2.1–2.3: Atoms form molecules through bonding.

Classical atomic models: a nucleus as a packed cluster of protons and neutrons, orbited by electrons.

A little bit of chemistry goes a long way in the study of biology and in understanding a great deal about your everyday life. Will eating those beans in your soup keep you up all night in gastrointestinal distress? Just knowing whether the beans are lentils or lima beans—each contains slightly different types of sugar molecules—will give you an answer. Will the butter you spread on your toast sabotage your efforts to lose weight? Understanding something about the carbon-hydrogen connections in the fat molecules can help you decide.

The chemistry that is most important in biology revolves around a few important elements, which are introduced at the end of this section. An element is a substance that cannot be broken down chemically into any other substances. Gold, carbon, and copper are elements you might be familiar with. Whatever the element, if you keep cutting it into ever smaller pieces, each of the pieces behaves exactly the same as any other piece. The smallest piece of pure gold will still have the softness, reflectivity, and malleability characteristic of that element (FIGURE 2-1).

Figure 2.1: Familiar elements.

If you could continue cutting, you would eventually separate the gold into tiny pieces that could no longer be divided without losing their gold-like properties. These individual component pieces of an element are called atoms. An atom is a bit of matter that cannot be subdivided any further without losing its essential properties. The word “atom” is from the Greek for “indivisible.”

Everything around us, living or not, can be reduced to atoms. All atoms—whether of the element gold or some other element such as oxygen or aluminum or calcium—have the same basic structure. At the center of an atom is a nucleus, which is usually made up of two types of particles, called protons and neutrons. Protons are particles that have a positive electrical charge and neutrons are particles that have no electrical charge. The amount of matter in a particle is its mass; protons and neutrons have approximately the same mass (FIGURE 2-2).

Figure 2.2: The atom.

Whirling in a cloud around the nucleus of every atom are negatively charged particles called electrons. An electron weighs almost nothing—less than one-twentieth of one percent of the weight of a proton. (The mass of an atom—its atomic mass—is made up of the combined mass of all of its protons and neutrons; for our purposes here, electrons are so light that their mass can be ignored.)

Particles that have the same charge repel each other; those with opposite charges are attracted to each other. Because all electrons have the same charge, the electrons in an atom repel each other. But because they are negatively charged, they are attracted to the positively charged protons in the nucleus. This attraction holds electrons close enough to the nucleus to keep them from flying away, while the energy of their fast movement keeps them from collapsing into the nucleus. When the number of protons and electrons is equal, the charges in the atom are balanced.

Atoms are tiny. Enlarge an atom by a billion times and it would only be the size of a grapefruit. Paradoxically, most of the space taken up by an atom is empty. That is, because the nucleus is very small and compact, the electrons zip about relatively far from the nucleus. If the nucleus were the size of a golf ball, the electrons would be anywhere from half a mile to six miles away.

Elements differ in their number of protons What distinguishes one element, such as chlorine, from another, such as neon or oxygen? The number of protons in an atom’s nucleus determines what element it is. As a rule, atoms of different elements have a different number of protons in the nucleus. A chlorine atom has 17 protons, a neon atom has 10 protons, and an oxygen atom has 8 protons. Each element is given a name (and an abbreviation, such as O for oxygen and C for carbon) and an atomic number that corresponds to how many protons it has (FIGURE 2-3).

Figure 2.3: The vital statistics of atoms. This key explains how to read a periodic table of elements. A full periodic table of the elements is at the back of the book.

The mass of an atom is often about double the element’s atomic number. This is the case when the number of neutrons in the nucleus is equal to the number of protons, because protons and neutrons have approximately the same mass. The element oxygen, for example, has the atomic number 8 reflecting that it has 8 protons, and because it has 8 neutrons it has an atomic mass of 16, simply the mass of the 8 protons and the mass of the 8 neutrons added together. (As noted above, we can ignore the mass of the electrons.)

Atoms of the same element don’t always have the exact same number of neutrons in their nucleus and electrons circling around it. They sometimes acquire or lose components. For example, an atom may have extra neutrons or fewer neutrons than the number of protons. Atoms with the same number of protons but different numbers of neutrons are called isotopes. An atom’s charge doesn’t change in an isotope, because neutrons have no electrical charge, but the atom’s mass changes with the loss or addition of another particle in the nucleus (FIGURE 2-4). Carbon, for example, has six protons and so usually has a mass of 12. Occasionally, though, a rare carbon atom has an extra neutron or two and an atomic mass of 13 or 14. These isotopes are called carbon-13 (¹³C) and carbon-14 (¹⁴C) and are referred to as “heavy” carbon. In nature, we frequently see mixtures of several isotopes for a given element. So although a sample of pure carbon is predominantly ¹²C atoms (with 6 protons and 6 neutrons), some ¹³C and ¹⁴C atoms are present in the sample, too.

Figure 2.4: Isotopes are atoms that have the same number of protons but a different number of neutrons.

Most elements and their isotopes have perfectly stable nuclei that remain unchanged virtually forever, never losing or gaining neutrons, protons, or electrons. A few atomic nuclei are not so stable, however, and break down spontaneously sometime after they are created. These atoms are radioactive, and in the process of decomposition they release, at a constant rate, a tiny, high-speed particle carrying a lot of energy. (The particle may be a proton, neutron, or electron; sometimes just energy is released and no particle.) For example, uranium-238 (which has 92 protons and 146 neutrons in its nucleus) is a radioactive element. It spontaneously loses a particle containing 2 protons and 2 neutrons, turning it into an isotope of a different element altogether—thorium-236, in this case. Thorium is radioactive as well, one in a long chain of isotopes, each decaying into another radioactive element, until finally producing the stable element lead (with an atomic mass of 206). Radioactive atoms turn out to be useful in determining the age of fossils (see Section 8-18), in medical imaging and cancer treatment, and in generating vast amounts of energy.

All the known elements can be arranged in a scheme, in the order of their atomic number, called the periodic table (see Figure 2-3). A copy of the periodic table is provided at the back of the book. So far, about 90 elements have been discovered that are present in nature, and about 25 others can be made in the laboratory. Everything you see around you is made up of some combination of the 90 or so naturally occurring elements.

Of all the elements found on earth, only 25 are found in your body. The “Top 10” most common of these make up 99.9% of your body mass, but they are not present in equal measures (FIGURE 2-5). The “Big 4” elements—oxygen, carbon, hydrogen, and nitrogen—predominate and make up more than 96% of your body mass. With knowledge about the Big 4, you can understand a huge amount about nutrition and physiology (how your body works), so we’ll focus on the properties of these four elements later in this chapter.

Figure 2.5: The body’s chemistry.