oxyacid

/ok"see as'id/, n. Chem.

an inorganic acid containing oxygen. Also called oxygen acid.

[1830-40; OXY-² + ACID]

* * *

▪ chemical compound

Introduction

      any oxygen-containing acid. Most covalent nonmetallic oxides (oxide) react with water to form acidic oxides; that is, they react with water to form oxyacids that yield hydronium ions (H₃O⁺) in solution. There are some exceptions, such as carbon monoxide, CO, nitrous oxide, N₂O, and nitric oxide, NO.

      The strength of an oxyacid is defined by the extent to which it dissociates in water (i.e., its ability to form H⁺ ions). In general, the relative strength of oxyacids can be predicted on the basis of the electronegativity and oxidation number of the central nonmetal atom. The acid strength increases as the electronegativity of the central atom increases. For example, because the electronegativity of chlorine (Cl) is greater than that of sulfur (S), which is in turn greater than that of phosphorus (P), it can be predicted that perchloric acid, HClO₄, is a stronger acid than sulfuric acid, H₂SO₄, which should be a stronger acid than phosphoric acid, H₃PO₄. For a given nonmetal central atom, the acid strength increases as the oxidation number of the central atom increases. For example, nitric acid, HNO₃, in which the nitrogen (N) atom has an oxidation number of +5, is a stronger acid than nitrous acid, HNO₂, where the nitrogen oxidation state is +3. In the same manner, sulfuric acid, H₂SO₄, with sulfur in its +6 oxidation state, is a stronger acid than sulfurous acid, H₂SO₃, where a +4 oxidation number of sulfur exists.

      The salt of an oxyacid is a compound formed when the acid reacts with a base: acid + base → salt + water. This type of reaction is called neutralization, because the solution is made neutral.

Oxyacids of nitrogen

nitric acid and nitrate salts

      Nitric acid, HNO₃, was known to the alchemists of the 8th century as “aqua fortis” (strong water). It is formed by the reaction of both dinitrogen pentoxide (N₂O₅) and nitrogen dioxide (NO₂) with water. Small amounts of nitric acid are found in the atmosphere after thunderstorms, and its salts, called nitrates (nitrate), occur widely in nature. Enormous deposits of sodium nitrate, NaNO₃, also known as Chile saltpetre, are found in the desert region near the boundary of Chile and Peru. These deposits can be 3 km (2 miles) wide, 300 km (200 miles) long, and up to 2 metres (7 feet) thick. Potassium nitrate, KNO₃, sometimes called Bengal saltpetre, is found in India and other countries in East Asia. Nitric acid can be prepared in the laboratory by heating a nitrate salt, such as those mentioned above, with concentrated sulfuric acid; for example,

NaNO₃ + H₂SO₄ + heat → NaHSO₄ + HNO₃.

Since HNO₃ boils at 86 °C (187 °F) and H₂SO₄ boils at 338 °C (640 °F) and NaNO₃ and NaHSO₄ are nonvolatile salts, nitric acid is easily removed by distillation.

      Commercially, nitric acid is produced by the Ostwald process. This process involves oxidation of ammonia, NH₃, to nitric oxide, NO, further oxidation of the NO to nitrogen dioxide, NO₂, and then conversion of the NO₂ to nitric acid (HNO₃). This is a flow process in which a mixture of ammonia and excess air is heated to 600 to 700 °C (1,100 to 1,300 °F) and passed through a platinum-rhodium catalyst. (A catalyst increases the rate of a reaction without itself being consumed in the reaction.) As the oxidation to NO occurs, this gaseous mixture literally burns with a flame. Additional air is added to oxidize the NO to NO₂. The NO₂, excess oxygen, and the unreactive nitrogen from the air are passed through a water spray, where HNO₃ and NO form as the NO₂ disproportionates. The gaseous NO is recycled through the process with more air, and the liquid HNO₃ is drawn off and concentrated. About 7 billion kg (16 billion pounds) of HNO₃ are produced commercially in the United States each year, with the bulk of it made by the Ostwald process.

      When pure, nitric acid is a colourless liquid that boils at 86 °C (187 °F) and freezes at −42 °C (−44 °F). Upon being exposed to light or heat, it decomposes to produce oxygen, water, and a mixture of nitrogen oxides (primarily NO₂).

