Chemistry

Chemistry

By Alan J. Rocke, Melvyn C. Usselman

Chemistry, the science that deals with the properties, composition, and structure of substances (defined as elements and compounds), the transformations they undergo, and the energy that is released or absorbed during these processes. Every substance, whether naturally occurring or artificially produced, consists of one or more of the hundred-odd species of atoms that have been identified as elements. Although these atoms, in turn, are composed of more elementary particles, they are the basic building blocks of chemical substances; there is no quantity of oxygen, mercury, or gold, for example, smaller than an atom of that substance. Chemistry, therefore, is concerned not with the subatomic domain but with the properties of atoms and the laws governing their combinations and how the knowledge of these properties can be used to achieve specific purposes.
 
The great challenge in chemistry is the development of a coherent explanation of the complex behaviour of materials, why they appear as they do, what gives them their enduring properties, and how interactions among different substances can bring about the formation of new substances and the destruction of old ones. From the earliest attempts to understand the material world in rational terms, chemists have struggled to develop theories of matter that satisfactorily explain both permanence and change. The ordered assembly of indestructible atoms into small and large molecules, or extended networks of intermingled atoms, is generally accepted as the basis of permanence, while the reorganization of atoms or molecules into different arrangements lies behind theories of change. Thus chemistry involves the study of the atomic composition and structural architecture of substances, as well as the varied interactions among substances that can lead to sudden, often violent reactions.
 
Chemistry also is concerned with the utilization of natural substances and the creation of artificial ones. Cooking, fermentation, glass making, and metallurgy are all chemical processes that date from the beginnings of civilization. Today, vinyl, Teflon, liquid crystals, semiconductors, and superconductors represent the fruits of chemical technology. The 20th century has seen dramatic advances in the comprehension of the marvelous and complex chemistry of living organisms, and a molecular interpretation of health and disease holds great promise. Modern chemistry, aided by increasingly sophisticated instruments, studies materials as small as single atoms and as large and complex as DNA (deoxyribonucleic acid), which contains millions of atoms. New substances can even be designed to bear desired characteristics and then synthesized. The rate at which chemical knowledge continues to accumulate is remarkable. Over time more than 8,000,000 different chemical substances, both natural and artificial, have been characterized and produced. The number was less than 500,000 as recently as 1965.
 
Intimately interconnected with the intellectual challenges of chemistry are those associated with industry. In the mid-19th century the German chemist Justus von Liebig commented that the wealth of a nation could be gauged by the amount of sulfuric acid it produced. This acid, essential to many manufacturing processes, remains today the leading chemical product of industrialized countries. As Liebig recognized, a country that produces large amounts of sulfuric acid is one with a strong chemical industry and a strong economy as a whole. The production, distribution, and utilization of a wide range of chemical products is common to all highly developed nations. In fact, one can say that the “iron age” of civilization is being replaced by a “polymer age,” for in some countries the total volume of polymers now produced exceeds that of iron.
 
The scope of chemistry
 
The days are long past when one person could hope to have a detailed knowledge of all areas of chemistry. Those pursuing their interests into specific areas of chemistry communicate with others who share the same interests. Over time a group of chemists with specialized research interests become the founding members of an area of specialization. The areas of specialization that emerged early in the history of chemistry, such as organic, inorganic, physical, analytical, and industrial chemistry, along with biochemistry, remain of greatest general interest. There has been, however, much growth in the areas of polymer, environmental, and medicinal chemistry during the 20th century. Moreover, new specialities continue to appear, as, for example, pesticide, forensic, and computer chemistry.
 
Analytical chemistry
 
Most of the materials that occur on Earth, such as wood, coal, minerals, or air, are mixtures of many different and distinct chemical substances. Each pure chemical substance (e.g., oxygen, iron, or water) has a characteristic set of properties that gives it its chemical identity. Iron, for example, is a common silver-white metal that melts at 1,535° C, is very malleable, and readily combines with oxygen to form the common substances hematite and magnetite. The detection of iron in a mixture of metals, or in a compound such as magnetite, is a branch of analytical chemistry called qualitative analysis. Measurement of the actual amount of a certain substance in a compound or mixture is termed quantitative analysis. Quantitative analytic measurement has determined, for instance, that iron makes up 72.3 percent, by mass, of magnetite, the mineral commonly seen as black sand along beaches and stream banks. Over the years, chemists have discovered chemical reactions that indicate the presence of such elemental substances by the production of easily visible and identifiable products. Iron can be detected by chemical means if it is present in a sample to an amount of 1 part per million or greater. Some very simple qualitative tests reveal the presence of specific chemical elements in even smaller amounts. The yellow colour imparted to a flame by sodium is visible if the sample being ignited has as little as one-billionth of a gram of sodium. Such analytic tests have allowed chemists to identify the types and amounts of impurities in various substances and to determine the properties of very pure materials. Substances used in common laboratory experiments generally have impurity levels of less than 0.1 percent. For special applications, one can purchase chemicals that have impurities totaling less than 0.001 percent. The identification of pure substances and the analysis of chemical mixtures enable all other chemical disciplines to flourish.
 
The importance of analytical chemistry has never been greater than it is today. The demand in modern societies for a variety of safe foods, affordable consumer goods, abundant energy, and labour-saving technologies places a great burden on the environment. All chemical manufacturing produces waste products in addition to the desired substances, and waste disposal has not always been carried out carefully. Disruption of the environment has occurred since the dawn of civilization, and pollution problems have increased with the growth of global population. The techniques of analytical chemistry are relied on heavily to maintain a benign environment. The undesirable substances in water, air, soil, and food must be identified, their point of origin fixed, and safe, economical methods for their removal or neutralization developed. Once the amount of a pollutant deemed to be hazardous has been assessed, it becomes important to detect harmful substances at concentrations well below the danger level. Analytical chemists seek to develop increasingly accurate and sensitive techniques and instruments.
 
Sophisticated analytic instruments, often coupled with computers, have improved the accuracy with which chemists can identify substances and have lowered detection limits. An analytic technique in general use is gas chromatography, which separates the different components of a gaseous mixture by passing the mixture through a long, narrow column of absorbent but porous material. The different gases interact differently with this absorbent material and pass through the column at different rates. As the separate gases flow out of the column, they can be passed into another analytic instrument called a mass spectrometer, which separates substances according to the mass of their constituent ions. A combined gas chromatograph–mass spectrometer can rapidly identify the individual components of a chemical mixture whose concentrations may be no greater than a few parts per billion. Similar or even greater sensitivities can be obtained under favourable conditions using techniques such as atomic absorption, polarography, and neutron activation. The rate of instrumental innovation is such that analytic instruments often become obsolete within 10 years of their introduction. Newer instruments are more accurate and faster and are employed widely in the areas of environmental and medicinal chemistry.
 
Inorganic chemistry
 
Modern chemistry, which dates more or less from the acceptance of the law of conservation of mass in the late 18th century, focused initially on those substances that were not associated with living organisms. Study of such substances, which normally have little or no carbon, constitutes the discipline of inorganic chemistry. Early work sought to identify the simple substances—namely, the elements—that are the constituents of all more complex substances. Some elements, such as gold and carbon, have been known since antiquity, and many others were discovered and studied throughout the 19th and early 20th centuries. Today, more than 100 are known. The study of such simple inorganic compounds as sodium chloride (common salt) has led to some of the fundamental concepts of modern chemistry, the law of definite proportions providing one notable example. This law states that for most pure chemical substances the constituent elements are always present in fixed proportions by mass (e.g., every 100 grams of salt contains 39.3 grams of sodium and 60.7 grams of chlorine). The crystalline form of salt, known as halite, consists of intermingled sodium and chlorine atoms, one sodium atom for each one of chlorine. Such a compound, formed solely by the combination of two elements, is known as a binary compound. Binary compounds are very common in inorganic chemistry, and they exhibit little structural variety. For this reason, the number of inorganic compounds is limited in spite of the large number of elements that may react with each other. If three or more elements are combined in a substance, the structural possibilities become greater.
 
After a period of quiescence in the early part of the 20th century, inorganic chemistry has again become an exciting area of research. Compounds of boron and hydrogen, known as boranes, have unique structural features that forced a change in thinking about the architecture of inorganic molecules. Some inorganic substances have structural features long believed to occur only in carbon compounds, and a few inorganic polymers have even been produced. Ceramics are materials composed of inorganic elements combined with oxygen. For centuries ceramic objects have been made by strongly heating a vessel formed from a paste of powdered minerals. Although ceramics are quite hard and stable at very high temperatures, they are usually brittle. Currently, new ceramics strong enough to be used as turbine blades in jet engines are being manufactured. There is hope that ceramics will one day replace steel in components of internal-combustion engines. In 1987 a ceramic containing yttrium, barium, copper, and oxygen, with the approximate formula YBa2Cu3O7, was found to be a superconductor at a temperature of about 100 K. A superconductor offers no resistance to the passage of an electrical current, and this new type of ceramic could very well find wide use in electrical and magnetic applications. A superconducting ceramic is so simple to make that it can be prepared in a high school laboratory. Its discovery illustrates the unpredictability of chemistry, for fundamental discoveries can still be made with simple equipment and inexpensive materials.
 
Many of the most interesting developments in inorganic chemistry bridge the gap with other disciplines. Organometallic chemistry investigates compounds that contain inorganic elements combined with carbon-rich units. Many organometallic compounds play an important role in industrial chemistry as catalysts, which are substances that are able to accelerate the rate of a reaction even when present in only very small amounts. Some success has been achieved in the use of such catalysts for converting natural gas to related but more useful chemical substances. Chemists also have created large inorganic molecules that contain a core of metal atoms, such as platinum, surrounded by a shell of different chemical units. Some of these compounds, referred to as metal clusters, have characteristics of metals, while others react in ways similar to biologic systems. Trace amounts of metals in biologic systems are essential for processes such as respiration, nerve function, and cell metabolism. Processes of this kind form the object of study of bioinorganic chemistry. Although organic molecules were once thought to be the distinguishing chemical feature of living creatures, it is now known that inorganic chemistry plays a vital role as well.
 
