would cause the fire to blaze up alarmingly. There is really nothing in a name. There are no sodium properties or chlorine properties in sodium chloride. Chemical behavior is determined by the alternatives available to the peripheral electrons of atoms, and all atoms are very much alike.
This is what made the Chinese discovery of nitrates around 1000 CE so remarkable. To purify nitrates then meant to separate one white powder from another with no hint of why they were different. This was the critical part of the discovery of gunpowder, but it was not gunpowder. The Chinese made use of nitrates for pyrotechnics, but explosive mixtures were developed in the West as propellants. These efforts depended on the magic Chinese white powder, so it was a cooperative effort. It is a common misconception that simply mixing saltpeter, charcoal and sulphur will give gunpowder. The necessary process is much more complicated.
Much confusion has resulted from the fact that the Greek nitron, nitron, refers to sodium carbonate, not to nitrate. Sodium carbonate was the strongest alkali known in the ancient world, and had many uses. However, it was very rare and was prized. Nitrates probably did get their name by a misapplication of nitron to another expensive and rare salt when it appeared in trade with China.
When salt is placed in water, the polar water molecules can cluster around an ion, effectively spreading the charge over a larger area and reducing the energy. This hydration can dismantle the salt crystal, producing what is called a salt solution. There is always a tendency for hydrated ions to stick to a crystal and become de-hydrated, and this tendency increases as the concentration of the solution increases. Therefore, if we do not run out of salt, an equilibrium is eventually reached in which the rate of hydration of ions is balanced by reattachment to the crystal. The solution is then called saturated. At 0°C, 35.7 g of NaCl dissolves in 100g of water. At 100°C, 39.8 g dissolves, and at 20°C, 36.0 g. Temperature does not affect the solubility of salt greatly. Conversely, dissolving salt has little thermal effect. An aqueous solution of salt is electrically conductive, because the hydrated ions are mobile. Chlorine and hydrogen gases are released at the electrodes whan a current is passed through the solution, while the liquid becomes sodium hydroxide solution.
Sea water contains around 35 parts per 1000 of dissoved salts. I presume this means about 3.6 g per 100 g of water, so sea water is quite far from saturation. Most of the dissoved matter is chlorine and sodium, with smaller amounts of magnesium, potassium, sulphate, carbonate and bicarbonate. Each ion has a certain lifetime between release into the ocean and its precipitation that determines its concentration. Ocean salinity is probably stable with time, but this is difficult to determine. It may be shocking to state that sea water contains no salt, but this is technically true. The solid salt is only formed by evaporating the water. Sea water will furnish something that is mostly salt. A little carbonate first precipitates, followed by gypsum, CaSO4·2H2O, at 10% salinity. Salt precipitates when the volume is about 1/10 the original volume, then anhydrite, CaSO4, and last of all the very soluble salts of K and Mg. These are the typical evaporites. The Piano del Sale of Eritrea, and the Gulf of Kara Bogaz by the Caspian, are two examples of evaporite deposits on the surface today. All thick deposits of evaporites are the result of continuous inflow and evaporation of saline water.
Salt for commercial use comes from three sources. First, the evaporation of sea water; second, the evaporation of brines in ground water; third, mining of rock salt. Even where rock salt is mined, it is dissolved in water for purification. The evaporation of water requires large amounts of energy, so that historically natural evaporation from pans was common. This is relatively easily done in dry, warm climates. In cooler and moister climates, the main drawback is not the lower temperature, but the unfortunate raining on the pans, which undoes the evaporation. For this reason, pans in moise regions were provided with roofs which could be pulled over the pans in case of rain. Natural evaporation could be hastened by heating, but this was economic only where a cheap fuel supply was at hand. At the current time, the much more efficient vacuum evaporator is used, especially for culinary salt. The crystal size depends on the rate of evaporation, rapid evaporation favoring small crystal size. The vacuum evaporator produces small, uniform crystals. If you grow salt crystals, they will usually be "hopper" crystals due to more rapid growth along the cube edges. Forms other than cubes are very rare for salt.
Well-digging technology was first developed for wells producing brine that was evaporated to make salt. One early location was Zigong in China, where natural gas was also available for heating the evaporation pans. Bamboo was used for pipes. The wells had wooden derricks, and were emptied by baling. The percussion drills were equipped with "jars" to improve the impact. This was very similar to the cable tools used extensively in the United States. The drilling methods were taken from the Kanawha Valley in Virginia to Pennsylvania by Edwin Clark to drill for oil instead. Rotary tools were imported from Russia around 1900 to drill in the soft sediments of the Gulf Coast, and rotary tools were also applied to brine wells. The invention of the rock bit by Hughes allowed rotary tools to dominate, since now they could attack hard rock even better than cable tools could. Most oil wells end their economic lives producing large quantities of salt water. Kanawha salt was favored by the economies of slave labor until after the Civil War. After that, salt from western New York State dominated the market for many years.
