The
GALE
ENCYCLOPEDIA
of
Science
The
GALE
ENCYCLOPEDIA
of
Science
THIRD EDITION
K. Lee Lerner and
Brenda Wilmoth Lerner,
Editors
VOLUME 2
Gale Encyclopedia of Science, Third Edition K. Lee Lerner and Brenda Wilmoth Lerner, Editors
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LIBRARY OF CONGRESS CATALOGING-IN-PUBLICATION DATA
Gale encyclopedia of science / K. Lee Lerner & Brenda Wilmoth Lerner, editors.— 3rd ed. p. cm.
Includes index.
ISBN 0-7876-7554-7 (set) — ISBN 0-7876-7555-5 (v. 1) — ISBN 0-7876-7556-3 (v. 2) — ISBN 0-7876-7557-1 (v. 3) — ISBN 0-7876-7558-X (v. 4) — ISBN 0-7876-7559-8 (v. 5) — ISBN 0-7876-7560-1 (v. 6) 1. Science—Encyclopedias. I. Lerner, K. Lee. II. Lerner, Brenda Wilmoth. Q121.G37 2004
503—dc22 2003015731
Disclaimer:
CONTENTS
Topic List. . . vii
Organization of the Encyclopedia. . . xxvii
Advisory Board . . . xxix
Contributors . . . xxxi
Entries
Volume 1 (Aardvark–Chaos). . . 1–818
Volume 2 (Charge-coupled device–Eye). . . 819–1572
Volume 3 (Factor–Kuru) . . . 1573–2254
Volume 4 (Lacewings–Pharmacogenetics) . . 2255–3036
Volume 5 (Pheasants–Star). . . 3037–3800
Volume 6 (Star cluster–Zooplankton) . . . 3801–4378
A
Anesthesia
Coffee plant
Formula, structural
Monoculture
Rate
Wren-warblers Wrens
Wrynecks
X
X-ray astronomy X-ray crystallography X rays
Xenogamy
Z
Zebras Zero
Zodiacal light Zoonoses Zooplankton
Y
Y2K Yak Yam Yeast Yellow fever Yew Yttrium
The Gale Encyclopedia of Science, Third Edition has been designed with ease of use and ready reference in mind.
• Entries are alphabetically arranged across six volumes, in a single sequence, rather than by scientific field • Length of entries varies from short definitions of one or
two paragraphs, to longer, more detailed entries on more complex subjects.
• Longer entries are arranged so that an overview of the subject appears first, followed by a detailed discussion conveniently arranged under subheadings.
• A list of key terms is provided where appropriate to de-fine unfamiliar terms or concepts.
• Bold-faced terms direct the reader to related articles. • Longer entries conclude with a “Resources” section,
which points readers to other helpful materials (includ-ing books, periodicals, and Web sites).
• The author’s name appears at the end of longer entries. His or her affiliation can be found in the “Contributors” section at the front of each volume.
• “See also” references appear at the end of entries to point readers to related entries.
• Cross references placed throughout the encyclopedia direct readers to where information on subjects without their own entries can be found.
• A comprehensive, two-level General Index guides readers to all topics, illustrations, tables, and persons mentioned in the book.
AVAILABLE IN ELECTRONIC FORMATS
ACADEMIC ADVISORS
Marcelo Amar, M.D.
Senior Fellow, Molecular Disease Branch National Institutes of Health (NIH) Bethesda, Maryland
Robert G. Best, Ph.D. Director
Divison of Genetics, Department of Obstetrics and Gynecology
University of South Carolina School of Medicine Columbia, South Carolina
Bryan Bunch Adjunct Instructor
Department of Mathematics Pace University
New York, New York Cynthia V. Burek, Ph.D.
Environment Research Group, Biology Department Chester College
England, UK David Campbell Head
Department of Physics
University of Illinois at Urbana Champaign Urbana, Illinois
Morris Chafetz
Health Education Foundation Washington, DC
Brian Cobb, Ph.D.
Institute for Molecular and Human Genetics Georgetown University
Washington, DC Neil Cumberlidge Professor
Department of Biology
Northern Michigan University Marquette, Michigan
Nicholas Dittert, Ph.D.
Institut Universitaire Européen de la Mer University of Western Brittany
France
William J. Engle. P.E.
Exxon-Mobil Oil Corporation (Rt.) New Orleans, Louisiana
Bill Freedman Professor
Department of Biology and School for Resource and Environmental Studies
Dalhousie University
Halifax, Nova Scotia, Canada Antonio Farina, M.D., Ph.D.
Department of Embryology, Obstetrics, and Gynecology
University of Bologna Bologna, Italy
G. Thomas Farmer, Ph.D., R.G.
Earth & Environmental Sciences Division Los Alamos National Laboratory
Los Alamos, New Mexico Jeffrey C. Hall
Lowell Observatory Flagstaff, Arizona Clayton Harris Associate Professor
Department of Geography and Geology Middle Tennessee State University Murfreesboro, Tennesses
Lyal Harris, Ph.D.
Tectonics Special Research Centre Department of Geology & Geophysics
ADVISORY BOARD
The University of Western Australia Perth, Australia
Edward J. Hollox, Ph.D. Queen’s Medical Centre University of Nottingham Nottingham, England
Brian D. Hoyle, Ph.D. (Microbiology) Microbiologist
Square Rainbow Nova Scotia, Canada Alexander I. Ioffe, Ph.D. Senior Scientist
Geological Institute of the Russian Academy of Sciences
Moscow, Russia Jennifer L. McGrath Northwood High School Nappannee, Indiana David T. King Jr., Ph.D. Professor
Department of Geology Auburn University Auburn, Alabama Danila Morano, M.D.
Department of Embryology, Obstetrics, and Gynecology
University of Bologna Bologna, Italy
Abdel Hakim Ben Nasr, Ph.D. Department of Genetics
Molecular Oncology and Development
Program/Boyer Center for Molecular Medicine Yale University School of Medicine
New Haven, Connecticut William S. Pretzer Curator
Henry Ford Museum and Greenfield Village Dearborn, Michigan
Judyth Sassoon, Ph.D., ARCS
Department of Biology and Biochemistry University of Bath
Bath, England, U.K. Yavor Shopov, Ph.D.
Professor of Geology & Geophysics University of Sofia
Bulgaria
Theodore Snow
Professor, Department of Astrophysical and Planetary Sciences
Fellow, Center for Astrophysics and Space Astronomy University of Colorado at Boulder
Boulder, Colorado
Michael J. Sullivan, M.D., Ph.D., FRACP Cancer Genetics Laboratory
University of Otago Dunedin, New Zealand Constance K. Stein, Ph.D.