Organic chemistry
 
Organic compounds are based on the chemistry of carbon. Carbon is unique in the variety and extent of structures that can result from the three-dimensional connections of its atoms. The process of photosynthesis converts carbon dioxide and water to oxygen and compounds known as carbohydrates. Both cellulose, the substance that gives structural rigidity to plants, and starch, the energy storage product of plants, are polymeric carbohydrates. Simple carbohydrates produced by photosynthesis form the raw material for the myriad organic compounds found in the plant and animal kingdoms. When combined with variable amounts of hydrogen, oxygen, nitrogen, sulfur, phosphorus, and other elements, the structural possibilities of carbon compounds become limitless, and their number far exceeds the total of all nonorganic compounds. A major focus of organic chemistry is the isolation, purification, and structural study of these naturally occurring substances. Many natural products are simple molecules. Examples include formic acid (HCO2H) in ants, ethyl alcohol (C2H5OH) in fermenting fruit, and oxalic acid (C2H2O4) in rhubarb leaves. Other natural products, such as penicillin, vitamin B12, proteins, and nucleic acids, are exceedingly complex. The isolation of pure natural products from their host organism is made difficult by the low concentrations in which they may be present. Once they are isolated in pure form, however, modern instrumental techniques can reveal structural details for amounts weighing as little as one-millionth of a gram. The correlation of the physical and chemical properties of compounds with their structural features is the domain of physical organic chemistry. Once the properties endowed upon a substance by specific structural units termed functional groups are known, it becomes possible to design novel molecules that may exhibit desired properties. The preparation, under controlled laboratory conditions, of specific compounds is known as synthetic chemistry. Some products are easier to synthesize than to collect and purify from their natural sources. Tons of vitamin C, for example, are synthesized annually. Many synthetic substances have novel properties that make them especially useful. Plastics are a prime example, as are many drugs and agricultural chemicals. A continuing challenge for synthetic chemists is the structural complexity of most organic substances. To synthesize a desired substance, the atoms must be pieced together in the correct order and with the proper three-dimensional relationships. Just as a given pile of lumber and bricks can be assembled in many ways to build houses of several different designs, so too can a fixed number of atoms be connected together in various ways to give different molecules. Only one structural arrangement out of the many possibilities will be identical with a naturally occurring molecule. The antibiotic erythromycin, for example, contains 37 carbon, 67 hydrogen, and 13 oxygen atoms, along with one nitrogen atom. Even when joined together in the proper order, these 118 atoms can give rise to 262,144 different structures, only one of which has the characteristics of natural erythromycin. The great abundance of organic compounds, their fundamental role in the chemistry of life, and their structural diversity have made their study especially challenging and exciting. Organic chemistry is the largest area of specialization among the various fields of chemistry.
 
Biochemistry
 
As understanding of inanimate chemistry grew during the 19th century, attempts to interpret the physiological processes of living organisms in terms of molecular structure and reactivity gave rise to the discipline of biochemistry. Biochemists employ the techniques and theories of chemistry to probe the molecular basis of life. An organism is investigated on the premise that its physiological processes are the consequence of many thousands of chemical reactions occurring in a highly integrated manner. Biochemists have established, among other things, the principles that underlie energy transfer in cells, the chemical structure of cell membranes, the coding and transmission of hereditary information, muscular and nerve function, and biosynthetic pathways. In fact, related biomolecules have been found to fulfill similar roles in organisms as different as bacteria and human beings. The study of biomolecules, however, presents many difficulties. Such molecules are often very large and exhibit great structural complexity; moreover, the chemical reactions they undergo are usually exceedingly fast. The separation of the two strands of DNA, for instance, occurs in one-millionth of a second. Such rapid rates of reaction are possible only through the intermediary action of biomolecules called enzymes. Enzymes are proteins that owe their remarkable rate-accelerating abilities to their three-dimensional chemical structure. Not surprisingly, biochemical discoveries have had a great impact on the understanding and treatment of disease. Many ailments due to inborn errors of metabolism have been traced to specific genetic defects. Other diseases result from disruptions in normal biochemical pathways.
 
Frequently, symptoms can be alleviated by drugs, and the discovery, mode of action, and degradation of therapeutic agents is another of the major areas of study in biochemistry. Bacterial infections can be treated with sulfonamides, penicillins, and tetracyclines, and research into viral infections has revealed the effectiveness of acyclovir against the herpes virus. There is much current interest in the details of carcinogenesis and cancer chemotherapy. It is known, for example, that cancer can result when cancer-causing molecules, or carcinogens as they are called, react with nucleic acids and proteins and interfere with their normal modes of action. Researchers have developed tests that can identify molecules likelyto be carcinogenic. The hope, of course, is that progress in the prevention and treatment of cancer will accelerate once the biochemical basis of the disease is more fully understood.
 
The molecular basis of biologic processes is an essential feature of the fast-growing disciplines of molecular biology and biotechnology. Chemistry has developed methods for rapidly and accurately determining the structure of proteins and DNA. In addition, efficient laboratory methods for the synthesis of genes are being devised. Ultimately, the correction of genetic diseases by replacement of defective genes with normal ones may become possible.
 
Polymer chemistry
 
The simple substance ethylene is a gas composed of molecules with the formula CH2CH2. Under certain conditions, many ethylene molecules will join together to form a long chain called polyethylene, with the formula (CH2CH2)n, where n is a variable but large number. Polyethylene is a tough, durable solid material quite different from ethylene. It is an example of a polymer, which is a large molecule made up of many smaller molecules (monomers), usually joined together in a linear fashion. Many naturally occurring substances, including cellulose, starch, cotton, wool, rubber, leather, proteins, and DNA, are polymers. Polyethylene, nylon, and acrylics are examples of synthetic polymers. The study of such materials lies within the domain of polymer chemistry, a specialty that has flourished in the 20th century. The investigation of natural polymers overlaps considerably with biochemistry, but the synthesis of new polymers, the investigation of polymerization processes, and the characterization of the structure and properties of polymeric materials all pose unique problems for polymer chemists.
 
Polymer chemists have designed and synthesized polymers that vary in hardness, flexibility, softening temperature, solubility in water, and biodegradability. They have produced polymeric materials that are as strong as steel yet lighter and more resistant to corrosion. Oil, natural gas, and water pipelines are now routinely constructed of plastic pipe. In recent years, automakers have increased their use of plastic components to build lighter vehicles that consume less fuel. Other industries such as those involved in the manufacture of textiles, rubber, paper, and packaging materials are built upon polymer chemistry.
 
Besides producing new kinds of polymeric materials, researchers are concerned with developing special catalysts that are required by the large-scale industrial synthesis of commercial polymers. Without such catalysts, the polymerization process would be very slow in certain cases.
 
Physical chemistry
 
Many chemical disciplines, such as those already discussed, focus on certain classes of materials that share common structural and chemical features. Other specialties may be centred not on a class of substances but rather on their interactions and transformations. The oldest of these fields is physical chemistry, which seeks to measure, correlate, and explain the quantitative aspects of chemical processes. The Anglo-Irish chemist Robert Boyle, for example, discovered in the 17th century that at room temperature the volume of a fixed quantity of gas decreases proportionally as the pressure on it increases. Thus, for a gas at constant temperature, the product of its volume V and pressure P equals a constant number—i.e., PV = constant. Such a simple arithmetic relationship is valid for nearly all gases at room temperature and at pressures equal to or less than one atmosphere. Subsequent work has shown that the relationship loses its validity at higher pressures, but more complicated expressions that more accurately match experimental results can be derived. The discovery and investigation of such chemical regularities, often called laws of nature, lie within the realm of physical chemistry. For much of the 18th century the source of mathematical regularity in chemical systems was assumed to be the continuum of forces and fields that surround the atoms making up chemical elements and compounds.
 
Developments in the 20th century, however, have shown that chemical behaviour is best interpreted by a quantum mechanical model of atomic and molecular structure. The branch of physical chemistry that is largely devoted to this subject is theoretical chemistry. Theoretical chemists make extensive use of computers to help them solve complicated mathematical equations. Other branches of physical chemistry include chemical thermodynamics, which deals with the relationship between heat and other forms of chemical energy, and chemical kinetics, which seeks to measure and understand the rates of chemical reactions. Electrochemistry investigates the interrelationship of electric current and chemical change. The passage of an electric current through a chemical solution causes changes in the constituent substances that are often reversible—i.e., under different conditions the altered substances themselves will yield an electric current. Common batteries contain chemical substances that, when placed in contact with each other by closing an electrical circuit, will deliver current at a constant voltage until the substances are consumed. At present there is much interest in devices that can use the energy in sunlight to drive chemical reactions whose products are capable of storing the energy. The discovery of such devices would make possible the widespread utilization of solar energy.
 
There are many other disciplines within physical chemistry that are concerned more with the general properties of substances and the interactions among substances than with the substances themselves. Photochemistry is a specialty that investigates the interaction of light with matter. Chemical reactions initiated by the absorption of light can be very different from those that occur by other means. Vitamin D, for example, is formed in the human body when the steroid ergosterol absorbs solar radiation; ergosterol does not change to vitamin D in the dark.
 
A rapidly developing subdiscipline of physical chemistry is surface chemistry. It examines the properties of chemical surfaces, relying heavily on instruments that can provide a chemical profile of such surfaces. Whenever a solid is exposed to a liquid or a gas, a reaction occurs initially on the surface of the solid, and its properties can change dramatically as a result. Aluminum is a case in point: it is resistant to corrosion precisely because the surface of the pure metal reacts with oxygen to form a layer of aluminum oxide, which serves to protect the interior of the metal from further oxidation. Numerous reaction catalysts perform their function by providing a reactive surface on which substances can react.
 
Industrial chemistry
 
The manufacture, sale, and distribution of chemical products is one of the cornerstones of a developed country. Chemists play an important role in the manufacture, inspection, and safe handling of chemical products, as well as in product development and general management. The manufacture of basic chemicals such as oxygen, chlorine, ammonia, and sulfuric acid provides the raw materials for industries producing textiles, agricultural products, metals, paints, and pulp and paper. Specialty chemicals are produced in smaller amounts for industries involved with such products as pharmaceuticals, foodstuffs, packaging, detergents, flavours, and fragrances. To a large extent, the chemical industry takes the products and reactions common to “bench-top” chemical processes and scales them up to industrial quantities.
 
