Chemical Bond

Chemical Bond

History

The origin of bond symbolism

Development of the modern theory of bonding

Bond types

Electronegativity

Bond polarity

Coordination compounds

Multiple bonds

Other types of bonds

Resources

A chemical bond is any force of attraction that holds two atoms or ions together. In most cases, that force of attraction is between one or more electrons held by one of the atoms and the positively charged nucleus of the second atom. Chemical bonds vary widely in their stability, ranging from relatively strong covalent bonds to very weak hydrogen bonds.

History

The concept of bonding as a force that holds two particles together is as old as the concept of ultimate particles of matter itself. As early as 100 BC, for example, Asklepiades of Prusa speculated about the existence of ;“clusters of atoms,” a concept that implies the existence of some force of attraction holding the particles together. At about the same time, the Roman poet Lucretius in his monumental work De Rerum Natura(On the nature of things) pictured atoms as tiny spheres to which were attached fishhook-like appendages. Atoms combined with each other, according to Lucretius, when the appendages from two adjacent atoms became entangled with each other.

Relatively little progress occurred in the field of bonding theory until the concept of an atom itself was clarified. When John Dalton proposed modern atomic theory in 1803, he specifically hypothesized that atoms would combine with each other to form “compound atoms.” Dalton’s concept of bonding was essentially nonexistent, however, and he imagined that atoms simply sat adjacent to each other in their compound form.

The real impetus to further speculation about bonding was provided by the evolution of the concept of a molecule, originally proposed by Amedeo Avogadro in 1811 and later refined by Stanislao Cannizzaro more than four decades later.

The origin of bond symbolism

Some of the most vigorous speculation about chemical bonding took place in the young field of organic chemistry. In trying to understand the structure of organic compounds, for example, Friedrich Kekule´ (1829–1896) suggested that the carbon atom is tetravalent; that is, it can bond to four other atoms. He also hypothesized that carbon atoms could bond with each other almost endlessly in long chains.

Kekulé had no clear notion as to how atoms bonded to each other, but he did develop an elaborate system for showing how those bonds might be arranged in space. That system was too cumbersome for everyday use by chemists, however, and it was quickly replaced by another system suggested earlier by the Scottish chemist Archibald Scott Couper (1831–1892), who proposed that the bond between two atoms (what the real physical nature of that bond might be) should be represented by a short dashed line. Thus, a molecule of water could be represented by the structural formula: H-O-H.

That system is still in existence today. The arrangement of atoms in a molecule is represented by the symbols of the elements present joined by dashed lines that show how the atoms of those elements are bonded to each other. Thus, the term chemical bond refers not only to the force of attraction between two particles, but also to the dashed line used in the structural formula for that substance.

Development of the modern theory of bonding

The discovery of the electron by J. J. Thomson (1856–1940) in 1897 was, in the long run, the key needed to solve the problem of bonding. In the short run, however, it was a serious hindrance to resolving that issue. The question that troubled many chemists at first was how two particles with the same electrical charge (as atoms then seemed to be) could combine with each other.

An answer to that dilemma began to evolve slowly, beginning with the work of the young German chemist Richard Abegg (1869–1910). In the early 1900s, Abegg came to the conclusion that inert gases are stable elements because their outermost shell of electrons always contains eight electrons. Abegg theorized that atoms combine with each other when they exchange electrons in such a way that they all end up with eight electrons in their outer orbit. In a simplistic way, Abegg had laid out the principle of ionic bonding. Ionic bonds are formed when one atom completely gives up one or more electrons, and a second atom takes on those electrons.

Since Abegg was killed in 1910 at the age of 41 in a balloon accident, he was prevented from improving upon his original hypothesis. That work was taken up in the 1910s, however, by a number of other scientists, most prominently the German chemist Walther Kossel (1888–1956) and the American chemists Irving Langmuir (1881–1957) and Gilbert Newton Lewis (1875–1946).