4HNO₃ + light (or heat) → 4ΝΟ₂ + 2H₂O + O₂

Consequently, nitric acid is often yellow or brown in colour because of the NO₂ that forms as it decomposes. Nitric acid is stable in aqueous solution, and 68 percent solutions of the acid (i.e., 68 grams of HNO₃ per 100 grams of solution) are sold as concentrated HNO₃. It is both a strong oxidizing agent and a strong acid. Nonmetallic elements such as carbon (C), iodine (I), phosphorus (P), and sulfur (S) are oxidized by concentrated HNO₃ to their oxides or oxyacids with the formation of NO₂; for example,

S + 6HNO₃ → H₂SO₄ + 6NO₂ + 2H₂O.

In addition, many compounds are oxidized by HNO₃. hydrochloric acid, aqueous HCl, is readily oxidized by concentrated HNO₃ to chlorine, Cl₂, and chlorine dioxide, ClO₂. aqua regia (“royal water”), a mixture of one part concentrated HNO₃ and three parts concentrated HCl, reacts vigorously with metals (metal). The use of this mixture by alchemists to dissolve gold is documented as early as the 13th century.

      The action of nitric acid on a metal usually results in reduction of the acid (i.e., a decrease in the oxidation state of the nitrogen). The products of the reaction are determined by the concentration of HNO₃, the metal involved (i.e., its reactivity), and the temperature. In most cases, a mixture of nitrogen oxides, nitrates, and other reduction products is formed. Relatively unreactive metals such as copper (Cu), silver (Ag), and lead (Pb) reduce concentrated HNO₃ primarily to NO₂. The reaction of dilute HNO₃ with copper produces NO, whereas more reactive metals, such as zinc (Zn) and iron (Fe), react with dilute HNO₃ to yield N₂O. When extremely dilute HNO₃ is used, either nitrogen gas (N₂) or the ammonium ion (NH₄⁺) may be formed. Nitric acid reacts with proteins (protein), such as those in human skin, to produce a yellow material called xanthoprotein.

      Nitrates, which are salts of nitric acid, are produced when metals or their oxides, hydroxides (hydroxide), or carbonates (carbonate) react with nitric acid. Most nitrates are soluble in water, and a major use of nitric acid is to produce soluble metal nitrates. All nitrates decompose when heated and may do so explosively. For example, when potassium nitrate (KNO₃) is heated, a nitrite (a compound containing NO₂⁻) is formed, and oxygen gas is evolved.

2KNO₃ + heat → 2KNO₂ + O₂

When heavy metal nitrates are heated, the metal oxide is produced, as in, for example,

2Cu(NO₃)₂ + heat → 2CuO + 4NO₂ + O₂.

ammonium nitrate, (NH₄)₂NO₃, produces nitrous oxide, N₂O, and is especially dangerous to heat or detonate.

      Nitric acid is heavily used in the laboratory and in chemical industries as a strong acid and as an oxidizing agent. The manufacture of explosives (explosive), dyes (dye), plastics (plastic), and drugs (drug) makes extensive use of the acid. Nitrates are valuable as fertilizers (fertilizer). gunpowder is a mixture of potassium nitrate, sulfur, and charcoal. Ammonal, an explosive, is a mixture of ammonium nitrate and aluminum powder.

nitrous acid and nitrite salts

       nitrous acid (HNO₂), a weak acid, is very unstable and exists only in aqueous solution. A pale blue solution of HNO₂ is obtained when dinitrogen trioxide (N₂O₃) is added to water, and it is also easy to prepare HNO₂ by adding acid to a solution of a nitrite.

NO₂⁻ + H₃O⁺ → HNO₂ + H₂O

It decomposes slowly at room temperature—and more rapidly at elevated temperatures—to nitric acid and nitric oxide. Nitrous acid is oxidized to nitric acid by active oxidizing agents and acts as an oxidizing agent with strong reducing agents. Sodium nitrite, NaNO₂, is an important example of a nitrite—that is, a salt of nitrous acid. It is typically prepared by reducing molten sodium nitrate with elemental lead.

NaNO₃ + Pb → NaNO₂ + PbO

This salt is added to meats, such as hot dogs, for two reasons. It prolongs the meat's retention of a red colour, and it inhibits the growth of bacteria that can cause food poisoning. The addition of sodium nitrite to meat is controversial because nitrous acid, which is produced in the human body when stomach acid reacts with the ingested nitrite ion, is known to react with certain organic compounds (organic compound) to form nitrosamines. Some of the compounds in the nitrosamine class are known to cause cancer in laboratory animals. Consequently, the United States Food and Drug Administration limits the amount of sodium nitrite that can be legally added to foods.

