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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

chemical bond: see chemical bond.

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