Atoms are bound together by electron glue.
The properties of all materials, be they sticky, magnetic, brittle, or green, arise from the infinite different ways atoms can be arranged and linked together. Two atoms form a chemical bond when their electrons rearrange themselves so that each atom feels an attractive force. The attraction of positive and negative electrical charges holds everything together.
Elements in Combination
The distinctive materials that enrich our world can be described in architectural terms. Various atomic building blocks—the elements —are joined in structures using chemical bonds. Elements can be combined with each other in countless ways to form chemical compounds that have properties and applications very different from their elemental raw materials. Nature and chemists have created millions of different chemicals. Some are very simple, like H2O (water) with two hydrogen atoms for every oxygen, or NaCl (table salt) with a one-to-one ratio of sodium and chlorine. Other compounds are exceedingly complex, combining a dozen or more different elements.
With fewer than one hundred elements to play with, it might seem that the chemist’s job of making new chemical combinations would be quickly exhausted. Not so. A million chemists could each make a new sample every day for a million years and still not come close to running out of things to try. There are about 70,000,000 possible four-element combinations in one-to-one-to-one-to-one ratios, to say nothing of the countless variations in element ratios. For each possible composition there are hundreds of ways to mix and treat the chemicals, each resulting in a different arrangement of the same set of constituent atoms, and therefore in a material with different properties. Just as in cooking, the temperature and time of baking can make all the difference between a culinary masterpiece and an unmitigated disaster. The successful chemist, like a great chef, must apply skill and intuition in preparing new recipes.
Chemists make their living by trying to create new and useful mixtures of atoms, or by trying to manufacture established useful chemicals in new ways. This is not an idle pursuit. Almost every aspect of modern life—food and clothing, transportation and communications, sports and entertainment—depends on the discoveries of chemistry.
THE ELECTRON GLUE
Chemistry is first and foremost the science of electrons and their interactions. That may seem an odd assertion, for chemistry is usually portrayed as the science of test tubes and bubbling beakers and odd mixtures of stuff that turns blue. But all those chemical reactions are the results of electrons shifting between atoms.
When two atoms approach each other, their outermost electrons come in close contact; the two negative electric charges repel each other. In most instances, as in the collision of gases in the atmosphere, the atoms just careen off to another chance rendezvous. Occasionally, however, colliding atoms stick together by an exchange or sharing of electrons.
The first step in understanding the extraordinary variety of materials that surround us is to look at the way that two or more atoms can link together: the chemical bond. There are several different kinds of bonds, each producing dramatically different properties. The general rule that dictates what kind of linkage, if any, will be formed when two atoms come near each other is simple: the two-atom system will try to reach a state of lowest possible energy Just as a ball rolls down a hill to get to a position of lowest gravitational potential energy, an atom’s electrons rearrange themselves to reach a state of lowest electrical potential and kinetic energy.
Three types of chemical bonds—ionic, covalent, and metallic—hold almost all materials together. Each type involves a different strategy to rearrange electrons among atoms, but the secret to forming most chemical bonds is the extreme stability of atoms with filled electron shells. Atoms with exactly 2 or 10 or 18 or 36 electrons (corresponding to 1, 2, 3, or 4 filled shells) are much happier (if you will excuse the anthropomorphism) than those with a few electrons more or less—11 or 17 electrons, for example. We think of 2, 10, 18, and 36 as the magic numbers of chemical bonding. In fact, the only elements that don’t readily engage in any kind of chemical bonding are found in the extreme right-hand column of the periodic table—elements 2, 10, 18, and 36, with filled outer electron shells. These are the so-called inert gases, including lighter-than-air helium gas and the neon gas of colorful lights.
