Ionic, Covalent & Metallic Bonding

Three ways atoms stick together — give electrons away, share them, or pool them — and why the choice decides everything about a material.

High schoolIntro Gen ChemUni Year 1
⏱️ About 20 min

Table salt is brittle and melts near 800 °C. Copper wire bends without breaking and carries electricity. Water is a soft liquid at room temperature. Same periodic table, wildly different behaviour — and the reason traces back to one question: what did the atoms do with their outer electrons when they joined?

💡
The big idea: Atoms bond to reach a more stable, lower-energy electron arrangement — usually a full outer shell. There are three ways to get there: TRANSFER electrons (ionic), SHARE them (covalent), or POOL them into a shared sea (metallic). Which one happens depends on the atoms involved.
🎯 By the end, you'll be able to
  • Explain why atoms bond in terms of reaching a stable electron arrangement
  • Distinguish ionic (transfer), covalent (sharing) and metallic (electron sea) bonding
  • Predict the likely bond type from where the elements sit (metal vs nonmetal)
  • Explain why ionic compounds form giant lattices rather than discrete molecules
📎 Helpful to know first

Why atoms bond at all

Left alone, most atoms are restless: their outer (valence) shell isn't full, and a full outer shell — the arrangement of a noble gas — is a low-energy, stable place to be. Atoms bond because joining up lets them reach that stability. Nature rolls downhill in energy, and bonding is the hill.

There are three ways atoms can rearrange their valence electrons to get there, and each produces a completely different kind of substance.

🔑 The three bond types in one line each
Ionic — one atom gives electrons to another (metal → nonmetal). Covalent — two atoms share a pair of electrons (nonmetal + nonmetal). Metallic — many atoms pool their electrons into a shared 'sea' (metal + metal).

Ionic bonding: a full electron transfer

When a metal meets a nonmetal, the metal (which holds its outer electrons loosely) hands them over to the nonmetal (which grabs them tightly). Sodium gives up its single 3s electron; chlorine accepts it to complete its shell. Both end up with a noble-gas arrangement.

After the transfer, the atoms are no longer neutral — they are ions. Sodium becomes Na⁺ (it lost a negative electron), chlorine becomes Cl⁻ (it gained one). Opposite charges attract, and that electrostatic pull is the ionic bond.

e⁻ e⁻ n=1 (2 e⁻) e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ n=2 (8 e⁻) e⁻ n=3 (1 e⁻) 11 p⁺ 12 n Sodium · Na · Z=11 · mass number 23
Sodium has one lone electron in its outer shell (2, 8, 1). Losing it leaves a full shell underneath — and a Na⁺ ion. Generated from Z = 11.
e⁻ e⁻ n=1 (2 e⁻) e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ n=2 (8 e⁻) e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ e⁻ n=3 (7 e⁻) 17 p⁺ 18 n Chlorine · Cl · Z=17 · mass number 35
Chlorine's outer shell holds 7 electrons (2, 8, 7) — one short of full. Accepting sodium's electron completes it, forming Cl⁻. Generated from Z = 17.
\[ \ce{Na + Cl -> Na+ + Cl^- } \]
Sodium transfers one electron to chlorine. The resulting Na⁺ and Cl⁻ ions attract one another.
⚠️ The #1 misconception: 'NaCl molecules'
There is no such thing as a single 'NaCl molecule' in solid salt. Each Na⁺ is surrounded by six Cl⁻, and each Cl⁻ by six Na⁺, repeating in a giant 3-D lattice of billions of ions. The formula NaCl just gives the ratio (1:1), not a discrete particle. This is why ionic solids are hard, brittle, and melt at high temperatures — you'd have to pull an entire lattice of attractions apart.

Covalent bonding: sharing a pair

Two nonmetals both want to gain electrons, so neither will simply hand them over. Instead they share: each contributes an electron to a shared pair that 'counts' toward both atoms' shells. That shared pair, held between the two nuclei, is a covalent bond.

