Difficulty: Introductory | Prerequisites: Basic atomic structure, periodic table layout, electron configurations
This topic covers why and how atoms form bonds, which is the foundation for nearly everything else in chemistry. You need to understand what holds molecules together before you can make sense of reaction mechanisms, thermodynamics, or molecular biology. The material assumes you already know what valence electrons are, how to read the periodic table for group and period information, and what electron configurations look like. If those feel shaky, revisit atomic structure notes first.
Atoms bond to reach a stable electron arrangement, usually eight valence electrons (the octet rule). Ionic bonds transfer electrons between metals and nonmetals; covalent bonds share electrons between nonmetals. Whether a covalent bond is polar or nonpolar depends on the electronegativity difference between the two atoms.
Octet rule
The tendency of atoms to bond so that each has eight electrons in its outermost shell, achieving a stable noble-gas configuration. In simple terms, atoms "want" a full outer shell of eight electrons, so they share or transfer electrons to get there.
Ionic bond
A bond formed when one or more electrons are transferred from an atom with low ionisation energy (typically a metal) to an atom with high electron affinity (typically a nonmetal), creating oppositely charged ions held together by electrostatic attraction. Think of it as one atom handing its electron to another, and the resulting positive and negative charges gluing them together.
Covalent bond
A bond in which two atoms (typically nonmetals) share one or more pairs of electrons. In simple terms, instead of giving electrons away, the atoms share them so both can fill their outer shells.
Electronegativity
A measure of how strongly an atom attracts the shared electrons in a bond. Think of it as each atom's "pull" on the electron pair. The bigger the difference in pull between two atoms, the more polar the bond.
Cation
A positively charged ion, formed when an atom loses one or more electrons. In simple terms, the atom gave away electrons and now has more protons than electrons, so it carries a positive charge.
Anion
A negatively charged ion, formed when an atom gains one or more electrons. Think of it as the atom that accepted the extra electrons and now carries a negative charge.
Coordinate covalent bond (dative bond)
A covalent bond in which both shared electrons are contributed by a single atom, rather than one from each. In simple terms, one atom donates the whole electron pair to the bond instead of each atom chipping in one electron.
Lattice structure (crystalline lattice)
A repeating three-dimensional arrangement of alternating positive and negative ions in an ionic compound. Think of it as a perfectly ordered grid of alternating charges, which is why ionic compounds form hard, brittle crystals.
Bond order
The number of shared electron pairs between two bonded atoms. Single = 1, double = 2, triple = 3. In simple terms, bond order tells you how many "sticks" connect the two atoms in a Lewis structure.
Bond length
The average distance between the nuclei of two bonded atoms. Think of it as how close the two atoms sit. Higher bond order means shorter bond length: triple bonds pull atoms closest together.
Bond energy (bond dissociation energy)
The energy required to break one mole of a particular bond in the gas phase, separating the atoms completely. In simple terms, it measures how strong the bond is. More shared electrons means more energy needed to pull the atoms apart.
Dipole moment (P = qd)
A vector quantity measuring the separation of partial charges in a polar bond, equal to the charge magnitude (q) multiplied by the displacement distance (d). Measured in Debye units. Think of it as a number that tells you how "lopsided" the electron sharing is in a bond.
Nonpolar covalent bond
A covalent bond in which electrons are shared equally between two atoms with identical or nearly identical electronegativities (difference of 0 to 0.5). In simple terms, neither atom pulls harder, so the electrons sit evenly between them.
Polar covalent bond
A covalent bond in which electrons are shared unequally because of a moderate electronegativity difference (0.5 to 1.7), creating partial positive and partial negative charges. Think of it as a tug-of-war where one atom pulls the electrons a bit closer to itself.
