Science Matters (15 page)

The energy of chemical explosions results when dense solids or liquids are quickly converted into hot expanding gas. Bullets, cannonballs, and rockets accelerate because of this force, which is nothing more than the cumulative effect of individual atomic collisions.

At extremely high temperatures, like those in the sun, gas takes on a different character. In this gas-like state of matter, called plasma, electrons are stripped off atoms. At low temperatures, only a few electrons per atom may be separated, but at very high temperatures—typically above 100,000 degrees—electrons are
completely stripped from gas molecules, yielding a complete plasma. Plasmas exhibit unusual properties not encountered in typical gases. For example, plasmas conduct electricity and can be confined in magnetic fields called magnetic bottles.

Even though you’ve never actually felt a plasma, they are by far the commonest state of matter in the universe. Every star, the sun included, is composed primarily of dense hydrogen-and helium-rich plasma, while tenuous plasma-like gas occurs in the outer atmospheres of several planets, including Earth. There is even plasma (though one in which only a few electrons have been stripped from each atom) in your fluorescent lightbulb.

Like gases, liquids have no fixed shape, but they differ from gases in having a fixed volume. At the atomic level liquids behave something like a bowl full of marbles. Like marbles, the liquid molecules slide over each other, easily shifting while retaining their available volume. Individual molecules do not stick tightly together, however, so the entire mass is free to change shape (or spill on the floor).

Solids are materials that have a more or less fixed shape. Atoms in solids bind together with sufficient strength to stay in place. Crystals, glasses, and plastics, three of the most important varieties of solids, differ primarily in the regularity of their atomic structures.

Metals, gemstones, bones, and computer chips consist of crystals. They contain regular three-dimensional arrays of atoms, repeated over and over again in a sort of Tinkertoy arrangement. You can imagine a crystal as something like a huge stack of boxes, with each box the same size and shape and all the boxes holding exactly the same atomic contents. The size of each “box” in a real crystal is usually no more than a ten millionth of an inch, and each box might contain up to a few dozen atoms. The structure of a crystal is orderly, with row upon row of exactly the same kinds of atoms; travel 100 boxes in any direction and you know exactly what you’ll find.

Three different kinds of solids are distinguished by the regularity of their atomic structures. Crystals (A) possess row after row of identical atom boxes; trillions of boxes stack to form one tiny crystal. Glasses (B) have random structures; there is no way to predict what atom might reside a short distance from any starting point. Plastics (C) form from long chains of carbon atoms called polymers; atoms are ordered in one direction (along the chain) but are spaced randomly elsewhere
.

Plastics, invented in the twentieth century, are among the commonest
solids we encounter in daily life. All plastics in common use are synthetic chemicals not found in nature, but built from long, chain-like molecules linked together by carbon atoms. Like any chain, they have a predictable structure in one direction; travel along any single chain and you’ll keep running into carbon atoms like beads on a string. The chains themselves, however, form a hopelessly tangled mass of interwoven strands that give plastics their rigidity and strength. When heated, the chains separate slightly and slide apart, softening the plastic and allowing it to be poured, rolled, or otherwise reshaped into new forms. The long list of common plastics includes nylon and polyester (in clothing and fabrics), Lucite (in sculptures), vinyl (floor tiles and furniture), and epoxy (in cements and glue).

Glass structures differ considerably from both crystals and plastics. Atoms in a glass are jumbled together almost at random, but usually with some regularities in the atomic architecture. Common window glass, for example, contains mostly silicon and oxygen, with almost every silicon atom surrounded by four oxygen atoms and almost every oxygen atom next to two silicons. But there aren’t any “boxes” of structure neatly stacked in window glass. Travel just a few atoms in any direction from any point and there is no way to guess whether you will find an oxygen or a silicon atom. Beyond the nearest neighbor atoms, the structure is random.

If an ice cube falls on the floor it melts, leaving a puddle for you to clean up. If you boil water on the stove, you produce a cloud of steam. If you let your paintbrush dry without cleaning it, the brush becomes hard and useless. These are examples of processes in which matter changes phase—from solid to liquid in the first case, from liquid to gas in the second, and from liquid to solid in the
third. The atoms and molecules are the same before and after the change of phase, but their relationship to one another is different.

In the ice cube, for example, the molecules of water in the crystal lattice vibrate faster and faster as heat from the outside air moves in. Eventually they vibrate so fast that they tear loose and begin to move around freely. When this happens, the ice turns to water and we say that the cube has melted. In just the same way, molecules in boiling water move faster and faster until they lose contact with one another and enter the air as a gas. As paint hardens, on the other hand, atoms link up to form long chains where before there were separate molecules. Physicists know melting, boiling, and solidification as phase transitions.

All phase transitions involve shifts of energy. Materials soak up heat while they melt or boil, but their temperature stays the same—all of the energy input goes into breaking up the old atomic arrangement. This is why you use ice cubes to keep your drink cold. When you take an ice cube from the refrigerator, it is at a temperature well below freezing. It starts to warm up as it absorbs heat from your drink, cooling the drink as it does so. When the ice gets to 32°F, however, it remains at that temperature until it has absorbed enough energy to break all the bonds between the water molecules in its crystal structure. Only after the cube has melted completely and all the water is in liquid form does the water temperature start to rise again.

Chemistry involves a lot more than bonding and changes of state. In fact, our world would be pretty boring if it weren’t for chemical reactions, the processes by which atoms and small molecules combine to form larger molecules, and larger molecules
break up into smaller fragments. When we take a bite of food, light a match, wash our hands, or drive a car, we initiate chemical reactions. Every moment of every day, from the digesting of your food to the clotting of your blood, countless chemical reactions sustain life in every cell of your body.