      In general, the salts of all oxyacids are more stable than the acids themselves; such is the case with nitrites. They are much more stable than nitrous acid. Most nitrites are soluble in water and in concentrated forms, like nitrates, can explode upon heating or detonation.

Oxyacids of phosphorus

Orthophosphoric acid (phosphoric acid) and phosphate salts

      Orthophosphoric acid, H₃PO₄, is usually called simply phosphoric acid. When pure, it is a colourless crystalline solid that melts at 42 °C (108 °F). It rapidly absorbs moisture from the air and liquefies. It is typically available commercially as syrupy phosphoric acid, which is an 85 percent solution in water. Pure H₃PO₄ is produced by dissolving phosphorus pentoxide (P₄O₁₀) in water, although it is more commonly prepared by treating calcium phosphate, Ca₃(PO₄)₂, with concentrated sulfuric acid, H₂SO₄.

Ca₃(PO₄)₂ + 3H₂SO₄ → 2H₃PO₄ + 3CaSO₄

The products are diluted with water, and the insoluble CaSO₄ is removed by filtration. The dilute acid produced is contaminated with calcium dihydrogen phosphate, Ca(H₂PO₄)₂, and other compounds found with naturally occurring Ca₃(PO₄)₂.

      Orthophosphoric acid is a triprotic acid—i.e., it can donate all three of its hydrogen atoms as protons (proton) in aqueous solution. Thus, it can form three series of salts: dihydrogen phosphates, containing the H₂PO₄⁻ ion; hydrogen phosphates, containing the HPO₄²⁻ ion; and orthophosphates, containing the PO₄³⁻ ion. When dissolved in water, soluble dihydrogen phosphate salts form solutions that are weakly acidic, because H₂PO₄⁻ is a weak acid. Aqueous solutions of hydrogen phosphates are basic, because the HPO₄²⁻ ion is stronger as a base (i.e., a proton acceptor) than as an acid. The PO₄³⁻ ion is a moderately strong base, so orthophosphate salts form strongly basic solutions. The hydrogen phosphate salts, as well as H₃PO₄, decompose with loss of water when heated to form compounds containing P−O−P bonds. The ease of polymerization via P−O−P linkages and the possibility of the formation of P−P and P−H bonds allow innumerable oxyacids and their salts to be formed. These acids are termed the lower oxyacids of phosphorus.

phosphorous acid and phosphite salts

      Pure phosphorous acid, H₃PO₃, is best prepared by hydrolysis of phosphorus trichloride, PCl₃.

PCl₃ + 3H₂O → H₃PO₃ + 3HCl

The resulting solution is heated to drive off the HCl, and the remaining water is evaporated until colourless crystalline H₃PO₃ appears on cooling. The acid can also be obtained by the action of water on P₄O₆, PBr₃, or PI₃. Colourless crystalline H₃PO₃ melts at 70.1 °C (158.2 °F), is very soluble in water, and has an odour similar to that of garlic. Heating phosphorous acid to about 200 °C (392 °F) causes it to disproportionate into phosphine, PH₃, and orthophosphoric acid.

4H₃PO₃ + heat → PH₃ + 3H₃PO₄

Phosphorous acid and its salts are active reducing agents, because of their easy oxidation to phosphoric acid and phosphate salts, respectively. For example, phosphorous acid reduces the silver ion (Ag⁺) to elemental silver (Ag), mercury(II) salts to mercury(I) salts, and sulfurous acid, H₂SO₃, to elemental sulfur. Whereas H₃PO₄ has three hydrogen atoms bound to oxygen and is triprotic, H₃PO₃ is diprotic, owing to its structure in which only two hydrogens are bonded to oxygen and are acidic. The third hydrogen is bonded directly to phosphorus and is not very acidic. For this reason, H₃PO₃ forms only two series of salts, one containing the dihydrogen phosphite ion, H₂PO₃⁻, and the other containing the hydrogen phosphite ion, HPO₃²⁻.

