Chemical Bonding: Ionic, Covalent, and Metallic Bonds Explained

Why Bonds Form

Atoms form chemical bonds because doing so lowers their potential energy. A sodium atom and a chlorine atom sitting far apart hold more total energy than the sodium chloride ion pair they become when they bond. That drop in energy is what drives bond formation and what makes compounds more stable than their isolated constituent atoms.

Every chemical bond is fundamentally electrostatic — opposite charges attract. What differs between bond types is how electrons are distributed between the atoms involved.

There are three main categories: ionic bonds (electrons transferred), covalent bonds (electrons shared), and metallic bonds (electrons delocalized across a lattice). Intermolecular forces are a separate, weaker class that acts between molecules rather than within them — more on those later.

Ionic Bonds

Ionic bonds form when one atom transfers one or more electrons to another atom. This happens most readily between metals (which release electrons easily) and nonmetals (which accept them readily).

How Transfer Works

Sodium (Na) has one valence electron and a low ionization energy — it gives that electron up without much resistance. Chlorine (Cl) has seven valence electrons and a high electron affinity — it strongly attracts one more to complete its octet. When sodium transfers its valence electron to chlorine:

Sodium becomes Na⁺ (lost an electron, now has more protons than electrons)
Chlorine becomes Cl⁻ (gained an electron, now has more electrons than protons)

The oppositely charged ions attract each other. In a bulk crystal of NaCl, each Na⁺ is surrounded by six Cl⁻ ions and vice versa, forming a three-dimensional lattice. The energy released when gaseous ions come together to form this lattice is called lattice energy. NaCl has a lattice energy of about −787 kJ/mol — a large negative value that reflects how stable the crystal is.

Properties of Ionic Compounds

High melting and boiling points. Breaking an ionic lattice requires overcoming many strong electrostatic attractions simultaneously. NaCl melts at 801 °C.
Brittle. Shift the layers of a crystal slightly and like charges align — the repulsion shatters it.
Conduct electricity only when dissolved or melted. In solid form the ions are locked in place. Dissolved in water or melted, the ions are free to move and carry charge.
Soluble in polar solvents. Water molecules can stabilize individual ions through ion-dipole interactions, pulling the lattice apart.

Covalent Bonds

Covalent bonds form when two atoms share one or more pairs of electrons rather than transferring them. This occurs most commonly between nonmetals.

Sharing Electrons

In a hydrogen molecule (H₂), neither hydrogen atom has a stronger pull on electrons than the other. Each atom contributes one electron to a shared pair. Both nuclei are attracted to the same electron pair sitting between them — that shared attraction is the bond.

Atoms can share more than one pair:

Single bond — one shared pair (H–H, C–H)
Double bond — two shared pairs (O=O, C=O)
Triple bond — three shared pairs (N≡N, C≡C)

More shared pairs means a shorter, stronger bond.

Polar vs. Nonpolar Covalent Bonds

When two identical atoms share electrons (H₂, N₂, Cl₂), the sharing is perfectly equal — the bond is nonpolar covalent.

When two different atoms share electrons, the one with higher electronegativity pulls the shared electrons closer to itself. This unequal sharing creates a partial negative charge (δ−) on the more electronegative atom and a partial positive charge (δ+) on the other. The bond is polar covalent.

In water (H₂O), oxygen (electronegativity 3.44) pulls the shared electrons much more strongly than hydrogen (electronegativity 2.20). Each O–H bond is polar, with oxygen carrying partial negative charge and each hydrogen carrying partial positive charge.

Properties of Covalent Compounds

Variable melting points. Small nonpolar molecules like H₂ have very low boiling points (−253 °C). Large network covalent solids like diamond have extremely high melting points (>3500 °C). The range is wide because it depends on intermolecular forces, not bond strength directly.
Poor electrical conductors. No free charged particles — electrons are localized in bonds.
Often soluble in nonpolar solvents. "Like dissolves like."

Electronegativity Difference as a Predictor

Electronegativity (EN) measures how strongly an atom in a bond attracts electrons toward itself. The Pauling scale runs from about 0.7 (cesium) to 4.0 (fluorine).

