Atoms are restless — nearly every one of them is chasing a stable, full outer shell of electrons, and the different ways they achieve it, by giving, taking, or sharing, create every material in the world around you.
A lone chlorine atom and a lone sodium atom are both dangerously reactive — sodium bursts into flame on contact with water, and chlorine gas is toxic enough to have been used as a weapon. Bond them together, though, and they become table salt, something you sprinkle on your dinner without a second thought. That transformation, from reactive individuals to stable compounds, is what chemical bonding is all about.
Atoms bond because isolated atoms are usually unstable. The most stable electron arrangement, the one found naturally in the unreactive noble gases, is a full outer shell of eight electrons — a pattern chemists call the octet rule, first proposed by American chemist Gilbert N. Lewis in 1916. Almost every atom "wants" this same full outer shell, and chemical bonding is simply the different strategies atoms use to get there: giving electrons away, taking them from another atom, or sharing them.
The electrons that actually take part in bonding are called valence electrons — the electrons sitting in an atom's outermost shell. Conveniently, an element's group number on the periodic table (covered in the previous article) reveals how many valence electrons it has: elements in Group 1 have just one, Group 2 has two, and Group 18, the noble gases, already have a full eight (or two, for helium) and rarely bond with anything at all.
Sodium (Group 1) has just 1 valence electron — one more than a stable, full shell. Rather than gain seven electrons, it is far easier for sodium to simply lose that one electron and achieve stability that way.
Ionic bonding happens when one atom transfers an electron entirely to another. This typically occurs between a metal, which has few valence electrons and loses them easily, and a non-metal, which has nearly a full shell and readily accepts more. Losing or gaining electrons leaves both atoms electrically charged: the atom that lost an electron becomes a positively charged cation, and the atom that gained one becomes a negatively charged anion. Opposite charges attract, and that electrostatic attraction is the bond itself.
Sodium transfers its single valence electron to chlorine, forming a positive Na⁺ ion and a negative Cl⁻ ion — together, sodium chloride (table salt).
Ionic compounds don't form loose pairs; the resulting ions stack together in a rigid, repeating pattern called a crystal lattice. This is why ionic compounds like salt tend to be hard, brittle solids with notably high melting points — pulling the tightly bonded lattice apart takes a great deal of energy.
When two non-metal atoms meet, neither one easily gives up electrons — so instead, they compromise by sharing a pair of electrons between them. This is covalent bonding, and it's how molecules like water (H₂O), carbon dioxide (CO₂), and oxygen gas (O₂) are held together. Each shared pair counts toward both atoms' outer shell simultaneously, letting both approach a full, stable octet at once.
In a water molecule, the oxygen atom shares one electron pair with each of two hydrogen atoms. Oxygen ends up with a full octet (its own 6 valence electrons plus 2 shared electrons), while each hydrogen reaches its own stable pair.
Not all covalent bonds share electrons equally. Electronegativity measures how strongly an atom pulls on shared electrons; American chemist Linus Pauling developed the standard numerical scale for it in the 1930s, running from about 0.7 up to 4.0 for fluorine, the most electronegative element of all. When two bonded atoms have very similar electronegativity, electrons are shared evenly, forming a nonpolar covalent bond. When one atom pulls noticeably harder, the electrons spend more time near it, creating a polar covalent bond with a slightly negative end and a slightly positive end.
| Bond Type | Electronegativity Difference | Example |
|---|---|---|
| Nonpolar covalent | Small (roughly 0 to 0.4) | O₂, Cl₂ (same element, shared equally) |
| Polar covalent | Moderate (roughly 0.4 to 1.8) | H₂O, HCl |
| Ionic | Large (greater than about 1.8) | NaCl, MgO |
Water's polarity, caused by oxygen pulling harder on the shared electrons than hydrogen does, is exactly why water molecules cling to each other so well and why water dissolves so many other substances — a property explored further in a later article on mixtures and solutions.
Metal atoms take a completely different approach. Rather than pairing off, metal atoms release their valence electrons into a shared, freely moving pool often called a "sea of electrons," while the resulting positive metal ions arrange themselves in a regular lattice throughout that sea. This unusual arrangement explains several familiar properties of metals at once.
Because the electrons in metallic bonding are free to move rather than locked to one pair of atoms, metals like copper conduct electricity extremely well — the moving electrons themselves carry the electric current. That same freely shifting sea of electrons also lets metal ions slide past one another without breaking their bonds, which is why metals like gold or aluminium can be hammered into new shapes or drawn into wire instead of shattering, unlike a brittle ionic crystal.
| Bond Type | Formed Between | Electron Behaviour | Typical Properties |
|---|---|---|---|
| Ionic | Metal + Non-metal | Transferred completely | Hard, brittle, high melting point |
| Covalent | Non-metal + Non-metal | Shared between atoms | Often gases or liquids, lower melting point |
| Metallic | Metal + Metal | Delocalised in a shared pool | Conductive, malleable, shiny |
Every material you can touch is, at its core, a story about electrons finding stability — sodium giving one away, oxygen sharing a pair, copper releasing its electrons into a communal sea. The octet rule may sound like a small, abstract idea, but it is quietly responsible for why salt is brittle, why water is wet, and why a copper wire can carry electricity across a city. The next article in this series builds on these same bonds to explore what happens when substances react and rearrange entirely — chemical reactions and equations.
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