Represented by the symbol Cl, chlorine has an atomic number of 17 and an atomic weight of 35.435. Like all halogens, chlorine forms a negative ion, receiving one electron into its outermost shell to acquire the highly stable electronic configuration of a noble gas. The naturally occurring isotopes of chlorine are 35 and 37. Because it is highly reactive, chlorine is found in nature only in compound form, primarily in salts. The greenish yellow diatomic gas has a noxious odor and is poisonous to humans if inhaled even in small quantities. Chlorine was used as a poison gas in the trenches of Europe during World War I.

The artificial production of pure chlorine gas is accomplished by the electrolysis of salt water, or sodium chloride (NaCl) brine. Because the oceans provide a practically inexhaustible supply of salt water, the supply of chlorine is thought to be inexhaustible as well. There are two processes used in the commercial production of chlorine. In one, the cathode in the electrolysis cell consists of a flowing bed of mercury. The sodium released in the electrolysis of NaCl amalgamates with the mercury. The mercury-sodium amalgam is then dissociated in a separate cell to form sodium hydroxide (NaOH) and the mercury returns to the main cell to be reused. At the anode chlorine gas is released, cooled with water, dried with sulfuric acid, then collected and compressed. Due to the health concerns regarding mercury, this method is being phased out.

The second method is called the diaphragm cell process. In this process, rather than removing the sodium from solution with a flowing bed of mercury, an asbestos diaphragm (or in another variation, a cation exchange membrane) is used to keep the fresh feed brine at the anode separate from the sodium (in the form of NaOH) that accumulates around the cathode. The chlorine gas is collected at the anode in the same way as in the mercury cell.

One of the most important uses of chlorine is for disinfection of drinking water. In the early part of the 20th century, unsafe water supplies were responsible for significant numbers of death and illnesses due to cholera, dysentery, and typhoid fever. Abel Wolman, a scientist with the Maryland State Department of Health, began testing chlorination of drinking water in the early 1900s. By the 1920s, the practice was widely accepted and resulted in an 85 percent drop in the rate of typhoid deaths nationwide.

However, the use of chlorine to benefit human health is not without environmental tradeoffs. Treating municipal and industrial water supplies (not to mention the local pool) consumes large amounts of chlorine. Because of these uses, excess chlorine often enters waterways where it mixes with organic compounds. These organic compounds then react with chlorine to form chloro-organic compounds, such as chloroform and other known carcinogens. The U.S. EPA has issued regulations setting limits on the allowable amounts of chlorine by-products in drinking water.

Currently, the largest use of chlorine is in the chemical industry, principally to produce various organic compounds such as chlorinated solvents and PVC plastic (polyvinylchloride). PVC pipe has replaced iron pipe as the most predominant material for sewage and water supply lines, since it does not rust or otherwise decompose. Chlorine is also used to make household bleach and other compounds used to bleach wood pulp in the making of paper. Pulp mills consume about 13 percent of U.S. chlorine production. The chlorine used in the making of paper does not end up in the final product, but rather in the wastes generated by the mills. Some of the chlorine can be reclaimed and reprocessed, but some cannot, and the disposal of pulp mill waste is a major source of excess chlorine in the environment. Insecticides made with chlorinated hydrocarbons break down very slowly. These are called persistent organic pollutants, a class of pollutants identified in 1997 by the United Nations Environment Program as a major threat to the environment and human health. Such pollutants accumulate in the fatty tissues of animals and, at high dosages, may cause cancer, neurological disease, and reproductive harm. The potential harm caused by insecticides and other chlorinated organic compounds must be weighed against their benefits. There is now discussion of permitting the use of DDT (the most notorious chlorinated hydrocarbon, long banned) once again to control mosquitoes in countries ravaged by malaria, a common cause of death and illness in the developing world.

From the 1940s through the 1990s, a significant amount of chlorine was used in the manufacture of chlorofluorocarbons (CFCs). CFCs such as freon were used as a heat transfer medium in refrigeration and air conditioning systems, as the propellant in aerosol cans, to foam plastics in the production of insulation and packing materials, and to make solvents for use in the electronics industry. Among other benefits, use of CFCs improved the energy efficiency of many industrial operations. Worldwide CFC production peaked at about one billion kilograms in 1973.

In 1974 it was discovered that CFCs damage the ozone layer in the Earth’s atmosphere. The ozone layer absorbs solar radiation in the ultraviolet range, protecting the Earth (and humans) from its damaging effects. Because they are chemically very stable, CFCs emitted at the earth’s surface migrate into the stratosphere unchanged. There they are broken down by UV radiation, liberating highly reactive chlorine atoms which then go on to react with ozone (O3), yielding O2. Through a pair of reactions, the chlorine atom is liberated once again, producing a chain reaction of ozone destruction in which a single chlorine atom can convert a million ozone molecules before reacting to form a molecule that is stable and non-reactive. The process is catalyzed by ice particles, and because there are vast clouds of such particles in the stratosphere at the Earth’s poles (especially the South Pole), the greatest ozone depletion is over Antarctica.

When the chlorine atoms eventually migrate back down to the lower atmosphere and form stable compounds, they react with methane to form hydrogen chloride, which then reacts with water vapor to form hydrochloric acid, which falls to Earth as acid rain. In 1987, more than 90 nations signed the United Nations Environment Program’s Montreal Protocol to ban CFCs, and their use was completely phased out by 1995.

Recommended Resources

Chem4Kids: Chlorine
Andrew Rader, a scientist and web-developer, maintains this free website explaining chemistry basics to kids. The site provides a breakdown on chlorine’s position in the periodic table, tidbits about chlorine in every day life, and the compounds it forms.

National Drinking Water Clearinghouse Tech Brief: Disinfection
Hosted by West Virginia University, the National Drinking Water Clearinghouse provides information about drinking water treatment. This fact sheet on disinfection spotlights relevant regulations and compares the typical methods used to rid drinking water of unsafe microorganisms.

Washington Post: Lead in D.C. Water Slashed (May 21, 2004)
Journalist D’Vera Cohn describes the impact on lead levels measured in Washington, D.C.’s water supply after authorities replaced chlorine with chloramines as a disinfectant. The article is a good example of the inherent risks and trade-offs of environmental policy-making at the community level.

Chlorine Chemistry Council: The Science Center: Building Block Chemistry
Sponsored by an association of U.S. chlorine producers, the Science Center provides teachers with classroom chemistry resources, most on chlorine. This two-day lesson plan includes a description of chlorine’s role in public health and daily life, along with two lab activities, and articles on different aspects of the chemistry of chlorine. [Grades 5-9]