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Chemical Bond

Chemical bond

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 negatively charged electrons held by one of the atoms and the positively charged nucleus of the second atom. Chemical bonds vary widely in their strength, ranging from relatively strong covalent bonds (in which electrons are shared between atoms) to very weak hydrogen bonds. The term chemical bond also refers to the symbolism used to represent the force of attraction between two atoms or ions. For example, in the chemical formula HOH, the short dashed lines are known as chemical bonds.

History

Theories of chemical bonds go back a long time. One of the first was developed by Roman poet Lucretius (c. 95c. 55 b.c.), author of De Rerum Natura (title means "on the nature of things"). In this poem, Lucretius described atoms as tiny spheres with fishhook-like arms. Atoms combined with each other, according to Lucretius, when the hooked arms of two atoms became entangled with each other.

Words to Know

Covalent bond: A chemical bond formed when two atoms share one or more pairs of electrons with each other.

Double bond: A covalent bond consisting of two pairs of electrons.

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

Hydrogen bond: A chemical bond formed between two atoms or ions with opposite charges.

Ionic bond: A chemical bond formed when one atom gains and a second atom loses electrons. An ion is a molecule or atom that has lost one or more electrons and is, therefore, electrically charged.

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.

Triple bond: A covalent bond consisting of three pairs of electrons.

Such theories were pure imagination, however, for many centuries, since scientists had no true understanding of an atom's structure until the beginning of the twentieth century. It was not until then that anything approaching a modern theory of chemical bonding developed.

Covalent bonding

Today, it is widely accepted that most examples of chemical bonding represent a kind of battle between two atoms for one or more electrons. Imagine an instance, for example, in which two hydrogen atoms are placed next to each other. Each atom has a positively charged nucleus and one electron spinning around its nucleus. If the atoms are close enough to each other, then the electrons of both atoms will be attracted by both nuclei. Which one wins this battle?

The answer may be obvious. Both atoms are exactly identical. Their nuclei will pull with equal strength on both electrons. The only possible result, overall, is that the two atoms will share the two electrons with each other equally. A chemical bond in which two electrons are shared between two atoms is known as a covalent bond.

Ionic bonding

Consider now a more difficult situation, one in which two different atoms compete for electrons. One example would be the case involving a sodium atom and a chlorine atom. If these two atoms come close enough to each other, both nuclei pull on all electrons of both atoms. In this case, however, a very different result occurs. The chlorine nucleus has a much larger charge than does the sodium nucleus. It can pull on sodium's electrons much more efficiently than the sodium nucleus can pull on the chlorine electrons. In this case, there is a winner in the battle: chlorine is able to pull one of sodium's electrons away. It adds that electron to its own collection of electrons. In a situation in which one atom is able to completely remove an electron from a second atom, the force of attraction between the two particles is known as an ionic bond.

Electronegativity

Most cases of chemical bonding are not nearly as clear-cut as the hydrogen and the sodium/chlorine examples given above. The reason for this is that most atoms are more nearly matched in their ability to pull electrons than are sodium and chlorine, although not as nearly matched as two identical atoms (such as two hydrogen atoms).

A method for expressing the pulling ability of two atoms was first suggested by American chemist Linus Pauling (19011994). Pauling proposed the name "electronegativity" for this property of atoms. Two atoms with the same or similar electronegativities will end up sharing electrons between them in a covalent bond. Two atoms with very different electronegativities will form ionic bonds.

Polar and nonpolar bonds

In fact, most chemical bonds do not fall into the pure covalent or pure ionic bond category. The major exception occurs when two atoms of the same kindsuch as two hydrogen atomscombine with each other. Since the two atoms have the same electronegativities, they must share electrons equally between them.

Consider the situation in which aluminum and nitrogen form a chemical bond. The electronegativity difference between these two atoms is about 1.5. (For comparison's sake, the electronegativity difference between sodium and chlorine is 2.1 and between hydrogen and hydrogen is 0.0.) A chemical bond formed between aluminum and nitrogen, then, is a covalent bond, but electrons are not shared equally between them. Instead, electrons that make up the bond spend more of their time with nitrogen (which pulls more strongly on electrons) than with aluminum (which pulls less strongly). A covalent bond in which electrons spend more time with one atom than with the other is called a polar covalent bond. In contrast, a bond in which electrons are shared equally (as in the case of hydrogen) is called a nonpolar covalent bond.

