Sodium carbonate

• 1. Introduction

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  Sodium carbonate goes by a crowd of names. Washing soda. Soda ash. Sal soda. Soda crystals. Each name points to a different life this single white salt has led. It is odorless, dissolves in water, and turns that water alkaline. Long before any factory made it, people pulled it from the ground and from fire. Deposits of a mineral called natron were mined from dry lake bottoms in Egypt since ancient times. There it helped prepare mummies and shaped some of the earliest glass. So how did a salt scraped from evaporated lakes become something produced in vast tonnages from rock and brine? Why does the same compound sit in a noodle kitchen, a swimming pool, a photographic darkroom, and a power station smokestack? And what drove chemists across three countries to keep reinventing the way it is made?

• 2. A Salt Of Many Hydrates

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  Na2CO3 rarely arrives alone. It clings to water, forming three distinct hydrates alongside the dry, anhydrous salt. The decahydrate, known as natron, carries ten molecules of water of crystallization. It crystallizes from solution between minus 2.1 and plus 32.0 degrees Celsius, then readily effloresces, shedding water to become the monohydrate. Between 32.0 and 35.4 degrees Celsius, a narrow window, the heptahydrate forms with seven waters, though it has never been found as a mineral. Above that temperature the monohydrate, thermonatrite, takes over, and chemists also call it crystal carbonate. The anhydrous form, natrite, also called calcined soda, appears when the hydrates are heated. It also forms when sodium hydrogencarbonate is calcined, as in the final step of the Solvay process. Even the count of waters can surprise. One reported hydrate holds 2.5 water units per sodium carbonate unit, a so-called penta hemihydrate. The decahydrate is the most common of these forms, and it carries the household name that follows soda ash into the laundry.

• 3. The Chemistry Of Cleaning

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  Washing soda earns its name in the laundry room, and it is the decahydrate doing the work. Dissolve soda ash in water, crystallize it, and you get washing soda, a component of many dry soap powders. Its cleaning power runs through saponification, a process that converts fats and grease into water-soluble salts, which are soaps. Hard water, the kind that fights against suds, usually carries dissolved calcium or magnesium ions. Sodium carbonate strips these out by offering a water-soluble source of carbonate. The calcium and magnesium ions meet that carbonate and drop out as insoluble solid precipitates. What remains is softened water, free of the ions that caused the trouble. It is one of the few metal carbonates that dissolves in water at all, and that rare solubility is exactly what makes it useful far beyond the wash.

• 4. Fire And Glass

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  Silica melts at 1,713 degrees Celsius, a temperature that defeats ordinary furnaces. Sodium carbonate acts as a flux for that silica, dragging the melting point down to something achievable without special materials. The result is soda glass, but soda glass has a flaw. It is mildly water-soluble. Adding calcium carbonate to the melt fixes this, producing the soda-lime glass that fills bottles and windows, with a transition temperature around 570 degrees Celsius. As the mixture of sodium carbonate, calcium carbonate, and silica sand heats, the carbonates release carbon dioxide. Sodium carbonate thereby becomes a source of sodium oxide inside the glass. Soda-lime glass has been the most common form of glass for centuries, and it remains a key input for tableware. The same salt that ancient Egyptians drew from natron deposits still sits at the heart of the windows we look through.

• 5. Flavor And The Kitchen Table

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  Stronger than baking soda, weaker than lye, sodium carbonate occupies a useful middle ground in the kitchen. Its alkalinity changes how gluten develops in kneaded dough. It also improves browning by lowering the temperature at which the Maillard reaction begins. That chemistry gives Japanese ramen noodles their characteristic flavor and chewy texture, and a similar alkaline solution shapes lamian in Chinese cuisine. Cantonese bakers reach for it as a substitute for lye-water, lending moon cakes their texture and browning. In German and Central European cooking, pretzels and lye rolls traditionally bathed in lye can be treated with sodium carbonate instead, safer to handle though it browns less strongly. The salt also powers sherbet powder. When saliva moistens the sherbet, sodium carbonate reacts with a weak acid, commonly citric acid, in an endothermic reaction that releases carbon dioxide and creates the cooling, fizzing sensation. As a recognized food additive it carries the European number E500, working as an acidity regulator, anticaking agent, raising agent, and stabilizer. It even stabilizes the pH of snus.