Director of Cytogenetics, Assistant Director of Molecular Diagnostics
SUNY Upstate Medical University Syracuse, New York
Robert Wolke Professor emeritus Department of Chemistry University of Pittsburgh Pittsburgh, Pennsylvania Richard Addison Wood Meteorological Consultant Tucson, Arizona
Diego F. Wyszynski, M.D., Ph.D. Department of Medicine, Epidemiology &
Biostatistics
Boston University School of Medicine Boston, Massachusetts
Rashmi Venkateswaran Undergraduate Lab Coordinator Department of Chemistry University of Ottawa Ottawa, Ontario, Canada
LIBRARIAN ADVISORS
Donna Miller Director
Craig-Moffet County Library Craig, Colorado
Judy Williams Media Center
Greenwich High School Greenwich, Connecticut Carol Wishmeyer
Science and Technology Department Detroit Public Library
Detroit, Michigan
Advisor
y Boar
Nasrine Adibe Professor Emeritus Department of Education Long Island University Westbury, New York Mary D. Albanese Department of English University of Alaska Juneau, Alaska
Margaret Alic Science Writer
Eastsound, Washington
James L. Anderson Soil Science Department University of Minnesota St. Paul, Minnesota
Monica Anderson Science Writer
Hoffman Estates, Illinois
Susan Andrew Teaching Assistant University of Maryland Washington, DC John Appel Director
Fundación Museo de Ciencia y Tecnología
Popayán, Colombia David Ball
Assistant Professor Department of Chemistry Cleveland State University Cleveland, Ohio
T. Parker Bishop Professor
Middle Grades and Secondary Education
Georgia Southern University Statesboro, Georgia
Carolyn Black Professor
Incarnate Word College San Antonio, Texas Larry Blaser Science Writer Lebanon, Tennessee Jean F. Blashfield Science Writer Walworth, Wisconsin Richard L. Branham Jr. Director
Centro Rigional de
Investigaciones Científicas y Tecnológicas
Mendoza, Argentina Patricia Braus Editor
American Demographics Rochester, New York David L. Brock Biology Instructor St. Louis, Missouri Leona B. Bronstein Chemistry Teacher (retired) East Lansing High School Okemos, Michigan Dana M. Barry
Editor and Technical Writer Center for Advanced Materials
Processing
Clarkston University Potsdam, New York Puja Batra
Department of Zoology Michigan State University East Lansing, Michigan Donald Beaty
Professor Emeritus College of San Mateo San Mateo, California Eugene C. Beckham
Department of Mathematics and Science
Northwood Institute Midland, Michigan Martin Beech Research Associate Department of Astronomy University of Western Ontario London, Ontario, Canada Julie Berwald, Ph.D. (Ocean
Sciences) Austin, Texas Massimo D. Bezoari Associate Professor Department of Chemistry Huntingdon College Montgomery, Alabama John M. Bishop III Translator
New York, New York
Sarah de Forest
David Goings, Ph.D. (Geology) Geologist
Amy Kenyon-Campbell University College of the Fraser
Valley Reuben H. Fleet Space Theater
and Science Center San Diego, California Adrienne Wilmoth Lerner Graduate School of Arts &
G. H. Miller
University of Arkansas at Little Rock
Vita Richman
University of Colorado at Boulder Boulder, Colorado
Department of Earth and Physical Sciences
Policy Analyst, Air Quality Issues U.S. General Accounting Office Raleigh, North Carolina
Laurie Toupin Science Writer
Pepperell, Massachusetts Melvin Tracy
Science Educator Appleton, Wisconsin Karen Trentelman Research Associate Archaeometric Laboratory University of Toronto Toronto, Ontario, Canada Robert K. Tyson Senior Scientist W. J. Schafer Assoc. Jupiter, Florida James Van Allen Professor Emeritus Department of Physics and
Astronomy University of Iowa Iowa City, Iowa Julia M. Van Denack Biology Instructor Silver Lake College Manitowoc, Wisconsin Kurt Vandervoort
Department of Chemistry and Physics
West Carolina University Cullowhee, North Carolina Chester Vander Zee Naturalist, Science Educator Volga, South Dakota
Pella, Iowa
Frederick R. West Astronomer
Hanover, Pennsylvania Glenn Whiteside Science Writer Wichita, Kansas John C. Whitmer Professor
Department of Chemistry Western Washington University Bellingham, Washington Donald H. Williams Department of Chemistry Hope College
Holland, Michigan Robert L. Wolke Professor Emeritus Department of Chemistry University of Pittsburgh Pittsburgh, Pennsylvania Xiaomei Zhu, Ph.D.
Postdoctoral research associate Immunology Department Chicago Children’s Memorial
Hospital, Northwestern University Medical School Chicago, Illinois
Jim Zurasky Optical Physicist
Nichols Research Corporation Huntsville, Alabama
Rashmi Venkateswaran Undergraduate Lab Coordinator Department of Chemistry University of Ottawa Ottawa, Ontario, Canada R. A. Virkar
Chair
Department of Biological Sciences
Kean College Iselin, New Jersey Kurt C. Wagner Instructor
South Carolina Governor’s School for Science and Technology
Hartsville, South Carolina Cynthia Washam Science Writer Jensen Beach, Florida Terry Watkins Science Writer Indianapolis, Indiana Joseph D. Wassersug Physician
Boca Raton, Florida Tom Watson
Environmental Writer Seattle, Washington Jeffrey Weld
Instructor, Science Department Chair
Pella High School
Charge-coupled device
Charge-coupled devices (CCDs) have made possible a revolution in image processing. They consist of a series of light-sensitive elements, called pixels, arranged in a
squareor rectangular array. When CCDs are exposed to
light, an image of the object being observed is formed; this image can be extracted from the CCD and stored on a computer for later analysis. CCDs are used in a variety of modern instruments, ranging from scanners and pho-tocopiers to video cameras and digital still cameras. They have transformed the way scientists measure and chart the universe. Because CCDs are available in a wide price range, they are accessible to amateurs as well as professionals, and enable both to make significant con-tributions to modern astronomy.
How the devices work
All CCDs work on the same principle. The CCD sur-face is a grid of pixels (pixel is a contraction for “picture element”). Small CCDs may have a grid of 256 x 256 pixels, while large CCDs may have 4,096 x 4,096 pixel grids. Although many CCD pixel grids are square, this is not always the case; scanners and photocopiers, for ex-ample, have a single line of pixels that passes over the picture or page of text being imaged. The pixels are tiny; some CCDs have pixels only 9 microns across, while oth-ers may have 27-micron pixels. The scale and resolution of the image a camera is able to form on the CCD de-pends both on the pixel size and the grid size. Regardless of the pixel or grid size, however, each pixel on the CCD has the ability to convert the light striking it into an elec-tric signal. The voltage accumulated by each pixel during an exposure is directly proportional to the amount of light striking it. When the CCD is exposed to light for a length of time, an image of whatever is being observed— whether a distant galaxyor cars in a parking lot—forms on the CCD as an array of differing electric voltages.