Multiple bonds

All covalent bonds, polar and nonpolar, always consist of two electrons. In some cases, both electrons come from one of the two atoms. In most cases, however, one electron comes from each of the two atoms joined by the bond.

In some cases, atoms may share more than two electrons. If so, however, they still share pairs only: two pairs or three pairs, for example. A bond consisting of two pairs of (that is, four) electrons is called a double bond. One containing three pairs of electrons is called a triple bond.

Other types of bonds

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

A hydrogen bond is a weak force of attraction that exists between two atoms or ions with opposite charges. For example, the hydrogen-oxygen bonds in water are polar bonds. The hydrogen end of these bonds are slightly positive, and the oxygen ends are 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 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. They are named in honor of the Dutch physicist Johannes Diderik van der Waals (18371923), who investigated the weak nonchemical bond forces between molecules. Such forces exist between particles that appear to be electrically neutral. The electrons in such particles shift back and forth very rapidly. That shifting of electrons means that some parts of the particle are momentarily charged, either positively or negatively. For this reason, very weak, short-term forces of attraction can develop between particles that are actually neutral.

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chemical bond

chemical bond, mechanism whereby atoms combine to form molecules. There is a chemical bond between two atoms or groups of atoms when the forces acting between them are strong enough to lead to the formation of an aggregate with sufficient stability to be regarded as an independent species. The number of bonds an atom forms corresponds to its valence. The amount of energy required to break a bond and produce neutral atoms is called the bond energy. All bonds arise from the attraction of unlike charges according to Coulomb's law; however, depending on the atoms involved, this force manifests itself in quite different ways. The principal types of chemical bond are the ionic, covalent, metallic, and hydrogen bonds. The ionic and covalent bonds are idealized cases, however; most bonds are of an intermediate type.

The Ionic Bond

The ionic bond results from the attraction of oppositely charged ions. The atoms of metallic elements, e.g., those of sodium, lose their outer electrons easily, while the atoms of nonmetals, e.g., those of chlorine, tend to gain electrons. The highly stable ions that result retain their individual structures as they approach one another to form a stable molecule or crystal. In an ionic crystal like sodium chloride, no discrete diatomic molecules exist; rather, the crystal is composed of independent Na+ and Cl- ions, each of which is attracted to neighboring ions of the opposite charge. Thus the entire crystal is a single giant molecule.

The Covalent Bond

A single covalent bond is created when two atoms share a pair of electrons. There is no net charge on either atom; the attractive force is produced by interaction of the electron pair with the nuclei of both atoms. If the atoms share more than two electrons, double and triple bonds are formed, because each shared pair produces its own bond. By sharing their electrons, both atoms are able to achieve a highly stable electron configuration corresponding to that of an inert gas. For example, in methane (CH4), carbon shares an electron pair with each hydrogen atom; the total number of electrons shared by carbon is eight, which corresponds to the number of electrons in the outer shell of neon; each hydrogen shares two electrons, which corresponds to the electron configuration of helium.

In most covalent bonds, each atom contributes one electron to the shared pair. In certain cases, however, both electrons come from the same atom. As a result, the bond has a partly ionic character and is called a coordinate link. Actually, the only purely covalent bond is that between two identical atoms.

Covalent bonds are of particular importance in organic chemistry because of the ability of the carbon atom to form four covalent bonds. These bonds are oriented in definite directions in space, giving rise to the complex geometry of organic molecules. If all four bonds are single, as in methane, the shape of the molecule is that of a tetrahedron. The importance of shared electron pairs was first realized by the American chemist G. N. Lewis (1916), who pointed out that very few stable molecules exist in which the total number of electrons is odd. His octet rule allows chemists to predict the most probable bond structure and charge distribution for molecules and ions. With the advent of quantum mechanics, it was realized that the electrons in a shared pair must have opposite spin, as required by the Pauli exclusion principle. The molecular orbital theory was developed to predict the exact distribution of the electron density in various molecular structures. The American chemist Linus Pauling introduced the concept of resonance to explain how stability is achieved when more than one reasonable molecular structure is possible: the actual molecule is a coherent mixture of the two structures.