The monitoring and control of bulk chemical processes, especially with regard to heat transfer, pose problems usually tackled by chemists and chemical engineers. The disposal of by-products also is a major problem for bulk chemical producers. These and other challenges of industrial chemistry set it apart from the more purely intellectual disciplines of chemistry discussed above. Yet, within the chemical industry, there is a considerable amount of fundamental research undertaken within traditional specialties. Most large chemical companies have research-and-development capability. Pharmaceutical firms, for example, operate large research laboratories in which chemists test molecules for pharmacological activity. The new products and processes that are discovered in such laboratories are often patented and become a source of profit for the company funding the research. A great deal of the research conducted in the chemical industry can be termed applied research because its goals are closely tied to the products and processes of the company concerned. New technologies often require much chemical expertise. The fabrication of, say, electronic microcircuits involves close to 100 separate chemical steps from start to finish. Thus, the chemical industry evolves with the technological advances of the modern world and at the same time often contributes to the rate of progress.
 
The methodology of chemistry
 
Chemistry is to a large extent a cumulative science. Over time the number and extent of observations and phenomena studied increase. Not all hypotheses and discoveries endure unchallenged, however. Some of them are discarded as new observations or more satisfying explanations appear. Nonetheless, chemistry has a broad spectrum of explanatory models for chemical phenomena that have endured and been extended over time. These now have the status of theories, interconnected sets of explanatory devices that correlate well with observed phenomena. As new discoveries are made, they are incorporated into existing theory whenever possible. However, as the discovery of high-temperature superconductors in 1986 illustrates, accepted theory is never sufficient to predict the course of future discovery. Serendipity, or chance discovery, will continue to play as much a role in the future as will theoretical sophistication.
 
Studies of molecular structure
 
The chemical properties of a substance are a function of its structure, and the techniques of X-ray crystallography now enable chemists to determine the precise atomic arrangement of complex molecules. A molecule is an ordered assembly of atoms. Each atom in a molecule is connected to one or more neighbouring atoms by a chemical bond. The length of bonds and the angles between adjacent bonds are all important in describing molecular structure, and a comprehensive theory of chemical bonding is one of the major achievements of modern chemistry. Fundamental to bonding theory is the atomic–molecular concept.
 
Atoms and Elements
 
As far as general chemistry is concerned, atoms are composed of the three fundamental particles: the proton, the neutron, and the electron. Although the proton and the neutron are themselves composed of smaller units, their substructure has little impact on chemical transformation. As was explained in an earlier section, the proton carries a charge of +1, and the number of protons in an atomic nucleus distinguishes one type of chemical atom from another. The simplest atom of all, hydrogen, has a nucleus composed of a single proton. The neutron has very nearly the same mass as the proton, but it has no charge. Neutrons are contained with protons in the nucleus of all atoms other than hydrogen. The atom with one proton and one neutron in its nucleus is called deuterium. Because it has only one proton, deuterium exhibits the same chemical properties as hydrogen but has a different mass. Hydrogen and deuterium are examples of related atoms called isotopes. The third atomic particle, the electron, has a charge of -1, but its mass is 1,836 times smaller than that of a proton. The electron occupies a region of space outside the nucleus termed an orbital. 
 
Some orbitals are spherical with the nucleus at the centre. Because electrons have so little mass and move about at speeds close to half that of light, they exhibit the same wave–particle duality as photons of light. This means that some of the properties of an electron are best described by considering the electron to be a particle, while other properties are consistent with the behaviour of a standing wave. The energy of a standing wave, such as a vibrating string, is distributed over the region of space defined by the two fixed ends and the up-and-down extremes of vibration. Such a wave does not exist in a fixed region of space as does a particle. Early models of atomic structure envisioned the electron as a particle orbiting the nucleus, but electron orbitals are now interpreted as the regions of space occupied by standing waves called wave functions. These wave functions represent the regions of space around the nucleus in which the probability of finding an electron is high. They play an important role in bonding theory, as will be discussed later.
 
Each proton in an atomic nucleus requires an electron for electrical neutrality. Thus, as the number of protons in a nucleus increases, so too does the number of electrons. The electrons, alone or in pairs, occupy orbitals increasingly distant from the nucleus. Electrons farther from the nucleus are attracted less strongly by the protons in the nucleus, and they can be removed more easily from the atom. The energy required to move an electron from one orbital to another, or from one orbital to free space, gives a measure of the energy level of the orbitals. These energies have been found to have distinct, fixed values; they are said to be quantized. The energy differences between orbitals give rise to the characteristic patterns of light absorption or emission that are unique to each chemical atom.
A new chemical atom—that is, an element—results each time another proton is added to an atomic nucleus. Consecutive addition of protons generates the whole range of elements known to exist in the universe. Compounds are formed when two or more different elements combine through atomic bonding. Such bond formation is a consequence of electron pairing and constitutes the foundation of all structural chemistry.
 
Ionic and Covalent Bonding
 
When two different atoms approach each other, the electrons in their outer orbitals can respond in two distinct ways. An electron in the outermost atomic orbital of atom A may move completely to an outer but stabler orbital of atom B. The charged atoms that result, A+ and B-, are called ions, and the electrostatic force of attraction between them gives rise to what is termed an ionic bond. Most elements can form ionic bonds, and the substances that result commonly exist as three-dimensional arrays of positive and negative ions. Ionic compounds are frequently crystalline solids that have high melting points (e.g., table salt).
 
The second way in which the two outer electrons of atoms A and B can respond to the approach of A and B is to pair up to form a covalent bond. In the simple view known as the valence-bond model, in which electrons are treated strictly as particles, the two paired electrons are assumed to lie between the two nuclei and are shared equally by atoms A and B, resulting in a covalent bond. Atoms joined together by one or more covalent bonds constitute molecules. Hydrogen gas is composed of hydrogen molecules, which consist in turn of two hydrogen atoms linked by a covalent bond. The notation H2 for hydrogen gas is referred to as a molecular formula. Molecular formulas indicate the number and type of atoms that make up a molecule. The molecule H2 is responsible for the properties generally associated with hydrogen gas. Most substances on Earth have covalently bonded molecules as their fundamental chemical unit, and their molecular properties are completely different from those of the constituent elements. The physical and chemical properties of carbon dioxide, for example, are quite distinct from those of pure carbon and pure oxygen.
 
The interpretation of a covalent bond as a localized electron pair is an oversimplification of the bonding situation. A more comprehensive description of bonding that considers the wave properties of electrons is the molecular-orbital theory. According to this theory, electrons in a molecule, rather than being localized between atoms, are distributed over all the atoms in the molecule in a spatial distribution described by a molecular orbital. Such orbitals result when the atomic orbitals of bonded atoms combine with each other. The total number of molecular orbitals present in a molecule is equal to the sum of all atomic orbitals in the constituent atoms prior to bonding. Thus, for the simple combination of atoms A and B to form the molecule AB, two atomic orbitals combine to generate two molecular orbitals. One of these, the so-called bonding molecular orbital, represents a region of space enveloping both the A and B atoms, while the other, the anti-bonding molecular orbital, has two lobes, neither of which occupies the space between the two atoms. The bonding molecular orbital is at a lower energy level than are the two atomic orbitals, while the anti-bonding orbital is at a higher energy level. The two paired electrons that constitute the covalent bond between A and B occupy the bonding molecular orbital. For this reason, there is a high probability of finding the electrons between A and B, but they can be found elsewhere in the orbital as well. Because only two electrons are involved in bond formation and both can be accommodated in the lower energy orbital, the anti-bonding orbital remains unpopulated. This theory of bonding predicts that bonding between A and B will occur because the energy of the paired electrons after bonding is less than that of the two electrons in their atomic orbitals prior to bonding. The formation of a covalent bond is thus energetically favoured. The system goes from a state of higher energy to one of lower energy.
 
Another feature of this bonding picture is that it is able to predict the energy required to move an electron from the bonding molecular orbital to the anti-bonding one. The energy required for such an electronic excitation can be provided by visible light, for example, and the wavelength of the light absorbed determines the colour displayed by the absorbing molecule (e.g., violets are blue because the pigments in the flower absorb the red rays of natural light and reflect more of the blue). As the number of atoms in a molecule increases, so too does the number of molecular orbitals. Calculation of molecular orbitals for large molecules is mathematically difficult, but computers have made it possible to determine the wave equations for several large molecules. Molecular properties predicted by such calculations correlate well with experimental results.
 
Isomerism
 
Many elements can form two or more covalent bonds, but only a few are able to form extended chains of covalent bonds. The outstanding example is carbon, which can form as many as four covalent bonds and can bond to itself indefinitely. Carbon has six electrons in total, two of which are paired in an atomic orbital closest to the nucleus. The remaining four are farther from the nucleus and are available for covalent bonding. When there is sufficient hydrogen present, carbon will react to form methane, CH4. When all four electron pairs occupy the four molecular orbitals of lowest energy, the molecule assumes the shape of a tetrahedron, with carbon at the centre and the four hydrogen atoms at the apexes. The C–H bond length is 110 picometres (1 picometre = 10-12 metre), and the angle between adjacent C–H bonds is close to 110°. Such tetrahedral symmetry is common to many carbon compounds and results in interesting structural possibilities. If two carbon atoms are joined together, with three hydrogen atoms bonded to each carbon atom, the molecule ethane is obtained. When four carbon atoms are joined together, two different structures are possible: a linear structure designated n-butane and a branched structure called iso-butane. These two structures have the same molecular formula, C4H10, but a different order of attachment of their constituent atoms. The two molecules are termed structural isomers. Each of them has unique chemical and physical properties, and they are different compounds. The number of possible isomers increases rapidly as the number of carbon atoms increases. There are five isomers for C6H14, 75 for C10H22, and 6.2 × 1013 for C40H82. When carbon forms bonds to atoms other than hydrogen, such as oxygen, nitrogen, and sulfur, the structural possibilities become even greater. It is this great potential for structural diversity that makes carbon compounds essential to living organisms.
 
Even when the bonding sequence of carbon compounds is fixed, further structural variation is still possible. When two carbon atoms are joined together by two bonding pairs of electrons, a double bond is formed. A double bond forces the two carbon atoms and attached groups into a rigid, planar structure. As a result, a molecule such as CHCl=CHCl can exist in two nonidentical forms called geometric isomers. Structural rigidity also occurs in ring structures, and attached groups can be on the same side of a ring or on different sides. Yet another opportunity for isomerism arises when a carbon atom is bonded to four different groups. These can be attached in two different ways, one of which is the mirror image of the other. This type of isomerism is called optical isomerism, because the two isomers affect plane-polarized light differently. Two optical isomers are possible for every carbon atom that is bonded to four different groups. For a molecule bearing 10 such carbon atoms, the total number of possible isomers will be 210 = 1,024. Large biomolecules often have 10 or more carbon atoms for which such optical isomers are possible. Only one of all the possible isomers will be identical to the natural molecule. For this reason, the laboratory synthesis of large organic molecules is exceedingly difficult. Only in the last few decades of the 20th century have chemists succeeded in developing reagents and processes that yield specific optical isomers. They expect that new synthetic methods will make possible the synthesis of ever more complex natural products.
 

