Chemical bonding

CONCEPT

Almost everything a person sees or touches in daily life—the air we breathe, the food we eat, the clothes we wear, and so on—is the result of a chemical bond, or, more accurately, many chemical bonds. Though a knowledge of atoms and elements is essential to comprehend the subjects chemistry addresses, the world is generally not composed of isolated atoms; rather, atoms bond to one another to form molecules and hence chemical compounds. Not all chemical bonds are created equal: some are weak, and some very strong, a difference that depends primarily on the interactions of electrons between atoms.

HOW IT WORKS

Early Ideas of Bonding

The theory that all of matter is composed of atoms did not originate in modern times: the atomic model actually dates back to the fifth century b.c. in Greece. The leading exponent of atomic theory in ancient times was Democritus (c. 460-370 b.c.), who proposed that matter could not be infinitely subdivided. At its deepest substructure, Democritus maintained, the material world was made up of tiny fragments he called atomos, a Greek term meaning “no cut” or “indivisible”
Forward-thinking though it was, Democritus’s idea was not what modern scientists today would describe as a proper scientific hypothesis. His “atoms” were not purely physical units, but rather idealized philosophical constructs, and thus, he was not really approaching the subject from the perspective of a scientist. In any case,there was no way for Democritus to test his hypothesis even if he had wanted to: by their very nature, the atoms he described were far too small to observe. Even today, what scientists know about atomic behavior comes not from direct observation, but indirect means.
Hence, Democritus and the few other ancients who subscribed to atomic theory went more on instinct than by scientific methods. Yet, some of them were remarkably prescient in their description of the bonding of atoms, in view of the primitive scientific methods they had at their disposal. No other scientist came close to the accuracy of their theory for about 2,000 years.

ASCLEPIADES AND LUCRETIUS DISCUSS BONDS

The physician Asclepiades of Prusa (c.130-40 b.c.) drew on the ideas of the Greek philosopher Epicurus (341270 b.c., another proponent of atomism. Asclepiades speculated on the ways in which atoms interact, and discussed “clusters of atoms,” though, of course, he had no idea what force attracted the atoms to one another.
A few years after Asclepiades, the Roman philosopher and poet Lucretius (c.95-c.55 b.c. espoused views that combined atomism with the idea of the “four elements”—earth, air, fire, and water. In his great work De rerum natura (“On the Nature of Things”), Lucretius described atoms as tiny spheres attached to each other by fishhook-like appendages that became entangled with one another.

LACK OF PROGRESS UNTIL 1800

Unfortunately, the competing idea of the four elements, handed down by the great philosopher Aristotle (384-322 b.c.), prevailed over the atomic model. As the Roman Empire
Electron-dot diagrams of some elements in the periodic table.
Electron-dot diagrams of some elements in the periodic table.
began to decline after a.d. 200, the pace of scientific inquiry slowed and—in Western Europe at least—eventually came to a virtual halt. Hence, the four elements theory, which had its own fanciful explanations as to why certain “elements” bonded with one another, held sway in Europe until the beginning of the modern era.
During the seventeenth century, a mounting array of facts from the realms of astronomy and physics collectively disproved the Aristotelian model. In the area of chemistry, English physicist and chemist Robert Boyle (1627-1691) showed that the four elements were not elements at all, because they could be broken down into simpler substances. Yet, no one really understood what constituted an element until the very beginning of the nineteenth century, and until that question was addressed, it was difficult to move on to the mystery of why certain atoms bonded with one another.

Early Modern Advances in Bonding Theory

Dalton’s atomic theory

The birth of atomic theory in modern times occurred in 1803, when English chemist John Dalton (1766-1844) formulated the idea that all elements are composed of tiny, indestructible particles. These he called by the name Democritus had given them nearly 23 centuries earlier: atoms. All known substances, he said, are composed of some combination of atoms, which differ from one another only in mass.
Though Dalton’s theory paved the way for enormous advances in the years that followed, there were a number of flaws in it. Mass alone, for instance, is not really what differentiates one atom from another: differences in mass reflect the presence of subatomic particles—protons and neutrons—of whose existence scientists were unaware at the time.
Furthermore, the properties of atoms that cause them to bond relate to a third subatomic particle, the electron, which, though it contributes little to the mass of the atom, is all-important to the energy it possesses. As for how atoms bond to one another, Dalton had little to say: in his conception of the atomic model, atoms simply sit adjacent to one another without forming true bonds, as such.

