oxyacids and their salts

Most covalent nonmetallic oxides react with water to form acidic oxides. In other words, they react with water to form oxyacids that yield hydronium ions (H3O¿) in solution. Some exceptions include carbon monoxide, CO, nitrous oxide, N2O, and nitric oxide, NO. An oxyacid’s strength is defined by the extent to which it dissociates in water (i.e., its ability to form H¿ ions). In general, the electronegativity and oxidation number of the central nonmetal atom is the basis for predicting the relative strength of oxyacids. For a given nonmetal central atom, the acid strength increases as the electronegativity of the central atom increases. Because the electronegativity of chlorine (Cl), for example, 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. Similarly, 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, a type of reaction called neutralization, because the solution is made neutral. In general, the salts of all oxyacids are more stable than the acids themselves. Nitric acid, known to the alchemists of the 8th century as "aqua fortis" (strong water), is formed by the reaction of both dinitrogen pentoxide (N2O5) and nitrogen dioxide (NO2) with water. Small amounts of nitric acid are found in atmosphere after thunderstorms, and its salts, called nitrates, occur widely in nature. Vast deposits of sodium nitrate, NaNO3, also known as Chile saltpeter, are found in the desert region near the boundary of Chile and Peru. Potassium nitrate, KNO3, sometimes called Bengal saltpeter, is found in India and other countries in East Asia. Pure nitric acid, a colorless liquid, boils at 86° C and freezes at -42° C. Exposure to light or heat causes it to decompose to produce oxygen, water, and a mixture of nitrogen oxides (primarily NO2). Because of the NO2 that forms as it decomposes, nitric acid is often yellow or brown in color. Nonmetallic elements such as carbon (C), iodine (I), phosphorus (P), and sulfur (S), as well as many compounds, are oxidized by concentrated HNO3 to their oxides or oxyacids with the formation of NO2. 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, and was used by alchemists to dissolve gold as early as the 13th century. Such relatively unreactive metals as copper (Cu), silver (Ag), and lead (Pb) reduce concentrated HNO3 primarily to NO2. The reaction of dilute HNO3 with copper produces NO, while 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, important in the manufacture of explosives, dyes, plastics, and drugs, is used extensively in the laboratory and in chemical industries as a strong acid and as an oxidizing agent. It reacts with proteins, 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, or carbonates react with nitric acid. A major use of nitric acid is to produce soluble metal nitrates. All nitrates decompose when heated and may do so explosively. Nitrates are valuable as fertilizers. Gunpowder is a mixture of potassium nitrate, sulfur, and charcoal. Ammonal, an explosive, is a mixture of ammonium nitrate and aluminum powder. Nitrous acid, a weak acid, is very unstable and exists only in aqueous solution, decomposing slowly at room temperature—and more rapidly at elevated temperatures—to nitric acid and nitric oxide. Oxidized to nitric acid by active oxidizing agents, HNO2 acts as an oxidizing agent with strong reducing agents. Sodium nitrite (NaNO2), typically prepared by reducing molten sodium nitrate with elemental lead, is an important example of a nitrite—that is, a salt of nitrous acid. Added to meats, such as hot dogs, it prolongs the meat's retention of a red color, and it inhibits the growth of bacteria that can cause food poisoning. However, 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 to form nitrosamines. Because some of the compounds in the nitrosamine class are known to cause cancer in laboratory animals, the United States Food and Drug Administration limits the amount of sodium nitrite that can be legally added to foods. Most nitrites, which are much more stable than nitrous acid, are soluble in water and, like nitrates, can explode upon heating or detonation. Orthophosphoric acid (H3PO4), usually called simply phosphoric acid, is a colorless, crystalline solid when pure that melts at 42° C. It rapidly absorbs moisture from the air and liquefies. Produced in pure form by dissolving phosphorus pentoxide (P4O10) in water, it is typically available commercially as syrupy phosphoric acid, which is an 85 percent solution in water. More commonly, H3PO4 is prepared by treating calcium phosphate, Ca3(PO4)2, with concentrated sulfuric acid, H2SO4. A triprotic acid—i.e., it can donate all three of its hydrogen atoms as protons in aqueous solution—it can thus form three series of salts: dihydrogen phosphates, containing the H2PO4 - ion; hydrogen phosphates, containing the HPO42- ion; and orthophosphates, containing the PO43- ion. Because H2PO4- is a weak acid, soluble dihydrogen phosphate salts form solutions that are weakly acidic. Because the HPO42- ion is stronger as a base (i.e., a proton acceptor) than as an acid, aqueous solutions of hydrogen phosphates are basic. Because the PO43- ion is a moderately strong base, orthophosphate salts form strongly basic solutions. Pure phosphorous acid, H3PO4, is best prepared by hydrolysis of phosphorus trichloride, PCl3, but can also be obtained by the action of water on P4O6, PBr3, or PI3. Colorless crystalline H3PO3 melts at 70.1° C, is very soluble in water, and has an odor similar to that of garlic. Heating phosphorous acid to about 200° C causes it to disproportionate into phosphine, PH3, and orthophosphoric acid. Phosphorous acid and its salts are active