What holds atoms together: ionic, covalent and metallic bonds

One rule for every bond

Every chemical bond is electrical, an attraction between opposite charges. The rule behind all of them fits in one line: a bond forms when electrons can be close to two or more nuclei at the same time. In the simplest case, the H₂⁺ ion, two hydrogen nuclei share one electron. Between the two protons, that electron attracts both towards itself, and so towards each other.

Off to one side, the electron pulls the nearer proton much harder, because Coulomb’s law weakens with the square of the distance, and it helps push the nuclei apart. Two protons with no electron repel at every distance and cannot form a molecule. Three models rest on this rule: ionic, covalent and metallic bonding, each useful within its limits.

An atom’s outermost-shell electrons are the ones that meet other atoms. Atoms in Groups 1, 2 and 13 usually give them up, forming ions such as Na⁺, Mg²⁺ and Al³⁺. Groups 15 to 17 tend to take electrons, forming ions such as Cl⁻, while Group 14 shares instead. Sodium’s shells hold 2, 8 and 1 electrons, chlorine’s 2, 8 and 7. Sodium’s figure draws the neutral atom, whose single outer electron is the one that leaves when sodium becomes an ion.

Ionic: the electron handed over, the ions in a lattice

The ionic idea began with salt solutions, which conduct electricity: opposite charges attract, so perhaps the solids behind them were built from charged atoms. In the ionic model, sharing is so unequal that one atom loses an electron completely to the other, leaving a pair of ions. Sodium loses its single outer electron as Na⁺; chlorine gains one as Cl⁻.

The model fits best the compounds of Group 1 and 2 metals with a halogen or oxygen. There the two elements differ so much in their pull on electrons that attraction between opposite charges usually dominates.

Each ion is attracted by opposite neighbours and repelled by like ones from all directions, which fixes it in place. With forces that have no direction, geometry alone sets the structure. In sodium chloride each ion has six neighbours of the opposite charge: (6,6) coordination.

Sodium chloride contains no discrete NaCl units. It is an extended solid, written only by its simplest formula: one Na⁺ with one Cl⁻ makes a formula unit, a ratio kept for bookkeeping, not a particle. Alkali halides form separate molecules only in the rarefied gas phase. Sodium chloride’s drawing shows one ion of each kind and no bond between them.

That three-dimensional structure puts more electrons closer to more nuclei than the same number of separate gaseous atoms, which is why the solid is the more stable of the two. Sodium chloride melts at 801 °C. Its fixed ions make the solid an electrical insulator; in the melt or a solution the ions themselves carry the current.

Covalent: a pair shared, and a molecule with edges

In 1916 G. N. Lewis proposed that two atoms sharing a pair of electrons attract each other through it, and that this is a bond. One shared pair holds the two nuclei of H₂ only 74 pm apart. In hydrogen chloride, hydrogen’s single electron pairs with one of chlorine’s seven outer electrons, one electron from each atom. Chlorine, which takes an electron from sodium in salt, shares one with hydrogen here instead. Hydrogen chloride’s viewer shows two atoms and one bond, and that one stick stands for the single shared pair.

A true molecule exists as an independent particle, while an extended solid can be written only as a formula. Methane, CH₄, is a molecule of that kind. Its central atom is carbon, from Group 14, which shares electrons rather than forming ions. A methane molecule is a separate particle of five atoms; in sodium chloride no group of ions stands apart from its neighbours as a unit.

A molecule’s atoms also sit at definite angles to one another. How that shape arises is the work of the valence shell electron pair repulsion model, VSEPR, set out in the VSEPR article.

Between the two: polar bonds and the continuum

Two identical atoms share their electrons exactly evenly, as in H₂, Cl₂ and O₂. Different atoms do not. Differences in size and nuclear charge make one atom pull harder on the bonding pair, and the shared electrons shift towards it, giving the bond a polarity. That pull is electronegativity, how strongly an atom in a bond draws electrons, compared with other atoms. An atom with a small one is electropositive.

On Pauling’s scale fluorine holds the top value, 3.98, and caesium sits at the low end, 0.79; lithium stands at 0.98 and sodium at 0.93. A higher nuclear charge and a smaller atom attract an outer electron more strongly.

In hydrogen chloride, hydrogen at 2.2 meets chlorine at 3.16, and the electronegativity rule puts the pair nearer chlorine. By that rule the bond is polar, with partial ionic character, which the viewer cannot show. The greater the electronegativity difference between two atoms, the more polar their bond.

Even in the most ionic compounds, both electrons stay close to both nuclei, and no bond is purely ionic. Lithium fluoride, which Chem1 calls the most ionic of the alkali halides, shows it. The electron lithium hands to fluorine does not simply leave. It comes within fluorine’s grasp while staying close to lithium’s nucleus as well, so calling it “lost” is only half true. The name “ionic bond” still serves for LiF and NaCl.

Filing every bond under ionic, covalent or metallic oversimplifies grossly, since many substances show some of all three; ionic and covalent are the two ends of one scale. A polar bond and a polar molecule are different things, and water and carbon dioxide are the cases that separate them, taken up in polarity.

Metals: ions in an electron fluid

Metal atoms show no tendency to form negative ions, which rules out the bonding of an ionic solid. They also never have enough outer-shell electrons to place eight around each atom. These two facts point to a third model.

In the third model, the outer electrons lose their tie to single atoms and form a mobile electron fluid. The metal’s strength comes from the attraction between its lattice of positive ions and that fluid. An applied electric field shifts the fluid, which is how a metal conducts. Chem1 ranks silver first among metals for conduction, then copper, gold and aluminium.

The fluid surrounds every ion on all sides, so metallic bonding has no preferred direction. That makes a metal malleable: it can be hammered into shape without breaking. An ionic crystal behaves otherwise: with enough force one layer of ions slides along the next, the slip spreads through a plane of the crystal, and the crystal splits. The electron-fluid picture is itself an oversimplification, and it fails to explain metals quantitatively.