Hypophosphorous acid and hypophosphite salts

 Free hypophosphorous acid, H₃PO₂, is prepared by acidifying aqueous solutions of hypophosphite ions, H₂PO₂⁻. For example, the solution remaining when phosphine is prepared from the reaction of white phosphorus and a base contains the H₂PO₂⁻ ion. If barium hydroxide (BaOH) is used as the base and the solution is acidified with sulfuric acid, barium sulfate, BaSO₄, precipitates, and an aqueous solution of hypophosphorous acid results.

Ba²⁺ + 2H₂PO₂⁻ + 2H₃O⁺ + SO₄²⁻ → BaSO₄ + 2H₃PO₂ + 2H₂O

The pure acid cannot be isolated merely by evaporating the water, however, because of the easy oxidation of the hypophosphorous acid to phosphoric acids (and elemental phosphorus) and its disproportionation to phosphine and phosphorous acid. The pure acid can be obtained by extraction of its aqueous solution by diethyl ether (ether), (C₂H₅)₂O. Pure hypophosphorous acid forms white crystals that melt at 26.5 °C (79.7 °F). The electronic structure of hypophosphorous acid is such that it has only one hydrogen atom bound to oxygen, and it is thus a monoprotic oxyacid. It is a weak acid and forms only one series of salts, the hypophosphites. Hydrated sodium hypophosphite, NaH₂PO₂ · H₂O, is used as an industrial reducing agent, particularly for the electroless plating of nickel onto metals and nonmetals.

Oxyacids of sulfur

 There are many oxyacids of sulfur. The most important of these acids are sulfuric acid, H₂SO₄, and sulfurous acid, H₂SO₃.

sulfuric acid

      Sulfuric acid is sometimes referred to as the “king of chemicals” because it is produced worldwide in such large quantities. In fact, per capita use of sulfuric acid has been taken as one index of the technical development of a country. Annual production in the United States, which is the world's leading producer, is well over 39 billion kg (86 billion pounds). It is the cheapest bulk acid.

Preparation

 Most sulfuric acid is produced by the modern contact process. First, elemental sulfur or sulfide ores (ore) are heated with oxygen to produce sulfur dioxide (SO₂). About 60 percent of the sulfur dioxide produced throughout the world comes from burning sulfur, and approximately 40 percent is derived from roasting sulfide minerals. (Roasting is the process by which ores are oxidized by heating in air.) Sulfur dioxide is then oxidized to sulfur trioxide, SO₃. This oxidation reaction is exothermic (i.e., releases energy in the form of heat) and reversible. Accordingly, a vanadium oxide catalyst is used on an inert support to increase the rate of the oxidation without decreasing the yield. Under optimum conditions, the feed gas consists of equimolar quantities of oxygen and sulfur dioxide (i.e., a 5:1 ratio of air to sulfur dioxide) that passes through a four-stage catalytic converter operating at various temperatures. After the gas mixture has passed over three of the catalyst beds and approximately 93 percent conversion to sulfur trioxide has occurred, it is cooled and absorbed into sulfuric acid in ceramic-packed towers. A final conversion of greater than 99 percent is achieved after passage through the final reaction bed. All three reactions used to produce sulfuric acid, as shown below, are exothermic. Efficient utilization of this energy to generate electricity, for example, is a key component in maintaining the inexpensive price of this heavily used acid.

S + O₂ → SO₂

2SO₂ + O₂ → 2SO₃

SO₃ + H₂O (in 98% H₂SO₄) → H₂SO₄

      Pure sulfuric acid is a colourless, oily, dense (1.83 grams per cc) liquid that freezes at 10.5 °C (50.9 °F). It fumes when heated because of its decomposition to water and sulfur trioxide. Because SO₃ has a lower boiling point than water, more SO₃ is lost during heating. When a concentration of 98.33 percent acid is reached, the solution boils at 338 °C without any further change in concentration. This is called a constant boiling solution, and it is this concentration that is sold as concentrated sulfuric acid. Anhydrous sulfuric acid mixes with water in all proportions in a very exothermic reaction. Adding water to concentrated acid can cause explosive spattering. Because it reacts with organic compounds in the skin, concentrated sulfuric acid can cause severe burns (burn). Thus, to decrease the risk of injury in the laboratory, sulfuric acid should always be added to water slowly and with stirring to distribute the heat.

