How Ionic and Covalent Bonds Hold Atoms Together

Electrons occupy regions around an atom's nucleus. The electrons farthest from the nucleus are called valence electrons, and they play the largest role in chemical bonding. Atoms often become more stable when their outer electron arrangement resembles that of a noble gas.

This idea is sometimes taught as the octet rule, because many atoms become stable with eight electrons in their outer level. It is a helpful starting model, but it is not a perfect rule for every atom and compound. Hydrogen is stable with two outer electrons, and several elements form structures that do not fit a simple set of eight.

Bonding is not driven by atoms thinking about neat electron diagrams. The arrangement forms when the combined system has lower energy and is more stable than the separated particles. The simple electron model helps you predict the result without pretending atoms hold committee meetings.

In ionic bonding, one atom transfers one or more electrons to another. The atom that loses electrons becomes a positively charged ion, while the atom that gains electrons becomes a negatively charged ion. Opposite charges attract, holding the ions together.

Sodium chloride is a common example. A sodium atom can lose one outer electron, and a chlorine atom can gain one. The resulting sodium and chloride ions attract in a repeating three-dimensional lattice. A crystal of salt is therefore not made from separate little sodium chloride pairs lined up politely. It is a large organised structure of alternating ions.

Ionic substances often have high melting points because many strong attractions must be overcome. When melted or dissolved in water, their ions can move and may carry electric current. In a solid crystal, the ions are fixed in place and cannot flow freely.

In covalent bonding, atoms share pairs of electrons. The shared electrons are attracted to more than one nucleus, helping hold the atoms together. Two hydrogen atoms can share a pair to form hydrogen gas, while oxygen atoms can share two pairs in an oxygen molecule.

Water contains covalent bonds between oxygen and hydrogen. The sharing is unequal because oxygen attracts the shared electrons more strongly. This gives different parts of the molecule slight charges and helps explain many of water's unusual properties.

Compare the basic patterns:

• •Ionic bonding transfers electrons.
• •Covalent bonding shares electrons.
• •Ionic bonds create charged ions.
• •Covalent bonds often create molecules.
• •Both arrangements can lower the system's energy.

Ionic and covalent bonding are best seen as two ends of a range. The key idea is electronegativity, which describes how strongly an atom attracts shared electrons. When two identical atoms bond, as in hydrogen gas, they share electrons equally. When the atoms differ a little, as oxygen and hydrogen do in water, the sharing becomes unequal and the bond is polar. When the difference is very large, as between sodium and chlorine, one atom effectively takes the electron and the bond is described as ionic. Many real compounds sit somewhere between these labels.

A common misconception is that boiling water breaks its bonds. It does not. The covalent bonds between oxygen and hydrogen inside each molecule stay intact, and steam is still made of water molecules. What boiling overcomes are the much weaker attractions between separate molecules. This explains why many covalent substances made of small molecules, such as oxygen or water, have low melting and boiling points compared with ionic compounds. Some covalent substances, such as diamond, are different. Their atoms form one giant network of strong bonds, so they melt only at extremely high temperatures.

The electron story lets you predict formulas instead of memorising them. Magnesium has two outer electrons and oxygen needs two more, so one magnesium gives both to one oxygen, forming Mg2+ and O2- ions and the formula MgO. Chlorine accepts only one electron, so magnesium needs two chlorine atoms, giving MgCl2. With aluminium, which forms 3+ ions, and oxygen, which forms 2- ions, you need a total charge of zero: two aluminium ions give +6 and three oxide ions give -6, so the formula is Al2O3.

For covalent molecules, count the bonds each atom usually forms. Hydrogen forms one, oxygen two, nitrogen three and carbon four. Carbon with four hydrogens gives methane, CH4. Nitrogen with three hydrogens gives ammonia, NH3, which still has one unshared pair of electrons on the nitrogen. In carbon dioxide, carbon cannot make four separate bonds to two oxygens, so it forms a double bond with each, written O=C=O. Each double bond uses two shared pairs. Checking that every atom has its usual number of bonds is a quick way to catch a wrong structure.

Ionic and covalent bonding are not the only ways atoms hold together. Metals use a third arrangement. In a piece of copper or iron, each atom releases one or more outer electrons into a shared pool that spreads through the whole solid. What remains is a regular lattice of positive metal ions surrounded by these delocalised electrons, which attract all the ions at once. No electron belongs to a particular pair of atoms. This is why a solid metal conducts electricity while a solid ionic crystal does not: the free electrons can drift through the metal when a voltage is applied.

The same model explains how materials respond to force. Hit a lump of rock salt with a hammer and it shatters, because shifting one layer of ions brings like charges side by side and they repel. Hit a piece of gold and it flattens, because layers of metal ions can slide over each other while the electron pool moves with them, so the bonding is never broken. Mixing in atoms of a different size, as in alloys such as bronze or steel, disturbs those neat layers and makes sliding harder, which is why many alloys are stronger than the pure metals they contain.

It is often said that sodium wants to lose an electron and chlorine wants to gain one, but the energy figures tell a different story. Removing the outer electron from a sodium atom costs about 496 kilojoules per mole. Adding an electron to a chlorine atom releases about 349 kilojoules per mole. On those two steps alone, making separate sodium and chloride ions from gas atoms costs energy overall, roughly 147 kilojoules per mole. If electron transfer were the whole story, salt would not form at all.

The payoff comes when the ions pack together. Bringing gaseous sodium and chloride ions into a crystal lattice releases around 787 kilojoules per mole, far more than the cost of transfer, and that is what makes the solid so stable. The attraction grows with the size of the charges and shrinks with the distance between ions. Magnesium oxide, built from small doubly charged ions, has a lattice energy of nearly 3,800 kilojoules per mole, which easily pays for removing magnesium's second electron. This is why magnesium oxide melts at around 2,800 degrees Celsius, compared with 801 degrees for sodium chloride.

Chemical bonds form through the behaviour of outer electrons. Ionic bonding creates attracting positive and negative ions, while covalent bonding holds atoms together through shared electrons. Learn the electron story first, and formulas begin to describe a pattern instead of a pile of symbols.