When two or more atoms share or exchange valence electrons with each other, the resulting electromagnetic forces cause the atoms to stick together. This is called a chemical bond. There are three types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds.
- In ionic bonds, atoms give away or accept electrons on “permanent loan,” so to speak. As a result, each of the atoms becomes an ion—an atom with more or fewer electrons than protons. Atoms that give away electrons become positively charged ions; atoms that take extra electrons become negatively charged ions. The two types of ions stick together, because opposite charges attract. Ionic bonds usually form when metals near the left side of the periodic table (especially alkali and alkaline metals) combine with halogens or other non-metals near the right side. The reason is that elements near the left side are eager to get rid of their valence electrons, as explained on the previous page, whereas elements near the right side are eager to fill their outermost shells (except for the noble gases, which already have full outer shells).
Sodium (atomic number 11) is eager to give its valence electron away, and chlorine (atomic number 17) wants just one more. They’re a perfect match! When sodium reacts with chlorine, each sodium atom gives its valence electron to a chlorine atom. The sodium ions have a positive charge, while the chlorine ions have a negative charge, so they stick together to form a substance called sodium chloride—better known as table salt.
- In covalent bonds, electrons are shared between clusters of neighboring atoms, which are called molecules. The electrons sort of “orbit” two or more atoms at once, in a way.
Two hydrogen atoms will share their electrons with a single oxygen atom to form a molecule of water (H2O). Hydrogen atoms can also share electrons with each other to form pure hydrogen molecules (H2). Oxygen can do the same, forming oxygen molecules (O2). Oxygen atoms can also team up in groups of three to form a slightly unstable molecule known as ozone (O3).
- In metallic bonds, electrons are shared by many atoms, and can flow freely. Metallic bonds form between metal atoms, which—as you may recall from the discussion on the previous page—prefer to give away their valence electrons. Some non-metals (e.g. hydrogen) can also form metallic bonds at sufficiently high pressure. So, when a bunch of metal atoms get together, nobody wants to babysit the valence electrons! Those poor, unwanted electrons roam aimlessly among the metal atoms like homeless vagrants. That’s why metals tend to be good conductors of electricity.
Metals are a diverse group of elements. Not all of them are solid, tough, and shiny like the ones most familiar in everyday experiences. Many are soft enough to cut with a knife. Some are so highly reactive that they burst into flame when exposed to air or explode when exposed to water. And some are just downright cool. My personal favorite is gallium (atomic number 31), which has a melting point of 29.76 °C (about 86 °F). That’s higher than room temperature, but lower than the temperature of the human body. For this reason, gallium will melt in your hand!
Substances comprised of two or more elements bonded together are called chemical compounds (or simply compounds). If a substance contains different elements that are not bonded together, it is called a mixture. Metal alloys are regarded as mixtures rather than compounds, even though metallic bonds form between the various types of metal atoms.
In all three types of bonds, atoms are held together by electromagnetic forces. In an ionic bond, the ions stick together because they have opposite electromagnetic charges, and hence attract each other. (Remember Coulomb’s law
.) The sharing of electrons in metallic and covalent bonds works similarly. Think of it this way: the electrons spend part of the time orbiting one atom, and part of the time orbiting another. So, sometimes one atom will have more than its “fair share” of electrons, and sometimes it will have less. Either way, its charge will be opposite the charge of the other atoms, so it will be attracted to them.
Chemical compounds are represented by chemical formulas. There are different types of chemical formulas:
- Structural formulas show how atoms are arranged in a molecule. For example, here’s the structural formula for water:
- Molecular formulas show how many of each type of atom are contained in a molecule, without showing how they are arranged. When there is more than one atom of a given element, subscripts are used to indicate how many atoms of that element are in the molecule. For example, the molecular formula for water is H2O. The subscript “2” shows that there are two atoms of hydrogen in this molecule. No subscript is written after “O,” because there is only one atom of oxygen in the molecule.
- Empirical formulas show the proportions of each type of atom, without regard to how many atoms are in each molecule. The empirical formula for water is exactly the same as the molecular formula, but the two types of formulas differ for some compounds. For example, the molecular formula for hydrogen peroxide is H2O2, while the empirical formula is simply HO.
I won’t be using structural formulas or empirical formulas in this book, so in what follows I’ll simply use the word “formula” to mean molecular formula.
Elements closer to the left side of the periodic table are usually written first in chemical formulas. For example, hydrogen peroxide is written H2O2 rather than O2H2, because hydrogen is closer to the left side of the table than oxygen is. There are exceptions to this rule. The order is reversed for some polyatomic ions—charged clusters of atoms that form parts of larger molecules. For instance, the formula for sodium hydroxide is NaOH. The negatively-charged hydroxide ion is written “OH,” even though H appears to the left of O in the periodic table. Similarly, the formula for sodium chlorite is NaClO2. The chlorite ion is written “ClO2,” even though Cl appears to the right of O in the periodic table.