After an image has been recorded on the CCD, the device can be “read out,” meaning that the voltages are extracted from the CCD for storage on a computer. The analogy that is almost universally used to describe this process is the “bucket brigade” analogy. Picture each pixel on the CCD as a bucket with a certain amount of
waterin it. When the CCD is read out, the water in each row of buckets is emptied into the adjacent row. The water in the first row goes into a special row of storage buckets, the water in each bucket in the second row goes into its neighbor bucket in the first row, and so on across the whole CCD. Then, the amount of water in each of these buckets is emptied, measured, and stored in a com-puter’s memory. This process is repeated until all of the
rows have been shifted into the storage buckets, emptied, and measured. If you now replace the water with electric voltages, and replace the measurement of water with the digital measurement of the analog electric signal, you have the basic process by which an image is extracted from the CCD. The actual process of reading out the CCD is performed by fairly complicated and exquisitely synchronized electronics that move all the electric charges between the “buckets,” convert the analog volt-ages into digital numbers, and make the data available for storage on a computer.
Once the pixel outputs have been measured and stored on a computer, they can be used in a variety of ways. For simple line drawings, the image processing software may render the data from the CCD in black and white. For pictures, a 256-level grayscale may be appro-priate. In either case, a grid of numbers, corresponding to the original light intensity, is present and can be ana-lyzed in any way the person studying the image desires.
From the description above, it may seem that CCDs cannot be used for colorimaging, since they respond only to light intensity. In fact, color CCDs are available, although they are used in video equipment such as cam-corders and almost never in astronomy. If an astronomer wanted to create a color image using a CCD, the old practice of taking three images through three different color filters is still the usual way to go. True color CCDs have pixels with built-in filters, alternating red, green, and blue. They can produce real-time color images, but they are undesirable for scientific work because they in-troduce significant difficulties into the data analysis process, as well as reducing the effective resolution of the CCD by a factor of three.
Applications in astronomy
Astronomers began using charge-coupled devices in their work in the early 1980s, when the increasing power and clock speed of semiconductors, and the computers needed to drive the hardware and analyze the data be-came both fast and affordable. Almost every field of as-tronomy was directly impacted by CCDs: for observa-tions of asteroids, galaxies, stars, and planets, whether by direct imaging or the recording of spectra, the CCD rapidly became the detector of choice.
CCDs are also useful to astronomers because an av-erage, CCDs are about ten times more light-sensitive than film. Astronomers are notorious for finding desper-ately faint objects to observe, so the CCD gave them the ability not only to see fainter objects than they could be-fore, but to reduce the amount of time spent tracking and observing a given object. A CCD camera can record in a 15 minute exposure the same information that would
Charge-coupled de
take a standard camera loaded with film two hours or more. While film typically records only 2–3% of the light that strikes it, charge-coupled device cameras can record between 50–80% of the light they detect. Further-more, CCDs can capture light outside the visible spec-trum, which film cannot do. The devices operate with-out darkrooms or chemicals, and the results can be re-constructed as soon as the information is loaded into an image processing program.
However, CCD cameras do have some drawbacks. The small size of the most affordable arrays results in a much smaller field of view. Large celestial bodies such as the moon, which are easily photographed with a 35mm camera, become very difficult to reproduce as a single image with a CCD camera. Although larger arrays are coming to the market, they remain pricy and beyond the resources of the amateur astronomer. They require com-plicated systems to operate, any many of them have to be cooled to typical temperatures of -112°F (-80°C) to re-duce their background electronic noise to an acceptable level. Finally, color images for astronomical CCD cam-eras (unlike commercially-available video and digital still cameras) require three separate exposures for each filter used. The final image has to be created by combining the data from each exposure within the computer.
CCDs, professionals, and amateurs
With web-based starcatalogues and other Internet and electronic resources, such as the Hubble Guide Star Catalog and the Lowell Observatory Asteroid Database, professional and amateur astronomers have begun shar-ing resources and comparshar-ing data in hopes of creatshar-ing a more accurate and complete picture of the heavens. Or-ganizations such as the Amateur Sky Survey help indi-viduals coordinate and share data with others. Thanks to CCDs, amateurs have often contributed as significantly to these projects as professional astronomers have. Paul Comba, an amateur based in Arizona, discovered and registered some 300 previously unknown asteroids in 1996–97, after adding a digital camera to his telescope. In 1998,astrophysicsstudent Gianluca Masi recorded the existence of an unknown variable star, discovered with the use of his Kodak KAF-0400 CCD, mounted in a Santa Barbara Instrument Group ST-7 camera. CCDs help level the playing field in the science of astrometry, drastically reducing the equipment barrier between the amateur and the professional.
Resources
Periodicals
di Cicco, Dennis. “Measuring the Sky with CCDs.”Sky & Tele-scope94 (December 1997): 115-18.
Gombert, Glenn, and Tom Droege. “Boldness: The Amateur Sky Survey.”Sky & Telescope95 (February 1998): 42- 45. Hannon, James. “Warming Up to Digital Imaging.”Sky &
Tele-scope97 (March 1999): 129.
Masi, Gianluca. “CCDs, Small Scopes, and the Urban Ama-teur.”Sky & Telescope95 (February 1998): 109-12. Terrance, Gregory. “Capture the Sky on a CCD: Digital
Imag-ing with a CCD Camera Is RevolutionizImag-ing the Way Ama-teur Astronomers Record Planets and Galaxies.” Astrono-my28 (February 2000): 72.
Kenneth R. Shepherd
Charles’s law
see
Gases, properties of
Cheetah
see
Cats
Chelate
A chelate is a type of coordination compoundin which a single metallic ion is attached by coordinate co-valent bonds to a moleculeor an ion called a ligand. The term chelate comes from the Greek word chela, meaning “crab’s claw.” The term clearly describes the appearance of many kinds of chelates, in which the ligand surrounds the central atom in a way that can be compared to the grasping of food by a crab’s claw.
Bonding in a chelate occurs because the ligand has at least two pairs of unshared electrons. These unshared pairs of electrons are regions of negative electrical charge to which are attracted cations such as the cop-per(I) and copper(II), silver, nickel, platinum, and alu-minumions. A ligand with only two pairs of unshared electrons is known as a bidentate (“two-toothed”) ligand; one with three pairs of unshared electrons, a tridentate (“three-toothed”) ligand, and so on.
The geometric shape of a chelate depends on the number of ligands involved. Those with bidentate lig-ands form linear molecules, those with four liglig-ands form planar or tetrahedral molecules, and those with six lig-ands form octahedral molecules.
One of the most familiar examples of a chelate is hemoglobin, the molecule that transports oxygen
through the blood. The “working part” of a hemoglobin molecule is heme, a complex molecule at whose core is an iron(II) ion bonded to four nitrogen atomswith coor-dinate covalent bonds.
Among the most common applications of chelates is in watersoftening and treatment of poisoning. In the former instance, a compound such as sodium
tripolyphosphate is added to water. That compound forms chelates with calciumand magnesiumions, ions
responsible for the hardness in water. Because of their ability to “tie up”metalions in chelates, compounds like sodium tripolyphosphate are sometimes referred to as se-questering agents.