• 6. The Quiet Workhorse

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  Cheaper than sodium hydroxide and far safer to handle, sodium carbonate serves as a strong base across many trades. Photographic film developing needs stable alkaline conditions, and sodium carbonate regulates the pH so most developing agents can do their job. In swimming pools and aquariums it holds a desired pH and carbonate hardness, the KH. Dyers who work with fiber-reactive dyes use it under names like soda ash fixative or soda ash activator, a mordant that bonds the dye to cellulose plant fiber. In froth flotation it conditions the float by keeping a favorable pH alongside CaO and other mildly basic compounds. The compound is also a launching point for other chemicals. Baking soda, sodium bicarbonate, is often generated from it, and that bicarbonate also appears in fire extinguishers. It makes sodium bisulfite for the sulfite method of separating lignin from cellulose, a reaction that doubles as a way to scrub sulfur dioxide from power station flue gases under stringent emission controls. The cotton industry uses it to neutralize the sulfuric acid of acid delinting. Brickmakers wet their clay with it to extrude using less water. It binds wet alginate to gelled alginate in casting, foams and abrades in toothpaste, and neutralizes acids in tanning animal hides. Its integral enthalpy of solution is minus 26.7 kilojoules per mole, and the monohydrate registers just 1.3 on the Mohs hardness scale, soft enough to scratch with a fingernail.

• 7. Mined, Erupted, And Extraterrestrial

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  Ol Doinyo Lengai, Tanzania's unique volcano, erupts sodium carbonate from its vents. Other volcanoes are presumed to have done the same in the past, but the minerals are unstable at the Earth's surface and likely erode away. Sodium carbonate occurs naturally in arid regions, especially in evaporite deposits left when seasonal lakes dry up. The anhydrous mineral natrite is quite rare. All three mineral forms, along with trona, trisodium hydrogen dicarbonate dihydrate, also turn up in ultra-alkaline pegmatitic rocks, such as those on the Kola Peninsula in Russia. The salt reaches beyond Earth too. Deposits identified as sodium carbonate are the source of the bright spots on Ceres, interior material brought to the surface. Mars carries carbonates, and sodium carbonate is expected among them, yet deposits remain unconfirmed. Some explain that absence by a global dominance of low pH in the planet's once-aqueous soil. Closer to home, trona supplies nearly all US consumption of sodium carbonate. Large deposits found in 1938, including one near Green River, Wyoming, made mining more economical than industrial production in North America. Turkey holds important reserves as well, and two million tons of soda ash have been extracted from the deposits near Ankara.

• 8. Inventing The Process

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  Until the early 19th century, Europe made soda ash by burning plants. Salt-tolerant halophytes like glassworts and saltworts, or seaweeds such as Fucus species, were harvested, dried, and burned. The ashes were lixivated, washed with water, into an alkali solution, then boiled dry. The very name soda ash traces to the Arabic word soda, applied to Salsola soda, one of the seashore plants used. Concentration varied wildly, from 2 to 3 percent in seaweed-derived kelp to 30 percent in the best Spanish barilla. As these sources fell short by the end of the 18th century, chemists raced to synthesize soda ash from salt. In 1792 the French chemist Nicolas Leblanc patented a route using salt, sulfuric acid, limestone, and coal, producing an intermediate called black ash from which soda ash was extracted with water. The Leblanc process spewed hydrochloric acid as air pollution and left calcium sulfide as troublesome waste, yet it dominated until the late 1880s. In 1861 the Belgian industrial chemist Ernest Solvay devised a cleaner method, reacting sodium chloride, ammonia, water, and carbon dioxide, recycling its ammonia and leaving only calcium chloride as waste. By 1900-90 percent of sodium carbonate came from the Solvay process, and the last Leblanc plant closed in the early 1920s. In the 1930s the Chinese chemist Hou Debang coupled the idea to the Haber process. His method, lianhe zhijian fa, the coupled manufacturing alkali method, eliminated calcium chloride and turned its byproduct ammonium chloride into a sellable fertilizer.