Investigations of chemical transformations

 
Basic Factors
 
The structure of ionic substances and covalently bonded molecules largely determines their function. As noted above, the properties of a substance depend on the number and type of atoms it contains and on the bonding patterns present. Its bulk properties also depend, however, on the interactions among individual atoms, ions, or molecules. The force of attraction between the fundamental units of a substance dictate whether, at a given temperature and pressure, that substance will exist in the solid, liquid, or gas phase. At room temperature and pressure, for example, the strong forces of attraction between the positive ions of sodium (Na+) and the negative ions of chlorine (Cl−) draw them into a compact solid structure. The weaker forces of attraction among neighbouring water molecules allow the looser packing characteristic of a liquid. Finally, the very weak attractive forces acting among adjacent oxygen molecules are exceeded by the dispersive forces of heat; oxygen, consequently, is a gas. Interparticle forces thus affect the chemical and physical behaviour of substances, but they also determine to a large extent how a particle will respond to the approach of a different particle. If the two particles react with each other to form new particles, a chemical reaction has occurred.
 
Not with standing the unlimited structural diversity allowed by molecular bonding, the world would be devoid of life if substances were incapable of change. The study of chemical transformation, which complements the study of molecular structure, is built on the concepts of energy and entropy.
 
Energy and The First Law of Thermodynamics
 
The concept of energy is a fundamental and familiar one in all the sciences. In simple terms, the energy of a body represents its ability to do work, and work itself is a force acting over a distance.
 
Chemical systems can have both kinetic energy (energy of motion) and potential energy (stored energy). The kinetic energy possessed by any collection of molecules in a solid, liquid, or gas is known as its thermal energy. Since liquids expand when they have more thermal energy, a liquid column of mercury, for example, will rise higher in an evacuated tube as it becomes warmer. In this way a thermometer can be used to measure the thermal energy, or temperature, of a system. The temperature at which all molecular motion comes to a halt is known as absolute zero.
 
Energy also may be stored in atoms or molecules as potential energy. When protons and neutrons combine to form the nucleus of a certain element, the reduction in potential energy is matched by the production of a huge quantity of kinetic energy. Consider, for instance, the formation of the deuterium nucleus from one proton and one neutron. The fundamental mass unit of the chemist is the mole, which represents the mass, in grams, of 6.02 × 1023 individual particles, whether they be atoms or molecules. One mole of protons has a mass of 1.007825 grams and one mole of neutrons has a mass of 1.008665 grams. By simple addition the mass of one mole of deuterium atoms (ignoring the negligible mass of one mole of electrons) should be 2.016490 grams. The measured mass is 0.00239 gram less than this. The missing mass is known as the binding energy of the nucleus and represents the mass equivalent of the energy released by nucleus formation. By using Einstein’s formula for the conversion of mass to energy (E = mc2), one can calculate the energy equivalent of 0.00239 gram as 2.15 × 108 kilojoules. This is approximately 240,000 times greater than the energy released by the combustion of one mole of methane. Such studies of the energetics of atom formation and interconversion are part of a specialty known as nuclear chemistry.
 
The energy released by the combustion of methane is about 900 kilojoules per mole. Although much less than the energy released by nuclear reactions, the energy given off by a chemical process such as combustion is great enough to be perceived as heat and light. Energy is released in so-called exothermic reactions because the chemical bonds in the product molecules, carbon dioxide and water, are stronger and stabler than those in the reactant molecules, methane and oxygen. The chemical potential energy of the system has decreased, and most of the released energy appears as heat, while some appears as radiant energy, or light. The heat produced by such a combustion reaction will raise the temperature of the surrounding air and, at constant pressure, increase its volume. This expansion of air results in work being done. In the cylinder of an internal-combustion engine, for example, the combustion of gasoline results in hot gases that expand against a moving piston. The motion of the piston turns a crankshaft, which then propels the vehicle. In this case, chemical potential energy has been converted to thermal energy, some of which produces useful work. This process illustrates a statement of the conservation of energy known as the first law of thermodynamics. This law states that, for an exothermic reaction, the energy released by the chemical system is equal to the heat gained by the surroundings plus the work performed. By measuring the heat and work quantities that accompany chemical reactions, it is possible to ascertain the energy differences between the reactants and the products of various reactions. In this manner, the potential energy stored in a variety of molecules can be determined, and the energy changes that accompany chemical reactions can be calculated.
 
Entropy and The Second Law of Thermodynamics
 
Some chemical processes occur even though there is no net energy change. Consider a vessel containing a gas, connected to an evacuated vessel via a channel wherein a barrier obstructs passage of the gas. If the barrier is removed, the gas will expand into the evacuated vessel. This expansion is consistent with the observation that a gas always expands to fill the volume available. When the temperature of both vessels is the same, the energy of the gas before and after the expansion is the same. The reverse reaction does not occur, however. The spontaneous reaction is the one that yields a state of greater disorder. In the expanded volume, the individual gas molecules have greater freedom of movement and thus are more disordered. The measure of the disorder of a system is a quantity termed entropy. At a temperature of absolute zero, all movement of atoms and molecules ceases, and the disorder—and entropy—of such perfectly compacted substances is zero. (Zero entropy at zero temperature is in accord with the third law of thermodynamics.) All substances above absolute zero will have a positive entropy value that increases with temperature. When a hot body cools down, the thermal energy it loses passes to the surrounding air, which is at a lower temperature. As the entropy of the cooling body decreases, the entropy of the surrounding air increases. In fact, the increase in entropy of the air is greater than the decrease in entropy of the cooling body. This is consistent with the second law, which states that the total entropy of a system and its surroundings always increases in a spontaneous reaction. Thus the first and second laws of thermodynamics indicate that, for all processes of chemical change throughout the universe, energy is conserved but entropy increases.
 
Application of the laws of thermodynamics to chemical systems allows chemists to predict the behaviour of chemical reactions. When energy and entropy considerations favour the formation of product molecules, reagent molecules will act to form products until an equilibrium is established between products and reagents. The ratio of products to reagents is specified by a quantity known as an equilibrium constant, which is a function of the energy and entropy differences between the two. What thermodynamics cannot predict, however, is the rate at which chemical reactions occur. For fast reactions an equilibrium mixture of products and reagents can be established in one millisecond or less; for slow reactions the time required could be hundreds of years.
 
Rates of Reaction
 
When the specific rates of chemical reactions are measured experimentally, they are found to be dependent on the concentrations of reacting species, temperature, and a quantity called activation energy. Chemists explain this phenomenon by recourse to the collision theory of reaction rates. This theory builds on the premise that a reaction between two or more chemicals requires, at the molecular level, a collision between two rapidly moving molecules. If the two molecules collide in the right way and with enough kinetic energy, one of the molecules may acquire enough energy to initiate the bond-breaking process. As this occurs, new bonds may begin to form, and ultimately reagent molecules are converted into product molecules. The point of highest energy during bond breaking and bond formation is called the transition state of the molecular process. The difference between the energy of the transition state and that of the reacting molecules is the activation energy that must be exceeded for a reaction to occur. Reaction rates increase with temperature because the colliding molecules have greater energies, and more of them will have energies that exceed the activation energy of reaction. The modern study of the molecular basis of chemical change has been greatly aided by lasers and computers. It is now possible to study short-lived collision products and to better determine the molecular mechanisms that fix the rate of chemical reactions. This knowledge is useful in designing new catalysts that can accelerate the rate of reaction by lowering the activation energy. Catalysts are important for many biochemical and industrial processes because they speed up reactions that ordinarily occur too slowly to be useful. Moreover, they often do so with increased control over the structural features of the product molecules. A rhodium phosphine catalyst, for example, has enabled chemists to obtain 96 percent of the correct optical isomer in a key step in the synthesis of L-dopa, a drug used for treating Parkinson’s disease.
 
Chemistry and society
 
For the first two-thirds of the 20th century, chemistry was seen by many as the science of the future. The potential of chemical products for enriching society appeared to be unlimited. Increasingly, however, and especially in the public mind, the negative aspects of chemistry have come to the fore. Disposal of chemical by-products at waste-disposal sites of limited capacity has resulted in environmental and health problems of enormous concern. The legitimate use of drugs for the medically supervised treatment of diseases has been tainted by the growing misuse of mood-altering drugs. The very word chemicals has come to be used all too frequently in a pejorative sense. There is, as a result, a danger that the pursuit and application of chemical knowledge may be seen as bearing risks that outweigh the benefits.
 
It is easy to underestimate the central role of chemistry in modern society, but chemical products are essential if the world’s population is to be clothed, housed, and fed. The world’s reserves of fossil fuels (e.g., oil, natural gas, and coal) will eventually be exhausted, some as soon as the 21st century, and new chemical processes and materials will provide a crucial alternative energy source. The conversion of solar energy to more concentrated, useful forms, for example, will rely heavily on discoveries in chemistry. Long-term, environmentally acceptable solutions to pollution problems are not attainable without chemical knowledge. There is much truth in the aphorism that “chemical problems require chemical solutions.” Chemical inquiry will lead to a better understanding of the behaviour of both natural and synthetic materials and to the discovery of new substances that will help future generations better supply their needs and deal with their problems.
 
Progress in chemistry can no longer be measured only in terms of economics and utility. The discovery and manufacture of new chemical goods must continue to be economically feasible but must be environmentally acceptable as well. The impact of new substances on the environment can now be assessed before large-scale production begins, and environmental compatibility has become a valued property of new materials. For example, compounds consisting of carbon fully bonded to chlorine and fluorine, called chlorofluorocarbons (or Freons), were believed to be ideal for their intended use when they were first discovered. They are nontoxic, nonflammable gases and volatile liquids that are very stable. These properties led to their widespread use as solvents, refrigerants, and propellants in aerosol containers. Time has shown, however, that these compounds decompose in the upper regions of the atmosphere and that the decomposition products act to destroy stratospheric ozone. Limits have now been placed on the use of chlorofluorocarbons, but it is impossible to recover the amounts already dispersed into the atmosphere.
 