The difference in electronegativity between two bonded atoms (ΔEN) is a practical guide to bond type:

| ΔEN | Bond Type | |-----|-----------| | < 0.5 | Nonpolar covalent | | 0.5 – 1.7 | Polar covalent | | > 1.7 | Ionic |

For NaCl: EN(Cl) = 3.16, EN(Na) = 0.93, ΔEN = 2.23 → ionic. For HCl: EN(Cl) = 3.16, EN(H) = 2.20, ΔEN = 0.96 → polar covalent. For Cl₂: ΔEN = 0 → nonpolar covalent.

Treat these cutoffs as guidelines, not hard lines. Bond character is a continuum — as ΔEN increases, the bond shifts gradually from perfectly equal sharing toward complete electron transfer. The 1.7 threshold is a rule of thumb, not a law of physics.

Metallic Bonds

In metals, valence electrons are not attached to any single atom. They detach and move freely throughout the entire solid, forming what chemists call an electron sea. The metal cations (positive nuclei) are held in a lattice, surrounded and stabilized by this mobile cloud of electrons.

Why This Model Explains Metal Properties

Electrical conductivity. Delocalized electrons flow easily under an applied voltage — no bonds need to break.
Thermal conductivity. Those same mobile electrons transfer kinetic energy quickly across the material.
Malleability and ductility. When stress deforms the lattice, the electron sea adjusts around the new arrangement. Layers slide past each other without the repulsion that shatters ionic crystals.
Metallic luster. Free electrons absorb and re-emit light across many frequencies, giving metals their characteristic shine.

Intermolecular Forces (Brief)

Intermolecular forces (IMFs) act between molecules, not within them. They are weaker than covalent or ionic bonds but determine physical properties like boiling point and solubility.

London dispersion forces — present in all molecules. Temporary fluctuations in electron density create instantaneous dipoles that induce dipoles in neighboring molecules. Larger molecules with more electrons have stronger dispersion forces.
Dipole-dipole interactions — occur between polar molecules. The δ+ end of one molecule attracts the δ− end of another.
Hydrogen bonds — a special, stronger case of dipole-dipole. Occur when H is bonded directly to N, O, or F, and that H interacts with a lone pair on an N, O, or F atom of a neighboring molecule. Hydrogen bonds explain why water has an unusually high boiling point for its molecular weight.

Do not confuse IMFs with the intramolecular bonds (ionic, covalent, metallic) that hold atoms together within a substance. Breaking IMFs accounts for phase changes; breaking intramolecular bonds accounts for chemical reactions.

Bond Polarity and Molecular Shape

A polar bond does not automatically produce a polar molecule. Molecular shape determines whether individual bond dipoles cancel or reinforce.

VSEPR theory (Valence Shell Electron Pair Repulsion) predicts shape: electron pairs around a central atom repel each other and spread as far apart as possible.

Consider two molecules with polar bonds:

CO₂ — carbon forms two double bonds with oxygen. ΔEN ≈ 1.0 per bond, so each C=O bond is polar. But CO₂ is linear; the two bond dipoles point in exactly opposite directions and cancel. Net dipole moment = 0. CO₂ is a nonpolar molecule.
H₂O — oxygen forms two single bonds with hydrogen and has two lone pairs. VSEPR gives water a bent shape (≈104.5° bond angle). The two O–H dipoles point in different directions but do not cancel — they add together. Water has a significant net dipole moment and is a polar molecule.

This distinction matters: polar molecules have stronger IMFs, higher boiling points, and greater solubility in water than nonpolar molecules of similar size.

Summary Comparison

| Property | Ionic | Covalent (Molecular) | Metallic | |---|---|---|---| | Electron behavior | Transferred | Shared | Delocalized (sea) | | Typical elements | Metal + nonmetal | Nonmetals | Metals | | Melting point | High | Low to very high | Moderate to high | | Electrical conductivity | Only when molten/dissolved | Poor | Excellent | | Mechanical properties | Brittle | Varies | Malleable, ductile | | Example | NaCl, MgO | H₂O, CO₂, glucose | Fe, Cu, Au |

Bond type depends on the identities and electronegativities of the atoms involved. Knowing ΔEN, the electron sea model, and VSEPR gives you the tools to predict not just what type of bond forms, but what properties the resulting substance will have.