Metallic and Hydrogen Bonds

Unlike the ionic and covalent bonds, which are found in a great variety of molecules, the metallic and hydrogen bonds are highly specialized. The metallic bond is responsible for the crystalline structure of pure metals. This bond cannot be ionic because all the atoms are identical, nor can it be covalent, in the ordinary sense, because there are too few valence electrons to be shared in pairs among neighboring atoms. Instead, the valence electrons are shared collectively by all the atoms in the crystal. The electrons behave like a free gas moving within the lattice of fixed, positive ionic cores. The extreme mobility of the electrons in a metal explains its high thermal and electrical conductivity.

Hydrogen bonding is a strong electrostatic attraction between two independent polar molecules, i.e., molecules in which the charges are unevenly distributed, usually containing nitrogen, oxygen, or fluorine. These elements have strong electron-attracting power, and the hydrogen atom serves as a bridge between them. The hydrogen bond, which plays an important role in molecular biology, is much weaker than the ionic or covalent bonds. It is responsible for the structure of ice.

Bibliography

See L. Pauling, The Nature of the Chemical Bond (3d ed. 1960); A. L. Companion, Chemical Bonding (2d ed. 1979).

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chemical bond

chemical bond A strong force of attraction holding atoms together in a molecule or crystal. In general, atoms combine to form molecules by sharing or transferring electrons in their outer shells. Typically chemical bonds have energies of about 1000 kJ mol–1 and are distinguished from the much weaker forces between molecules. See also covalent bond; electrovalent bond; hydrogen bond.

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chemical bond

chemical bond Mechanism that holds together atoms to form molecules. There are several types which arise either from the attraction of unlike charges, or from the formation of stable configurations through electron-sharing. The number of bonds an atom can form is governed by valence. The main types are ionic, covalent, metallic, and hydrogen bonds.

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bond, chemical

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Chemical Bond

Chemical bond


A chemical bond is any force of attraction between two atoms strong enough to hold the atoms together for some period of time. At least five primary types of chemical bonds are known, ranging from very strong to very weak. They are covalent, ionic, metallic, and hydrogen bonds, and London forces.

In all cases, a chemical bond ultimately involves forces of attraction between the positively-charged nucleus of one atom and the negatively-charged electron of a second atom. Understanding the nature of chemical bonds has practical significance since the type of bonding found in a substance explains to a large extent the macroscopic properties of that substance.

An ionic bond is one in which one atom completely loses one or more electrons to a second atom. The first atom becomes a positively charged ion and the second, a negatively charged ion. The two ions are attracted to each other because of their opposite electrical charges.

In a covalent bond, two atoms share one or more pairs of electrons. For example, a hydrogen atom and a fluorine atom each donate a single electron to form a shared pair that constitutes a covalent bond between the two atoms. Both electrons in the shared pair orbit the nuclei of both atoms.

In most cases, covalent and ionic bonding occur in such a way as to satisfy the Law of Octaves. Essentially that law states that the most stable configuration for an atom is one in which the outer energy level of the atom contains eight electrons or, in the case of smaller atoms, two electrons.

Ionic and covalent bonds might appear to represent two distinct limits of electron exchange between atoms, one in which electrons are totally gained and lost (ionic bonding) and one in which electrons are shared (covalent bonding). In fact, most chemical bonds fall somewhere between these two extreme cases. In the hydrogen-fluorine example mentioned above, the fluorine nucleus is much larger than the hydrogen nucleus and, therefore, exerts a greater pull on the shared electron pair. The electrons spend more time in the vicinity of the fluorine nucleus and less time in the vicinity of the hydrogen nucleus. For this reason, the fluorine end of the bond is more negative than the hydrogen end, and the bond is said to be a polar covalent bond. A non-polar covalent bond is possible only between two atoms with equal attraction for electrons as, for example, between two atoms of the same element.

