What holds atoms together? Ionic, covalent and metallic bonds

Why atoms bond

Atoms bond because it lets them reach a more stable arrangement of outer electrons, often a full outer shell like that of a noble gas. Only the outer electrons, the valence electrons, take part, and atoms can give them away, take them or share them. Each choice leads to a different kind of bond.

Ionic bonds: giving and taking electrons

An ionic bond usually forms between a metal and a nonmetal. The metal atom loses electrons and becomes a positive ion, and the nonmetal atom gains them and becomes a negative ion. Opposite charges attract strongly, and that attraction is the ionic bond.

Table salt is the classic example. A sodium atom gives its one outer electron to a chlorine atom, forming Na⁺ and Cl⁻. A salt crystal contains no separate NaCl pairs. Instead, each Na⁺ is surrounded by six Cl⁻ ions and each Cl⁻ by six Na⁺ ions, in a repeating cube-shaped pattern called a crystal lattice.

Ionic compounds are usually hard, brittle solids with high melting points: salt melts at 801 °C. They do not conduct electricity as solids, but they do when melted or dissolved in water, because the ions are then free to move.

Covalent bonds: sharing electrons

A covalent bond usually forms between two nonmetal atoms. Neither atom pulls hard enough to take electrons from the other, so they share a pair of electrons instead. The shared pair is attracted to both nuclei, which holds the atoms together. A group of atoms joined this way is called a molecule.

In a water molecule, H₂O, the oxygen atom shares one pair of electrons with each of the two hydrogen atoms. Oxygen started with 6 outer electrons and now has 8 around it, and each hydrogen now has 2, a full first shell. Atoms can also share two pairs, making a double bond as in O₂, or three pairs, making a triple bond as in N₂.

Many covalent substances have low melting and boiling points, so they are often gases or liquids at room temperature; water boils at 100 °C. A few, such as diamond, form giant networks of covalent bonds and are extremely hard.

Metallic bonds: a sea of electrons

In a metal, the atoms release their outer electrons into a shared pool that moves freely through the whole metal. What remains is a regular arrangement of positive metal ions, held together by this “sea” of electrons. The attraction between the ions and the moving electrons is the metallic bond.

The free electrons carry electric current, which is why copper is used for wires, and they carry heat, which is why a metal spoon in hot soup warms up quickly. Because the bonds are not fixed between particular atoms, layers of ions can slide past one another, so metals bend and can be hammered into sheets instead of shattering.

Using electronegativity to predict the bond

Electronegativity measures how strongly an atom pulls on shared electrons. To compare two atoms, subtract the smaller value from the larger one. As a rough guide, a difference below about 0.4 means the electrons are shared almost equally (nonpolar covalent), a difference between about 0.4 and 1.7 means they are shared unequally (polar covalent), and a difference above about 1.7 usually means an ionic bond.

Sodium chloride: 3.16 − 0.93 = 2.23, so the bond is ionic. Water: oxygen (3.44) and hydrogen (2.20) differ by 1.24, so the O–H bonds are polar covalent. Oxygen pulls the shared electrons closer and carries a small negative charge, while the hydrogen atoms are slightly positive. Chlorine gas, Cl₂: 3.16 − 3.16 = 0, so the electrons are shared perfectly equally.

These cutoffs are rough guidelines, and textbooks use slightly different numbers, because bonding is really a sliding scale. A shortcut works most of the time: metal + nonmetal gives ionic, nonmetal + nonmetal gives covalent, and metal + metal gives metallic bonds.