Formation of sulfate salts

      Sulfuric acid has its two hydrogen atoms bonded to oxygen, ionizes in two stages, and is a strong diprotic acid. In aqueous solution, loss of the first hydrogen (as a hydrogen ion, H⁺) is essentially 100 percent. The second ionization takes place to an extent of about 25 percent, but HSO₄⁻ is nonetheless considered a moderately strong acid. Because it is a diprotic acid, H₂SO₄ forms two series of salts: hydrogen sulfates, HSO₄⁻, and sulfates, SO₄²⁻. The sulfates of the alkaline-earth metals (alkaline-earth metal)— calcium (Ca), strontium (Sr), and barium (Ba)—as well as that of lead (Pb) are virtually insoluble, and these salts are found as naturally occurring minerals. These important minerals include gypsum (CaSO₄ · 2H₂O), celestine (SrSO₄), barite (BaSO₄), and anglesite (PbSO₄). These insoluble salts can be prepared in the laboratory by metathesis reactions. A metathesis reaction is one in which compounds exchange anion- cation partners. For example, if a solution of barium nitrate, Ba(NO₃)₂, is added to a solution of sodium sulfate, Na₂SO₄, a precipitation of barium sulfate, BaSO₄, occurs. This is an important reaction because it can be used as both a qualitative and quantitative test for the sulfate ion and the barium ion. (Qualitative tests are used to determine the presence or absence of a substance, whereas quantitative tests are used to measure the amount of a constituent.) In addition to metathesis reactions, sulfate salts can generally be prepared by dissolution of metals in aqueous H₂SO₄, neutralization of aqueous H₂SO₄ with metal oxides (oxide) or hydroxides (hydroxide), oxidation of metal sulfides (a sulfide contains S²⁻) or sulfites (SO₃²⁻), or decomposition of salts of volatile acids, such as carbonates, with aqueous H₂SO₄. Some important soluble sulfate salts are Glauber's salt, Na₂SO₄ · 10H₂O; Epsom salt, MgSO₄ · 7H₂O; blue vitriol, CuSO₄ · 5H₂O; green vitriol, FeSO₄ · 7H₂O; and white vitriol, ZnSO₄ · 7H₂O.

Reactions and uses

      Pure H₂SO₄ undergoes extensive self- ionization (sometimes called autoprotolysis).

2H₂SO₄ → H₃SO₄⁺ + HSO₄⁻

This autoprotolysis reaction is, however, only one of the equilibrium reactions that occur in pure H₂SO₄ to give it an extremely high electrical conductivity. There are three additional equilibrium reactions that take place because of the ionic self-dehydration of sulfuric acid.

2HSO₄ ⇌ H₃O⁺ + HS₂O₇⁻

H₂O + H₂SO₄ ⇌ H₃O⁺ + HSO₄⁻

H₂S₂O₇ + H₂SO₄ ⇌ H₃SO₄⁺ + HS₂O₇

Thus, there are at least seven well-defined species that exist in “pure” H₂SO₄. The value of the dielectric constant of the acid is also quite high (ε = 100).

      Concentrated sulfuric acid is not a very strong oxidizing agent unless it is hot. When it acts as an oxidizing agent, however, it can be reduced to several different sulfur species, including SO₂, HSO₃⁻, SO₃²⁻, elemental sulfur (S₈), hydrogen sulfide (H₂S), and the sulfide anion (S²⁻). Concentrated sulfuric acid is a good dehydrating agent, as it reacts with many organic materials to remove the elements of water.

      The amount of sulfuric acid used in industry exceeds that of any other manufactured compound. In the United States approximately 67 percent of the acid is utilized to convert phosphate rock to phosphoric acid. The phosphoric acid is then converted to phosphate fertilizers (fertilizer). Other major uses include the refining of petroleum, the removal of impurities from gasoline and kerosene, the pickling of steel (the cleaning of its surface), and the manufacture of other chemicals, such as nitric and hydrochloric acids. It also is utilized in lead storage batteries (battery) and in the production of paints (paint), plastics (plastic), explosives (explosive), and textiles (textile).

Sulfurous acid and sulfite salts

      When sulfur dioxide is dissolved in water, an acidic solution results. This has long been loosely called a sulfurous acid, H₂SO₃, solution. However, pure anhydrous sulfurous acid has never been isolated or detected, and an aqueous solution of SO₂ contains little, if any, H₂SO₃. Studies of these solutions indicate that the predominant species are hydrated SO₂ molecules, SO₂ · nH₂O. The ions present in these solutions are dependent on concentration, temperature, and pH and include H₃O⁺, HSO₃⁻, S₂O₅²⁻, and perhaps SO₃²⁻. However, “sulfurous acid” has two acid dissociation constants. It acts as a moderately strong acid with an apparent ionization of about 25 percent in the first stage and much less in the second stage. These ionizations produce two series of salts—sulfites, containing SO₃²⁻, and hydrogen sulfites, containing HSO₃⁻. Only with large cations, such as Rb⁺ ( rubidium) or Cs⁺ ( cesium), have solid HSO₃⁻ salts been isolated. Attempts to isolate these salts with smaller cations tend to yield disulfites as a product of dehydration.