Working independently, these researchers came up with a second method by which atoms might bond to each other. Rather than completely losing or gaining electrons, they hypothesized that perhaps atoms could share electrons with each other. One might imagine, for example, that in a molecule of methane (CH₄ ), each of the four valence electrons in carbon is shared with the single electron available from each of the four hydrogen atoms. Such an arrangement could provide carbon with a full outer shell of eight electrons and each hydrogen atom with a full outer shell of two. Chemical bonds in which two atoms share pairs of electrons with each other are known as covalent bonds.

In trying to illustrate this concept, Lewis developed another system for representing chemical bonds. In the Lewis system (also known as the electron-dot system), each atom is represented by its chemical symbol with the number of electrons in its outermost orbit, its bonding or valence electrons. The formula of a compound, then, is to be represented by showing how two or more atoms share electrons with each other.

Bond types

Credit for the development of the modern theory of chemical bonding belongs largely to the great American chemist Linus Pauling (1901–1994). Early in his career, Pauling learned about the revolution in physics that was taking place largely in Europe during the 1920s. That revolution had come about with the discovery of the relativity theory, quantum mechanics, the uncertainty principle, the duality of matter and energy, and other new and strikingly different concepts in physics.

Most physicists recognized the need to reformulate the fundamental principles of physics because of these discoveries. Relatively few chemists, however, saw the relevance of the revolution in physics to their own subject. Pauling was the major exception. By the late 1920s, he had already begun to ask how the new science of quantum mechanics could be used to understand the nature of the chemical bond.

In effect, the task Pauling undertook was to determine the way in which any two atoms might react with each other to put them in the lowest possible energy state. Among the many discoveries he made was that, for most cases, atoms form neither a purely ionic nor purely covalent bond. That is, atoms typically do not completely lose, gain, or share equally the electrons that form the bond between them. Instead, the atoms tend to form hybrid bonds in which a pair of shared electrons spend more time with one atom and less time with the second atom.

Electronegativity

The term that Pauling developed for this concept is electronegativity. This, in a general sense, is the tendency of an atom to attract the electrons in a covalent bond. The numerical values for the electronegativities of the elements range from a maximum of 4.0 for fluorine to a minimum of about 0.7 for cesium. A bond formed between fluorine and cesium would tend to be ionic because fluorine has a much stronger attraction for electrons than does cesium. On the other hand, a bond formed between cobalt (electronegativity = 1.9) and silicon (electronegativity = 1.9) would be a nearly pure covalent bond since both atoms have an equal attraction for electrons.

The modern concept of chemical bonding, then, is that bond types are not best distinguished as purely ionic or purely covalent. Instead, they lie somewhere along a continuum between those two extremes. The position of any particular bond can be predicted by calculating the difference between the two electronegativities of the atoms involved. The greater that difference, the more ionic the bond; the smaller the difference, the more covalent.

Bond polarity

The preceding discussion suggests that most chemical bonds are polar; that is, one end of the bond is more positive than the other end. In the bond formed between hydrogen (electronegativity = 2.2) and sulfur (electronegativity = 2.6), for example, neither atom has the ability to take electrons completely from the other. Neither is equal sharing of electrons likely to occur. Instead, the electrons forming the hydrogen-sulfur bond will spend somewhat more time with the sulfur atom and somewhat less time with the hydrogen atom. Thus, the sulfur end of the hydrogen-sulfur bond is somewhat more negative (represented as δ–), and the hydrogen end is somewhat more positive (δ+).

Coordination compounds

Some chemical bonds are unique in that both electrons forming the bond come from a single atom. The two atoms are held together, then, by the attraction between the pair of electrons from one atom and the positively charged nucleus of the second atom. Such bonds have been called coordinate covalent bonds.

An example of this kind of bonding is found in the reaction between copper(II) ion and ammonia. The nitrogen atom in ammonia has an unshared pair of electrons that is often used to bond with other atoms. The copper(II) ion is an example of such an anion. It is positively charged and tends to surround itself with four ammonia molecules to form the cupric ammonium ion, Cu(NH₃ )₄²⁺. The bonding in this ion consists of coordinate covalent bonds with all bonding electrons supplied by the nitrogen atom.