The Ionic Bond
An ionic bond forms when one atom of a pair gives up an electron while the other acquires it on permanent loan, so that both atoms achieve a magic number. In common table salt, for example, a sodium atom with 11 electrons transfers its outermost electron to a chlorine atom, which has 17 electrons. When this shuffle occurs, the sodium atom winds up with only 10 electrons, while the chlorine atom increases to 18 electrons—both magic numbers. What’s more, by shifting an electron, sodium becomes a positive ion and chlorine becomes a negative ion, so the atoms form a strong electrostatic attraction—an ionic bond.
A glance at the periodic table (see page 75) reveals some elements engage in ionic bonding. All the elements in the two farthest left-hand columns of the table have only one or two electrons in their outer shell. These elements would be much happier to get rid of their outermost electrons and become positively charged ions. By contrast, elements in columns near the right-hand side of the table are one or more electrons shy of a full outer shell; they are eager to add a few more electrons and become negatively charged ions. Positive attracts negative and so, voilà, an ionic bond forms.
Hundreds of everyday materials are held together by ionic bonds. Table salt is the classic textbook example of ionic bonding, but compounds that combine silicon and oxygen, two of the Earth’s most abundant elements, are by far the most common ionic compounds. Most beach sand is a mineral called quartz, formed when silicon atoms, which yield four electrons, bond to pairs of oxygen atoms, each of which collects two electrons, leaving both silicon and oxygen with closed shells. The large electrostatic charges of +4 on every silicon and—2 on every oxygen results in a strong electrostatic force between adjacent atoms. That is why quartz, which can scratch steel, is such a tough and resistant mineral. Similar ionic silicon-oxygen bonds provide the strength and hardness of window glass, china and other ceramics, and most common rocks and minerals.
The Metallic Bond
Atoms in metals generally have filled electron shells plus one or more electrons left over. Metallic sodium, a soft silvery metal with 11 electrons per atom, has one electron in its outer shell; magnesium (element 12) has 2; aluminum (element 13) has 3; and so on. When metal atoms combine, some or all of these overflow electrons leave their homes and wander freely throughout the metal, leaving behind positively charged ions as they swim in a sea of negative charge. Each nucleus creates a local island of positive charge, and electrostatic forces hold the whole metal system together. You can think of the metallic bond as a system in which outer electrons are shared by all the atoms in the system, in contrast to the ionic bond, in which one atom permanently donates an electron to another.
Metallic bonds differ dramatically from ionic bonds. In an ionic crystal there are always two very different kinds of atoms, some with positive charge and some with negative charge. Each ion is surrounded by those of opposite charge. Obviously, a pure element (with only one kind of atom) can never have ionic bonds. In a metal, however, all atoms play exactly the same role as their neighbors. Every metal atom is surrounded by similar metal atoms. It’s probably not surprising, therefore, that about three quarters of all known pure elements, including iron, aluminum, copper, and gold, form metallic bonds. Metallurgists often mix two or more of these elements to create metal alloys with distinctive properties. Brass combines copper and zinc, bronze forms from copper and tin, and steel alloys can incorporate a dozen metallic elements along with the essential iron and carbon components.
The Covalent Bond
A curious dilemma faces two carbon atoms as they approach each other. Each carbon atom has 6 electrons, including 4 in the outer shell—halfway between 2 and
10. Are these atoms supposed to donate or accept electrons to achieve a magic number? In fact, neither atom does anything of the sort. The two carbons actually end up sharing their outer electrons—in effect creating a filled outer shell between both atoms. Electrons, constantly flipping back and forth between adjacent atoms, do not really belong to either carbon. This sharing produces a force that holds the two atoms together—a covalent bond. Although the covalent bond is most often seen holding carbon atoms together, many other elements, including silicon, sulfur, and nitrogen, participate in this sort of bonding as well.
The covalent bond is the basis of all life. It holds the tissues of your body together and keeps your DNA from falling apart. It also can be seen in plastics, nylon, diamonds, and superglue. Look around you. Carbon-carbon bonding is an essential part of almost everything you see, except for metals, water, glass, and ceramics. Carbon-carbon bonding is so important that tens of thousands of specialists, known as organic chemists, devote their research lives exclusively to the study of carbon compounds.