Avogadro and the molecule

Though Dalton recognized that the structure of atoms in a particular element or compound is uniform, he maintained that compounds are made up of compound atoms: thus,
Hydrogen bonding in HF, H2O, and NH3.
Hydrogen bonding in HF, H2O, and NH3.
water is a compound of “water atoms” However, water is not an element, and therefore, there had to be some structure—still very small, but larger than the atom—in which atoms coalesced to form the basic materials of a compound.
That structure was the molecule, first described by Italian physicist Amedeo Avogadro (1776-1856). For several decades, Avogadro, who originated the idea of the mole as a means of comparing large groups of atoms or molecules, remained a more or less unsung hero. Only in 1860, four years after his death, was his idea of
the molecule resurrected by Italian chemist Stanislao Cannizzaro (1826-1910). Cannizzaro’s work was occasioned by disagreement among scientists regarding the determination of atomic mass; however, the establishment of the molecular model had far-reaching implications for theories of bonding.

Symbolizing atomic bonds

In 1858, German chemist Friedrich August Kekule (1829-1896) made the first attempt to define the concept of valency, or the property an atom of one element possesses that determines
"Pure" water from a mountain stream is actually filled with traces of the rocks over which it has flowed. in fact, water is almost impossible to find in pure form except by purifying it in a laboratory.
“Pure” water from a mountain stream is actually filled with traces of the rocks over which it has flowed. in fact, water is almost impossible to find in pure form except by purifying it in a laboratory.
its ability to bond with atoms of other elements. A pioneer in organic chemistry, which deals with chemical structures containing carbon, Kekule described the carbon atom as tetravalent, meaning that it can bond to four other atoms. (The Latin prefix tetra- means “four.”) He also speculated that carbon atoms are capable of bonding with one another in long chains.
This was one of the first attempts to examine the subject of bonding using modern scientific terminology, complete with hypotheses that could be tested by experimentation. Kekule also recognized that in order to discuss bonds understandably, there needed to be some means of representing those bonds with symbols. He even went so far as to develop a system for showing the arrangement of bonds in space; however, his system was so elaborate that it was replaced in favor of a simpler one developed by Scottish chemist Archibald Scott Couper (1831-1892).
Couper, who also studied valency and the tetravalent carbon bond—he is usually given equal credit with Kekule for these ideas—created an extremely straightforward schematic representation still in use by chemists today. In Couper’s system, short dashed lines serve to designate chemical bonds. Hence, the bond between two hydrogen atoms and an oxygen atom in a water molecule would be represented thus: H-O-H. As the understanding of bonds progressed in modern times, this system was modified to take into account multiple bonds, discussed below.

REAL-LIFE APPLICATIONS

Atoms, Electrons, and Ions

Today, chemical bonding is understood as the joining of atoms through electromagnetic force. Before that understanding could be achieved, however, scientists had to unlock the secret of the electromagnetic interactions that take place within an atom.
The key to bonding is the electron, discovered in 1897 by English physicist J. J. Thomson (1856-1940). Atomic structure in general, and the properties of the electron in particular, are discussed at length elsewhere in this volume. However, because these specifics are critical to bonding, they will be presented here in the shortest possible form.
At the center of an atom is a nucleus, consisting of protons, with a positive electrical charge; and neutrons, which have no charge. These form the bulk of the atom’s mass, but they have little to do with bonding. In fact, the neutron has nothing to do with it, while the proton plays only a passive role, rather like a flower being pollinated by a bee. The “bee” is the electron,and, like a bee, it buzzes to and fro, carrying a powerful “sting”—its negative electric charge, which attracts it to the positively charged proton.

Electrons and ions

Though the electron weighs much, much less than a proton, it possesses enough electric charge to counterbalance the positive charge of the proton. All atoms have the same number of protons as electrons, and hence the net electric charge is zero. However, as befits their highly active role, electrons are capable of moving from one atom to another under the proper circumstances. An atom that loses or acquires electrons has an electric charge, and is called an ion.
The atom that has lost an electron or electrons becomes a positively charged ion, or cation. On the other hand, an atom that gains an electron or electrons becomes a negatively charged ion, or anion. As we shall see, ionic bonds, such as those that join sodium and chlorine atoms to form NaCl, or salt, are extremely powerful.