The chlorofluorocarbon problem illustrates how difficult it is to anticipate the overall impact that new materials can have on the environment. Chemists are working to develop methods of assessment, and prevailing chemical theory provides the working tools. Once a substance has been identified as hazardous to the existing ecological balance, it is the responsibility of chemists to locate that substance and neutralize it, limiting the damage it can do or removing it from the environment entirely. The last years of the 20th century will see many new, exciting discoveries in the processes and products of chemistry. Inevitably, the harmful effects of some substances will outweigh their benefits, and their use will have to be limited. Yet, the positive impact of chemistry on society as a whole seems beyond doubt.
 
The history of chemistry
 
Chemistry has justly been called the central science. Chemists study the various substances in the world, with a particular focus on the processes by which one substance is transformed into another. Today, chemistry is defined as the study of the composition and properties of elements and compounds, the structure of their molecules, and the chemical reactions that they undergo. Rather than starting with such modern concepts, though, a fuller appreciation of the subject requires an examination of the historical processes that led to these concepts.
 
Philosophy of matter in antiquity
 
Indeed, the philosophers of antiquity could have had no notion that all matter consists of the combinations of a few dozen elements as they are understood today. The earliest critical thinking on the nature of substances, as far as the historical record indicates, was by certain Greek philosophers beginning about 600 bce. Thales of Miletus, Anaximander, Empedocles, and others propounded theories that the world consisted of varieties of earth, water, air, fire, or indeterminate “seeds” or “unbounded” matter. Leucippus and Democritus propounded a materialistic theory of invisibly tiny irreducible atoms from which the world was made. In the 4th century bce, Plato (influenced by Pythagoreanism) taught that the world of the senses was but the shadow of a mathematical world of “forms” beyond human perception.
 
In contrast, Plato’s student Aristotle took the world of the senses seriously. Adopting Empedocles’s view that the terrestrial region consisted of earth, water, air, and fire, Aristotle taught that each of these materials was a combination of qualities such as hot, cold, moist, and dry. For Aristotle, these “elements” were not building blocks of matter as they are thought of now; rather, they resulted from the qualities imposed on otherwise featureless prime matter. Consequently, there were many different kinds of earth, for instance, and nothing precluded one element from being transformed into another by appropriate adjustment of its qualities. Thus, Aristotle rejected the speculations of the ancient atomists and their irreducible fundamental particles. His views were highly regarded in late antiquity and remained influential throughout the Middle Ages.
 
For thousands of years before Aristotle, metalsmiths, assayers, ceramists, and dyers had worked to perfect their crafts using empirically derived knowledge of chemical processes. By Hellenistic and Roman times, their skills were well advanced, and sophisticated ceramics, glasses, dyes, drugs, steels, bronze, brass, alloys of gold and silver, foodstuffs, and many other chemical products were traded. Hellenistic Alexandria in Egypt was a centre for these arts, and it was apparently there that a group of ideas emerged that later became known as alchemy.
 
Alchemy
 
Three different sets of ideas and skills fed into the origin of alchemy. First was the empirical sophistication of jewelers, gold- and silversmiths, and other artisans who had learned how to fashion precious and semiprecious materials. Among their skills were smelting, assaying, alloying, gilding, amalgamating, distilling, sublimating, painting, and lacquering. The second component was the early Greek theory of matter, especially Aristotelian philosophy, which suggested the possibility of unlimited transformability of one kind of matter into another. The third of alchemy’s roots consisted of a complex combination of ideas derived from Asian philosophies and religions, Hellenistic mystery religions, and what became known as the Hermetic writings (a body of pseudonymous Greek writings on magic, astrology, and alchemy ascribed to the Egyptian god Thoth or his Greek counterpart Hermes Trismegistos). It is important to note, however, that Hellenistic Egypt is only one of several candidates for the homeland of alchemy; at about the same time, similar ideas were developing in Persia, China, and elsewhere.
 
In general, alchemists sought to manipulate the properties of matter in order to prepare more valuable substances. Their most familiar quest was to find the philosopher’s stone, a magical substance that would transmute ordinary metals such as copper, tin, iron, or lead into silver or gold. Important materials in this craft included sulfur, mercury, and electrum (a gold-silver alloy). However, many other alchemists spurned alchemical transmutation (aurifaction), devoting their efforts instead to a pharmaceutical preparation known as the “elixir of life” that would cure any disease, including the ultimate disease, death. The philosopher’s stone and the elixir of life could be considered parallel quests, for each would “cure” metallic or human bodies, respectively, yielding immortal perfection. There was a parallel religious dimension to all this as well. Finally, some alchemists spurned material manipulations entirely, devoting themselves to meditation with the goal of achieving spiritual purity and ultimate redemption.
 
After the rise of Islam, Arabic-speaking scholars of the 9th century translated Greek scientific and philosophical works into their own language. Thereafter, philosophers in the Islamic world pursued chemical and alchemical ideas with enthusiasm and success. The sizable number of modern chemical words derived from Arabic—alcohol, alkali, alchemy, zircon, elixir, natron, and others—suggests the importance of this period for the history of chemistry. One of the leading ideas of medieval Arabic alchemy was the theory that all metals were formed of sulfur and mercury in various proportions and that altering those proportions could transform the metal under study—even to produce silver or gold from lead or iron. Not every alchemist, however, believed in the possibility of such transmutations.
 
Later, scholars in Christian western Europe learned of ancient Greek and early medieval Arabic philosophy by translating these books into Latin. Thus, the alchemical tradition, along with the rest of the Greco-Arabic philosophical and scientific corpus, passed to the West in the course of the 12th century. Well-known Scholastic philosophers of the 13th century, such as Roger Bacon in England and Albertus Magnus in Germany and France, wrote on alchemy. Alongside this learned literature, the empirical chemical arts continued to flourish and comprised a largely separate realm of expertise among artisans, engineers, and mechanics.
 
An important Western alchemist of the late 13th century was the pseudonymous Latin writer who called himself Geber in homage to the 8th-century Arab alchemist Jābir ibn Ḥayyān. Geber was the first to record methods for the preparation and use of sulfuric acid, nitric acid, and hydrochloric acid; the earliest clear evidence for widespread familiarity with distilled alcohol also does not much predate his day. These substances could only have been produced by novel stills that were more robust and efficient than their predecessors, and the appearance of these remarkable new materials produced dramatic changes in the repertoire of chemists.
 
The Renaissance saw even stronger interest in the science. The German-Swiss physician Paracelsus practiced alchemy, Kabbala, astrology, and magic, and in the first half of the 16th century he championed the role of mineral rather than herbal remedies. His emphasis on chemicals in pharmacy and medicine was influential on later figures, and lively controversies over the Paracelsian approach raged around the turn of the 17th century. Gradually the Hermetic influence declined in Europe, however, as certain celebrated feats of putative aurifaction were revealed as frauds.
 
It would be a mistake to think that open-minded empirical investigation that is well integrated with theory (which is how one might define science) was absent from the history of alchemy. Alchemy had many quite scientific practitioners through the centuries, notably including Britain’s Robert Boyle and Isaac Newton—heroes of the scientific revolution of the 17th century—who applied systematic and quantitative method to their (mostly secret) alchemical studies. Indeed, as late as the end of the 17th century there was little to distinguish alchemy from chemistry, either substantively or semantically, since both words were applied to the same set of ideas. It was only in the early 18th century that chemists conferred different definitions on the two words, banishing alchemy to the ashbin of discredited occult pseudosciences.
 
Phlogiston theory
 
This shift was partly simple self-promotion by chemists in the new environment of the Enlightenment, whose vanguard glorified rationalism, experiment, and progress while demonizing the mystical. However, it was also becoming ever clearer that certain central ideas of alchemy (especially metallic transmutation) had never been demonstrated. One of the leaders in this regard was the German physician and chemist Georg Ernst Stahl, who vigorously attacked alchemy (after dabbling in it himself) and proposed an expansive new chemical theory. Stahl noted parallels between the burning of combustible materials and the calcination of metals—the conversion of a metal into its calx, or oxide. He suggested that both processes consisted of the loss of a material fluid, contained within all combustibles, called phlogiston.
 
Phlogiston became the centrepiece of a broad-ranging theory that dominated 18th-century chemical thought. Phlogiston, in short, was thought to be a material substance that defined combustibility. When metallic iron becomes red rust, it loses its phlogiston, just as a burning log does. The ashes of the log and the red rust “ashes” (calx) of iron can no longer burn because they no longer contain the principle of combustibility, or phlogiston. But iron calx can be converted back to the metal if it is strongly heated in the presence of a phlogiston-rich substance such as charcoal. The charcoal donates its phlogiston (becoming ashes itself), while the calx turns into molten metallic iron. Thus, smelting (reduction) of metallic ores could also be understood in phlogistic terms. Later phlogistonists added respiration to the number of phenomena that the theory could elucidate. An animal breathes air, emitting phlogiston in an analogy to a slow fire, fueled by the phlogiston-rich food it consumes. Earth’s atmosphere avoids excess accumulation of phlogiston because plants incorporate it into combustible plant tissues that can then be used as animal food. Combustion, calcination, or respiration eventually cease in an enclosed space because air has a limited capacity to absorb the phlogiston emitted from the burning, calcining, or respiring entity.
 
The phlogiston theory became popular both because of its great success in explaining phenomena and guiding further investigation and because of a certain Enlightenment predilection for materialistic physical theories (the putative fluid of heat became known as caloric, and there were other suggested fluids of electricity, light, and so on). This materialist-mechanist trend can also be seen in the diffuse but powerful influence of Newton and René Descartes on chemists of the 18th century. Enlightenment chemists established distinctive scientific communities and a well-defined discipline (closely allied, to be sure, with medical and artisanal studies) in the major countries of Europe. The chemist’s workplace or laboratory (the word itself had been coined in the Renaissance to apply to the chemical arts) was now closely associated with the field, and a standardized repertoire of operations was taught there.
 
