Three ways to satisfy the octet
In the previous guide you met the atom — a tiny nucleus wrapped in shells of electrons, with the outer valence electrons doing all the chemistry — and you traced the bonding-energy curve: attraction and repulsion fighting to a truce at an equilibrium spacing, the bottom of a well whose depth is the bond energy. That curve tells you THAT two atoms bind and how strongly. This guide answers the next question — by what MEANS do they bind, and how does that choice quietly pick the whole crystal structure.
Almost every atom behaves as if it wants a filled outer shell — the octet tendency, roughly eight electrons outside (last guide's rule of thumb). There are exactly three ways for atoms to rearrange their valence electrons to get there, and each is a distinct primary bond: give electrons away or take them (ionic), share them in a fixed direction (covalent), or pool them all into a common cloud (metallic). Which one nature picks is set almost entirely by one number from the last guide — electronegativity, each atom's pull on shared electrons.
The ionic bond: transfer, then pack by charge and size
When one atom pulls far harder than the other — a big electronegativity gap, typically a metal meeting a nonmetal, like sodium meeting chlorine — the greedy atom simply takes the loose electron. Sodium loses its lone valence electron to become Na+, chlorine gains one to become Cl-, and now two oppositely charged ions attract. That is the ionic bond: not sharing at all, but a completed transfer followed by plain electrostatic attraction.
The crucial structural fact is that an ion's electric field is the same in every direction — a charged sphere pulls equally north, south, east, and west. So the ionic bond is non-directional: it does not care about angles, only about two things. First, charge must balance (equal plus and minus, so the crystal comes out neutral). Second, the ions must fit — a small cation nestled among larger anions, packed as tightly as their sizes allow. Structure here is a geometry puzzle of stacking charged spheres, governed by ionic size (the ionic radius from earlier).
Take common salt. In the rock-salt structure the larger Cl- ions form a face-centred array and every Na+ drops into an octahedral hole — a gap surrounded by six chlorides. So each Na+ touches six Cl- and vice versa: a coordination number of 6. Why six and not more? The radius ratio. Na+ is about 1.02 angstrom, Cl- about 1.81 angstrom, a ratio near 0.56 — exactly the window where six neighbours fit snugly but eight would rattle around a too-small cation. Size, not angle, chooses the structure — a theme guide 5 develops in full.
The covalent bond: sharing in a fixed direction
When two atoms pull on electrons about equally hard — two nonmetals of similar, high electronegativity, like two carbons — neither can win the tug-of-war, so they compromise: they SHARE a pair of electrons, one contributed by each, sitting in the space between the two nuclei. Both atoms count that shared pair toward their octet. This is the covalent bond, and its defining feature is completely different from the ionic bond: it points.
A shared electron pair lives in a lobe of space pointing from one atom toward its partner, so a covalent bond has a definite direction and a definite angle to its neighbours. Carbon forms four bonds aimed at the corners of a tetrahedron, locked at 109.5 degrees. You cannot squeeze more neighbours in just by pushing harder — the angles are fixed by the electron geometry. This is why covalent solids feel 'stubborn': the bonds behave like rigid struts of set length and set angle.
Follow those rigid struts and you get an open, low-coordination structure. In the diamond cubic structure every carbon bonds to just four others at tetrahedral angles, giving a coordination number of only 4 and a wide-open framework — its atomic packing factor is a mere 0.34, filling barely a third of space, against 0.74 for a dense metal. All that empty room is the price of insisting on angles. The reward is ferocious strength and stiffness: diamond is the hardest natural material precisely because those directional bonds resist being bent or sheared.
The metallic bond: an electron sea
The third route appears when every atom present is a weak holder of electrons — metals, with low electronegativity all round. No atom can grab the electrons and none can pin a shared pair, so the valence electrons simply let go of their home atoms and roam freely through the whole solid. What remains is a regular array of positive ion cores bathed in a common electron sea — the metallic bond. The glue is the attraction between that negative sea and all the positive cores at once.
Because the sea surrounds every core evenly, the metallic bond — like the ionic one — is non-directional; it too has no favoured angle. But unlike ions, metal atoms carry no charge-balance constraint and are all essentially the same size, so nothing stops them from packing as densely as identical spheres possibly can. The winning move is close packing: stack layers of spheres the way a grocer stacks oranges into a pyramid.
Copper is a textbook case. Its atoms stack in the face-centred cubic pattern, layers repeating ABCABC, and each atom touches twelve neighbours — a coordination number of 12, the maximum possible for equal spheres, filling 74 percent of space (packing factor 0.74). Non-directional bonding buys density: high coordination and tight packing. The mobile electron sea also explains two everyday facts — metals conduct electricity, and they bend instead of shattering, because the sea flows and re-forms as the ion cores slide past one another.
Directionality is the master idea
bond electrons directional? coordination packing example -------- ---------------- ------------ ------------ --------- ----------- ionic transferred +/- no ~ 6 to 8 medium NaCl (6) covalent shared, aimed YES (angles) low ~ 4 open 0.34 diamond (4) metallic pooled 'sea' no high ~ 8-12 dense 0.74 copper (12)
Step back and one pattern organises everything. Compare the two non-directional bonds (ionic, metallic) against the one directional bond (covalent) and the rule leaps out: bond directionality decides whether a structure is open or dense. Directional covalent bonds fix angles and force open, low-coordination frameworks like diamond's 4. Non-directional ionic and metallic bonds care only about packing spheres, giving high coordination (6, 8, 12) and dense structures. This single idea — angles versus packing — is the hinge of guide 4, coming next.
No bond is pure — reading the bond
Real materials rarely sit at a pure corner. Bonding is a spectrum, and most compounds show mixed bonding. Zinc blende (ZnS) shares electrons like a covalent solid yet with a partly-ionic tilt, because zinc and sulphur differ in electronegativity; quartz (SiO2) is famously part covalent, part ionic. And directional covalent sheets can be stacked by weak, non-directional forces: graphite's carbon layers are covalently bonded within each sheet but held to one another only by feeble secondary bonds — which is why pencil 'lead' slides off in flakes. Those weak stackers are the whole subject of the next guide.
- Find each atom's electronegativity — the dial you built in guide 1.
- Big gap (metal meets nonmetal, e.g. Na and Cl): electrons transfer, so expect an ionic bond — non-directional, packed by size and charge balance.
- Both high and similar (two nonmetals, e.g. C and C, or Si and O): electrons are shared, so expect a covalent bond — directional, open, low coordination.
- Both low (metal atoms all round): electrons pool into a sea, so expect a metallic bond — non-directional, dense, high coordination.
- In between: expect mixed bonding — read the material as a blend, not a pure type, and don't be surprised by exceptions.