Sulfuric Acid

Sulfuric Acid, corrosive, oily, colorless liquid, with a specific gravity of 1.85. It melts at 10.36° C (50.6° F), boils at 340° C (644° F), and is soluble in all proportions in water. When sulfuric acid is mixed with water, considerable heat is released. Unless the mixture is well stirred, the added water may be heated beyond its boiling point and the sudden formation of steam may blow the acid out of its container (see Acids and Bases). The concentrated acid destroys skin and flesh, and can cause blindness if it gets into the eyes. The best treatment is to flush away the acid with large amounts of water. Despite the dangers created by careless handling, sulfuric acid has been commercially important for many years. The early alchemists prepared it in large quantities by heating naturally occurring sulfates to a high temperature and dissolving in water the sulfur trioxide thus formed. About the 15th century a method was developed for obtaining the acid by distilling hydrated ferrous sulfate, or iron vitriol, with sand. In 1740 the acid was produced successfully on a commercial scale by burning sulfur and potassium nitrate in a ladle suspended in a large glass globe partially filled with water.

Sulfuric acid is a strong acid, that is, in aqueous solution it is largely changed to hydrogen ions (H+) and sulfate ions. Each molecule gives two H+ ions, thus sulfuric acid is dibasic. Dilute solutions of sulfuric acid show all the behavior characteristics of acids. They taste sour, conduct electricity, neutralize alkalies, and corrode active metals with formation of hydrogen gas. From sulfuric acid one can prepare both normal salts containing the sulfate group and acid salts containing the hydrogen sulfate group.

Smog

Smog, mixture of solid and liquid fog and smoke particles formed when humidity is high and the air so calm that smoke and fumes accumulate near their source. Smog reduces natural visibility and often irritates the eyes and respiratory tract. In dense urban areas, the death rate usually goes up considerably during prolonged periods of smog, particularly when a process of heat inversion creates a smog-trapping ceiling over a city.

Smog prevention requires control of smoke from furnaces; reduction of fumes from metal-working and other industrial plants; and control of noxious emissions from automobiles, trucks, and incinerators. In the U.S. internal-combustion engines are regarded as the largest contributors to the smog problem, emitting large amounts of contaminants, including unburned hydrocarbons and oxides of nitrogen. The number of undesirable components in smog, however, is considerable, and the proportions highly variable. They include ozone, sulfur dioxide, hydrogen cyanide, and hydrocarbons and their products formed by partial oxidation. Fuel obtained from fractionation of coal and petroleum produces sulfur dioxide, which is oxidized by atmospheric oxygen, forming sulfur trioxide. Sulfur trioxide is in turn hydrated by the water vapor in the atmosphere to form sulfuric acid.

The so-called photochemical smog, which irritates sensitive membranes and damages plants, is formed when nitrogen oxides in the atmosphere undergo reactions with the hydrocarbons energized by ultraviolet and other radiations from the sun. See Air Pollution.

Camphor

Camphor, volatile, white, crystalline compound,with a characteristic aromatic odor. Ordinary camphor is obtained from the camphor tree, Cinnamomum camphora, which grows in Asia and Brazil. The camphor is distilled by steaming chips of the root, stem, or bark. The leaves of certain plants, such as tansy and feverfew, contain a second form of camphor, which is not used commercially. A racemic form is present in the oil of an Asian chrysanthemum and is also produced synthetically for most commercial uses. Camphor is used in the manufacture of celluloid and explosives and medicinally in liniments and other preparations for its mild antiseptic and anesthetic qualities. It is poisonous if ingested in large amounts.

Camphor is insoluble in water, soluble in organic solvents, and melts at 176° C (349° F) and boils at 209° C (405° F).

Ozone

Ozone (Greek ozein, “to smell”), pale blue, highly poisonous gas with a strong odor. Ozone is considered a pollutant at ground level, but the ozone layer of the upper atmosphere protects life on Earth from the Sun’s harmful ultraviolet radiation.

Ozone is one of three forms, called allotropes, of the element oxygen. Ozone is triatomic, meaning that it has three atoms in each molecule (formula O3). Ordinary, or diatomic, oxygen (O2) is more stable than ozone and accounts for the bulk of oxygen in the atmosphere. Electrical sparks and ultraviolet light can cause ordinary oxygen to form ozone. The presence of ozone sometimes causes a detectable odor near electrical outlets.