Electron configuration

Even in covalent bonding, which does not involve ions, the configurations of electrons in two atoms are highly important. The basics of electron configuration are explained in the Electrons essay, though even there, this information is presented with the statement that the student should consult a chemistry text topic for a more exhaustive explanation.
In the simplest possible terms, electron configuration refers to the distribution of electrons at various positions in an atom. However, because the behavior of electrons cannot be fully predicted, this distribution can only be expressed in terms of probability. An electron moving around the nucleus of an atom can be compared to a fly buzzing around some form of attractant (e.g., food or a female fly, if the moving fly is male) at the center of a sealed room. We can state positively that the fly is in the room, and we can predict that he will be most attracted to the center, but we can never predict his location at any given moment.
As one moves along the periodic table of elements, electron configurations become ever more complex. The reason is that with an increase in atomic number, there is an increase in the energy levels of atoms. This indicates a greater range of energies that electrons can occupy, as well as a greater range of motion. Electrons occupying the highest energy level in an atom are called valence electrons, and these are the only ones involved in chemical bonding. By contrast, the core electrons, or the ones closest to the nucleus, play no role in the bonding of atoms.

Ionic and Covalent Bonds

the goal of eight valence electrons. The above discussion of the atom, and the electron’s place in it, refers to much that was unknown at the time Thomson discovered the electron. Protons were not discovered for several more years, and neutrons several decades after that. Nonetheless, the electron proved the key to solving the riddle of how substances bond, and not long after Thomson’s discovery, German chemist Richard Abegg (1869-1910) suggested as much.
While studying noble gases, noted for their tendency not to bond, Abegg discovered that these gases always have eight valence electrons. His observation led to one of the most important principles of chemical bonding: atoms bond in such a way that they achieve the electron configuration of a noble gas. This has been shown to be the case in most stable chemical compounds.

Two different types of bonds

Perhaps, Abegg hypothesized, atoms combine with one another because they exchange electrons in such a way that both end up with eight valence electrons. This was an early model of ionic bonding, which results from attractions between ions with opposite electric charges: when they bond, these ions “complete” one another.
Ionic bonds, which occur when a metal bonds with a nonmetal are extremely strong. As noted earlier, salt is an example of an ionic bond: the metal sodium loses an electron, forming a cation; meanwhile, the nonmetal chlorine gains the electron to become an anion. Their ionic bond results from the attraction of opposite charges.
Ionic bonding, however, could not explain all types of chemical bonds for the simple reason that not all compounds are ionic. A few years after Abegg’s death, American chemist Gilbert Newton Lewis (1875-1946) discovered a very different type of bond, in which nonionic compounds share electrons. The result, once again, is eight valence electrons for each atom, but in this case, the nuclei of the two atoms share electrons.
In ionic bonding, two ions start out with different charges and end up forming a bond in which both have eight valence electrons. In the type of bond Lewis described, a covalent bond, two atoms start out as atoms do, with a net charge of zero. Each ends up possessing eight valence electrons, but neither atom “owns” them; rather, they share electrons.

Lewis structures

In addition to discovering the concept of covalent bonding, Lewis developed the Lewis structure, a means of showing schematically how valence electrons are arranged among the atoms in a molecule. Also known as the electron-dot system, Lewis structures represent the valence electrons as dots surrounding the chemical symbols of the atoms involved. These dots, which look rather like a colon, may be above or below, or on either side of, the chemical symbol. (The dots above or below the chemical symbol are side-by-side, like a colon turned at a 90°-angle.)
To obtain the Lewis structure representing a chemical bond, it is first necessary to know the number of valence electrons involved. One pair of electrons is always placed between elements, indicating the bond between them. Sometimes this pair of valence electrons is symbolized by a dashed line, as in the system developed by Couper. The remaining electrons are distributed according to the rules by which specific elements bond.