A typical sequestering agent used to treat poison victims is ethylenediaminetetraacetic acid, commonly known as EDTA. Suppose that a person has swallowed a significant amount of lead and begins to display the symptoms of lead poisoning. Giving the person EDTA allows that molecule to form chelates with lead ions, re-moving that toxic material from the bloodstream.
Chemical bond
A chemical bond is any forceof attraction that holds two atomsor ions together. In most cases, that force of attraction is between one or more electrons held by one of the atoms and the positively charged nucleus of the second atom. Chemical bonds vary widely in their stabil-ity, ranging from relatively strong covalent bonds to very weak hydrogenbonds.
History
The concept of bonding as a force that holds two particles together is as old as the concept of ultimate particles of matteritself. As early as 100 B.C., for ex-ample, Asklepiades of Prusa speculated about the exis-tence of “clusters of atoms,” a concept that implies the existence of some force of attraction holding the parti-cles together. At about the same time, the Roman poet Lucretius in his monumental work De Rerum Natura
(“On the nature of things”) pictured atoms as tiny spheres to which were attached fishhook-like ap-pendages. Atoms combined with each other, according to Lucretius, when the appendages from two adjacent atoms became entangled with each other.
Relatively little progress could occur in the field of bonding theory, of course, until the concept of an atom itself was clarified. When John Dalton proposed the modern atomic theoryin 1803, he specifically hypothe-sized that atoms would combine with each other to form “compound atoms.” Dalton’s concept of bonding was es-sentially non-existent, however, and he imagined that atoms simply sit adjacent to each other in their com-pound form.
The real impetus to further speculation about bond-ing was provided by the evolution of the concept of a molecule, originally proposed by Amedeo Avogadro in 1811 and later refined by Stanislao Cannizzaro more than four decades later.
The origin of bond symbolism
Some of the most vigorous speculation about chemi-cal bonding took place in the young field of organic chem-istry. In trying to understand the structure of organic com-pounds, for example, Friedrich Kekulé suggested that the carbonatom is tetravalent; that is, it can bond to four other atoms. He also hypothesized that carbon atoms could bond with each other almost endlessly in long chains.
Kekulé had no very clear notion as to how atoms bond to each other, but he did develop an elaborate sys-tem for showing how those bonds might be arranged in space. That system was too cumbersome for everyday use by chemists, however, and it was quickly replaced by another system suggested earlier by the Scottish chemist Archibald Scott Couper. Couper proposed that the bond between two atoms (what the real physical nature of that bond might be) be represented by a short dashed line. Thus, a molecule of watercould be represented by the structural formula: H-O-H.
That system is still in existence today. The arrange-ment of atoms in a molecule is represented by the symbols of the elements present joined by dashed lines that show how the atoms of those elements are bonded to each other. Thus, the term chemical bond refers not only to the force of attraction between two particles, but also to the dashed line used in the structural formula for that substance.
Development of the modern theory of bonding
The discovery of the electronby J. J. Thomson in 1897 was, in the long run, the key needed to solve the problem of bonding. In the short run, however, it was a serious hindrance to resolving that issue. The question that troubled many chemists at first was how two parti-cles with the same electrical charge (as atoms then seemed to be) could combine with each other.
An answer to that dilemma slowly began to evolve, beginning with the work of the young German chemist Richard Abegg. In the early 1900s, Abegg came to the conclusion that inert gases are stable elements because their outermost shell of electrons always contain eight electrons. Perhaps atoms combine with each other, Abegg said, when they exchange electrons in such as way that they all end up with eight electrons in their outer orbit. In a simplistic way, Abegg had laid out the principle of ionic bonding. Ionic bonds are formed when one atom completely gives up one or more electrons, and a second atom takes on those electrons.
Since Abegg was killed in 1910 at the age of 41 in a balloon accident, he was prevented from improving upon his original hypothesis. That work was taken up in
the 1910s, however, by a number of other scientists, most prominently the German chemist Walther Kossel and the American chemists Irving Langmuir and Gilbert Newton Lewis.
Working independently, these researchers came up with a second method by which atoms might bond to each other. Rather than completely losing or gaining electrons, they hypothesized, perhaps atoms can share electrons with each other. One might imagine, for exam-ple, that in a molecule of methane (CH4), each of the
four valenceelectrons in carbon is shared with the single electron available from each of the four hydrogen atoms. Such an arrangement could provide carbon with a full outer shell of eight electrons and each hydrogen atom with a full outer shell of two. Chemical bonds in which two atoms share pairs of electrons with each other are known as covalent bonds.
In trying to illustrate this concept, Lewis developed another system for representing chemical bonds. In the Lewis system (also known as the electron-dot system), each atom is represented by its chemical symbol with the number of electrons in its outermost orbit, its bonding or valence electrons. The formula of a compound, then, is to be represented by showing how two or more atoms share electrons with each other.
Bond types
Credit for the development of the modern theory of chemical bonding belongs largely to the great American chemist Linus Pauling. Early in his career, Pauling learned about the revolution in physicsthat was taking place largely in Europeduring the 1920s. That revolu-tion had come about with the discovery of the relativity theory,quantum mechanics, the uncertainty principle, the duality of matter and energy, and other new and strikingly different concepts in physics.
Most physicists recognized the need to reformulate the fundamental principles of physics because of these discoveries. Relatively few chemists, however, saw the relevance of the revolution in physics for their own sub-ject. Pauling was the major exception. By the late 1920s, he had already begun to ask how the new science of quantum mechanics could be used to understand the na-ture of the chemical bond.
In effect, the task Pauling undertook was to deter-mine the way in which any two atoms might react with each other in such a way as to put them in the lowest possible energy state. Among the many discoveries he made was that, for most cases, atoms form neither a purely ionic nor purely covalent bond. That is, atoms typically do not completely lose, gain, or share equally the electrons that form the bond between them. Instead,
the atoms tend to form hybrid bonds in which a pair of shared electrons spend more time with one atom and less time with the second atom.
Electronegativity
The term that Pauling developed for this concept is electronegativity. Electronegativity is, in a general sense, the tendency of an atom to attract the electrons in a covalent bond. The numerical values for the elec-tronegativities of the elements range from a maximum of 4.0 for fluorine to a minimum of about 0.7 for ce-sium. A bond formed between fluorine and cesium would tend to be ionic because fluorine has a much stronger attraction for electrons than does cesium. On the other hand, a bond formed between cobalt (elec-tronegativity = 1.9) and silicon (elec(elec-tronegativity = 1.9) would be a nearly pure covalent bond since both atoms have an equal attraction for electrons.
The modern concept of chemical bonding, then, is that bond types are not best distinguished as purely ionic or purely covalent. Instead, they can be envisioned as lying somewhere along a continuum between those two extremes. The position of any particular bond can be pre-dicted by calculating the difference between the two electronegativities of the atoms involved. The greater that difference, the more ionic the bond; the smaller the difference, the more covalent.