reducing agents, because of their easy oxidation to phosphoric acid and phosphate salts, respectively. Thus, 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. Because in phosphorous acid only two hydrogens are bonded to oxygen and are acid, H3PO3 is diprotic, whereas H3PO4 has three hydrogen atoms bound to oxygen and is triprotic. Because the third hydrogen, bonded directly to phosphorus, is not very acidic, H3PO3 forms only two series of salts, one containing the dihydrogen phosphite ion, H2PO3- , and the other containing the hydrogen phosphite ion, HPO32-. Free hypophosphorous acid, H3PO2, is prepared by acidifying aqueous solutions of hypophosphite ions, H2PO2-. Pure hypophosphorous acid forms white crystals that melt at 26.5° C. The electronic structure of H3PO2 is such that it has only one hydrogen atom bound to oxygen, and is thus a monoprotic oxyacid. A weak acid, it forms only one series of salts, the hypophosphites. Hydrated sodium hypophosphite, NaH2PO2 9 H2O, is used as an industrial reducing agent, particularly for the electroless plating of nickel onto metals and nonmetals. Sulfuric acid, sometimes referred to as the "king of chemicals" because it is produced worldwide in such large quantities, is the cheapest bulk acid. Pure sulfuric acid, a colorless, oily, dense (1.83 grams per cubic centimeter) liquid that freezes at 10.5° C, 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. Called a constant boiling solution, this concentration is sold as concentrated sulfuric acid. A strong diprotic acid, sulfuric acid has its two hydrogen atoms bonded to oxygen, and ionizes in two stages. In aqueous solution, loss of the first hydrogen (as a hydrogen ion, H¿) is essentially 100 percent, whereas the second ionization takes place to an extent of about 25 percent. 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 virtually insoluble sulfates of the alkaline earth metals—calcium (Ca), strontium (Sr), and barium (Ba)—as well as that of lead (Pb) —are found as naturally occurring minerals, including gypsum (CaSO4 9 2H2O), celestite (SrSO4), barite (BaSO4), and anglesite (PbSO4). Some important soluble sulfate salts are Glauber's salt, Na2SO4 9 10H2O; Epsom salt, MgSO4 9 7H2O; blue vitriol, CuSO4 9 5H2O; green vitriol, FeSO4 9 7H2O; and white vitriol, ZnSO4 9 7H2O. Approximately 67 percent of the sulfuric acid produced in the United States is utilized to convert phosphate rock to phosphoric acid, which is then converted to phosphate fertilizers. 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 employed in lead storage batteries and in the production of paints, plastics, explosives, and textiles. The dissolution of sulfur dioxide in water results in an acidic solution that 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. The predominant species appear to be hydrated SO2 molecules, SO2 9 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- . Both salts are strong reducing agents. Sodium sulfite is used in the paper pulp industry and as a reducing agent in photographic film development. Carbonic acid (H2CO3) is formed in small amounts when its anhydride, carbon dioxide (CO2), dissolves in water. It can be considered to be a diprotic acid from which two series of salts can be formed—namely, hydrogen carbonates, containing HCO3- , and carbonates, containing CO32-. These salts can be prepared by the reaction of carbon dioxide with metal oxides and metal hydroxides, respectively. When an aqueous solution of sodium hydroxide (NaOH) is saturated with carbon dioxide, sodium hydrogen carbonate, NaHCO3, is formed in solution. 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. No reaction occurs as long as the soda is dry, but 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. Carbonates react with acids to form 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). Because the hydrogen carbonate anion is actually stronger as a base than it is as an acid, aqueous solutions of salts of hydrogen carbonates are weakly alkaline (basic) and are active ingredients in many antacids. If equivalent amounts of sodium hydroxide and a solution of sodium hydrogen carbonate are combined and the solution is then evaporated, the result is crystals of a hydrated form of sodium carbonate (Na2CO3 9 10H2O), a compound sometimes called washing soda. It can be used as a water softener because it forms insoluble carbonates—for example, calcium carbonate—that can then be filtered from the water. Gentle heating of the hydrated sodium carbonate produces the anhydrous compound Na2CO3, called by the chemical industry soda ash or simply soda. This important industrial chemical is used extensively in the manufacture of other chemicals, glass, soap, paper and pulp, cleansers, and water softeners and in the refining of petroleum. Small doses of lithium carbonate, Li2CO3, orally administered, are an effective treatment for manic-depressive psychoses. 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—better known as limestone, a mineral second in abundance only to the silicate-forming minerals in the Earth's crust—is usually composed of calcite, which is the low-temperature form of calcium carbonate. Calcite results when CaCO3 is precipitated below 30° C., whereas 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; i.e., 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. Another carbon-containing acid that is sometimes referred to as a carbonic acid is formic acid (HCOOH), the acid that formally has carbon monoxide (CO) as its acid anhydride. This acid has a low solubility in water. Carbon suboxide, C3O2, the acid anhydride of malonic acid, CH2(COOH)2, is considered by some to be a carbonic acid.