2HSO₃⁻ ⇌ S₂O₅²⁻ + H₂O

      With the exception of the alkali metal sulfites, these salts are relatively insoluble. The HSO₃⁻ ion has an interesting structure in that the hydrogen atom is bonded to the sulfur atom and not to the oxygen atom, as might be expected. There is some suggestion that in solution both the sulfur-hydrogen and oxygen-hydrogen structures may exist in equilibrium with one another, but there is no concrete evidence for this phenomenon. Heating solid hydrogen sulfite salts (shown by the equation above) or passing gaseous sulfur dioxide into their aqueous solutions produces disulfites.

HSO₃⁻(aq) + SO₂ → HS₂O₅⁻(aq)

Disulfite ions possess a sulfur-sulfur bond and are therefore unsymmetrical. Addition of acid to the solution of HS₂O₅⁻ above does not produce “disulfurous acid” (H₂S₂O₅) but instead regenerates HSO₃⁻ and SO₂. “Sulfurous acid” solutions can be oxidized by strong oxidizing agents, and oxygen in the air slowly oxidizes the solution to the more stable sulfuric acid.

2H₂SO₃ + O₂ + 4H₂O → 4H₃O⁺ + 2SO₄²⁻

Likewise, solutions of sulfites are susceptible to air oxidation to produce solutions of sulfates. Sulfites and hydrogen sulfites are moderately strong reducing agents. For example, the reaction with iodine (I₂) is quantitative (i.e., proceeds nearly to completion) and can be used in volumetric analysis.

HSO₃⁻ + I₂ + H₂O → HSO₄⁻ + 2H⁺ + 2I⁻

Sodium sulfite is used in the paper-pulp (paper pulp) industry and as a reducing agent in photographic film development.

Carbonic acid and carbonate salts

      Carbonic acid (H₂CO₃) is formed in small amounts when its anhydride, carbon dioxide (CO₂), dissolves in water.

CO₂ + H₂O ⇌ H₂CO₃

The predominant species are simply loosely hydrated CO₂ molecules. Carbonic acid can be considered to be a diprotic acid from which two series of salts can be formed—namely, hydrogen carbonates (carbonate), containing HCO₃⁻, and carbonates, containing CO₃²⁻.

H₂CO₃ + H₂O ⇌ H₃O⁺ + HCO₃⁻

HCO₃⁻ + H₂O ⇌ H₃O⁺ + CO₃²⁻

However, the acid-base behaviour of carbonic acid depends on the different rates of some of the reactions involved, as well as their dependence on the pH of the system. For example, at a pH of less than 8, the principal reactions and their relative speed are as follows:

CO₂ + H₂O ⇌ H₂CO₃ (slow)

H₂CO₃ + OH⁻ ⇌ HCO₃⁻ + H₂O (fast)

Above pH 10 the following reactions are important:

CO₂ + OH⁻ ⇌ HCO₃⁻ (slow)

HCO₃⁻ + OH⁻ ⇌ CO₃²⁻ + H₂O (fast)

Between pH values of 8 and 10, all the above equilibrium reactions are significant.

Carbonate and hydrogen carbonate salts

      These salts can be prepared by the reaction of carbon dioxide with metal oxides and metal hydroxides, respectively.

CO₂ + O₂ → CO₃²⁻

CO₂ + OH⁻ → HCO₃⁻

For example, when an aqueous solution of sodium hydroxide (NaOH) is saturated with carbon dioxide, sodium hydrogen carbonate, NaHCO₃, is formed in solution.

