Chem1 General Chemistry Virtual Textbook → Acid-base concepts → Gallery
A gallery of some acids and bases
you should know
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Acids and bases are of interest not only to the chemically inclined; they play a major role in our modern industrial society — so anyone who participates in it, or who is interested in its history and development, needs to know something about them. Five of the major acids and bases fall into the "Top 20" industrial chemicals manufactured in the world. The following table shows year-2000 figures for the U.S:
chemical and rank | Sulfuric acid - 1 |
Lime (CaO) |
Phosphoric acid - 4 |
Ammonia 5 |
Sodium hydroxide - 9 |
Nitric acid - 11 |
production in 109 kg | 40 | 20 | 16 | 15 | 11 | 8 |
major use | chemicals | cement | fertilizers | fertilizers | chemicals | chemicals |
In this unit, we will survey some of the major acids and bases that anyone who is engaged in chemistry, business, or economics should know about.
This term refers to any inorganic acid, but its common use is usually limited to the major strong acids plus phosphoric acid.
The major mineral acids— sulfuric, nitric, and hydrochloric— have been known since medieval times. Their discovery is usually credited to the Persian alchemist Abu Musa Jabir ibn Hayyan, known in the West by his Latinized name Geber. Jabir also invented aqua regia, the mixture of nitric and hydrochloric acids that has the unique ability to dissolve gold.
More sulfuric acid is manufactured than any other industrial chemical, and it is the cheapest industrial acid worldwide. It has been continuously manufactured in the U.S. since 1793 and in Europe for much longer.
2 H2SO4 → H3SO4+ + HSO4–
H2SO4 + 2 NaCl → 2 Na+ + SO42– + HCl(g)
C12H22O11(s) → 12 C(s) + 11 H2O
H2SO4 + H2SO4 → HS2O7– + H3O+
Sulfur trioxide, the anhydride of sulfuric acid, is the immediate precursor. Gaseous SO3 reacts vigorously with water, liberating much heat in the process:
SO3(g) + H2O(l) → H2SO4(l)
Industrial manufacture of the acid starts with sulfur dioxide, prepared from burning elemental sulfur or obtained as a byproduct from roasting sulfide ores. The oxidation of SO2 to SO3 looks simple
SO2(g) + ½ O2(g) → SO3(g)
but there are several complications:
H2SO4(l) + SO3(g) → H2S2O7(l)
H2S2O7(l) + H2O(l) → 2 H2SO4(l)
Sulfuric acid has a broad spectrum of industrial uses, and the annual tonnage follows the economic cycle quite closely.
Acid Rain - Combusion of fossil fuels which contain organic sulfur compounds releases SO2 into the atmosphere. Photochemical oxidation of this compound to SO3, which rapidly takes up moisture, leads to the formation of H2SO4, a major component of acid rain.
Acid mine drainage results when sediments of the very common iron pyrite, FeS2, are exposed to air and are oxidized:
FeS2(s) + 7/2 O2 + H2O → Fe2+ 2 SO42– + 2 H+
further oxidation of the iron to Fe3+ results in additional reactions. The resulting drainage liquid is often orange-brown in color and can a have a pH of below zero.
Anhydrous HNO3 is a colorless liquid boiling at 82.6°C, but "pure" HNO3 only exists as the solid which melts at –41.6°C. In its liquid and gaseous states, the acid is always partially decomposed into nitrogen dioxide:
2 HNO3 → 2 NO2 + ½ O2 + H2O
This reaction, which is catalyzed by light, accounts for the brownish color of HNO3 solutions.
The simplest method, which was used industrially before 1900, was by treatment of sodium nitrate ("Chile saltpetre", NaNO3) with sulfuric acid.
If it were not for the high activation energy required to sustain this reaction, all of the oxygen in the atmosphere would be consumed and the oceans would be a dilute solution of nitric acid.
