Chemical Bonding
Why atoms bond, the octet rule that predicts how, and how to classify a bond as ionic or covalent using electronegativity.
Almost everything you encounter — molecules, crystals, the compounds that make up living tissue — consists of atoms held together by chemical bonds. In fact, the only elements that exist stably as separate, unbonded atoms are the Group 18 noble gases. This subtopic covers why atoms bond, the octet rule that predicts how they bond, and how to classify a bond as ionic or covalent using electronegativity — with a close look at ionic bonds, crystal lattices, and lattice energy.
Key Takeaways
Atoms bond through valence electron interactions, most often following the octet rule (eight valence electrons, a noble-gas-like configuration).
Octet rule exceptions: incomplete octet (H, He, Li, Be, B — including electron-deficient molecules like BeCl₂ and BCl₃), expanded octet (period 3+ elements using d-orbitals, e.g., P, S, Cl), and odd-electron species (e.g., NO).
Electronegativity difference (ΔEN) classifies bond type: less than 0.5 is nonpolar covalent, 0.5–1.7 is polar covalent, and greater than 1.7 is ionic.
Ionic bonds form primarily between metals (which lose electrons to form cations) and nonmetals (which gain electrons to form anions); no electrons are shared.
Ionic compounds arrange into an ordered 3D crystal lattice (e.g., NaCl's 6:6 Na⁺/Cl⁻ coordination), giving them high melting and boiling points.
Lattice energy — the energy needed to separate one mole of ionic crystal into gas-phase ions — increases with greater ionic charge and smaller ionic radius, per Coulomb's law, with charge having the larger effect.
Why Atoms Bond: Valence Electrons and the Octet Rule
When atoms come together to form molecules, they're held together by chemical bonds formed through the interaction of their valence electrons — the outermost electrons available for bonding. The resulting compound's properties are usually very different from the properties of the individual atoms that formed it.
For most elements, this bonding follows the octet rule: an atom tends to bond with other atoms until it has eight electrons in its outermost shell, achieving a stable electron configuration like that of the noble gases. Argon, for example, already has a full octet of eight valence electrons and is stable on its own.
Exceptions to the Octet Rule
The octet rule is really more of a rule of thumb — there are enough exceptions that it's worth knowing them by name.
Incomplete octet. Some elements are stable with fewer than eight valence electrons:
Hydrogen and helium, stable with 2 electrons (matching helium's configuration)
Lithium, stable with 2
Beryllium, stable with 4
Boron, stable with 6
When beryllium or boron sits as the central atom in a molecule, that molecule is often electron deficient — it has fewer than eight electrons around the central atom. BeCl₂ and BCl₃ are classic examples.
Expanded octet. Any element in period 3 or below can hold more than eight electrons, because these elements have empty d-orbitals in their valence shell that can accommodate the extra electrons. Phosphorus can hold 10, sulfur 12, and chlorine 14, among others.
Odd numbers of electrons. A molecule with an odd total number of valence electrons can't distribute them so that every atom gets a full octet. Nitric oxide (NO) is the classic example — nitrogen's 5 valence electrons plus oxygen's 6 add up to 11, an odd number.
If you want a shortcut, the elements that almost always obey the octet rule cleanly are carbon, nitrogen, oxygen, fluorine, sodium, and magnesium. As a general pattern, nonmetals gain electrons and metals lose electrons to reach a complete octet.
Classifying Bonds: Electronegativity and ΔEN
Chemical bonds fall into two broad categories: ionic and covalent.
Ionic bonds are electrostatic attractions between oppositely charged particles, usually between a metal and a nonmetal, resulting from one atom transferring electrons to another.
Covalent bonds involve an electron pair shared between two atoms — typically nonmetals with relatively similar electronegativities. If the pair is shared equally, the bond is nonpolar covalent; if it's shared unequally, the bond is polar covalent.
The variable that determines which kind of bond forms is electronegativity — a measure of how strongly an atom attracts electrons in a bond. Every atom has a characteristic electronegativity value, and the difference between two bonding atoms' electronegativities (ΔEN) predicts the bond type:
ΔEN | Bond Type |
|---|---|
Less than 0.5 | Nonpolar covalent |
0.5 to 1.7 | Polar covalent |
Greater than 1.7 | Ionic |
Ionic Bonds
Ionic bonds form between atoms with significantly different electronegativities. The atom that loses electrons becomes a cation (more protons than electrons); the atom that gains electrons becomes an anion (more electrons than protons). No electrons are shared — they're fully transferred, and for that transfer to occur, ΔEN must exceed 1.7 on the Pauling scale.
You're not usually expected to memorize specific electronegativity values, just to recognize that ionic bonds generally form between a metal and a nonmetal. A useful way to remember it: Metals lose electrons to become caTions — posiTive ions. Nonmetals gain electrons to become aNions — Negative ions.