Bond polarity
The preceding discussion suggests that most chemi-cal bonds are polar; that is, one end of the bond is more positive than the other end. In the bond formed between hydrogen (electronegativity = 2.2) and sulfur (elec-tronegativity = 2.6), for example, neither atom has the ability to take electrons completely from the other. Nei-ther is equal sharing of electrons likely to occur. Instead, the electrons forming the hydrogen-sulfur bond will spend somewhat more time with the sulfur atom and somewhat less time with the hydrogen atom. Thus, the sulfur end of the hydrogen-sulfur bond is somewhat more negative(represented as ␦-) and the hydrogen end, somewhat more positive (␦+).
Coordination compounds
Some chemical bonds are unique in that both elec-trons forming the bond come from a single atom. The two atoms are held together, then, by the attraction be-tween the pair of electrons from one atom and the posi-tively charged nucleus of the second atom. Such bonds have been called coordinate covalent bonds.
that takes place within such molecules means that some parts of the molecule are momentarily charged, either positively or negatively. For this reason, very weak, tran-sient forces of attraction can develop between particles that are actually neutral.
Resources
Books
Bynum, W.F., E.J. Browne, and Roy Porter. Dictionary of the
History of Science.Princeton, NJ: Princeton University
Press, 1981, pp. 433-435.
Kotz, John C., and Paul Treichel. Chemistry and Cehmical
Re-activity.Pacific Grove, CA: Brooks/Cole, 1998.
Lide, D.R., ed. CRC Handbook of Chemistry and Physics.
Boca Raton: CRC Press, 2001. An example of this kind of bonding is found in the
reaction between copper(II) ion and ammonia. The ni-trogenatom in ammonia has an unshared pair of elec-trons that is often used to bond with other atoms. The copper(II) ion is an example of such an anion. It is posi-tively charged and tends to surround itself with four am-monia molecules to form the cupric ammonium ion, Cu(NH3)42+. The bonding in this ion consists of
coordi-nate covalent bonds with all bonding electrons supplied by the nitrogen atom.
Multiple bonds
The bonds described thus far can all be classified as single bonds. That is, they all consist of a single pair of electrons. Not uncommonly, two atoms will combine with each other by sharing two pairs of electrons. For ex-ample, when leadand sulfur combine to form a com-pound, the molecules formed might consist of two pairs of electrons, one electron from lead and one electron from sulfur in each of the pairs. The standard shorthand for a double bond such as this one is a double dashed line (=). For example, the formula for a common double-bonded compound, ethylene, is: H2C=CH2.
Compounds can also be formed by the sharing of three pairs of electrons between two atoms. The formula for one such compound, acetylene, shows how a triple bond of this kind is represented: HC-CH.
Other types of bonds
Other types of chemical bonds also exist. The atoms that make up a metal, for example, are held together by a metallic bond. A metallic bond is one in which all of the metal atoms share with each other a cloud of electrons. The electrons that make up that cloud originate from the outermost energy levels of the atoms.
A hydrogen bond is a weak force of attraction that exists between two atoms or ions with opposite charges. For example, the hydrogen-oxygen bonds in water are polar bonds. The hydrogen end of these bonds are slight-ly positive and the oxygenends, slightly negative. Two molecules of water placed next to each other will feel a force of attraction because the oxygen end of one mole-cule feels an electrical force of attraction to the hydrogen end of the other molecule. Hydrogen bonds are very common and extremely important in biological systems. They are strong enough to hold substances together, but weak enough to break apart and allow chemical changes to take place within the system.
Van der Waals forcesare yet another type of chem-ical bond. Such forces exist between particles that appear to be electrically neutral. The rapid shifting of electrons
Chemical bond
KEY TERMS
. . . .
Coordinate covalent bond—A type of covalent bond in which all shared electrons are donated by only one of two atoms.
Covalent bond—A chemical bond formed when two atoms share a pair of electrons with each other.
Double bond—A covalent bond consisting of two pairs of shared electrons that hold the two atoms together.
Electronegativity—A quantitative method for indi-cating the relative tendency of an atom to attract the electrons that make up a covalent bond.
Ionic bond—A chemical bond formed when one atom gains and a second atom loses electrons.
Lewis symbol—A method for designating the structure of atoms and molecules in which the chemical symbol for an element is surrounded by dots indicating the number of valence electrons in the atom of that element.
Molecule—A collection of atoms held together by some force of attraction.
Multiple bond—A double or triple bond.
Polar bond—A covalent bond in which one end of the bond is more positive than the other end.
Structural formula—The chemical representation of a molecule that shows how the atoms are arranged within the molecule.
Triple bond—A triple bond is formed when three pairs of electrons are shared between two atoms..
Oxtoby, David W., et al. The Principles of Modern Chemistry.
5th ed. Pacific Grove, CA: Brooks/Cole, 2002.
Pauling, Linus. The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to
Modern Structural Chemistry.3rd edition. Ithaca, NY:
Cornell University Press, 1960.
Chemical compound
see
Compound,
chemical
Chemical element
see
Element, chemical
Chemical equilibrium
see
Equilibrium,
chemical
Chemical evolution
Chemical evolution describes chemical changes on the primitive Earththat gave rise to the first forms of life. The first living things on Earth were prokaryotes with a type of cell similar to present-day bacteria. Prokaryotefossils have been found in 3.4-million-year-old rock in the southern part of Africa, and in even older rocksin Australia, including some that appear to be photosynthetic. All forms of life are theorized to have evolved from the original prokaryotes, probably 3.5-4.0 billion years ago.
The primitive Earth
The chemical and physical conditions of the prim-itive Earth are invoked to explain the origin of life, which was preceded by chemical evolution of organic chemicals. Astronomers believe that 20-30 billion years ago, all matter was concentrated in a single mass, and that it blew apart with a “big bang.” In time, a disk-shaped cloud of dust condensed and formed the Sun, and the peripheral matter formed its planets. Heatproduced by compaction,radiation, and impact-ing meteorites melted Earth. Then, as the planet cooled, Earth’s layers formed. The first atmosphere was made up of hot hydrogengas, too light to be held by Earth’s gravity. Watervapor,carbon monoxide, carbon dioxide,nitrogen, and methane replaced the hydrogen atmosphere. As Earth cooled, water vapor condensed and torrential rains filled up its basins, thereby forming the seas. Also present were lightning, volcanic activity, and ultraviolet radiation. It was in this setting that life began.
According to one theory, chemical evolution oc-curred in four stages.