The main reason that the covalent bond is seen so often in nature is that covalent compounds seldom stop at a single pair of atoms. If we start with two carbon atoms sharing a pair of electrons, for example, other carbon atoms can start to share other electrons, forming ever longer and more complex structures. Our world is full of giant atomic structures with thousands of carbon atoms in long chains or intricate branches or interlocking ring patterns. There is literally no limit to the variety of potential organic compounds. In all of these substances, no matter how fancy, each pair of adjacent carbon atoms shares some electrons in a covalent bond.
Chemical Bonds and the Real World
Chemical bonds in lots of materials don’t exactly fit the neat definitions of ionic, covalent, or metallic. Electrons are always on the move, and it is not always possible to say exactly where a given particle spends its time. If outer electrons remain near one ion then the bonding is ionic; if they are shared equally between a pair of atoms then the bonding is covalent; if they are free to roam the crystal then the bonding is metallic. But in many substances electrons divide their time unevenly between two or more atoms. Such divided loyalties result in bonds with a mixed character.
The common mineral pyrite, known to miners as fool’s gold, owes its unusual combination of properties to iron-sulfur bonds that have a complex mixture of covalent, metallic, and ionic character. Fool’s gold’s shiny luster confused thousands of novice California prospectors during the 1849 gold rush, but the mineral breaks with rough, brittle edges, characteristic of materials with ionic or covalent bonding.
The van der Waals and Hydrogen Bonds
In thousands of compounds groups of atoms bond without ever giving up an electron, a situation very different from ionic, metallic, or covalent bonding. This type of attractive force arises because when two atoms come near each other, their electron clouds distort as a result of the repulsive force between electrons. The electrical forces between two distorted atoms are (1) repulsion between the two nuclei, (2) attraction between the electrons of one atom and the nucleus of the other, and (3) repulsion between the electrons. Although it isn’t obvious, it turns out that in this situation the attractive forces can win, in which case a weak bond forms between the atoms. We call this bond the van der Waals force, after Dutch physicist Johannes van der Waals (1837–1923). Many soft materials, from candle wax to talcum powder, owe their distinctive properties to weak van der Waals bonds.
The “hydrogen bond,” found in every living thing, is an important variant of van der Waals bonding. Hydrogen atoms, with just one proton and one electron, tend to bond to just one other atom at a time—often oxygen or carbon. The lone hydrogen electron is shifted to that other atom, leaving the proton as a somewhat exposed positive bump on the resulting molecule. This positively charged region can attract other atoms by ordinary electrical force, and the bond formed in this way is called the hydrogen bond. You can picture it as a distorted hydrogen atom forming the glue that holds two other atoms together.
A molecule of water, for example, has two hydrogens attached to one oxygen, and looks very much like Mickey Mouse’s head with small ears: the oxygen is Mickey’s negatively charged head, while the positive protons form the ears. When solid ice forms by hydrogen bonding, the molecules all arrange themselves so that oppositely charged ends line up opposite each other. The behavior of hydrogen also explains why water dissolves so many things, from salt to sugar. Positive and negative ends of water molecules exert powerful forces on ions—ripping apart ionic crystals like table salt, ion by ion. Positive sodium ions congregate near the oxygen ends of water molecules, while negative chlorine is greeted by hydrogens.
Each boomerang-shaped molecule of water (H2O) contains a central atom of oxygen flanked by two smaller atoms of hydrogen.
ELECTRICAL CONDUCTIVITY
AND THE CHEMICAL BOND
The most important electrical property of any material is its ability to conduct electricity—to allow electrons to flow through it. Our technological society depends on a wide range of materials with different electrical properties: efficient conductors to carry current, insulators that protect the user from harm, and semiconductors, the backbone of the microelectronics industry, which can be controlled to conduct electricity under special circumstances.