Multiple bonds

Hydrogen bonds according to what is known as the duet rule, meaning that a hydrogen atom has only two valence electrons. In most other elements—there are exceptions, but these will not be discussed here—atoms end up with eight valence electrons, and thus are said to follow the octet rule. If the bond is covalent, the total number of valence electrons will not be a multiple of eight, however, because the atoms share some electrons.
When carbon bonds to two oxygen atoms to form carbon dioxide (CO2), it is represented in the Couper system as O-C-O. The Lewis structure also uses dashed lines, which stand for two valence electrons shared between atoms. In this case, then, the dashed line to the left of the carbon atom indicates a bond of two electrons with the oxygen atom to the left, and the dashed line to the right of it indicates a bond of two electrons with the oxygen atom on that side.
The non-bonding valence electrons in the oxygen atoms can be represented by sets of two dots above, below, and on the outside of each atom, for a total of six each. Combined with the two dots for the electrons that bond them to carbon, this gives each oxygen atom a total of eight valence electrons. So much for the oxygen atoms, but something is wrong with the representation of the carbon atom, which, up to this point, is shown only with four electrons surrounding it, not eight.
In fact carbon in this particular configuration forms not a single bond, but a double bond, which is represented by two dashed lines—a symbol that looks like an equals sign. By showing the double bonds joining the carbon atom to the two oxygen atoms on either side, the carbon atom has the required number of eight valence electrons. The carbon atom may also form a triple bond (represented by three dashed lines, one above the other) with an oxygen atom, in which case the oxygen atom would have only two other valence electrons.

Electronegativity and Polar Covalent Bonds

Today, chemists understand that most bonds are neither purely ionic nor purely covalent; rather, there is a wide range of hybrids between the two extremes. Credit for this discovery belongs to American chemist Linus Pauling (1901-1994), who, in the 1930s, developed the concept of electronegativity—the relative ability of an atom to attract valence electrons.
Elements capable of bonding are assigned an electronegativity value ranging from a minimum of 0.7 for cesium to a maximum of 4.0 for fluorine. Fluorine is capable of bonding with some noble gases, which do not bond with any other elements or each other. The greater the electronegativity value, the greater the tendency of an element to draw valence electrons to itself.
If fluorine and cesium bond, then, the bond would be purely ionic, because the fluorine exerts so much more attraction for the valence electrons. But if two elements have equal electronegativity values—for instance, cobalt and silicon, both of which are rated at 1.9—the bond is purely covalent. Most bonds, as stated earlier, fall somewhere in between these two extremes.

Polar covalent bonding

When substances of differing electronegativity values form a covalent bond, this is described as polar covalent bonding. Sometimes these are simply called “polar bonds,” but that is not as accurate: all ionic bonds, after all, are polar, due to the extreme differences in electronegativity. The term “polar covalent bond” is much more specific, describing a bond, for instance, between hydrogen (2.1) and sulfur (2.6). Because sulfur has a slightly greater electronegativity value, the valence electrons will be slightly more attracted to the sulfur atom than to the hydrogen atom.
Another example of a polar covalent bond is the one that forms between hydrogen and oxygen (3.5) to form H2O or water, which has a number of interesting properties. For instance, the polar quality of a water molecule gives it a great attraction for ions, and thus ionic substances such as salt dissolve easily in water. “Pure” water from a mountain stream is actually filled with traces of the rocks over which it has flowed. In fact, water—sometimes called the “universal sol-vent”—is almost impossible to find in pure form, except when it is purified in a laboratory.
By contrast, molecules of petroleum (CH2) tend to be nonpolar, because carbon and hydrogen have almost identical electronegativity values—2.5 and 2.1 respectively. Thus, an oil molecule offers no electric charge to bond it with a water molecule, and for this reason, oil and water do not mix. It is a good thing that water molecules attract each other so strongly, because this means that a great amount of energy is required to change water from a liquid to a gas. If this were not so, the oceans and rivers would vaporize, and life on Earth could not exist as it does.

Bond energy

The last two paragraphs allude to attractions between molecules, which is not the same as (nor is it as strong as) the attraction between atoms within a molecule. In fact, the bond energy—the energy required to pull apart the atoms in a chemical bond—is low for water. This is due to the presence of hydrogen atoms, with their two (rather than eight) valence electrons. It is thus relatively easy to separate water into its constituent parts of hydrogen and oxygen, through a process known as electrolysis.
Covalent bonds that involve hydrogen are among the weakest bonds between atoms. (Again, this is different from bonds between molecules.) Stronger than hydrogen bonds are regular, octet-rule covalent bonds: as one might expect, double covalent bonds are stronger than single ones, and triple covalent bonds are stronger still. Strongest of all are ionic bonds, involved in the bonding of a metal to a metal, or a metal to a nonmetal, as in salt. The strength of the bond energy in salt is reflected by its boiling point of 1,472°F (800°C), much higher than that of water, at 212°F (100°C).