Alkali metals and alkaline earth metals (Groups 1 and 2) readily form ionic bonds with halogens (Group 17): the active metals hold their valence electrons loosely, while halogens are eager to gain one more electron to complete their octet. A few worked examples of ΔEN for common ionic compounds:
Cesium chloride: ΔEN ≈ 2.3
Potassium iodide: ΔEN ≈ 1.7
Sodium fluoride: ΔEN ≈ 3.1
The Ionic Crystal Lattice
In an ionic compound, oppositely charged ions are held together in an ordered, three-dimensional structure called a crystal lattice. Because the electrostatic forces holding the ions together are non-directional, each ion ends up surrounded by several ions of the opposite charge, arranged to maximize attraction and minimize repulsion.
Sodium chloride is the classic example: in a crystal of NaCl, each Na⁺ cation is surrounded by 6 Cl⁻ anions, and each Cl⁻ anion is surrounded by 6 Na⁺ cations. Because these electrostatic attractions are strong and extend in three dimensions throughout the whole crystal, ionic compounds tend to have very high melting and boiling points.
Lattice Energy and Coulomb's Law
One way to quantify the strength of an ionic bond is lattice energy — the energy required to separate one mole of a solid ionic crystal into gas-phase ions far removed from each other. A larger lattice energy means a stronger ionic bond.
Lattice energy can be predicted using Coulomb's law, which describes the electrostatic force of attraction (or repulsion) between two charged particles:
F = k(q₁q₂) / r²
where q₁ and q₂ are the charges of the two ions, r is the distance between their centers, and k is a proportionality constant. Coulomb's law implies that lattice energy increases as ionic charge increases and as ionic radius decreases. Both charge and size matter, but charge has the larger effect.
Worked example (original, for illustration): Compare NaCl and MgO. Sodium and chloride ions each carry a single charge (Na⁺, Cl⁻), while magnesium and oxide ions each carry a double charge (Mg²⁺, O²⁻). Since q₁q₂ appears in the numerator of Coulomb's law, MgO's charge product (2 × 2 = 4) is four times larger than NaCl's (1 × 1 = 1). Even though Mg²⁺ and O²⁻ are also somewhat smaller than Na⁺ and Cl⁻ — which would further increase lattice energy by shrinking r — charge alone already predicts that MgO has a substantially larger lattice energy than NaCl. This matches the real-world trend: MgO's melting point (about 2,800 °C) is dramatically higher than NaCl's (about 801 °C), consistent with a much stronger ionic bond.
Common MCAT Mistakes
Treating incomplete-octet elements as rule violations rather than named exceptions. H, He, Li, Be, and B are stable with fewer than eight valence electrons — BeCl₂ and BCl₃ aren't "broken" molecules, they're standard electron-deficient examples.
Misreading the ΔEN cutoffs. The boundaries are less than 0.5 (nonpolar covalent), 0.5 to 1.7 (polar covalent), and greater than 1.7 (ionic) — mixing up which side of 1.7 is ionic is a common slip.
Describing ionic bonds as "sharing" electrons. Ionic bonding is a full transfer of electrons from one atom to another, producing a cation and an anion — it's covalent bonds that involve a shared electron pair.
Assuming ionic size matters more than ionic charge for lattice energy. Per Coulomb's law, both charge and radius affect lattice energy, but charge (in the q₁q₂ numerator) has the larger effect — which is why MgO's lattice energy dwarfs NaCl's despite Mg²⁺ and O²⁻ not being dramatically smaller than Na⁺ and Cl⁻.
MCAT-Style Concept Check
Question: Based on Coulomb's law, which pair of ions would be predicted to form the ionic compound with the highest lattice energy?
A) Na⁺ and Cl⁻
B) K⁺ and Br⁻
C) Mg²⁺ and Cl⁻
D) Mg²⁺ and O²⁻
Answer: D
Explanation: Lattice energy scales with the charge product (q₁q₂) in Coulomb's law, and charge has a larger effect on lattice energy than ionic radius does. Mg²⁺ and O²⁻ each carry a charge of 2, giving a charge product of 4 — larger than any of the other pairs, which involve at least one singly-charged ion (charge product of 1 or 2). This is the same logic that makes MgO's lattice energy, and its melting point, dramatically higher than NaCl's.
FAQ
What determines whether a bond is ionic or covalent?
The electronegativity difference (ΔEN) between the two bonding atoms. A ΔEN less than 0.5 gives a nonpolar covalent bond, 0.5 to 1.7 gives a polar covalent bond, and greater than 1.7 gives an ionic bond, where electrons are fully transferred rather than shared.
What is the octet rule, and what are its main exceptions?
The octet rule says atoms tend to bond until they have eight valence electrons, matching a noble gas configuration. The main exceptions are incomplete octets (H, He, Li, Be, B), expanded octets (period 3 and below elements like P, S, and Cl, which use empty d-orbitals), and odd-electron species like NO.
Why do ionic compounds have such high melting and boiling points?
Ionic compounds form an ordered, three-dimensional crystal lattice where each ion is surrounded by several oppositely charged ions, held together by strong, non-directional electrostatic attractions extending throughout the whole crystal. Breaking apart that lattice takes a lot of energy, which shows up as high melting and boiling points.
What is lattice energy, and what affects it?
Lattice energy is the energy required to separate one mole of a solid ionic crystal into gas-phase ions far removed from each other. Per Coulomb's law, it increases with greater ionic charge and smaller ionic radius, with charge having the larger effect of the two.