In the first stage of chemical evolution, molecules in the primitive environment formed simple organic sub-stances, such as amino acids. This concept was first pro-posed in 1936 in a book entitled, “The Origin of Life on Earth,” written by the Russian scientist, Aleksandr Ivanovich Oparin. He considered hydrogen,ammonia, water vapor, and methane to be components in the early atmosphere. Oxygenwas lacking in this chemically- re-ducing environment. He stated that ultraviolet radiation from the Sun provided the energyfor the transformation of these substances into organic molecules. Scientists today state that such spontaneous synthesis occurred only in the primitive environment. Abiogenesis became impos-sible when photosynthetic cells added oxygen to the at-mosphere. The oxygen in the atmosphere gave rise to the ozonelayer which then shielded Earth from ultraviolet radiation. Newer versions of this hypothesis contend that the primitive atmosphere also contained carbon monox-ide, carbon dioxmonox-ide, nitrogen, hydrogen sulfmonox-ide, and hy-drogen. Present-day volcanoes emit these substances.
In 1957, Stanley Miller and Harold Urey provided laboratory evidence that chemical evolution as described by Oparin could have occurred. Miller and Urey created an apparatus that simulated the primitive environment. They used a warmed flask of water for the ocean, and an atmosphere of water, hydrogen, ammonia and methane. Sparks discharged into the artificial atmosphere represent-ed lightning. A condenser coolrepresent-ed the atmosphere, causing rain that returned water and dissolved compounds back to the simulated sea. When Miller and Urey analyzed the components of the solutionafter a week, they found vari-ous organic compounds had formed. These included some of the amino acids that compose the proteins of living things. Their results gave credence to the idea that simple substances in the warm primordial seas gave rise to the chemical building blocks of organisms.
In the second stage of chemical evolution, the sim-ple organic molecules (such as amino acids) that formed and accumulated joined together into larger structures (such as proteins). The units linked to each other by the process of dehydration synthesis to form polymers. The problem is that the abiotic synthesis of polymers had to occur without the assistance of enzymes. In addition, these reactions give off water and would, therefore, not occur spontaneously in a watery environment. Sydney Fox of the University of Miami suggested that waves or rain in the primitive environment splashed organic monomers on fresh lava or hot rocks, which would have allowed polymers to form abiotically. When he tried to do this in his laboratory, Fox produced proteinoids—abi-otically synthesized polypeptides.
The next step in chemical evolution suggests that polymers interacted with each other and organized into
Chemical e
Periodicals
Franklin, Carl. “Did Life Have a Simple Start?”New Scientist
(October 2, 1993).
Radetsky, Peter. “How Did Life Start?”Discover(November 1992).
Bernice Essenfeld
Chemical oxygen demand
Chemical oxygen demand (COD) is a measure of the capacity of waterto consume oxygen during the de-compositionof organic matterand the oxidation of in-organic chemicals such as ammoniaand nitrite. COD measurements are commonly made on samples of waste waters or of natural waters contaminated by domestic or industrial wastes. Chemical oxygen demand is measured as a standardized laboratory assay in which a closed water sampleis incubated with a strong chemical oxi-dant under specific conditions of temperatureand for a particular period of time. A commonly used oxidant in COD assays is potassium dichromate (K2Cr2O7) which is
used in combination with boiling sulfuric acid(H2SO4).
Because this chemical oxidant is not specific to oxygen-consuming chemicals that are organic or inorganic, both of these sources of oxygen demand are measured in a COD assay.
Chemical oxygen demand is related to biochemical oxygen demand(BOD), another standard test for assay-ing the oxygen-demandassay-ing strength of waste waters. However, biochemicaloxygen demand only measures the amount of oxygen consumed by microbial oxidation and is most relevant to waters rich in organic matter. It is im-portant to understand that COD and BOD do not neces-sarily measure the same types of oxygen consumption. For example, COD does not measure the oxygen-con-suming potential associated with certain dissolved or-ganic compounds such as acetate. However, acetate can be metabolized by microorganismsand would therefore be detected in an assay of BOD. In contrast, the oxygen-consuming potential of celluloseis not measured during a short-term BOD assay, but it is measured during a COD test.
Chemical reactions
Chemical reactions describe the changes between re-actants (the initial substances that enter into the reaction) and products (the final substances that are present at the aggregates, known as protobionts. Protobionts are not
ca-pable of reproducing, but had other properties of living things. Scientists have successfully produced protobionts from organic molecules in the laboratory. In one study, proteinoids mixed with cool water assembled into droplets or microspheres that developed membranes on their surfaces. These are protobionts, with semipermeable and excitable membranes, similar to those found in cells.
In the final step of chemical evolution, protobionts developed the ability to reproduce and pass genetic in-formation from one generation to the next. Some scien-tists theorize RNA to be the original hereditary molecule. Short polymers of RNA have been synthe-sized abiotically in the laboratory. In the 1980s, Thomas Cech and his associates at the University of Colorado at Boulder discovered that RNA molecules can function as enzymes in cells. This implies that RNA molecules could have replicated in prebiotic cells without the use of protein enzymes. Variations of RNA molecules could have been produced by mutations and by errors during replication. Natural selection, operating on the different RNAs would have brought about subsequent evolution-ary development. This would have fostered the survival of RNA sequences best suited to environmental parame-ters, such as temperatureand salt concentration. As the protobionts grew and split, their RNA was passed on to offspring. In time, a diversity of prokaryote cells came into existence. Under the influence of natural selection, the prokaryotes could have given rise to the vast variety of life on Earth.
See also Amino acid.
Resources
Books
Keeton, William T., and James L. Gould. Biological Science.
New York: W.W. Norton and Co., 1993.
Chemical r
eactions
KEY TERMS
. . . .
Abiogenesis—Origin of living organisms from nonliving material.
Autotroph—This refers to organisms that can syn-thesize their biochemical constituents using inor-ganic precursors and an external source of energy.
Heterotroph—Organism that requires food from the environment since it is unable to synthesize nutrients from inorganic raw materials.
end of the reaction). Describing interactions among chem-ical species, chemchem-ical reactions involve a rearrangement of the atomsin reactants to form products with new struc-tures in such a way as to conserve atoms. Chemical equa-tions are notaequa-tions that are used to concisely summarize and convey information regarding chemical reactions.
In a balanced chemical reaction all of the matter (i.e., atoms or molecules) that enter into a reaction must be accounted for in the products of a reaction. Accord-ingly, associated with the symbols for the reactants and products are numbers (stoichiometry coefficients) that represent the number of molecules, formula units, or moles of a particular reactant or product. Reactants and products are separated by additionsymbols (addition signs). The addition signs represent the interaction of the reactants and are used to separate and list the prod-ucts formed. The chemical equations for some reac-tions may have a lone reactant or a single product. The subscript numbers associated with the chemical formu-la designating individual reactants and products repre-sent the number of atoms of each element that are in each molecule(for covalently bonded substances) or formula unit (for ironically associated substances) of reactants or products.
For a chemical reaction to be balanced, all of the atoms present in molecules or formula units or moles of reactants to the left of the equation arrow must be present in the molecules, formula units and moles of product to the right of the equation arrow. The combinations of the atoms may change (indeed, this is what chemical reac-tions do) but the number of atoms present in reactants must equal the number of atoms present in products.