Whether a material will conduct electricity depends on how its electrons move. It is not surprising, therefore, that conductivity and chemical bonding are closely linked. Each bond type gives rise to distinctive electrical properties.
In general, materials in which electrons are loosely bound make good conductors. In such materials the presence of an outside force (supplied by a battery, for example) is enough to break the electrons loose from their moorings and start them moving. Their movement then constitutes an electrical current in the material. If, on the other hand, electrons are tightly locked in, an outside force can’t dislodge them and no current will flow.
Insulators
Insulators resist the flow of electrons, and so serve such essential functions as preventing short circuits and protecting electrical consumers from shocks. Materials with ionic bonds make excellent insulators. In ionic compounds there is a onetime transfer of an electron from one ion to the next, but from then on each electron is bound tightly to a given nucleus. In this situation electrons are not easily moved. Glass and ceramics, composed primarily of ionically bonded silicon and oxygen, have long been used for the most critical high-voltage insulators.
Many covalent compounds with long chains of carbon-carbon bonds—materials known collectively as plastics—also serve as cheap, reliable insulators. Covalently bonded electrons seldom stray far from their home atom, so they are often almost as good insulators as ionic compounds. Flexible plastics, furthermore, are much easier to fabricate and use than brittle ceramics, which explains why plastic wall sockets and light switches are the norm in homes and offices.
Conductors
The terms “metal” and “electrical conductor” are almost synonymous. The electrons in a metal’s negative sea respond easily to outside forces, and the best electrical conductors rely on this ready supply of mobile electrons. Silver is the best conductor at room temperature, but copper, almost as good, is much cheaper. Gold finds important uses as a noncorroding coating for contact surfaces, while lightweight (and cheap) aluminum is often used for carrying large currents in power lines.
Semiconductors
As the name suggests, a semiconductor is a material that will carry electricity (like a conductor), but won’t carry it very well. Semiconductors walk a tightrope between a conductor with lots of mobile electrons and an insulator with none. The key to a successful commercial semiconductor is to have just a few mobile carriers of electrical charge.
The abundant element silicon provides the starting material for almost all of today’s semiconductor devices. Silicon, with a half-filled outer shell of electrons, behaves very much like carbon. Silicon-silicon bon
ds are covalent, and pure silicon does not conduct electricity very well because the shared electrons are tightly bound. Atomic vibrations always shake some electrons loose, however, so that at any given moment there will be a few electrons wandering around in a silicon crystal. The result: silicon has a net conductivity much less than that of metals, but much more than insulators. That’s why it’s called a semiconductor.
The loose electrons don’t tell the whole story as far as electrical current in a semiconductor is concerned. Whenever an electron shakes loose, it leaves behind a vacancy—what physicists call a hole. As far as physicists are concerned, holes can conduct electricity just as well as electrons. To see why this is so, think of the moving holes as analogous to something you see in stop-and-go traffic at rush hour. As a space (a hole) opens up in front of your car, you advance to fill that space, the next car behind moves forward to fill the new space, and so on. Most people would say the cars have moved forward, but it is just as correct to say that the hole has moved in the opposite direction. As commuters, most of us couldn’t care less how fast the holes are moving, but if you are a physicist worried about the net movement of electrical charge, electrons or holes serve equally well. You can move electric charge by shifting negative electrons one way, or just as easily by shifting positive charge—the holes—in the opposite direction.
A slight impurity can alter the behavior of a semiconductor like silicon dramatically. For example, semiconductors can be made so that an atom of phosphorus replaces one out of every few million silicon atoms. Phosphorus has one more electron than silicon has in its outer shell. All but one of these electrons form the covalent bonds that hold the crystal together, leaving one electron free to wander around. Thus, the altered crystal can conduct electricity without having to depend on the electrons stolen from covalent silicon bonds. The process of adding a trace of phosphorus is called “doping” the crystal, and the slight excess of negative electrons yields an “n-type” (for negative charge carriers) semiconductor.
Science Matters Page 10