Metallic bonds are very different from ionic and covalent bonds in that they involve large numbers of atoms. The outer electrons of these atoms feel very little attraction to any one nucleus and are able, therefore, to move freely throughout the metal.

Hydrogen bonds are very weak forces of attraction between atoms with partial positive and negative charges. Hydrogen bonds are especially important in living organisms since they can be broken and reformed easily during biochemical changes.

London forces are the weakest of chemical bonds. They are forces of attraction between two uncharged molecules. The force appears to arise from the temporary shift of electrical charges within each molecule.

[David E. Newton ]


RESOURCES

BOOKS

Giddings, J. Calvin. Chemistry, Man, and Environmental Change: An Integrated Approach. San Francisco: Canfield Press, 1973.

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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. Daltons 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´ (18291896) 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 (18311892), 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 (18561940) 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 (18691910). 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 (18881956) and the American chemists Irving Langmuir (18811957) and Gilbert Newton Lewis (18751946).

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 (19011994). 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: HCCH.

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 bondone 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 elements 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 Eriks 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>.

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Chemical Bond

Chemical bond

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 b.c., 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 could occur in the field of bonding theory, of course, until the concept of an atom itself was clarified. When John Dalton proposed the modern atomic theory in 1803, he specifically hypothesized that atoms would combine with each other to form "compound atoms." Dalton's concept of bonding was essentially non-existent, however, and he imagined that atoms simply sit 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 Kekulé 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 very clear notion as to how atoms bond 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. Couper proposed that the bond between two atoms (what the real physical nature of that bond might be) be represented by a short dashed line. Thus, a molecule of 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 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 slowly began to evolve, beginning with the work of the young German chemist Richard Abegg. In the early 1900s, Abegg came to the conclusion that inert gases are stable elements because their outermost shell of electrons always contain eight electrons. Perhaps atoms combine with each other, Abegg said, when they exchange electrons in such as way that they all end up with eight electrons in their outer orbit . In a simplistic way, Abegg had laid out the principle of ionic bonding. Ionic bonds are formed when one atom completely gives up one or more electrons, and a second atom takes on those electrons.

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

Working independently, these researchers came up with a second method by which atoms might bond to each other. Rather than completely losing or gaining electrons, they hypothesized, perhaps atoms can share electrons with each other. One might imagine, for 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. 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 for 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 in such a way as to put them in the lowest possible energy state. Among the many discoveries he made was that, for most cases, atoms form neither a purely ionic nor purely covalent bond. That is, atoms typically do not completely lose, gain, or share equally the electrons that form the bond between them. Instead, the atoms tend to form hybrid bonds in which a pair of shared electrons spend more time with one atom and less time with the second atom.

Electronegativity

The term that Pauling developed for this concept is electronegativity. Electronegativity is, in a general sense, the tendency of an atom to attract the electrons in a covalent bond. The numerical values for the 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 can be envisioned as lying 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, 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)4 2+ . 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: H2C=CH2.

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


Other types of bonds

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

A hydrogen bond is a weak force of attraction that exists between two atoms or ions with opposite charges. For example, the hydrogen-oxygen bonds in water are polar bonds. The hydrogen end of these bonds are 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 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 Cehmical Reactivity. Pacific Grove, CA: Brooks/Cole, 1998.

Lide, D.R., ed. 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 edition. Ithaca, NY: Cornell University Press, 1960.

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.

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"Chemical Bond." The Gale Encyclopedia of Science. . Encyclopedia.com. 10 Sep. 2018 <http://www.encyclopedia.com>.

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Because each style has its own formatting nuances that evolve over time and not all information is available for every reference entry or article, Encyclopedia.com cannot guarantee each citation it generates. Therefore, it’s best to use Encyclopedia.com citations as a starting point before checking the style against your school or publication’s requirements and the most-recent information available at these sites:

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Notes:
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