An egg was often used as a hydrometer to test the concentration of salt solutions. If the egg floated, then the brine was strong enough for pickling. A fresh, whole egg I measured weighed 56 g and had a volume of 50 cc, so its density was about 1.12. This is the density of a 16% solution by weight, or 16 g salt dissolved in 84 g of water. 100 g of solution has a volume of 100/1.12 or 89 cc, so the solution expands slightly as salt is added.
The use of brine freezing mixtures is well known, especially in making ice cream at home. A 23.3% solution produces a temperature of -6.0°F (Fahrenheit thought 0°F was the lowest that could be reached). At greater concentrations, the hydrate NaCl·2H2O crystallizes out, and the brine maintains the eutectic composition. A 16% solution reaches -11.9°C. Magnesium chloride, now often used on roads, reaches lower temperatures, down to -27°F. Salt solutions are not suitable as antifreeze mixtures, because they are very corrosive and promote rusting (since they are electrically conductive). Nonconductive alcohols and glycols are used instead. Pure ice freezes out of salt solution, which becomes more concentrated.
Crystalline salt not only grows in right-angled crystals, but cleaves easily into perfect rectangular parallelepipeds. This is very clearly and easily explained by considering the crystal structure. The Na+ ion is about 0.098 nm in radius, while the Cl_ ion has a radius of 0.181 nm. One way of arranging the ions in space is alternately along rows in each dimension. This is not the only way the ions can be packed, but it is the way chosen by the salt crystal. The distance between Na+ and Cl- ions is 0.098 + 0.181 = 0.279 nm. Think of a cube with Na+ and Cl- ions on alternate corners. This cube, by itself, does not generate the crystal if displaced by the edge distance, but has to be supplemented by cubes rotated in various ways. It also does not have the rotational symmetry of a cube. Since each ion is shared by the 8 cubes that share the lattice point, the cube represents half of an NaCl unit.
If you assemble 8 small cubes rotated in various ways, a cube can be formed that has the rotational symmetry of a cube. This cube has a side of 0.558 nm. It can be chosen to have an Na+ ion at each of its 8 corners, and in the centre of each of its 6 faces. The Cl- ions will then occupy the centres of the 12 edges, as well as the centre of the cube. Considering sharing with adjacent cubes, this cube will contain 4 NaCl units. The Na+ and Cl- ions each occupy a face-centred cubic lattice. There is no NaCl molecule as such; the whole crystal is a "molecule" in a sense, though this term really has no meaning here. When you smash a salt crystal, each fragment is a molecule. This cube does generate the whole crystal if displaced by multiples of its edge length in the three dimensions, and so can be considered a unit cell. A unit cell that contains only one NaCl unit can be identified. Its axis is along cube body diagonals, and is not as easy to visualize as the cubit unit cell.
A mole, or 6.023 x 1023, of Na atoms weighs 22.997 g, and a mole of Cl- atoms 35.457 g. Therefore, a cube containing (1/2)NaCl has a weight of 29.227 / 6.023 x 1023 g, and a volume of (0.279)3 nm = 2.172 x 10-23 cm3. This gives a density of 2.23 g/cm3, close to the observed value of 2.16 g/cm3. The actual unit cell is slightly larger than what we have calculated from the ionic radii, for various reasons. The observed side of the fcc unit cell is 0.563 nm. Except for thermal vibrations and structual defects, the crystal is perfect, since all ions of any type are identical to each other.
Nothing compels the ions to form interpenetrating face-centred cubic lattices. This is simply the best solution for reducing the electrostatic energy of the charges, considering their sizes. It is quite a different problem from the packing of uncharged spheres. Alkali metal ions increase in radius from Na+ at 0.095 nm to Cs+ at 0.169 nm, which is nearly as large as Cl-, 0.181 nm. Cesium chloride prefers a lattice that can be built from cubic units with Cs- at each corner and a Cl- at the centre (or vice versa). This is interpenetrating simple cubic lattices of the ions. Rubidium chloride is fcc at lower temperatures, and simple cubic at higher temperatures; obviously, it teeters on the edge between one structure and the other. Calculations of the energy are not easy, since the result is the difference between many contributions of opposite signs. The energy can be calculated as a function of the lattice spacing, and the minimum found. The structure with the lowest (free) energy is the one that wins out. An excellent treatment of crystal binding will be found in Kittel (see References).
These considerations were verified in detail by X-ray analysis, though they were suspected much earlier. They would have been a great comfort to Epicurus, who taught the atomic structure of matter, but had no way to confirm his views. The crystal is held together by the strong electrostatic forces between the ions, though the crystal itself is electrically neutral. Salt melts at about 800°C. High melting points are typical of ionic crystals. Fused salt conducts electricity, since the charged ions are mobile. Liquid sodium and chlorine gas appear at the electrodes.
Pure salt is transparent, with an index of refraction of 1.5443 for the Na D line, very similar to that of glass. Unlike glass, it is also transparent in the infrared, out to about 18μ. As a prism for dispersion in a spectroscope, it is most useful in the range 5-15μ, and so very useful for familiar vibration-rotation spectra. Salt optics must be protected from moisture; they are also soft and easily scratched, and fragile. The high excitation energies of the closed shells of the ions is responsible for the transparency.