The direct synthesis of the acid from atmospheric nitrogen and oxygen is thermodynamically favorable
½ N2 + 5/4 O2 + ½ H2O → HNO3
but is kinetically hindered by an extremely high activation energy, a fact for which we can be most thankful (see sidebar.) The first industrial nitrogen fixation process, developed in 1903, used this reaction to produce nitric acid, but it required the use of an electric arc to supply the activation energy and was therefore too energy-intensive to be economical.
The modern Ostwald process involves the catalytic oxidation of ammonia to nitric oxide NO, which is oxidized in a further step to NO2; reaction of the latter with water yields HNO3. This route, first developed in 1901, did not become practical until the large-scale production of ammonia by the Haber-Bosch process in 1910.
The major industrial uses of nitric acid are for the production of ammonium nitrate fertilizer, and in the manufacture of explosives. On a much small scale, the acid is used in metal pickling, etching semiconductors and electronic circuit boards, and in the manufacture of acrylic fibers.
In the laboratory, the acid finds use in a wide variety of roles.
High-temperature combustion processes (in internal combustion engines, power plants, and incinerators) can oxidize atmospheric nitrogen to nitric oxide (NO) and other oxides ("NOx"); the NO is then photooxidized to NO2, which reacts with water to form HNO3 which is a major component of acid rain. NO2 is the major precursor of photochemical smog.
"Nitric Acid Acts Upon Trousers" - a delightful pre-1900 account of chemistry professor Ira Remsen's first encounter with nitric acid.
Hydrochloric acid is still sometimes sold under its older name muriatic acid for cleaning bricks and other household purposes. The name comes from the same root as marine, reflecting its preparation from salt.
Unlike the other major acids, there is no such substance as "pure" hydrochloric acid; what we call "hydrochloric acid" is just an aqueous solution of hydrogen chloride gas (bp –84°C). But in a sense it is more "pure" than the acids discussed above, since there is no autoprotolysis; hydronium and chloride ions are the only significant species in the solution.
Hydrochloric acid is usually sold as a 32-38% (12M) solution of HCl in water; concentrations greater than this are known as fuming hydrochloric acid.
The acid has been known to chemists (and alchemists), and used for industrial purposes since the middle ages. Its composition HCl was demonstrated by Humphrey Davy in 1816.
The uses of hydrochloric acid are far too many to enumerate individually, but the following stand out:
The ancient method of treating salt with sulfuric acid to release HCl has long since been supplanted by more efficient processes, including direct synthesis by "burning" hydrogen gas in chlorine:
H2(g) + Cl2(g) → 2 HCl(g)
Most hydrochloric acid production now comes from reclaiming byproduct hydrogen chloride gas from other processes, especially those associated with the production of industrial organic compounds.
The term alkali usually means a basic salt of a Group 1 or 2 ("alkali" or "alkaline earth") metal. All alkalies are of course bases, but the latter term is much more general, whether defined according to the Arrhenius, Brønsted-Lowry, or Lewis concepts.
The word alkali comes from the Arabic al-qali, which refers to the ashes from which sodium and potassium hydroxides (potash, "ashes remaining in the pot", and the origin of the element name potassium) were extracted as a step in the making of soap.
Pure sodium hydroxide is a white solid consisting of Na+ and OH– ions in a crystal lattice. Although it is widely thought of as an ionic solid, van der Waals forces make a substantial contribution to its stability.
Sodium hydroxide is now manufactured by the electrolysis of brine solutions, and along with chlorine, is one of the two major products of the chloralkali industry.
Electrolysis of aqueous NaCl produces Cl2 at the anode, but because H2O can be reduced more readily than Na+, the water is decomposed to H2 and OH– at the cathode, leaving a solution of NaOH. An older mercury cell process reduces the Na+ to Na within a mercury amalgam (alloy), and the metallic sodium is then combined with water to produce NaOH and hydrogen. The net reaction for the reduction step is the same for both methods:
2 Na+ + 2 H2O + 2e– → H2(g) + 2 NaOH
The resulting solution is usually evaporated to such a high concentration that it solidifies at ordinary temperatures. It is commonly shipped in rail cars or barges which can be heated with steam to liquify the mixture for removal. (It is obviously uneconomical to ship large quantities of water across the country!)