KEY TERMS

Anion: The negative ion that results when an atom gains one or more electrons. An anion (pronounced “AN-ie-un”) of an element is never called, for instance, the chlorine anion. Rather, for an anion involving a single element, it is named by adding the suffix -ide to the name of the original element—hence, “chloride.” Other rules apply for more complex anions. Atom: The smallest particle of an element. An atom can exist either alone or in combination with other atoms in a molecule.
Atomic number: The number of protons in the nucleus of an atom. Since this number is different for each element, elements are listed on the periodic table of elements in order of atomic number.
Bond energy: The energy required to pull apart the atoms in a chemical bond.
Cation: The positive ion that results when an atom loses one or more electrons. A cation (pronounced “KAT-ie-un”) is named after the element of which it is an ion and thus is called, for instance, the aluminum ion or the aluminum cation.
Chemical bonding: The joining, through electromagnetic force, of atoms representing different elements. The principal types of bonds are covalent bonding and ionic bonding, though few bonds are purely one or the other. Rather, there is a wide range of “hybrid” bonds, in accordance with the electronegativity values of the elements involved.
Chemical symbol: A one-or two-letter abbreviation for the name of an element.
Compound: A substance made up of atoms of more than one element. These atoms are usually joined in molecules.
Covalent bonding: A type of chemical bonding in which two atoms share valence electrons. Atoms may bond by single, double, or triple covalent bonds, which, in representations of Lewis structures, are shown by single, double, or triple dashed lines. (The double dashed line looks like an equals sign.) When atoms have differing values of electronegativity, they form polar covalent bonds.
Duet rule: A term describing the distribution of valence electrons when hydrogen atoms—which end up with only two valence electrons—experience chemical bonding with other atoms. Most other elements follow the octet rule.
Electron: Negatively charged particles in an atom. Electrons, which spin around the protons and neutrons that make up the atom’s nucleus, are essential to chemical bonding.
Electronegativity: The relative ability of an atom to attract valence electrons.
Element: A substance made up of only one kind of atom. Unlike compounds, elements cannot be broken down chemically into other substances.
Ion: An atom that has lost or gained one or more electrons, and thus has a net electrical charge. Ions may either be anions or cations.
Ionic bonding: A form of chemical bonding that results from attractions between ions with opposite electrical charges.
Lewis structure: A means of showing schematically how valence electrons are distributed among the atoms in a molecule. Also known as the electron-dot system, Lewis structure represents pairs of electrons with a symbol rather like a colon, which—depending on the situation—can be placed above, below, or on either side of the chemical symbol. In the Lewis structure, the pairs of electrons involved in chemical bonds are usually represented by a dashed line.
Molecule: A group of atoms, usually, but not always, representing more than one element, joined in a structure. Compounds are typically made of up molecules.
Neutron: A subatomic particle that has no electrical charge. Neutrons are found at the nucleus of an atom, alongside protons.
Nucleus: The center of an atom, a region where protons and neutrons are located, and around which electrons spin.
Octet rule: A term describing the distribution of valence electrons that takes place in chemical bonding for most elements, which end up with eight valence electrons. Hydrogen is an exception, and follows the duet rule. A few elements follow other rules, and some (most notably the noble gases) do not typically bond with other elements.
Periodic table of elements:A chart showing the elements arranged in order of atomic number. Vertical columns within the periodic table indicate groups or “families” of elements with similar chemical characteristics.
Polar covalent bonding: The type of chemical bonding between atoms that have differing values of electronegativity. If the difference is extreme, of course, the bond is not a covalent bond at all, but an ionic bond. Thus, although these are sometimes called polar bonds, they are more properly identified as polar covalent bonds.
Proton: A positively charged particle in an atom.
Valence electrons: Electrons that occupy the highest energy levels in an atom. These are the only electrons involved in chemical bonding. By contrast, the core electrons, or the ones at lower energy levels, play no role in the bonding of atoms.
Valency: The property of the atom of one element that determines its ability to bond with atoms of other elements.

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