Salt occurring as a mineral is scientifically called halite, from the Greek for salt, hals, and colloquially rock salt. It is usually impure, often containing magnesium and sulphate, and translucent rather than transparent, with yellow, red, blue or purple coloration. Its Mohs hardness is 2.5, as hard as the fingernail but not as hard as calcite. Halite can be confused with the much rarer KCl or sylvite, which is also cubic and has a salty taste. KCl is more bitter in taste than NaCl, and is more soluble. Octahedral faces are sometimes found on sylvite crystals, but not on halite crystals. Sylvite can be scratched by the finger nail, and is slightly less dense than halite (1.99). It gives the violet flame of potassium in a flame test, if the sodium yellow is filtered out with blue glass.
Beds of salt occur when the outlets of bodies of water are cut off and the waters evaporate. This has happened many times in the earth's history, and salt beds are widespread, often associated with gypsum and anhydrite, which has a similar origin. For example, this occurred in the region that is now the Gulf of Mexico in the Jurassic, when the Atlantic was just beginning to open. Sea water was supplied from the west to basins that subsequently dried up in many cycles, forming a thick bed of salt now called the Louann salt, which has been torn in half as South America separated from North America. Most other North American salt is of Permian age, as is European salt.
When later sediments were piled on the Louann salt and increased the pressure on it, the salt was squeezed upwards into diapirs (from the Greek for "through-piercers") that pierced and bowed upwards the overlying beds, forming salt domes. These can rise close to the surface and have surface expression. They are better described as plugs or stocks than as domes. Their radii are about equal to the thickness of the salt bed from which they arose, though the upper parts may be belled out. Petroleum was made by bacterial action in the oxygen-poor but warm seas of the area, and collected in the structures created by the domes around their edges. Sulphur was also made by bacteria at the same time, associated with the petroleum and derived from the gypsum of the cap rock. The cap rock is the residue left when salt has been dissolved. The salt had nothing directly to do with the petroleum or the sulphur, but created the conditions necessary for trapping them. Salt, though apparently a crystalline solid, can flow over geologic times. This came as a surprise to geologists, as did plate tectonics. The nature of salt domes was not appreciated until after the 1901 Spindletop discovery on the Texas Gulf coast. The Spindletop dome is about a mile in diameter. Hundreds of salt domes have been recognized in this region, both on land and underwater in the Gulf.
It should not be thought that salt domes occur only around the Gulf of Mexico. Another important region is southern Iran and the Persian Gulf, where the salt is Cambrian in age and the salt domes generally larger than those of the Gulf of Mexico. Salt mountains and salt glaciers are found here. Here, too, petroleum occurs. Salt domes are found in the North German Plain, around Hannover, for example, and in southwestern France. There are even salt domes in the Queen Elizabeth islands of arctic Canada. This salt is late Silurian, and is good evidence of continental drift, since it could not have formed in the Arctic.
The United States Strategic Petroleum Reserve of about 700 million barrels is stored in salt dome traps, since salt seals them well. Currently, the U.S. consumes about 20 million barrels a day, and produces 8.8 domestically (the second largest producer in the world). If all imports were halted, the reserve would last about two months. This is really not much of a reserve, and is certainly insufficient to affect the petroleum market.
Among the interesting salt lore collected by Kurlansky is the history of ketchup. The name comes from the Indonesian fermented fish sauce called kechap ikan, where kechap is pronounced ketchap. This name was applied to English salted anchovy sauce in the 17th century. All sauces of this type contained a good deal of salt as a preservative. They include Roman garum and Chinese soy sauce, used to flavor and salt food instead of adding salt directly. Tomatoes were made into a sauce around 1800 in the United States, which was called tomato ketchup, and has now become simply ketchup.
Lox is salt-cured Pacific salmon, made famous by Jewish delicatessen shops. It has now been replaced almost entirely by Nova Lox. Nova Lox is much less salty, which is practical due to refrigeration. The "Nova" refers to Nova Scotia, where it was first made. Since Atlantic salmon has vanished, it is now made from Pacific salmon. Pastrami is a Romanian salt beef, dried, spiced, salted, smoked and steamed. Salt cod, once an inexpensive staple and the basis of many famous recipes, is now very expensive, if available at all.
M. Kurlansky, Salt (New York: Penguin Books, 2003)
C. Kittel, Introduction to Solid State Physics, 3rd ed. (New York: Wiley, 1966). Chapter 3, pp. 89-99. Ionic Crystals.
A. Holmes, Principles of Physical Geology, 2nd ed. (New York: Ronald, 1965). Chapter 10, pp. 234-248. Salt diapirs.
N. A. Lange, Handbook of Chemistry, 10th ed. (New York: McGraw-Hill, 1961). Density of salt solutions is given on p. 1168 and on p. 1194, where the freezing point is also given.
Composed by J. B. Calvert
Created 12 April 2005
Last revised 14 September 2007