PROPERTIES

At normal temperatures and pressures ozone is a gas with a specific gravity of 2.144 (about 1.5 times the density of ordinary oxygen gas). Ozone accounts for only a tiny fraction of the atmosphere and is normally invisible, but high concentrations of ozone gas are pale blue. The gas condenses to a liquid at -111.9°C (-169.52°F) and freezes at -192.5°C (-314.5°F). Liquid ozone is deep blue, and is diamagnetic (repelled by magnetic fields). Solid ozone is dark purple. Ozone is much more active chemically than ordinary oxygen. It is used in purifying water, sterilizing air, and bleaching certain foods.

Learn more:

ENVIRONMENTAL EFFECTS

Chlorofluorocarbons (CFCs)

Chlorofluorocarbons (CFCs), family of synthetic chemicals that are compounds of the elements chlorine, fluorine, and carbon. CFCs are stable, nonflammable, noncorrosive, relatively nontoxic chemicals and are easy and inexpensive to produce. During the 1970s, scientists linked CFCs to the destruction of Earth’s ozone layer. The manufacture of CFCs has since been banned in most countries.

USES

Scientists developed the first CFCs during the late 1920s. The compounds subsequently became used in a wide range of industrial products in the United States, Europe, and Japan. Manufacturers used CFCs as refrigerants in refrigerators, freezers, air conditioners, and heat pumps, and as propellants in aerosols and medical inhalers. CFCs also served as insulating foams in packaging materials, furniture, bedding, and car seats. Cleaning agents for electronic circuit boards, metal parts, and dry cleaning processes also used CFCs.

Learn more:

• Harmful Effects of CFCs
• Regulation
• Extended Impact
  
  

Allotrope

Allotrope, two or more distinct physical forms of a chemical element in the same physical state. The term allotropy comes from the Greek allos tropos meaning “another shape.” Allotropes arise because of differing arrangements of an element’s atoms within its molecules or crystals.

One of the best-known examples of allotropy is carbon, which has multiple distinct allotropes including graphite, diamond, and buckminsterfullerene. Carbon atoms in diamond form a rigid, three-dimensional structure, with each carbon atom bonded to four other carbon atoms. In graphite the carbon atoms form stacks of flat honeycomb layers with only weak intermolecular forces between layers, while buckminsterfullerene forms balls and tubes with structures reminiscent of the geodesic domes designed by the architect Richard Buckminster Fuller.

Elements exhibiting allotropy include arsenic, antimony, iron, oxygen, phosphorus, selenium, sulfur, and tin.

There are two main kinds of allotropy, monotropy and enantiotropy. Monotropy (monotropic allotropy) occurs when one form of a substance is stable at all temperatures, while any other forms are metastable, and change (sometimes very slowly) into the stable form. For carbon, graphite is the stable monotrope, while diamond and buckminsterfullerene change, extremely slowly, into graphite. Phosphorus also exhibits monotropy: Red phosphorus is stable, while white (sometimes called yellow) phosphorus is metastable.

Enantiotropy (enantiotropic allotropy) occurs when one solid form of the substance changes into another solid form of the same substance when at a definite transition temperature. Tin, which changes from white tin to gray tin below 13°C (55°F), is an example of this type of allotropy. Gray tin is much more brittle than white tin. The Greek philosopher Aristotle recorded tin statues collapsing in the intense cold as long ago as the 4th century BC.

Infrared Radiation

Infrared Radiation, emission of energy as electromagnetic waves in the portion of the spectrum just beyond the limit of the red portion of visible radiation (see Electromagnetic Radiation). The wavelengths of infrared radiation are shorter than those of radio waves and longer than those of light waves. They range between approximately 10-6 and 10-3 (about 0.0004 and 0.04 in). Infrared radiation may be detected as heat, and instruments such as bolometers are used to detect it.

Infrared radiation is used to obtain pictures of distant objects obscured by atmospheric haze, because visible light is scattered by haze but infrared radiation is not. The detection of infrared radiation is used by astronomers to observe stars and nebulas that are invisible in ordinary light or that emit radiation in the infrared portion of the spectrum.

An opaque filter that admits only infrared radiation is used for very precise infrared photographs, but an ordinary orange or light-red filter, which will absorb blue and violet light, is usually sufficient for most infrared pictures. Developed about 1880, infrared photography has today become an important diagnostic tool in medical science as well as in agriculture and industry. Use of infrared techniques reveals pathogenic conditions that are not visible to the eye or recorded on X-ray plates. Remote sensing by means of aerial and orbital infrared photography has been used to monitor crop conditions and insect and disease damage to large agricultural areas, and to locate mineral deposits. In industry, infrared spectroscopy forms an increasingly important part of metal and alloy research, and infrared photography is used to monitor the quality of products.