Chemistry: Chemical Reactions and the Conservation of Mass and Energy
Chemistry: Chemical Reactions and the Conservation of Mass and Energy
Chemical reactions involve molecules, the smallest units of matter that retain a substance's unique properties. Molecules, in turn, are composed of atoms, the smallest unit in which an element can exist. Molecules that represent a chemical combination of different atoms are called compounds. Molecules can also be composed of only one kind of atom; oxygen molecules, for example, have two oxygen atoms bonded together.
When two different molecules combine to make a new product, the reaction is called a synthesis reaction. In a decomposition reaction one molecule splits into smaller molecules. More commonly, chemical reactions involve two or more molecules swapping atoms in what are called replacement reactions. In every reaction the molecular reactant's atoms (or mass) are conserved. This means that the atoms that go into the reaction are the same as those that are found after the reaction in the molecules of the products.
All molecules have some degree of heat energy. Exothermic reactions increase heat energy in their environments; endothermic reactions absorb heat energy from their environment. Heat is a form of kinetic energy; “kinetic” comes from the Greek kinetikos, which means “in motion,” because both the atoms in the molecules and the molecules themselves are moving. The degree of movement determines the temperature of the molecule; the collective total of the movements determines the heat content of the system.
In addition to having heat energy, molecules have stored or potential energy in the chemical bonds that hold them together. The potential energy of chemical bonds depends on the type of bond between the atoms of elements and on the unique characteristics of each element.
The total energy of the system, the kinetic energy of heat added to the potential energy stored in the chemical bonds at the beginning of a reaction, is the same as the total energy of the system at the end of the reaction. Like mass, the total energy involved in a chemical reaction is also conserved.
Historical Background and Scientific Foundations
Chemistry evolved from alchemy, a practice and philosophy that proposed, among other ideas, that the so-called base elements could be turned into gold. Alchemy combined the philosophical ideas of ancient Greeks with Arabic chemical arts. Greek philosophers introduced the term “element,” but believed there were only four: earth, air, fire, and water. They considered atoms the essential particles of matter, but thought they were indivisible and didn't understand that they were the smallest units of chemical elements. Only in the seventeenth century did the modern concepts of elements, atoms, and molecules begin to emerge.
The breakthrough that led to modern chemistry came from the study of gases by British natural philosopher Robert Boyle (1627–1691). Until this point alchemists were concerned only with solids and liquids. Boyle's studies and writings inspired many others to investigate gases, including the one we now call oxygen, which was isolated in 1774 by British scientist Joseph Priestley (1733–1804).
Oxygen's discovery raised questions about what exactly it was, questions that were answered by French chemist Antoine Laurent Lavoisier (1743–1794). He established the modern concept of a chemical element and identified Priestley's discovery as an element, which he named oxygen. Lavoisier published his ideas in 1789 in a treatise called Traité élémentaire de chimie (Elementary treatise on chemistry). Following Lavoisier, English meteorologist John Dalton (1766–1844) promoted a theory of atoms, chemical elements, molecules, and compounds.
Advances in chemistry happened quickly after Lavoisier's work was published. Swedish chemist Jöns Berzelius (1779–1848) set out to make systematic and very precise measurements of chemicals. He published a table of atomic weights in 1826 that still agrees with many of the values accepted today.
Electricity was discovered at about this time, and many chemists explored its effects. When English chemist Humphry Davy (1778–1829) passed a current through molten caustic soda (sodium hydroxide), a chemical favored for its reactive properties since ancient times, the result was pure elemental sodium. Humphry used the same method to isolate potassium.
The idea of ionic compounds (those that form ions [electrically charged particles] in solutions) emerged from Swedish chemist Svante Arrhenius's (1859–1927) studies of reactions in solutions that conduct electricity. He won a 1903 Nobel Prize for his contributions.
Like Berzelius, other nineteenth-century scientists saw vast differences between compounds that came from living beings and those that came from nonliving sources, although this distinction was challenged in 1828 when Friedrich Wöhler (1800–1882) was able to synthesize urea, a compound previously found only in the urine of most animals. This spurred the development of organic chemistry, the science of carbon compounds—whatever their source—as a separate discipline. Organic compounds, scientists learned, are based on the unique shape of the carbon atom, making their molecular bonds different from those in ionic compounds.
The questions of bonding were not easily answered until atomic structure was better understood. In 1896 French physicist Henri Becquerel (1852–1908) discovered radioactivity. The following year British physicist J.J. Thomson (1856–1940) discovered the electron, a negatively charged subatomic particle common to all atoms that is emitted during radioactive decay. Becquerel won the Nobel Prize in 1903 for his discovery; Thomson won the prize in 1906.
The radioactivity of certain atoms led to further speculation on the structure of atoms. Thomson's model, called the plum pudding model, suggested that electrons are scattered randomly through a positive mass, giving a neutral atom. That model was discarded in 1911 when New Zealand-born British physicist Ernest Rutherford (1871–1937) discovered atoms have a dense positive nucleus.
The final model of an atom was described by Danish physicist Niels Bohr (1885–1962) in 1913, work for which he won the Nobel Prize for physics in 1922. Bohr discovered that an atom's electrons surround the dense positive nucleus but can move only in very specific regions called shells or orbits. His basic ideas have been refined to a model that works well to explain chemical reactions.
Types of Chemical Reactions
Every chemical reaction involves the electrons in the outer orbits of atoms. In reactions involving ionic bonds, electrons are given or taken to make ions of opposite charge that attract to form a chemical bond. Ionic reactions occur between metals and nonmetals.
