Although the atom is the basic building block of matter, you will seldom find one existing in isolation. An oxygen atom does not waft through the air by its lonesome. It is always connected to at least one other atom, be it another oxygen atom or some other element such as carbon. Similarly, the atoms that make up your body, this book, and the floor beneath your feet all exist in closely bound groups.
In fact, only six out of more than 100 known elements exist in monoatomic, or single-atom, form. Those six are the noble gases: helium, neon, argon, krypton, xenon, and radon. To be fair, helium is the second-most-abundant element in the universe, thought to make up almost one-quarter of all matter in existence. But here on Earth, helium and its fellow noble gases are exceedingly rare, totaling less than 1 percent of the lower atmosphere.
The fact that virtually all atoms exist in groups is so important because the physical structure of these groups and the forces that bind them together determine how they will "behave." In other words, a basic understanding of atomic groups enables us to predict the physical properties of virtually every chemical we will ever encounter. Importantly, these substances include not only everything we touch, see, or smell, but also the chemicals in your body that make it possible for you to pick up this book, read its words, and grasp their meaning.COMING TOGETHER
Atoms join together by sharing, losing, or gaining electrons. As a general rule, they do so to achieve a stable set of eight electrons in their outer electron shell. The exception to this octet rule is hydrogen. It can either lose its single electron to become a bare proton (H+), or it can achieve a stable duet of electrons by sharing or stealing one.
When atoms steal or lose electrons, they become charged particles called ions. The metallic element sodium (Na), for example, readily gives up an electron to form the positively charged ion, or cation, Na+. The poisonous gas chlorine (Cl2) readily steals electrons in order to produce negatively charged ions, or anions (2Cl−).
Like all ions, sodium cations and chlorine anions have very different properties from the neutral atoms from which they form. They also show a great attraction for each other. Like positive and negative ends of two magnets, cations and anions are drawn together by an electromagnetic force. This force creates an ionic bond that joins Na+ and Cl− to form NaCl, otherwise known as sodium chloride, or table salt.IONIC COMPOUNDS
Such a group of two or more oppositely charged ions is called an ionic compound. It always contains equal amounts of positive and negative charge. In other words, its positive ions balance out the negative ones.
The way we represent an ionic compound—the formula NaCl, for example—gives somewhat misleading information about its structure. Ionic bonds seldom if ever exist between just two ions such as one Na+ and one Cl−. Instead, the attraction between oppositely charged ions causes them to group into a three-dimensional pattern (see Figure 1). So around every sodium cation (Na+), one would find six tightly packed chlorine anions (Cl−); and around every chlorine anion, one would find six tightly packed sodium cations. The resulting compound forms a crystal such as the tiny cube we know as a "grain" of salt.
Because of differences in size, shape, and charges of ions, not all ionic compounds look like cubic salt crystals. But they all share many important characteristics. For example, ionic compounds all form solids at room temperature. This is because strong attractive forces hold their ions together in a three-dimensional grid, or crystal lattice.
Heat does not easily disrupt this rigid lattice. So the melting points of such compounds are typically high—from several hundred degrees to more than 3,500° F (2,000° C).
Yet most ionic compounds can be easily cracked, or split. This, too, results from their unique structure. When a force strikes a portion of an ionic crystal, it pushes positive and negative ions out of alignment. As a result, ions of the same charge come into contact and repel one another. This causes the crystal lattice to cleave, or fracture cleanly, along a flat plane.
Despite their hardness in open air, many (but not all) ionic compounds easily dissolve in water. This is because the particles that make up water (H2O) exert strong positive and negative forces of their own, and so can pull the oppositely charged ions apart.
When in a water solution or in their molten (liquid) state, ionic compounds become very good conductors of electricity. The electricity is carried by their charged ions (such as Na+ and Cl−) as they move through the liquid.MOLECULES
Instead of losing or stealing electrons outright, many atoms share one or more pairs of electrons. When they do so, the shared electrons form a new orbit that encompasses both atoms and ties them together. Such a shared-electron connection is called a covalent bond.
Generally speaking, two atoms form a covalent bond instead of an ionic bond if their competition for electrons is not so lopsided that one or the other ends up stealing the electrons entirely. A group of two or more atoms joined by a covalent bond or bonds is called a molecule.
The simplest molecules comprise two atoms of the same element. Examples include the common gaseous elements hydrogen (H2), oxygen (O2), nitrogen (N2), and chlorine (Cl2). Ozone (O3) is an example of an elemental molecule that is made up of three atoms.
