Some acids and bases you should know

On this page:

• 1 - The mineral acids
• Sulfuric acid
• Nitric acid
• Hydrochloric acid
• 2 - The alkalies
• Sodium hydroxide
• Sodium carbonate
• Ammonia
• 3 - Some important organic acids
• Acetic acid
• Formic acid
• Oxalic acid
• Lactic acid
• Citric acid
• Ascorbic acid
• Salicylic acid
• Fatty acids
• Amino acids
• What you should be able to do
• Concept map

 

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) 3	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 chemical science, business, or economics should know about.

This term can refer to any inorganic acid, but its common use is usually limited to the major strong acids of industry, 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.

What you should know about it

• Pure anhydrous H₂SO₄ is a dense, viscous liquid which melts at 10.4°C and boils at about 300°C, decomposing back into its constituents, H₂O and SO₃.
• The acid undergoes autoprotolysis

2 H₂SO₄ → H₃SO₄⁺ + HSO₄^–

• Its high boiling point makes the acid ideal for making other acids, such as nitric and hydrochloric, which are more volatile; removal of the gaseous product drives the reaction to the right as predicted by the LeChâtelier Principle:

H₂SO₄ + 2 NaCl → 2 Na⁺ + SO₄^2– + HCl(g)

• Sulfuric acid has a voracious appetite for water, and thus is an excellent dehydrating agent. This is seen most spectacularly if some concentrated acid is poured onto a small pile of table sugar; after a short time, a vigorous reaction ensues resulting in a pile of porous, steaming carbon

C₁₂H₂₂O₁₁(s) → 12 C(s) + 11 H₂O

• Sulfuric acid can even dehydrate itself!

H₂SO₄ + H₂SO₄ → HS₂O₇^– + H₃O⁺

• The super-dehydrating ability of sulfuric acid is due to the stability of the many sulfuric acid hydrates H₂SO₄·(H₂O)_n , some of which have been detected on the surface of Jupiter's moon Europa.
• Owing to the autoprotolysis and self-dehydration reactions described above, "pure" sulfuric acid contains at least six minority species in addition to H₂SO₄.

How it is made

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(g)

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:

• All chemical reactions take place more rapidly at higher temperatures, but because this reaction is highly exothermic, raising the temperature decreases the yield.
• Because of the decrease in volume (1.5 moles of gases to 1 mole), raising the pressure will increase the yield, so the reaction is carried out at a temperature below 600°C but at very high pressure.
• Specialized catalysts are used to speed up the reaction at these lower temperatures.
• Dissolving the SO₃ directly in water would release large amounts of heat, creating a a mist of fine acid droplets that would escape into the atmosphere. The SO₃ is instead dissolved in sulfuric acid to form pyrosulfuric acid or oleum, sometimes known as fuming sufuric acid:

H₂SO₄(l) + SO₃(g) → H₂S₂O₇(l)

• The oleum is then treated with water to form industrial grade (96-98%) sulfuric acid:

H₂S₂O₇(l) + H₂O(l) → 2 H₂SO₄(l)

What it is used for

Sulfuric acid has a broad spectrum of industrial uses, and the annual tonnage follows the economic cycle quite closely.

• Sixty percent of worldwide product goes into the manufacture of phosphoric acid H₃PO₄ which is used to make phosphate fertilizers and phosphate-based household detergents.
• An important nitrogen fertilizer, ammonium sulfate (NH₄)₂SO₄, is made by reacting sulfuric acid with ammonia, the latter often obtained from the thermal decomposition of coal.
• Sulfuric acid is the major component of pickling acid that is used to remove surface oxide scale from steel before it is fabricated into steel product for the automobile and other industries.
• Aluminum sulfate (made from bauxite Al₂O₃ with H₂SO₄), is widely used in the papermaking industry to coagulate the cellulose fibers, producing a smooth, hard paper surface. Another major use is to make aluminum hydroxide which is used to filter out particulate matter in water treatment facilities.
• The familiar lead-acid storage battery employs sulfuric acid as an electrolyte. As the battery discharges, the concentraton of sulfuric acid in the electrolyte decreases as sulfate ions are taken up as PbSO₄. Owing to the high density of the acid, the state of charge of the battery can be measured by means of a hydrometer.

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 SO₄^2– + 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 slightly 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, produces brown nitrogen dioxide, which accounts for the color of HNO₃ solutions.

What you should know about it

• HNO₂ undergoes autoprotolysis to a greater extent than any other liquid. Further reactions of its conjugate acid H₂NO₃⁺ with HNO₃ lead to a complicated mixture of species in the liquid.
• Dilute nitric acid can be concentrated by distillation up to a maximum of 68%, at which point it forms a constant-boiling (azeotropic) mixture with water. Higher concentrations require dehydration with sulfuric acid; the result is fuming nitric acid.
• "Concentrated nitric acid" is sold as a 70% solution in water, corresponding to a concentration of about 16M.
• Nitric acid is a very strong oxidizing agent, which adds to its corrosive behavior with organic materials including, of course, skin, which it turns yellow owing to a reaction with the protein keratin. Reactions with many organic compounds are highly exothermic and often violent. The well-known reaction of nitric acid with metallic copper produces copious amounts of brown nitrogen dioxide gas.

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How it is made

The simplest method, which was used industrially before 1900, was by treatment of sodium nitrate ("Chile saltpetre", NaNO₃) with sulfuric acid.

The direct synthesis of the acid from atmospheric nitrogen and oxygen is thermodynamically favorable

½ N₂ + 5/4 O₂ + ½ H₂O → HNO₃

An interesting fact: 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.

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.

What it is used for

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.