Still unsettled were some fundamental issues relating to chemical composition. To a phlogistonist, a metallic calx was elemental, and the associated metal was a compound of calx plus phlogiston. This puzzled some, though, since the metal gained rather than lost weight when it supposedly lost phlogiston to become a calx. The issues were sharpened in the 1770s, when the virtuoso English chemist (and Unitarian minister) Joseph Priestley produced a new gas by heating certain minerals. A candle burned in this gas with extraordinary vigour, and in an enclosed space a mouse breathing it survived far longer than one could in ordinary air. Priestley’s explanation was that the new gas had been radically dephlogisticated and, hence, had much greater capacity than air for absorbing phlogiston.
 
Actually, gases (then usually known as airs) were a relatively novel object of chemical attention. In Scotland in 1756, Joseph Black studied the gas given off in respiration and combustion, characterizing it chemically and following its participation in certain chemical reactions. (Black, a physician, taught chemistry as a branch of medicine, as did most academic chemists of this era.) He called the new gas “fixed air,” since it was also found “fixed” in certain minerals such as limestone. His discovery that this gas was a normal component of common air (at a fraction of a percent, to be sure) was the first clear indication that atmospheric air was a mixture rather than a homogeneous element. In the following quarter century, many new gases were discovered and studied, by such workers as Priestley, the English physicist and chemist Henry Cavendish, and the Swedish pharmacist Carl Scheele.
 
The chemical revolution
 
The new research on “airs” attracted the attention of the young French aristocrat Antoine-Laurent Lavoisier. Lavoisier commanded both the wealth and the scientific brilliance to enable him to construct elaborate apparatuses to carry out his numerous ingenious experiments. In the course of just a few years in the 1770s, Lavoisier developed a radical new system of chemistry, based on Black’s methods and Priestley’s dephlogisticated air.
 
Lavoisier first determined that certain metals and nonmetals absorb a gaseous substance from the air in undergoing calcination or combustion and, in the process, increase in weight. Initially, he thought that this gas must be Black’s fixed air, for he knew of no other chemical species present in ordinary air; moreover, fixed air was known to be produced in smelting, so it seemed reasonable to think that it was present in the calx that was smelted. At this point (October 1774), Priestley communicated to Lavoisier his discovery of dephlogisticated air. Further experiments led Lavoisier to continuously modify his ideas, until it finally became clear to him that it was this new gas, and not fixed air, that was the active entity in combustion, calcination, and respiration. Moreover, he determined (or so he thought, at least) that this gas was contained in all acids. He renamed it oxygen, Greek for “acid producer.”
 
Lavoisier’s oxygen was in some respects the inverse of phlogiston. Rather than releasing anything, the combustible or metal absorbed (more precisely, chemically combined with) oxygen in the process that Lavoisier now called oxidation. He showed that atmospheric air was a mixture of two principal components, oxygen and a physiologically inert gas (known to Priestley) that he called azote or nitrogen. He also showed that water is a chemical compound of two substances, oxygen and what Cavendish had called “inflammable air.” The latter gas was now renamed hydrogen (“water producer”). Black’s fixed air proved to be a gaseous form of oxidized carbon, or carbon dioxide. The various parts of Lavoisier’s new system were beginning to fit together beautifully.
 
The keys to Lavoisier’s success were twofold. First, he carefully accounted for all the substances, including gases, entering into and emerging from the chemical reactions he studied by tracking their weights with the greatest possible precision. He knew to do this partly from Black’s example, but he proceeded with a mastery that the science had never before seen. Second, he established a simple operational definition of a chemical element—namely, a substance that could not be reduced in weight as the result of any chemical reaction that it undergoes. Oxygen, carbon, iron, and sulfur were now regarded as elements, along with close to 30 other substances. Lavoisier wrote a textbook to promote the new oxygenist chemistry, Traité élémentaire de chimie (1789), which appeared in the same year the French Revolution began. He and his associates also developed a new nomenclature—essentially the one used today for inorganic compounds—along with a new journal. As an aristocrat of the ancien régime and an investor in a tax-collection agency, Lavoisier was executed in the Reign of Terror, but by that time (1794) the chemical revolution that he had started had largely succeeded in replacing phlogistonist chemistry.
 
Atomic and molecular theory
 
Lavoisier’s set of chemical elements, and the new way of understanding chemical composition, proved to be invaluable for analytic and inorganic chemistry, but in a real sense the chemical revolution had only just begun. Around the turn of the century, the English Quaker schoolteacher John Dalton began to wonder about the invisibly small ultimate particles of which each of these elemental substances might be composed. He thought that if the atoms of each of the elements were distinct, they must be characterized by a distinct weight that is unique to each element. Although these atoms were far too small to weigh individually, he realized that he could deduce their weights relative to each other—the ratio of the weight of an atom of oxygen to one of hydrogen, for instance—by examining reacting weights of macroscopic quantities of these elements. In fact, the laws of stoichiometry (combining weights of elements) were just then being developed, and Dalton used these regularities to justify his inferences. His first discussion of these issues dates to 1803, and he presented his atomic theory in the multivolume New System of Chemical Philosophy (1808–27).
 
Dalton’s atomic theory was a landmark event in the history of chemistry, but it had a crucial flaw. His procedure required that one know the formulas of the simple compounds resulting from the combination of the elements. For example, analytical data of that day indicated that water resulted from the combination of seven parts by weight of oxygen with one part of hydrogen. If the resulting water molecule was HO (one atom of each element combining to form a molecule of water), then the weight ratio of the atoms of these elements must be the same, seven to one. However, if the formula were H2O, then the weight of an oxygen atom would have to be 14 times the weight of a hydrogen atom. There was simply no way to determine molecular formulas at that time, so Dalton made assumptions based on the simplicity of nature. He chose HO as his water formula and, therefore, seven as the relative atomic weight of oxygen.
 
In the following years, several leading chemists adopted essential elements of Dalton’s theory, but many objected to the hypothetical elements just described; some also doubted the very possibility of investigating the world of the invisibly small. In 1808 the French chemist Joseph-Louis Gay-Lussac discovered that when gases combine chemically, they do so in small integral multiples by volume. Three years later the Italian physicist Amedeo Avogadro argued that this fact suggested that equal volumes of gases contain equal numbers of constituent particles (Avogadro’s law), physical conditions being the same. This idea provided a physical method of determining certain molecular formulas. For instance, Gay-Lussac had pointed out that exactly two volumes of hydrogen combine with precisely one of oxygen to form water. If Avogadro was right, the formula for water had to be H2O. But this line of reasoning also led to the uncomfortable notion that elementary gases had polyatomic molecules (O2, H2, and so on), and therefore many chemists rejected Avogadro’s hypotheses.
 
By far the greatest of the early atomists was the Swede Jöns Jacob Berzelius, who accepted parts of Avogadro’s ideas and developed an elaborate version of chemical atomism by 1826. It was Berzelius who in 1813 had proposed the alphabetic system for denoting elements, atoms, and molecular formulas, and the use of formulas as an aid for studying chemical composition and reactions began to blossom about 1830. However, different chemists were still making different assumptions regarding the formulas of simple compounds such as water, and so, for decades, various inconsistent systems of atomic weights and formulas were in use in the various European countries.
 
Berzelius also developed a theory of chemical combination based on the electrochemical studies that the invention of the battery (1800) had spawned. He became convinced that all molecules were held together by the Coulomb force, the electrostatic attraction between oppositely charged objects. (Berzelius assumed that a molecule’s constituent atoms or groups of atoms were not neutral, and he called these charged components radicals.) This theory of electrochemical dualism worked well with inorganic compounds, but organic substances seemed anomalous. Particularly in the 1830s, when chemists learned how to replace the hydrogen of organic compounds with chlorine atoms, Berzelius’s theory appeared to be threatened—after all, hydrogen and chlorine had opposite electrochemical characteristics, yet the substitution seemed to make little difference in the properties of the compounds. In the 1840s and ’50s, extensive debates over rival systems of chemical atomism and over electrochemical dualism enlivened the journal literature.
 
Organic radicals and the theory of chemical structure
 
Both problems were finally resolved through the further development of organic chemistry. The leading organic chemists of the day were the German Justus von Liebig and the Frenchman Jean-Baptiste-André Dumas. In 1830 Liebig invented a device that made organic analysis rapid, convenient, and accurate, and his laboratory institute at the tiny University of Giessen in Hesse became the most famous chemical school in the world. Liebig taught an enormous number of chemists, and his students assisted in his research program. He was the leading figure in the rise of the research university and in the idea of a research group. As a professor at Giessen, and later at the University of Munich, he laid much emphasis on practical applications of chemistry, especially for physiology, agriculture, and consumer products. Dumas exerted a similar influence in France, training students and pursuing research at a private laboratory in Paris.
 
Both Liebig and Dumas initially accepted the Berzelian scheme and sought to understand organic molecules as composed of identifiable radicals held together electrochemically. The younger French chemists Auguste Laurent and Charles Gerhardt pursued chlorine substitution reactions and cast doubt on this simple model; sometime after 1840 Liebig and Dumas both retreated into positivism. In 1852 Liebig’s English former postdoctoral assistant Edward Frankland noticed a regularity in the combining capacity of the atoms of certain metals and semimetals. At about the same time, two former students of both Liebig and Dumas, Alexander Williamson in London and Charles-Adolphe Wurtz in Paris, were independently approaching the same idea from a different direction. Using a system of atomic weights and formulas developed by Gerhardt and Laurent—a modified version of Berzelius’s system that incorporated Avogadro’s ideas more consistently—they proposed that oxygen atoms could combine with two other simple atoms, such as hydrogen, or with two organic radicals and that nitrogen atoms could combine with three. This was the beginning of the concept of atomic valence.
 
In 1858 the young German theorist August Kekule then expanded this concept to carbon, not only proposing that carbon atoms were tetravalent but adding the idea that they could bond to each other to form chains, comprising a molecular “skeleton” to which other atoms could cling. Kekule’s theory of chemical structure clarified the compositions of hundreds of organic compounds and served as a guide to the synthesis of thousands more. (The self-chaining of carbon atoms was independently developed by the Scottish chemist Archibald Scott Couper.) This theory experienced dramatic expansion when Kekule successfully applied it to aromatic compounds (after 1865) and after Jacobus Henricus van ’t Hoff of the Netherlands and Joseph LeBel of France independently began to investigate molecular structures in three dimensions—later called stereochemistry.
 