Molecules with more than one element, or kind of atom, are called molecular compounds. Just as an element loses its original qualities when it becomes an ion, elements forming compounds lose their individual natures and gain a new, shared identity. Oxygen, for example, is not nearly as good to breathe when it combines with carbon to form deadly carbon monoxide (CO), or with hydrogen to form water (H2O). Two- and three-atom compounds are the simplest examples of molecular compounds. Some of the most complex include biological proteins and synthetic polymers with many thousands of atoms.
While elemental molecules such as oxygen (O2) or chlorine (Cl2) share their electrons equally, molecular compounds can show polarity, or differences in electrical charge between their atoms. This results from the different electronegativities, or electron pulls, of the different elements that comprise these compounds. An example of this type of uneven covalent bond exists in water (H2O), a very polar molecule. The oxygen atom tends to "hog" the electrons it shares with the two hydrogen atoms. A similarly lopsided bond exists between hydrogen and nitrogen in ammonia (NH3).
Importantly, the slightly positive hydrogen ends of these molecules tend to attract the slightly negative ends of their neighbor molecules. This type of attraction is called hydrogen bonding. It is a very weak sort of chemical bond, similar to the intermolecular force seen with ionic compounds, but not nearly as strong.
Because they are more or less neutral particles, most molecules show only weak attractions to each other. As a result, molecules do not cling together to form the strong crystal lattices seen with ionic compounds. Instead, they are commonly found as gases or liquids. A few familiar examples are gases made up of covalently bonded molecules, including oxygen (O2), chlorine (Cl2), carbon dioxide (CO2), carbon monoxide (CO), and methane (CH4), and liquids made up of polar molecules, such as water (H2O), ethanol (CH3CH2OH), and methylene chloride (CH2Cl2).
Larger or heavier molecules, such as table sugar (C12H22O11) and elemental iodine (I2), tend to settle into solid form at room temperature. Molecules can also exist in giant structures called macromolecules, in which an entire crystal consists of one huge molecule. Familiar examples include silicon dioxide (quartz) and the crystalline form of carbon (diamonds).Molecular Shapes. Molecules with more than two atoms have bonds that point in different directions. As a result, they have three-dimensional shapes. Understanding these shapes is important, because they influence how a given molecule will behave around other molecules.
We can represent these shapes in two dimensions with the symbols known as Lewis structures, as follows:
Here, water (H2O) is seen to have a "bent," or "V shape." (The double dots in this Lewis structure represent oxygen's remaining pair of unshared electrons.)
A molecule of carbon dioxide (CO2), however, has a linear shape:
and a molecule of formaldehyde (H2CO) has a triangular shape:
We can predict the shape of almost any molecule by remembering that each bond is made up of a paired set of electrons, and that each pair of electrons tries to stay as far away from the next as possible. This is just what we see in each of the above examples.
What shape will four covalent bonds assume around a central atom? We can try to represent this with a two-dimensional Lewis structure for methane (CH4), as follows:
But a three-dimensional model would work better to show the resulting shape of this molecule: a tetrahedron.
Another common molecular shape is the triangular pyramid, a rather flattened version of the tetrahedron shown above. It is what we see with three covalent bonds and a pair of free electrons around our central atom, as with ammonia (NH3):
Chemists and chemistry students often use "Tinkertoy" models of sticks and balls to represent and better "see" such three-dimensional molecules.
With these useful mental images of molecules and molecular compounds, we can better understand many of their qualities. The bent shape of a water molecule, for example—caused because the electrons of the oxygen atom “push away” the electrons of the two hydrogen atoms—shows us that this is a polar molecule:
That is, its electrons tend to cluster at one end of the highly electronegative oxygen atom. As a result, the two ends of a water molecule carry opposite charges. Indeed, the unique and fascinating properties of water have very much to do with its shape and resulting polarity.POLYATOMIC IONS
In addition to simple ions and molecules, there exists a third kind of hybrid compound called the polyatomic ion. Polyatomic ions can be thought of as molecules that bear either a positive or a negative charge. In fact, they consist of groups of covalently bonded atoms that bear an overall negative or positive charge. So, as a unit, the group behaves like an ion and forms ionic bonds with other ions, both simple and polyatomic.
Some familiar examples of polyatomic ions include phosphates (PO43−), commonly used to produce fertilizers, detergents, insecticides, and medicines; and sulfites (SO32−), used to preserve food and make wine.
Most polyatomic ions bear negative charges. An important exception is the ammonium ion (NH4+), which combines with a chlorine ion (Cl−) to form ammonium chloride (NH4Cl), the electrolyte in electric batteries, and with the sulfate ion (SO42−) to form the common fertilizer ammonium sulfate [(NH4)2SO4].
Having useful mental images of ions, molecules, and compounds will also help us understand the various physical states of matter. Chemicals take solid, liquid, and gaseous form for reasons very much related to their shape and attractive forces.