Na⁺ + OH⁻ + CO₂ → Na⁺ + HCO₃⁻

When the water is removed, the solid compound is also called sodium bicarbonate, or baking soda. When baking soda is used in cooking and, for example, causes bread or cake to rise, this effect is due to the reaction of the basic hydrogen carbonate anion (HCO₃⁻) with an added acid, such as potassium hydrogen tartrate (cream of tartar), KHC₄H₄O₆, or calcium dihydrogen phosphate, Ca(H₂PO₄)₂. As long as the soda is dry, no reaction occurs. When water or milk is added, the acid-base neutralization takes place, producing gaseous carbon dioxide and water. The carbon dioxide becomes trapped in the batter, and when heated the gas expands to create the characteristic texture of biscuits and breads.

      Carbonates are moderately strong bases. Aqueous solutions are basic because the carbonate anion can accept a hydrogen ion from water.

CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻

Carbonates react with acids, forming salts of the metal, gaseous carbon dioxide, and water. This is the reaction that occurs when an antacid containing the active ingredient calcium carbonate (CaCO₃) reacts with stomach acid ( hydrochloric acid).

CaCO₃ + 2HCl → CaCl₂ + CO₂ + H₂O

The hydrogen carbonate anion is also a base.

HCO₃⁻ + H₃O⁺ → H₂CO₃ + H₂O → CO₂ + 2H₂O

It is actually stronger as a base than it is as an acid. Because of this, aqueous solutions of salts of hydrogen carbonates are weakly alkaline (basic) and are also active ingredients in many antacids.

HCO₃⁻ + H₂O ⇌ H₂CO₃ + OH⁻

If equivalent amounts of sodium hydroxide and a solution of sodium hydrogen carbonate are combined and the solution is then evaporated, crystals of a hydrated form of sodium carbonate are formed. This compound, Na₂CO₃ · 10H₂O, is sometimes called washing soda. It can be used as a water softener because it forms insoluble carbonates—for example, calcium carbonate—which can then be filtered from the water. Gently heating the hydrated sodium carbonate produces the anhydrous compound Na₂CO₃, which is called soda ash or, simply, soda in the chemical industry. This is an important industrial chemical that is used extensively in the manufacture of other chemicals, glass, soap (soap and detergent), paper and pulp, cleansers, and water softeners and in the refining of petroleum.

      An interesting use of lithium carbonate, Li₂CO₃, stems from the discovery that small doses of the salt, orally administered, are an effective treatment for manic-depressive psychoses (psychosis). It is not entirely understood how this treatment works, but it is almost certainly related to the effect of the Li⁺ ion on the Na⁺:K⁺ or the Mg²⁺:Ca²⁺ balance in the brain.

  The mineral calcium carbonate is better known as limestone, a mineral second in abundance only to the silicate-forming minerals in the Earth's crust. Most limestone is composed of calcite, which is the low-temperature form of calcium carbonate. Calcite results when CaCO₃ is precipitated below 30 °C (86 °F). The calcium carbonate that precipitates above 30 °C (the high-temperature form) is known as aragonite. Transparent calcite, sometimes called Iceland spar, has the unusual property of birefringence, or double refraction. That is to say, when a beam of light enters a single crystal of calcite, the beam is broken into two beams, and two images of any object viewed through the crystal are produced.

      When water containing carbon dioxide comes in contact with limestone rocks, the rocks dissolve because Ca(HCO₃)₂, a water-soluble compound that forms aqueous Ca²⁺ and HCO₃⁻ ions, is formed. Thus, this reaction is responsible for the formation of the caves (cave) that are often found in limestone rock beds. On the other hand, if water containing Ca(HCO₃)₂ liberates carbon dioxide, calcium carbonate may again be deposited.

Ca(HCO₃)₂ (aqueous) → CaCO₃ + CO₂ + H₂O

These depositions of calcium carbonate are what are known as stalactites and stalagmites (stalactite and stalagmite), beautiful formations found in caves and caverns. When aqueous solutions of Ca(HCO₃)₂ (a form of hard water) are heated, precipitates of calcium carbonate in the form of scale (crust) are often observed in pots, boilers, and other vessels containing these solutions. Thus, one method for eliminating the hardness of water is to precipitate aqueous Ca²⁺ and HCO₃⁻ ions as solid CaCO₃, which can then be removed.

Other carbonic acids

      Two other carbon-containing acids are sometimes referred to as carbonic acids. formic acid (HCOOH) is the acid that formally has carbon monoxide (CO) as its acid anhydride. This acid has a low solubility in water. As noted previously, carbon suboxide, C₃O₂, is the acid anhydride of malonic acid, CH₂(COOH)₂, which is considered by some to be a carbonic acid.

Steven S. Zumdahl

* * *