Multiple bonds

The bonds described thus far can all be classified as single bonds. That is, they all consist of a single pair of electrons. Not uncommonly, two atoms will combine with each other by sharing two pairs of electrons. For example, when lead and sulfur combine to form a compound, the molecules formed might consist of two pairs of electrons, one electron from lead and one electron from sulfur in each of the pairs. The standard shorthand for a double bond such as this one is a double dashed line (=). For example, the formula for a common double-bonded compound, ethylene, is: H₂ C=CH₂.

Compounds can also be formed by the sharing of three pairs of electrons between two atoms. The formula for one such compound, acetylene, shows how a triple bond of this kind is represented: HC≡CH.

Other types of bonds

Other types of chemical bonds also exist. The atoms that make up a metal, for example, are held together by a metallic bond—one in which all of the metal atoms share a cloud of electrons with each other. The electrons that make up that cloud originate from the outermost energy levels of the atoms.

A hydrogen bond is a weak force of attraction that exists between two atoms or ions with opposite charges. For example, the hydrogen-oxygen bonds in water are polar bonds. The hydrogen ends of these bonds are slightly positive and the oxygen ends, slightly negative. Two molecules of water placed next to each other will feel a force of attraction because the oxygen end of one molecule feels an electrical force of attraction to the hydrogen end of the other molecule. Hydrogen bonds are very common and extremely important in biological systems. They are strong

KEY TERMS

Coordinate covalent bond —A type of covalent bond in which all shared electrons are donated by only one of two atoms.

Covalent bond —A chemical bond formed when two atoms share a pair of electrons with each other.

Double bond —A covalent bond consisting of two pairs of shared electrons that hold the two atoms together.

Electronegativity —A quantitative method for indicating the relative tendency of an atom to attract the electrons that make up a covalent bond.

Ionic bond —A chemical bond formed when one atom gains and a second atom loses electrons.

Lewis symbol —A method for designating the structure of atoms and molecules in which the chemical symbol for an element is surrounded by dots indicating the number of valence electrons in the atom of that element.

Molecule —A collection of atoms held together by some force of attraction.

Multiple bond —A double or triple bond.

Polar bond —A covalent bond in which one end of the bond is more positive than the other end.

Structural formula —The chemical representation of a molecule that shows how the atoms are arranged within the molecule.

Triple bond —A triple bond is formed when three pairs of electrons are shared between two atoms.

Valence electrons —The electrons in the outermost shell of an atom that determine an element’s chemical properties.

enough to hold substances together, but weak enough to break apart and allow chemical changes to take place within the system.

Van der Waals forces are yet another type of chemical bond. Such forces exist between particles that appear to be electrically neutral. The rapid shifting of electrons that takes place within such molecules means that some parts of the molecule are momentarily charged, either positively or negatively. For this reason, very weak, transient forces of attraction can develop between particles that are actually neutral.

Resources

BOOKS

Bynum, W. F., E. J. Browne, and Roy Porter. Dictionary of the History of Science. Princeton, NJ: Princeton University Press, 1981, pp. 433-435.

Kotz, John C., and Paul Treichel. Chemistry and Chemical Reactivity. Pacific Grove, CA: Brooks/Cole, 1998.

Lide, D. R., editor. CRC Handbook of Chemistry and Physics Boca Raton: CRC Press, 2001.

Oxtoby, David W., et al. The Principles of Modern Chemistry. 5th ed. Pacific Grove, CA: Brooks/Cole, 2002.

Pauling, Linus. The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry. 3rd ed. Ithaca, NY: Cornell University Press, 1960.

OTHER

Epp, Erik. “Chapter 8: Chemical Bonding” Erik’s Chemistry. (accessed November 13, 2006) <http://eppe.tripod.com/chembond.html>.

Oregon State University. “Linus Pauling and the Nature of the Chemical Bond: A Documentary” (accessed November 13, 2006) <http://osulibrary.oregonstate.edu/specialcollections/coll/pauling/bond/narrative/page34.html>.