Charge is also conserved in balanced chemical reac-tions and therefore there is a conservation of electrical charge between reactants and products.
Although chemical equations are usually concerned only with reactants and products chemical reactions may proceed through multiple intermediate steps. In such multi-step reactions the products of one reaction become the reactants (intermediary products) for the next step in the reaction sequence.
Reaction catalysts are chemical species that alter the energyrequirements of reactions and thereby alter the speed at which reactions run (i.e., control the rateof for-mation of products).
Combustion reactions are those where oxygen combines with another compound to form waterand carbon dioxide. The equations for these reactions usual-ly designate that the reaction is exothermic (heat pro-ducing). Synthesis reactions occur when two or more simple compounds combine to form a more complicated compound. Decompositionreactions reflect the reversal
of synthesis reactions (e.g., reactions where complex molecules are broken down into simpler molecules). The electrolysisof water to make oxygen and hydrogenis an excellent example of a decomposition reaction.
Equations for single displacement reactions, double displacement, and acid-base reactions reflect the appro-priate reallocation of atoms in the products.
In accord with the laws of thermodynamics, all chemical reactions change the energy state of the reac-tants. The change in energy results from changes in the in the number and strengths of chemical bonds as the re-action proceeds. The heat of rere-action is defined as the quantity of heat evolved or absorbed during a chemical reaction. A reaction is called exothermic if heat is re-leased or given off during a chemical transformation. Al-ternatively, in an endothermicreaction, heat is absorbed in transforming reactants to products. In endothermic re-actions, heat energy must be supplied to the system in order for a reaction to occur and the heat content of the products is larger than that of the reactants. For example, if a mixture of gaseous hydrogen and oxygen is ignited, water is formed and heat energy is given off. The chemi-cal reaction is an exothermic reaction and the heat con-tent of the product(s) is lower than that for the reactants. The study of energy utilization in chemical reactions is called chemical kinetics and is important in understand-ing chemical transformations.
A chemical reaction takes place in a vessel which can be treated as a system. If the heat “flows” into the vessel during reaction, the reaction is said to be “en-dothermic” (e.g., a decomposition process) and the amount of heat, say,q, provided to the system is taken as a positive quantity. On the other hand, when the system has lost heat to the outside world, the reaction is “exothermic” (e.g., a combustion process) and q is viewed as a negativenumber. Normally the heat change involved in a reaction can be measured in an adiabatic bomb calorimeter. The reaction is initiated inside a con-stant-volume container. The observed change in temper-atureand the information on the total heat capacityof the colorimeter are employed to calculate q. If the heat of reaction is obtained for both the products and reac-tants at the same temperature after reaction and also in their standard states, it is then defined as the “standard heat of reaction,” denoted by ⌬H°.
Both chemical kinetics and thermodynamics are crucial issues in studying chemical reactions. Chemical kinetics help us search for the factors that influence the rate of reaction. It provides us with the information about how fast the chemical reaction will take place and about what the sequence of individual chemical events is to produce observed reactions. Very often, a single
reac-Chemical r
have to first encounter each other so that they can ex-change atoms or groups of atoms. In gas phases, this step relies on collision, whereas in liquid and solid phases,diffusionprocess (masstransfer) plays a key role. However, even reactive species do encounter each other, and certain energy state changes are required to surmount the energy barrier for the reaction. Normally, this minimum energy requirement (e.g., used to break old chemical bonds and to form new ones) is varied with temperature, pressure, the use of catalysts, etc. In other words, the rate of chemical reaction depends heav-ily on encounter rates or frequencies and energy avail-ability, and it can vary from a value approaching infini-tyto essentially zero.
See also Catalyst and catalysis; Chemical bond; Chemistry; Conservation laws; Entropy; Enzyme; Equa-tion, chemical; Equilibrium, chemical; Molecular formu-la; Moles; Stereochemistry.
Resources
Books
Housecroft, Catherine E., et al. Inorganic Chemistry.Prentice Hall, 2001.
Incropera, Frank P., and David P. DeWitt. Fundamentals of Heat and Mass Transfer.5th ed. John Wiley & Sons, 2001. Moran, Michael J., and Howard N. Shapiro. Fundamentals of
Engineering Thermodynamics. 4th ed. John Wiley &
Sons, 2000.
K. Lee Lerner Pang-Jen Kung tion like A 씮B may take several steps to complete. In
other words, a chain reaction mechanism is actually in-volved which can include initiation, propagation, and termination stages, and their individual reaction rates may be very different. With a search for actual reaction mechanisms, the expression for overall reaction rate can be given correctly. As to determining the maximum ex-tent to which a chemical reaction can proceed and how much heat will be absorbed or liberated, we need to esti-mate from thermodynamics data. Therefore, kinetic and thermodynamic information is extremely important for reactor design.
As an example of a chemical reaction, Hydrogen (H2) and oxygen (O2) gases under certain conditions can
react to form water (H2O). Water then exists as solid
(ice), liquid, or vapor (steam); they all have the same composition, H2O, but exhibit a difference in how H2O
molecules are brought together due to variations in tem-perature and pressure.
Chemical reactions can take place in one phase alone and are termed “homogeneous.” They can also proceed in the presence of at least two phases, such as reduction of iron oreto iron and steel, which are nor-mally described as “heterogeneous” reactions. Quite fre-quently, the rate of chemical reaction is altered by for-eign materials, so-called catalysts, that are neither reac-tants nor products. Although usually used to accelerate reactions, reaction catalysts can either accelerate or hin-der the reaction process. Typical examples are found in Pt as the catalyst for oxidation of sulfur dioxide(SO2)
and iron promoted with Al2O3and K as the catalyst for ammonia(NH3) synthesis.
Chemical reactions can characterized as irreversible, reversible, or oscillating. In the former case, the equilib-rium for the reaction highly favors formation of the products, and only a very small amount of reactants re-mains in the system at equilibrium. In contrast to this, a reversible reaction allows for appreciable quantities of all reactants and products co-existing at equilibrium. H2O ⫹3NO2씯씮2HNO3⫹NO is an example of a
re-versible reaction. In an oscillating chemical reaction, the concentrations of the reactants and products change with
timein a periodic or quasi-periodic manner. Chemical oscillators exhibit chaotic behavior, in which concentra-tions of products and the course of a reaction depend on the initial conditions of the reaction.
Chemical reactions may proceed as a single reac-tion A 씮B, series reactions A 씮B 씮C, side-by-side
parallelreactions A 씮B and C 씮D, two competitive parallel reactions 씮B and A 씮C, or mixed parallel and series reactions A ⫹B 씮C and C ⫹ B 씮D. In order for chemical reactions to occur, reactive species
Chemical r
eactions
KEY TERMS
. . . .
Chemical kinetics—The study of the reaction mechanism and rate by which one chemical species is converted to another.