All chemical reactions arise from the rearrangement of atoms, as well as the shuffling of the atoms’ outer electrons to form new chemical bonds. So when a piece of soft silvery sodium metal comes into contact with the corrosive gas chlorine, the two elements react violently to form sodium chloride—table salt. When a piece of iron metal is exposed to oxygen in the atmosphere they react, albeit much more slowly, to form rust.

Countless millions of chemical reactions take place in the world around us. Some occur naturally and others occur as the result of human design or intervention. In your everyday world, however, you are likely to see a few types of chemical reactions over and over again. Perhaps the most distinctive chemical feature of our planet’s atmosphere is the abundance of the highly reactive gas oxygen, which drives many of the most familiar chemical reactions in our lives. Oxidation includes any chemical reaction in which oxygen accepts electrons while combining with other elements. Rusting is a common gradual oxidation reaction in which iron metal combines with oxygen to form the familiar reddish iron oxide, while burning is a much more rapid oxidation in which oxygen and carbon-rich fuel combines to produce carbon dioxide.

Reduction is the opposite of oxidation—that is, it involves the addition of electrons to an atom. Thousands of years before scientists discovered the element oxygen, primitive metalworkers had learned how to reduce metal ores by smelting. In the iron-smelting
process, ironmasters heat a mixture of iron oxide ore and lime (calcium oxide) in an extremely hot charcoal fire. The lime lowers the melting temperature of the entire mixture, which then reacts to produce iron metal and carbon dioxide.

Oxidation and reduction reactions are essential to life, and they define the principal difference between plants and animals (see Chapter 15). Animals eat carbon-rich food and obtain energy through oxidation in their cells, releasing carbon dioxide as a by-product. Plants take in carbon dioxide and use the energy in sunlight to reduce it, releasing oxygen as a by-product.

Acids are common corrosive chemicals that, when put into water, produce positively charged hydrogen ions in the solution. Lemon juice, orange juice, and vinegar are examples of common weak acids, while sulfuric acid (used in car batteries) and hydrochloric acid (used in industrial cleaning) are strong acids. Bases are another class of corrosive chemicals that, when put into water, produce negatively charged OH ions. Milk of magnesia and other antacids are weak bases, while ammonia cleaning fluids and most drain cleaners are stronger bases.

Acid-base chemical reactions occur when an acid and a base are brought together. When you swallow an antacid, your stomach acid (with excess H) and the antacid (with excess OH) mix together. When the H and OH ions react to form water (H
₂
O), we say that the stomach acid has been neutralized.

The molecular building blocks of most common biological structures, including simple sugars and amino acids (see Chapter
15), consist at most of a few dozen atoms. Yet proteins, DNA, and other essential biological molecules are huge, with millions of atoms in a single unit. How can small building blocks yield the large structures characteristic of living things? The answer lies in the process of polymerization.

A polymer is a large molecule formed by linking together many smaller molecules. In spiderwebs, clotted blood, muscle fibers, and a thousand other substances, living systems have mastered the art of combining small molecules into long chains through polymerization reactions. Synthetic polymers from plastics to paints usually begin in liquid form, with small molecules that move freely past their neighbors. Polymers form from the liquid when the ends of these molecules begin to link up.

The longevity of some polymers presents a growing problem in an age of diminishing landfills. Nevertheless, over time most polymers decompose into short segments through the process of depolymerization. Some of the most familiar depolymerization reactions occur in your kitchen. Polymers cause the toughness of uncooked meat and the stringiness of many raw vegetables, so we marinate and cook our food to break down these polymers.

As museum curators are painfully aware, not all depolymerization reactions are desirable. The breakdown process can affect leather, paper, textiles, and other historic artifacts made of organic materials. Storage in a cool, dry environment may slow the depolymerization process, but there is no known way to repolymerize old brittle objects.

When you go to the grocery store you test the food you’re about to buy. Are the tomatoes firm? Does the meat have good color? Is
the lettuce crisp, the bread soft? As you evaluate your potential purchases you are determining physical properties of the materials you buy.

A physical property is any aspect of a material that you can measure. Every measurement has three essential components: a sample, a source of energy, and a detector. The sample is the thing you want to measure. Samples vary widely in size and character—you can study single subatomic particles, a piece of fruit, the whole Earth, entire galaxies, and anything in between.

All measurements depend on an interaction of the sample with some form of energy. Color is the interaction of matter with light. Electrical conductivity is the interaction of matter with an electric field. Brittleness can be measured by the interaction of matter with the force of a moving hammer. Without energy, you can’t measure these properties.

The detector measures the interaction between the sample and the energy. The human senses provide marvelously sensitive and portable detectors such as the eyes and ears, and scientists have developed many detectors that extend our senses, such as photographic film, speedometers, thermometers, and radios. All of these detectors (and many others) record the interactions of matter and energy.

We are surrounded by countless materials, each with properties ideally suited to its function. You are reading a book with thin, white, flexible paper pages, fast-drying black ink, and strong glues for binding. You are probably sitting in a comfortable chair constructed from some combination of laminated woods, resilient plastics, light metal alloys, and synthetic fabrics. We wrote this book on PCs made of semiconductor integrated circuits, ultrathin LCD (liquid crystal display) screens, flexible copper wires, and colorful plastic insulation.