Mendeleyev’s periodic law
 
Kekule’s innovations were closely connected with a reform movement that gathered steam in the 1850s, seeking to replace the multiplicity of atomic weight systems with Gerhardt’s and Laurent’s proposal. Indeed, Kekule could not have succeeded with structure theory if he had not started with the reformed atomic weights. Kekule, Wurtz, and German chemist Carl Weltzien were organizers of the first international chemical conference, held at Karlsruhe in southwestern Germany in September 1860, which was intended to gain unity and understanding across the European chemical community. The Italian chemist Stanislao Cannizzaro played perhaps the most critical role at the conference. The reformers’ success was incomplete, but the Karlsruhe Congress can stand as an appropriate symbol of the era when chemistry attained a recognizably modern appearance.
 
The widespread adoption of a single reformed set of atomic weights for the 60-odd known elements appears to have prompted renewed speculation on the relationships of the elements to each other, and various proposals for systems of classification were developed in the 1860s. By far the most successful of these systems was that of the Russian chemist Dmitry Mendeleyev. In 1869 he announced that when the elements were arranged horizontally according to increasing atomic weight, and a new horizontal row was begun below the first whenever similar properties in the elements reappear, then the resulting semi-rectangular table revealed consistent periodicities. The vertical columns of similar elements were called groups or families, and the entire array was called the periodic table of the elements. Mendeleyev demonstrated that this manner of looking at the elements was more than mere chance when he was able to use his periodic law to predict the existence of three new elements, later named gallium, scandium, and germanium, which were discovered in the 1870s and ’80s.
 
To be sure, there were still many anomalies. For example, 15 chemically similar rare earth elements had been discovered by the end of the century. These elements were resistant to any periodic system; eventually they were grouped together in a separate category, the lanthanides (later called the lanthanoids; see transition element). Then in the 1890s British scientists William Ramsay and Lord Rayleigh discovered the inert, or rare, gases argon, helium, neon, krypton, and xenon. These were all clearly members of a single chemical family, but there were no vacant spaces in the table for them. Soon after the turn of the 20th century, chemists decided simply to create an extra group for them.
 
Structuralist ideas from organic chemistry, as well as the development of the periodic table, gave new impetus to the study of inorganic compounds in the late 19th century. The leading chemical field in the second half of the century, however, was clearly organic chemistry, and the leading country was Germany. It was the Germans who exploited the structure theory most aggressively, and their success was measured by the explosive growth of university institutes as well as by practical applications developed in commercial enterprises. Organic chemists such as August Wilhelm von Hofmann and Emil Fischer at the University of Berlin and Adolf von Baeyer at the University of Munich developed large research groups that turned out novel compounds, research publications, and doctoral dissertations by the score. By the late 19th century, German chemistry, both academic and industrial, dominated Europe and the world.
 
The rise of physical chemistry
 
This is not to say that other approaches to chemistry were neglected, nor that other countries failed to participate in the excitement. Physical studies of chemical compounds and reactions began early in the century, and the field of physical chemistry had achieved maturity by the 1880s. Michael Faraday in England, Hermann Kopp and Robert Bunsen in Germany, and Henri-Victor Regnault in France carried out investigations on the physical characteristics of substances in the period 1830–60. Studies of heat, work, and force led to the rise of thermodynamics around 1850; originally oriented almost entirely to the science of physics, figures such as the American Josiah Willard Gibbs, the Frenchmen Marcellin Berthelot and Pierre Duhem, and the Germans Hermann von Helmholtz and Wilhelm Ostwald then applied energy and entropy concepts to chemistry in the 1870s and ’80s. Electrochemistry, invented by the independent efforts of Berzelius and Humphry Davy in England at the beginning of the century, was pursued fruitfully by Faraday and others. Bunsen and Gustav Kirchhoff of Germany developed chemical spectroscopy in the late 1850s. Studies on the kinetics of chemical reactions began in the 1860s.
 
All this work culminated in the “official” establishment of the field of physical chemistry, traditionally considered to be when the Zeitschrift für Physikalische Chemie (“Journal of Physical Chemistry”) began publication in 1887. The editors were Ostwald and van ’t Hoff, with Svante Arrhenius of Sweden, a future Nobelist, an especially important member of its editorial board. Controversies over the reality of ionic dissociation and other issues connected with electrochemistry, the theory of solutions, and thermodynamics enlivened early issues of the journal.
 
Physical chemists were in increasing demand as universities turned to them for instruction in basic courses on general and theoretical chemistry. This was nowhere more true than in the United States, with its vigorously expanding educational structure, including both private and state (land-grant) universities and emerging German-influenced doctoral programs. Soon after the turn of the century, two chemists at the Massachusetts Institute of Technology (MIT) who had studied with Ostwald, Arthur Noyes and Gilbert Lewis, formed the nucleus of a rising American chemical community. Noyes continued his career at Throop Polytechnic in Pasadena (later renamed the California Institute of Technology, commonly known as Caltech), and Lewis went on to the University of California at Berkeley.
 
Physical chemistry was profoundly altered by what some have called the second scientific revolution—namely, the discoveries of the electron, X-rays, radioactivity, and new radioactive elements, the understanding of radioactive emissions and nuclear decay processes, and early versions of the theories of quantum mechanics and relativity. All of this happened in just 10 years, from 1895 to 1905, and the scientific bombshells continued in the following years. In 1911 the British physicist Ernest Rutherford proposed a nuclear model of the atom, but his orbiting electrons seemed to violate classical electromagnetic theory, and the model was not immediately embraced. However, two years later the Danish physicist Niels Bohr resolved some of these anomalies by applying spectroscopic data and the quantum theory of the German physicists Max Planck and Albert Einstein to Rutherford’s model (see figure). Bohr went on to head an international theoretical research group in Copenhagen that led in developing quantum mechanics during the 1920s. In the meantime, Rutherford revealed the existence of the proton and Einstein advanced his theory of general relativity.
 
Electronic theories of valence
 
So much for the physicists; but the chemists were not sitting on their hands through all of this. Since its discovery a half century earlier, one of the greatest puzzles in chemistry had been the central phenomenon of valence. It was as inexplicable as it was incontrovertibly true that oxygen atoms had exactly two valence “hooks” with which to form bonds and carbon normally had four (that is to say, oxygen is divalent, carbon tetravalent). Moreover, these bonds were not radially symmetrical like electrostatic charges or gravitation but seemed to be directed at distinct spatial angles around the atom. And the existence of highly stable elementary molecules such as H2 was downright embarrassing—for what could be the basis for the strong attraction of two identical atoms for each other? Some scientists, such as the great Swiss chemist Alfred Werner, used combinations of structural-organic and ionic theories to develop a scheme that brilliantly explained the structures of complex inorganic substances known as coordination compounds.
 
Others would take their cue from the discovery of the electron. As early as 1902, taking into account the work of the English physicist J.J. Thomson, Werner, and Ramsay and Rayleigh on the rare gases, Lewis privately drew casual sketches—depicting cubic atoms with outer electrons—that constituted the first step toward an electronic theory of chemical bonding. However, it was not until after Rutherford and Bohr had provided the early development of the nuclear theory of the atom that Lewis’s ideas gelled. (Simultaneously and independently, the German physicist Walther Kossel published a similar theory.) Lewis suggested that a chemical bond consisted of a pair of electrons that was shared between the combining atoms. By equal sharing of electrons (forming what the American physical chemist Irving Langmuir was soon to call a covalent bond), each atom could complete its outer electron shell and thus achieve stability. The normally complete outer shell, Lewis thought, contained eight electrons—the configuration of the notably stable (that is, inert) rare gases. This was the octet rule, and it helped to explain why Mendeleyev’s periodicities often came in multiples of eight.
 
The Lewis-Kossel-Langmuir electronic theory of valence (1916–23) was very incomplete, but was also extraordinarily fruitful for further developments, and essential elements of it survived for decades. In 1922 Bohr proposed electron configurations in the so-called K, L, M, and N shells. The theory was soon thereafter modified by breaking developments in quantum mechanics achieved by Bohr, German physicist Werner Heisenberg, Austrian physicist Erwin Schrödinger, and others. In 1927 two German researchers working with Schrödinger in Zürich, Fritz London and Walter Heitler, produced the first-ever quantum mechanical treatment of a chemical system, the hydrogen molecule.
 
The American physical chemist Linus Pauling (along with another American, John Slater) independently developed this approach into what he called the valence bond method of understanding chemical combination. The orbitals in the various electron shells (classified by the letters s, p, d, and f) could be mathematically “hybridized,” resulting in the directed bonds actually observed in chemical compounds. Pauling also made extensive use of the quantum mechanical resonance effect, especially for understanding aromatic compounds. All of this was summarized in his classic work The Nature of the Chemical Bond (1939). An alternative quantum mechanical method of understanding chemical bonding, called the molecular orbital method, was developed by the American chemist Robert Mulliken and the German physicist Friedrich Hund. Although mathematically more complex, this approach has largely replaced Pauling’s. In any case, ever since Lewis and Bohr, it has been understood that all chemical reactions and all chemical bonding involves the outer electron shells—the valence electrons—of participating atoms.
 
Organic chemists also incorporated electronic ideas into their theories. In the 1920s the Englishmen Robert Robinson and Christopher Ingold—bitter rivals then and later—led in the development of electronic theories of organic reaction mechanisms by focusing on rearranging electron pairs over the course of chemical reactions. Not only did this allow chemists to understand the intimate details of reactions in a way that had not previously been possible, but it also allowed them to successfully predict the reactivities of organic compounds in different chemical environments. Other studies of quantum mechanics applied to organic substances, combined with the kinetics of reactions, the nature of acids and bases, and instrumental methods of understanding compounds, led to a well-developed specialty field of physical organic chemistry.
 
Biochemistry, polymers, and technology
 
Organic chemistry, of course, looks not only in the direction of physics and physical chemistry but also, and even more essentially, in the direction of biology. Biochemistry began with studies of substances derived from plants and animals. By about 1800 many such substances were known, and chemistry had begun to assist physiology in understanding biological function. The nature of the principal chemical categories of foods—proteins, lipids, and carbohydrates—began to be studied in the first half of the century. By the end of the century, the role of enzymes as organic catalysts was clarified, and amino acids were perceived as constituents of proteins. The brilliant German chemist Emil Fischer determined the nature and structure of many carbohydrates and proteins. The announcement of the discovery (1912) of vitamins, independently by the Polish-born American biochemist Casimir Funk and the British biochemist Frederick Hopkins, precipitated a revolution in both biochemistry and human nutrition. Gradually, the details of intermediary metabolism—the way the body uses nutrient substances for energy, growth, and tissue repair—were unraveled. Perhaps the most representative example of this kind of work was the German-born British biochemist Hans Krebs’s establishment of the tricarboxylic acid cycle, or Krebs cycle, in the 1930s.
 