In reactions that involve carbon and nonmetal atoms, bonding is called “covalent,” because electrons in the outer orbits (valence electrons) are shared in pairs in which the electrons are spinning oppositely, giving them north-south magnetic properties that make them attract. The bonding paired electrons are not always equally shared between atoms.
Collections of covalently bonded atoms can form molecules or ions. For example, the sulfate ion is made of four oxygen atoms clustered around a sulfur atom. It
has a -2 charge because it has two extra electrons. Such ions can form ionic bonds with metals—one example is copper sulfate.
For a reaction to occur the atoms must collide in such a way that the valence electrons can be rearranged. Reactions usually happen in solutions in which molecules are moving randomly, making favorable collisions more probable. In addition, they must collide with enough force to break their existing bonds and collide at just the right angle to form new bonds. This takes energy, which is released when the product bonds are formed. Exothermic reactions release more energy in the formation of products than they use to break the bonds of reactants. If the reverse is true, the reaction is endothermic.
Many chemical reactions are reversible. At first the only reaction is the forward reaction but, as products accumulate in the system, the reverse reaction becomes more likely. In time, all observable properties, such as the color of the reaction mixture, appear to be constant. When the forward reaction is happening as fast as the reverse, a condition called equilibrium has been reached.
Equilibrium reactions were studied by French chemist Henry-Louis Le Chatelier (1850–1936). His observation, known still as Le Chatelier's principle, was that if an equilibrium reaction is disturbed by an external occurrence, either the forward or reverse reaction will increase until equilibrium is restored.
The speed at which a reaction proceeds depends on a number of factors, including what types of molecules are involved, their concentration, and the temperature of the reactants. Even with optimum concentration and temperature, the probability of collisions between atoms in molecules may not produce products directly. Very often some intermediate product forms first, which will then either form the products or go back to reform reactants. American chemist Henry Eyring (1901–1981) suggested the concept of intermediate products, which he called activated complexes. The energy needed to form the activated complex is called the activation energy.
In many reactions catalysts can aid in the formation of intermediate molecules that then form products faster. Catalysts generally work by lowering the activation energy required to make an intermediate product; they are not part of the final product. A catalyst speeds up a reaction but never produces a greater concentration of the products. However, speeding up the reaction can be very advantageous in some chemical reactions.
There are catalysts for a great many chemical reactions used in industrial processes. Catalysts, such as enzymes—proteins that are key to every living organism—are also continuously at work. Three American biochemists, Christian Anfinsen (1916–1992), Stanford Moore (1913–1982), and William Stein (1911–1980), shared the 1972 Nobel Prize for their contributions to the understanding of catalytic activity as it relates to RNA, an essential part of all living cells.
The observation that chemical reactions reach equilibrium and the fact that temperature and pressure can influence the rate of a reaction had a major impact on one famous chemical reaction: the Haber-Bosch process for synthesizing nitrogen and hydrogen into the soluble nitrogen product ammonia.
By manipulating Le Chatelier's principle, the German physical chemist Fritz Haber (1868–1934) found a catalyst that created favorable conditions for producing ammonia. Haber's process was scaled up to industrial production levels in 1908 by German chemist Carl Bosch (1874–1940), who worked for the German chemical company BASF.
The Haber-Bosch process is credited with prolonging World War I by increasing the German's ability to produce munitions and grow more food after the British navy tried to block Germany from receiving critical nitrogen compounds from Chile. It is also considered to have affected food production worldwide, thereby contributing to the rapid increase in world population since World War II. Haber won the Nobel Prize in 1918; Bosch was honored in 1931.
Late-twentieth century advances in chemistry have considerably improved the quality of modern life. Unfortunately chemistry has also contributed to such problems as pesticide poisonings, landfills full of plastics that do not biodegrade, and air pollution.
Modern Cultural Connections
Chemical reactions happen continuously in every living organism. They are also at the heart of all chemical industries. The products of chemical reactions are everywhere.
One of the more noteworthy chemical reactions is the synthesis of ethanol as a renewable alternative fuel to supplement the use of petroleum-based fuels like gasoline. Even more noteworthy is the development of enzymes that make possible the conversion of cellulose (such as waste plant and wood fibers) into ethanol. Presently ethanol is produced from corn and sugar cane. Since both are also food crops, the production of ethanol for fuel is competing with world food supply demands.
Yet another advance in the movement toward green chemistry is the use of environmentally friendly materials and processes to manufacture everything from pharmaceuticals to plastics. The production of biodegradable plastics may help relieve the waste disposal problem that developed from the widespread use of many types of plastics. Research and development of chemical synthesis processes has also produced biocompatible plastics that can be used in life-saving medical devices. Other plastics are being developed and used in all parts of vehicles to reduce their weight and so save fuel.
See Also Chemistry: Biochemistry: The Chemistry of Life; Chemistry: Chemical Bonds; Chemistry: Chemical Reactions and the Conservation of Mass and Energy; Chemistry: Molecular Structure and Stereochemistry; Chemistry: Organic Chemistry; Chemistry: States of Matter: Solids, Liquids, Gases, and Plasma.
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Moore, Walter J. Physical Chemistry. Englewood Cliffs, New Jersey: Prentice Hall, 1979.
Morris, Richard. The Last Sorcerers: The Path from Alchemy to the Periodic Table. Washington, DC, Joseph Henry Press, 2003.
Weeks, Mary Elvira. Discovery of the Elements. White-fish, Montana: Kessinger Publishing Company, 2003.
Miriam C. Nagel