The economic push and pull of caustic and chlorine
In contrast to the extremely diverse applications of sodium hydroxide which makes the demand for this commodity relatively immune to the ups and downs of the economic cycle, the consumption of chlorine is directly dependent on the economy as reflected in the demand for polyvinyl chloride products that are now widely used in the the construction and home furnishings industries. Because chlorine, being a gas, is expensive to store, the output of the chloralkali industry is governed largely by the demand for this commodity. When times are good this presents no problem; caustic is then largely a by-product and can easily be stockpiled if supply exceeds demand. But during an economic downturn, the demand for chlorine declines, limiting its production along with that of caustic. But because the demand for caustic tends to decline much less, it becomes scarce and its price rises, thus tending to drive the industrial economy into even deeper trouble.
This compound is known industrially as soda ash, and domestically as washing soda. The common form is the heptahydrate, Na2CO3·7 H2O. The white crystals of this substance spontaneously lose water (effloresce) when exposed to the air, forming the monohydrate.
Most of the world's sodium carbonate is made by the ammonia-soda ("Solvay") process developed in 1861 by the Belgian chemist Ernest Solvay (1838-1922) whose patents made him into a major industrialist and a rich philanthropist. This process involves a set of simple reactions that essentially converts limestone (CaCO3), ammonia NH3 and brine (NaCl) into sodium bicarbonate NaHCO3 and eventually Na2CO3, recycling several of the intermediate products in an ingenious way.
A minor source of soda ash (but quite significant in some countries, such as the U.S.) is the mining of natural evaporites (the remains of ancient lakes), such as the trona found in Southern California.
Tradition dies slowly:
a non-existent chemical
available in bottles!
Ammonia NH3 is of course not a true alkali, but it is conveniently included in this section for discussion purposes. Most people are familiar with the pungent odor of this gas, which can be detected at concentrations as low as 20-50 ppm.
Ammonia is made by direct synthesis from the elements:
N2(g) + 3 H2(g) → 2 NH3(g)
... a simple-looking reaction, but one that required some very creative work to implement; the Haber-Bosch process is considered to be the most important chemical synthesis of the 20th Century. See here for more on the theory, history, and social impact of this process.
Most acids are organic— there are millions of them. The acidic function is usually a hydroxyl group connected to a carbon that is bonded to an electron-withdrawing oxygen atom; the combination is the well-known carboxyl group, –COOH. Here are a few that are part of everyone's life.
This is next to formic acid in being the simplest of the organic acids, and in the form of vinegar (a 5-8% solution in water) its characteristic odor is known to everyone. The pure acid is a colorless liquid above 16.7°C ; below this temperature it forms a crystalline solid, hence the term "glacial acetic acid" that is commonly applied to the pure substance. The name of the acid comes from acetum, the Latin word for vinegar.
Slightly less than half of the world production of acetic acid goes into the production of polymers. The end product visible to most people would be the flexible plastic bottles in which drinking water is sold. Other uses are related mostly to the production of other chemicals, mainly acetic anhydride, but also including aspirin.