Nitrogen oxides in the environment

"Nitric Acid Acts Upon Trousers" - a delightful pre-1900 account of chemistry professor Ira Remsen's first encounter with nitric acid.

High-temperature combustion processes (in internal combustion engines, power plants, and incinerators) can oxidize atmospheric nitrogen to nitric oxide (NO) and other oxides ("NO_x"); 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.

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.

What you should know about it

• Hydrochloric acid is the least hazardous of the strong mineral acids to work with because unlike the others, it is not an oxidizing agent. It is usually the acid of choice for titrations and other operations in which the main requirement is simply a strong source of hydronium ions.
• The concentrated acid boils at 48°C. As boiling continues, it loses HCl and the boiling point rises to 109°C, at which point a constant-boiling (azeotropic) solution remains, consisting of 20.2% HCl.

What it is used for

The uses of hydrochloric acid are far too many to enumerate individually, but the following stand out:

• A major industrial use is to remove surface scale from iron or steel ("pickling") before it is processed into sheets or other forms, or galvanized or coated.
• Production of chlorinated organic chemicals, particularly vinyl chloride, polyurethanes, and other construction polymers, consumes huge amounts of HCl.
• The acid is widely used for pH control of water, including neutralization of alkaline wastwater streams, and for regenerating ion-exchange water softeners.

How it is made

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 metallic element of Group 1 or 2 ("alkali" or "alkaline earth".) 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.

What you should know about it

• Owing to its deliquescence (ability to take up moisture) and its tendency to react with carbon dioxide, the solid must be stored in a closed container.
• In industry, sodium hydroxide is commonly known as caustic soda or simply as caustic; NaOH sold for household purposes is usually known as lye.
• Sodium hydroxide slowly attacks glass to form sodium silicate. Glass vessels used to store concentrated solutions gradually develop a cloudy coating on the inside.
• Some metals, notably aluminum, zinc, and titanium, react with strongly alkaline solutions, but iron and copper are immune to this kind of attack.
• Highly alkaline solutions also soften and dissolve skin, accounting for the slippery feeling associated with strong bases. Sodium hydroxide was once used to dispose of animal carcasses, digesting them into an easily disposable liquid form.

How it is made

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:

[illustration link]

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!)

What it is used for

• Sodium hydroxide is one of the most diverse industrial chemicals in terms of its applications. Most householders know it as the active ingredient of drain cleaning agents.
• Huge quantities are consumed by the pulp and paper industry, which is probably its single largest specific industrial application. It is used to remove the liqnin component of wood pulp from the cellulose so that the latter can be processed into paper.
• About half of the NaOH output goes into the prouction of a wide variety of other industrial chemicals, and in degreasing steel drums and other industrial surfaces.
• "Lye" plays a role in the processing of many types of foods, including chocolate, olives, pretzels, and the "hominy" and "grits" corn products used in the Southern U.S. It also acts as a chemical peeling agent for fruits and vegetables.

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.

What it is used for

• Although carbonates are much weaker bases than hydroxides, a solution of sodium carbonate can still have a pH of 11 or so, sufficiently high to allow it to substitute for sodium hydroxide in many applications— especially when the price of caustic is high.
• The single most important use of soda ash is in the manufacture of glass, where it serves to lower the melting point of the principal component, SiO₂.
• Another emerging major use is to neutralize the SO₂ emissions of fossil fuel-burning power plants.
• The older use of sodium carbonate as a cleaning agent (hence the name washing soda) was based partly on the ability of its alkaline solutions to emulsify grease, but mainly as a means of precipitating the insoluble carbonates of calcium and magnesium before these ions (commonly present in hard water) could form undesirable preciptates with soaps. The use of modern detergents has largely eliminated this once important market.

How it is made

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.

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.

What you should know about it

• More moles of ammonia are manufactured than of any other industrial chemical.
• Ammonia is extremely soluble in water. An aqueous solution of ammonia is still sometimes referred to in commerce as "ammonium hydroxide", but this term is no longer favored by chemists because no such compound as NH₄OH has ever been shown to exist. At neutral pH, about 99% of the ammonia in water exists as NH₄⁺ ions.
• Ammonia is an end product of nitrogen metabolism in most organisms. One source that may be familiar to parents of infants is the bacterial decomposition of the contents of diapers.
• Liquid ammonia (bp –33°C) is often used as an ionizing laboratory solvent.

What it is used for

• The major use of ammonia (about 80%) is as a fertilizer, most commonly as anhydrous ammonia (the gas is injected directly into the soil) or after conversion to (NH₄)₂SO₄, NH₄NO₃, or urea O–C(NH₃)₂.
• Ammonia is used in the production of numerous polymers, including nylons and polyurethanes.
• Explosives manufacture accounts for about 5% of ammonia production.
• Beyond these, there are hundreds of minor uses, including as a household cleaning agent (aqua ammonia). a refrigerant, and as a laboratory reagent.

How it is made

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.

What you should know about it

• The pure acid, although quite weak in the proton-donor sense, is quite corrosive and its vapors are very irritating.
• A 1.0M solution of the acid has a pH of about 2.4, corresponding to only four out of every thousand CH₃COOH molecules being dissociated.

What it is used for

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.

How it is made

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

 

A brief look at some other interesting organic acids

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.

• Describe some of the special properties of sulfuric acid that make it especially important both in the laboratory and in industry.
• Name the major acids and bases that are important to the fertilizer industry
• Name the natural sources of any three of the major organic acids.
• What are fatty acids? In what major way do the physical properties of saturated and unsaturated fatty acids differ?
• Describe the general structure of an amino acid, and state why they are important.