Equilibrium—The conditions under which a sys-tem shows no tendency for a change in its state. At equilibrium the net rate of reaction becomes zero.
Phase—A homogeneous region of matter.
Standard state—The state defined in reaction ther-modynamics for calculation purposes in which the pure gas in the ideal-gas state at 1 atm and pure liquid or solid at 1 atm are taken for gas and liquid or solid, respectively.
Chemical warfare
Chemical warfare involves the use of natural or syn-thetic substances to incapacitate or kill an enemy or to deny them the use of resources such as agricultural prod-ucts or screening foliage. The effects of the chemicals may last only a short time, or they may result in perma-nent damage and death. Most of the chemicals used are known to be toxic to humans or plantlife. Other normal-ly benign (mild) chemicals have also been intentionalnormal-ly misused in more broadly destructive anti-environmental actions, called ecocide, and as a crude method of causing mayhem and damaging an enemy’s economic system. The deliberate dumping of large quantities of crude oil on the land or in the oceanis an example.
Chemical warfare dates back to the earliest use of weapons. Poisoned arrows and darts used for hunting were also used as weapons in intertribal conflicts (and primitive peoples still use them for these purposes today). In 431 B.C., the Spartans used burning sulfurand pitch to produce cloudsof suffocating sulfur dioxidein their sieges against Athenian cities. When the Romans defeated the Carthaginians of North Africain 146 B.C. during the last of a series of Punic Wars, they levelled the city of Carthage and treated the surrounding fields with salt to destroy the agricultural capability of the land, thereby preventing the rebuilding of the city.
The attraction of chemicals as agents of warfare was their ability to inflict masscasualties or damage to an enemy with only limited risk to the forces using the chemicals. Poisoning a town’s watersupply, for exam-ple, posed almost no threat to an attacking army, yet re-sulted in the death of thousands of the town’s defenders. In many cases, the chemicals were also not detectable by the enemy until it was too late to take action.
Chemical agents can be classified into several gen-eral categories. Of those that attack humans, some, like tear gas, cause only temporary incapacitation. Other agents cause violent skin irritation and blistering, and may result in death. Some agents are poisonous and are absorbed into the bloodstream through the lungs or skin to kill the victim. Nerve agents attack the nervous sys-temand kill by causing the body’s vital functions to cease. Still others cause psychological reactions includ-ing disorientation and hallucinations. Chemical agents which attack vegetation include defoliants that kill plant leaves,herbicidesthat kill the entire plant, and soil steri-lants that prevent the growth of new vegetation.
Antipersonnel agents—chemicals used against people
The first large-scale use of poisonous chemicals in
warfare occurred during World War I. More than 100,000 tons (90,700 metric tons) of lethal chemicals were used by both sides during several battles in an ef-fort to break the stalemate of endless trench warfare. The most commonly used chemicals were four lung-destroy-ing poisons: chlorine, chloropicrin, phosgene, and trichloromethyl chloroformate, along with a skin-blister-ing agent known as mustard gas, or bis (2-chloroethyl) sulfide. These poisons caused about 100,000 deaths and another 1.2 million injuries, almost all of which involved military personnel.
Despite the agreements of the Geneva Protocol of 1925 to ban the use of most chemical weapons, the Unit-ed States, Britain, Japan, Germany, Russia and other countries all continued development of these weapons during the period between World War I and World War II. This development included experimentation on ani-mals and humans. Although there was only limited use of chemical weapons during World War II, the opposing sides had large stockpiles ready to deploy against mili-tary and civilian targets.
During the war in Vietnam, the United States mili-tary used a nonlethal “harassing agent” during many op-erations. About 9,000 tons (8,167 tonnes) of tear gas, known as CS or o-chlorobenzolmalononitrile, were sprayed over 2.5 million acres (1.0 million ha) of South Vietnam, rendering the areas uninhabitable for 15-45 days. Although CS is classified as nonlethal, several hun-dred deaths have been reported in cases when CS has been used in heavy concentrations in confined spaces such as underground bunkers and bomb shelters.
Poisonous chemicals were also used during the Iran-Iraq War of 1981-1987, especially by Iran-Iraqi forces. During that war, both soldiers and civilians were targets of chemi-cal weapons. Perhaps the most famous incident was the gassing of Halabja, a town in northern Iraq that had been overrun by Iranian-supported Kurds. The Iraqi military at-tacked Halabja with two rapidly acting neurotoxins, known as sabin and tabun, which cause rapid death by in-terfering with the transmission of nerve impulses. About 5,000 people, mostly civilians, were killed in this incident.
Use of herbicides during the Vietnam War During the Vietnam War, the U.S. military used large quantities of herbicides to deny their enemies agri-cultural food production and forest cover. Between 1961 and 1971, about 3.2 million acres (1.3 million ha) of for-est and 247,000 acres (100,000 ha) of croplands were sprayed at least once. This is an area equivalent to about one-seventh of South Vietnam.
The most commonly used herbicide was called agent orange, a one-to-one blend of two phenoxy
herbi-Chemical w
arfar
Chemical w
arfar
e
Soldiers at Assaf Harofe Hospital washing “victims” in simulated chemical attack.Photograph by Jeffrey L. Rotman. Corbis.
Reproduced by permission.
There were also severe ecological effects of herbi-cide spraying in the extremely biodiverse upland forests of Vietnam, especially rain forests. Mature tropical forests in this region have many species of hardwood trees. Because this forested ecosystem has such a dense and complexly layered canopy, a single spraying of her-bicide killed only about 10% of the larger trees. Re-sprays of upland forests were often made, however, to achieve a greater and longer-lasting defoliation. To achieve this effect, about 34% of the area of Vietnam that was subjected to herbicides was treated more than once.
The effects on animals of the herbicide spraying in Vietnam were not well documented; however, there are many accounts of sparse populations of birds, mam-mals,reptiles, and other animals in the herbicide-treated mangrove forests and of large decreases in the yield of near-shore fisheries, for which an intact mangrove ecosystem provides important spawning and nursery
habitat. More than a decade after the war, Vietnamese ecologists examined an inland valley that had been con-verted by herbicide spraying from a rich upland tropical forest into a degraded ecosystem dominated by grasses
and shrubs. The secondary, degraded landscape only cides, 2,4-D and 2,4,5-T. Picloram and cacodylic acid
were also used, but in much smaller amounts. In total, this military action used about 25,000 tons of 2,4-D; 21,000 tons of 2,4,5-T; and 1,500 tons of picloram. Agent orange was sprayed at a rateof about 22.3 lb/acre (25 kg/ha), equivalent to about ten times the rate at which those same chemicals were used for plant control in forestry. The spray rate was much more intense dur-ing warfare, because the intention was to destroy the ecosystems through ecocide, rather than to manage them towards a more positive purpose.
The ecological damages caused by the military use of herbicides in Vietnam were not studied in detail; how-ever, cursory surveys were made by some visiting ecolo-gists. These scientists observed that coastal mangrove