Bacterial fermentation of sugars has been the source of vinegar since ancient times, and is still accounts for most food-grade acetic acid and vinegar, but it now amounts to only about 10% of total acetic acid production:
C6H12O6 → 3 CH3COOH
There are several important synthetic routes to acetic acid production, but the major one is by treating methanol with carbon monoxide:
CH3OH + CO → CH3COOH
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This, the simplest of the carboxylic acids, is the chemical weapon that Nature has given ants and bees (the Latin word for ant is formicus.) Known since the 15th Century, it was first distilled from ants, but is now made synthetically. Its main uses are as a preservative and antibacterial agent in livestock feed. Chemists find it a useful source of carbon monoxide (just add sulfuric acid.) |
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Just two carboxyl groups joined together, this acid has a special ability to grab up divalent metal ions and bind them into five-membered MO2C2 rings. Because many such metal ions (Ca2+, for example) are essential nutrients, this can be dangerous to your health. The acid occurs in many plants, notably rhubarb (it is what makes the leaves poisonous), parsley and spinach. If you have ever noticed a funny feeling in your mouth after drinking milk with a rhubarb desert, it is due to precipitation of calcium oxalate. This same solid is often a major component of kidney stones, and it contributes to the miseries of gout. |
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Structurally, lactic acid is both an alcohol and a carboxylic acid— not all that uncommon. We know it as the acid found in sour milk, yogurt, and in tired muscles. When the blood cannot deliver enough oxygen to oxidize glucose all the way to CO2 and water, your muscles go into a much-less-energy-efficient "anaerobic" mode that generates lactate. In milk products, lactic acid is produced in about the same way by acidophilus bacteria that eat the lactose ("milk sugar") and who don't know how to utilize oxygen. |
Citric acidKnown by Arabic alchemists in the 8th Century; first isolated by Scheele in 1784. Wikipedia article
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Although it is one part alcohol and three parts acid, citric acid is quite weak, but nevertheless strong enough to make it unpleasant to suck on a lemon, of which it can comprise as much as 8% of the dry weight of this fruit. Biochemists know it as part of the citric acid cycle, a sequence of reactions involved in extracting energy from the oxidative metabolism of foods. Its major industrial use is as a food flavoring and preservative agent; large quantities are used to make soft drink beverages. The acid also finds use in cleaning agents. |
Ascorbic acid |
The major industrial use of ascorbic acid is as an antioxidant, so it is often added to foods and other materials as a preservative. |
Salicylic acid
2-hydroxybenzoic acid is found in willow trees |
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Fatty acidsThis generic term refers to carboxylic acids built from a chain of 4 to 22 carbon atoms. Fatty acids, as the name implies, are derived from fats, in which they are bound to glycerol in the form of triglycerides. Fats, which occur in all animal tissues, are the most efficient means of storing metabolic energy. Vegetable oils are another source.
Saturated fatty acids consist of straight chains of carbon atoms with a methyl group on one end and a carboxyl (acidic) group on the other, with methylene groups -CH2- in between. Thus stearic (octadecanoic) acid is CH3(CH2)16COOH. The linear shape of these molecules allows them to pack together very efficiently, contributing to their ability to store energy in a small space; the higher saturated acids tend to be waxy solids. Unsaturated fatty acids contain one or more carbon-carbon double bonds, which introduces a complication: because free rotation around the axis of a double bond is not possible, the neighboring hydrogens can be either on the same side of the bond (cis) or on opposite sites (trans)— thus the cis- and trans-fatty acids. These double bonds introduce kinks into the carbon chain, especially in the case of the cis-acids. The bent and curved carbon chains that result cannot pack together compactly enough to interact very strongly with each other, so unsaturated fatty acids tend to be liquids. The human body is able to synthesize most of the fatty acids it needs, but two classes of unsaturated acids, known as the essential fatty acids, can be obtained only from foods. The compounds are known as ω-3 and ω-6 fatty acids; the ω (omega)-n notation means that the double bond is located n carbons away from the terminal CH3 group of the molecule.
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Amino acids
Representation of an enzyme protein showing its complex folding [link] |
All amino acids have the basic structure shown above; they differ in the nature of the group of atoms designated by "R" in the diagram. Although we call them "acids", the amino acids are really chemical hermaphrodites; you will recall that amines are weak bases. The balance between their acidic and basic properties can be shifted simply by changing the pH.
The polymeric chains that result are known as peptides if they are fairly short (2 to about 20-50 amino acid units); longer chains, or aggregate made up of multiple peptide units, are proteins. Owing to their very large size (500 amino acid residues is quite common, but some have as many as 1500), proteins are able to fold in various ways, so the amino acid sequence alone is not sufficient to determine their properties. |
Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially imortant that you know the precise meanings of all the highlighted terms in the context of this topic.
Page last modified: 09.09.2010