But the most dramatic discovery in the history of 20th-century biochemistry was surely the structure of DNA (deoxyribonucleic acid), revealed by American geneticist James Watson and British biophysicist Francis Crick in 1953—the famous double helix. The new understanding of the molecule that incorporates the genetic code provided an essential link between chemistry and biology, a bridge over which much traffic continues to flow. The individual “letters” that make the code—four nucleotides named adenine, guanine, cytosine, and thymine—were discovered a century ago, but only at the close of the 20th century could the sequence of these letters in the genes that make up DNA be determined en masse. In June 2000, representatives from the publicly funded U.S. Human Genome Project and from Celera Genomics, a private company in Rockville, Md., simultaneously announced the independent and nearly complete sequencing of the more than three billion nucleotides in the human genome. However, both groups emphasized that this monumental accomplishment was, in a broader perspective, only the end of a race to the starting line.
 
DNA is, of course, a macromolecule, and an understanding of this centrally important category of chemical compounds was a precondition for the events just described. Starch, cellulose, proteins, and rubber are other examples of natural macromolecules, or very large polymers. The word polymer (meaning “multiple parts”) was coined by Berzelius about 1830, but in the 19th century it was only applied to special cases such as ethylene (C2H4) versus butylene (C4H8). Only in the 1920s did the German chemist Hermann Staudinger definitely assert that complex carbohydrates and rubber had huge molecules. He coined the word macromolecule, viewing polymers as consisting of similar units joined head to tail by the hundreds and connected by ordinary chemical bonds.
 
Empirical work on polymers had long predated Staudinger’s contributions, though. Nitrocellulose was used in the production of smokeless gunpowder, and mixtures of nitrocellulose with other organic compounds led to the first commercial polymers: collodion, xylonite, and celluloid. The last of these was the earliest plastic. The first totally synthetic plastic was patented by Leo Baekeland in 1909 and named Bakelite. Many new plastics were introduced in the 1920s, ’30s, and ’40s, including polymerized versions of acrylic acid (a variety of carboxylic acid), vinyl chloride, styrene, ethylene, and many others. Wallace Carothers’s nylon excited extraordinary attention during the World War II years. Great effort was also devoted to develop artificial substitutes for rubber—a natural resource in especially short supply during wartime. Already by World War I, German chemists had substitute materials, though many were less than satisfactory. The first highly successful rubber substitutes were produced in the early 1930s and were of great importance in World War II.
 
During the interwar period, the leading role for chemistry shifted away from Germany. This was largely the result of the 1914–18 war, which alerted the Allied countries to the extent to which they had become dependent on the German chemical industries. Dyes, drugs, fertilizers, explosives, photochemicals, food chemicals (such as chemicals for food additives, food colouring, and food preservation), heavy chemicals, and strategic materiel of many kinds had been supplied internationally before the war largely by German chemical companies, and, when supplies of these vital materials were cut off in 1914, the Allies had to scramble to replace them. One particularly striking example is the introduction of chlorine gas and other poisons, starting in 1915, as chemical warfare agents. In any case, after the war ended, chemistry was enthusiastically promoted in Britain, France, and the United States, and the interwar years saw the United States rise to the status of a world power in science, including chemistry.
 
All this makes clear why World War I is sometimes referred to as “the chemists’ war,” in the same way that World War II can be called “the physicists’ war” because of radar and nuclear weapons. But chemistry was an essential partner to physics in the development of nuclear science and technology. Indeed, the synthesis of transuranium elements (atomic numbers greater than 92) was a direct consequence of the research leading to (and during) the Manhattan Project in World War II. This is all part of the legacy of the dean of nuclear chemists, American Glenn Seaborg, discoverer or codiscoverer of 10 of the transuranium elements. In 1997, element 106 was named seaborgium in his honour.
 
The instrumental revolution
 
As far as the daily practice of chemical research is concerned, probably the most dramatic change during the 20th century was the revolution in methods of analysis. In 1930 chemists still used “wet-chemical,” or test-tube, methods that had changed little in the previous hundred years: reagent tests, titrations, determination of boiling and melting points, elemental combustion analysis, synthetic and analytic structural arguments, and so on. Starting with commercial labs that provided an out-source for routine analyses and with pH meters that displaced chemical indicators, chemists increasingly began to rely on physical instrumentation and specialists rather than personally administered wet-chemical methods. Physical instrumentation provides the sharp “eyes” that can see to the atomic-molecular level.
 
In the 1910s J.J. Thomson and his assistant Francis Aston had developed the mass spectrograph to measure atomic and molecular weights with high accuracy. It was gradually improved, so that by the 1940s the mass spectrograph had been transformed into the mass spectrometer—no longer a machine for atomic weight research but rather an analytical instrument for the routine identification of complex unknown compounds (see mass spectrometry). Similarly, colorimetry had a long history, dating back well into the previous century. In the 1940s colorimetric principles were applied to sophisticated instrumentation to create a range of usable spectrophotometers, including visible, infrared, ultraviolet, and Raman spectroscopy. The later addition of laser and computer technology to analytical spectrometers provided further sophistication and also offered important tools for studies of the kinetics and mechanisms of reactions.
 
Chromatography, used for generations to separate mixtures and identify the presence of a target substance, was ever more impressively automated, and gas chromatography (GC) in particular experienced vigorous development. Nuclear magnetic resonance (NMR), which uses radio waves interacting with a magnetic field to reveal the chemical environments of hydrogen atoms in a compound, was also developed after World War II. Early NMR machines were available in the 1950s; by the 1960s they were workhorses of organic chemical analysis. Also by this time, GC-NMR combinations were introduced, providing chemists unexcelled ability to separate and analyze minute amounts of sample. In the 1980s NMR became well known to the general public, when the technique was applied to medicine—though the name of the application was altered to magnetic resonance imaging (MRI) to avoid the loaded word nuclear.
 
Many other instrumental methods have seen vigorous development, such as electron paramagnetic resonance and X-ray diffraction. In sum, between 1930 and 1970 the analytical revolution in chemistry utterly transformed the practice of the science and enormously accelerated its progress. Nor did the pace of innovation in analytical chemistry diminish during the final third of the century.
 
Organic chemistry in the 20th century
 
No specialty was more affected by these changes than organic chemistry. The case of the American chemist Robert B. Woodward may be taken as illustrative. Woodward was the finest master of classical organic chemistry, but he was also a leader in aggressively exploiting new instrumentation, especially infrared, ultraviolet, and NMR spectrometry. His stock in trade was “total synthesis,” the creation of a (usually natural) organic substance in the laboratory, beginning with the simplest possible starting materials. Among the compounds that he and his collaborators synthesized were alkaloids such as quinine and strychnine, antibiotics such as tetracycline, and the extremely complex molecule chlorophyll. Woodward’s highest accomplishment in this field actually came six years after his receipt of the Nobel Prize for Chemistry in 1965: the synthesis of vitamin B12, a notable landmark in complexity. Progress continued apace after Woodward’s death. By 1994 a group at Harvard University had succeeded in synthesizing an extraordinarily challenging natural product, called palytoxin, that had more than 60 stereocentres.
 
These total syntheses have had both practical and scientific spin-offs. Before the “instrumental revolution,” syntheses were often or even usually done to prove molecular structures. Today they are a central element of the search for new drugs. They can also illuminate theory. Together with a young Polish-born American chemical theoretician named Roald Hoffmann, Woodward followed up hints from the B12 synthesis that resulted in the formulation of orbital symmetry rules. These rules seemed to apply to all thermal or photochemical organic reactions that occur in a single step. The simplicity and accuracy of the predictions generated by the new rules, including highly specific stereochemical details of the product of the reaction, provided an invaluable tool for synthetic organic chemists.
 
Stereochemistry, born toward the end of the 19th century, received steadily increasing attention throughout the 20th century. The three-dimensional details of molecular structure proved to be not only critical to chemical (and biochemical) function but also extraordinarily difficult to analyze and synthesize. Several Nobel Prizes in the second half of the century—those awarded to Derek Barton of Britain, John Cornforth of Australia, Vladimir Prelog of the Soviet Union, and others—were given partially or entirely to honour stereochemical advances. Also important in this regard was the American Elias J. Corey, awarded the Nobel Prize for Chemistry in 1990, who developed what he called retrosynthetic analysis, assisted increasingly by special interactive computer software. This approach transformed synthetic organic chemistry. Another important innovation was combinatorial chemistry, in which scores of compounds are simultaneously prepared—all permutations on a basic type—and then screened for physiological activity.
 
Chemistry in the 21st century
 
Two more innovations of the late 20th century deserve at least brief mention, especially as they are special focuses of the chemical industry in the 21st century. The phenomenon of superconductivity (the ability to conduct electricity with no resistance) was discovered in 1911 at temperatures very close to absolute zero (0 K, −273.15 °C, or −459.67 °F). In 1986 two Swiss chemists discovered that lanthanum copper oxide doped with barium became superconducting at the “high” temperature of 35 K (−238 °C, or −397 °F). Since then, new superconducting materials have been discovered that operate well above the temperature of liquid nitrogen—77 K (−196 °C, or −321 °F). In addition to its purely scientific interest, much research focuses on practical applications of superconductivity.
 
In 1985 Richard Smalley and Robert Curl at Rice University in Houston, Tex., collaborating with Harold Kroto of the University of Sussex in Brighton, Eng., discovered a fundamental new form of carbon, possessing molecules consisting solely of 60 carbon atoms. They named it buckminsterfullerene (later nicknamed “buckyball”), after Buckminster Fuller, the inventor of the geodesic dome. Research on fullerenes has accelerated since 1990, when a method was announced for producing buckyballs in large quantities and practical applications appeared likely. In 1991 Science magazine named buckminsterfullerene their “molecule of the year.”
 
Two centuries ago, Lavoisier’s chemical revolution could still be questioned by the English émigré Joseph Priestley. A century ago, the physical reality of the atom was still doubted by some. Today, chemists can maneuver atoms one by one with a scanning tunneling microscope, and other techniques of what has become known as nanotechnology are in rapid development. The history of chemistry is an extraordinary story.
  • Recommend Us