Two hydrogen atoms share a pair to make H₂; sharing two pairs makes a double bond, three makes a triple. Because the sharing binds a fixed, small set of atoms, covalent bonding produces discrete molecules (H₂O, CO₂, O₂) — unlike the endless ionic lattice.

\[ \ce{H\bond{-}H}\qquad \ce{O=C=O}\qquad \ce{N#N} \]
Single, double and triple covalent bonds: one, two and three shared electron pairs between the atoms.

Metallic bonding: a sea of electrons

In a chunk of metal, atoms are packed together and each gives up its valence electrons to a shared pool. The result is a lattice of positive metal ions sitting in a mobile 'sea' of delocalised electrons that are free to drift.

That free-flowing sea explains a metal's signature properties: it conducts electricity and heat (the electrons carry charge and energy), it is malleable (layers of ions can slide past each other without snapping bonds), and it is shiny.

✨ It's a spectrum, not three boxes
Bond type really tracks the electronegativity difference between the atoms. A large difference (metal + nonmetal) pulls electrons all the way over → ionic. A tiny or zero difference (two similar nonmetals) shares them evenly → covalent. Metals, which barely hold their electrons, pool them → metallic. Real bonds sit on a continuum between these ideals.
📝 Worked example: Sodium (a metal) reacts with chlorine (a nonmetal). Predict the bond type and describe what happens to the electrons.
  1. Identify the elements: sodium is a metal, chlorine is a nonmetal — a metal + nonmetal pairing.
  2. Metal + nonmetal means a large electronegativity gap, so electrons transfer rather than share: this is ionic.
  3. Sodium loses its one outer electron to become Na⁺; chlorine gains it to become Cl⁻.
  4. The opposite charges attract and lock into a repeating lattice — not discrete NaCl molecules.
✓ Ionic bonding: sodium transfers one electron to chlorine, giving Na⁺ and Cl⁻ ions that build a giant lattice.
✏️ Practice: Calcium is in group 2, so a calcium atom has 2 valence electrons. To reach a full outer shell by transfer (forming Ca²⁺), how many electrons does each calcium atom give away?
electrons
Solution
  1. Group 2 metals have 2 valence electrons.
  2. Emptying that outer shell exposes the full shell beneath it — the stable arrangement.
  3. So calcium gives away 2 electrons, becoming Ca²⁺ (e.g. it forms CaCl₂ with two Cl⁻).

Check your understanding

1. A crystal of table salt (NaCl) is best described as…
Ionic compounds are giant lattices, not molecules. Every ion is surrounded by ions of opposite charge; 'NaCl' is only the 1:1 ratio, not a discrete particle.
2. Which property is explained by metallic bonding's 'sea of electrons'?
Delocalised electrons are free to move, so they carry electric current — and let layers of ions slide, which is why metals are also malleable.
3. Which pairing of elements most likely forms a covalent bond?
Two nonmetals both attract electrons strongly, so neither transfers — they share pairs, which is covalent. Metal + nonmetal tends to be ionic; metal + metal is metallic.
🔎 Related lessons
Next: Lewis Structures VSEPR & Molecular Geometry Electron Configuration
✅ Key takeaways
  • Atoms bond to reach a stable, lower-energy electron arrangement (usually a full outer shell).
  • Ionic = electron transfer (metal + nonmetal) → charged ions in a giant lattice, not molecules.
  • Covalent = sharing electron pairs (nonmetal + nonmetal) → discrete molecules.
  • Metallic = positive ions in a shared 'sea' of delocalised electrons → conductive, malleable.
  • Which bond forms tracks the electronegativity difference between the atoms.
➡️ Covalent molecules share electrons — but exactly how many pairs, and where the leftover electrons sit, controls a molecule's shape and reactivity. Next, we learn to map every valence electron with Lewis structures.
Want to test yourself on this? Try the Chemistry practice test →