Atoms bond to achieve eight electrons in their outermost shell, mimicking noble-gas stability
Elements that always obey the octet rule: C, N, O, F, Na, Mg
Incomplete octet: some elements are stable with fewer than eight valence electrons
Hydrogen (2), helium (2), lithium, beryllium, boron
Boron compounds such as BF₃ are stable with only six electrons around boron
Expanded octet: elements in period 3 and beyond can hold more than eight electrons by using their empty d orbitals
Phosphorus (up to 10), sulfur (up to 12), chlorine (up to 14)
This is why PCl₅ and SF₆ exist despite having more than four bonds
Odd-electron species: molecules with an odd total of valence electrons cannot give every atom a full octet
Example: NO has 11 valence electrons, so one atom is left with seven
Ionic bonds: electron transfer from a metal to a nonmetal, driven by large electronegativity differences (greater than 1.7)
Covalent bonds: electron sharing between nonmetals with similar electronegativities
Nonpolar covalent: electronegativity difference of 0 to 0.5, equal sharing
Polar covalent: electronegativity difference of 0.5 to 1.7, unequal sharing
Coordinate covalent: both shared electrons come from one atom
Form between atoms with significantly different electronegativities, typically a metal and a nonmetal
The metal loses electrons to become a cation (positive ion); the nonmetal gains electrons to become an anion (negative ion)
Held together by electrostatic attraction between opposite charges
Electronegativity difference threshold: greater than 1.7
Properties of ionic compounds:
Very high melting and boiling points (strong electrostatic forces throughout the lattice)
Many dissolve readily in water and other polar solvents
Conduct electricity when molten or dissolved in water (ions are free to move), but not as solids
Form crystalline lattices: repeating patterns of alternating positive and negative ions that minimise repulsion
Bond order, length, and energy relationships:
Single bond: bond order 1, longest, weakest
Double bond: bond order 2, shorter, stronger
Triple bond: bond order 3, shortest, strongest
The pattern: as you add more shared electron pairs, the bond gets shorter and requires more energy to break
Polarity and electronegativity thresholds:
0 to 0.5 difference: nonpolar covalent (equal sharing). Examples include all homonuclear diatomics: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂
0.5 to 1.7 difference: polar covalent (unequal sharing). The more electronegative atom pulls electron density toward itself, gaining a partial negative charge (δ−); the less electronegative atom acquires a partial positive charge (δ+)
Greater than 1.7 difference: ionic (electron transfer)
Coordinate covalent bonds:
Both electrons in the shared pair come from the same atom
The donor atom (Lewis base) provides the lone pair; the acceptor atom (Lewis acid) receives it
Typically formed during Lewis acid-base reactions
Once formed, a coordinate covalent bond is indistinguishable from any other covalent bond; you can only identify it by counting valence electrons
Dipole moment:
P = qd
Where P is the dipole moment (in Debye), q is the magnitude of the partial charge, and d is the displacement vector separating the two partial charges.
Electronegativity difference thresholds for bond classification:
Electronegativity Difference | Bond Type |
|---|---|
0 to 0.5 | Nonpolar covalent |
0.5 to 1.7 | Polar covalent |
Greater than 1.7 | Ionic |
Table salt (NaCl) is the classic ionic compound. Its high melting point (801 °C) and its ability to conduct electricity when dissolved in water both follow directly from the lattice structure and the behaviour of free ions in solution.
The polar covalent bonds in water molecules are what make water such a good solvent for ionic and polar substances, which is why your body can transport nutrients and waste through the bloodstream.
Students often think the octet rule is universal. It is not. Hydrogen only needs two electrons, and elements in period 3+ can exceed eight by using d orbitals.
Students frequently confuse electronegativity with electron affinity. Electronegativity is about pulling shared electrons within a bond; electron affinity is the energy change when a free atom gains an electron.
A common mistake is thinking ionic bonds involve sharing. They do not. In ionic bonding, electrons are fully transferred, not shared.
Students sometimes believe that a coordinate covalent bond is somehow weaker or different from a regular covalent bond. Once formed, it is identical in strength and character; the distinction is only in where the electrons originally came from.
⚠️ Expect questions asking you to classify a bond as ionic, polar covalent, or nonpolar covalent based on an electronegativity difference. Know the thresholds cold: 0 to 0.5, 0.5 to 1.7, greater than 1.7.
⚠️ The relationship between bond order, bond length, and bond energy is a favourite multiple-choice topic. Remember: higher bond order = shorter length = greater energy.
⚠️ Octet rule exceptions are commonly tested. Be ready to identify incomplete octets (H, He, Li, Be, B), expanded octets (period 3+), and odd-electron species (like NO).
⚠️ Properties of ionic compounds (high melting point, conductivity when dissolved, lattice structure) are frequently tested as a set. Know all of them together.
True or false: Boron trifluoride (BF₃) violates the octet rule because boron has only six electrons in its valence shell. (True)
Fill in the blank: An electronegativity difference greater than ______ indicates an ionic bond. (1.7)
True or false: A triple bond is longer than a single bond. (False, it is shorter)
Fill in the blank: The atom that loses electrons in an ionic bond becomes a ______. (cation)
True or false: Ionic compounds conduct electricity as solids. (False, only when molten or dissolved)
Q: What type of bond forms between sodium (Na) and chlorine (Cl), and why?
A: Ionic bond. Sodium has a low ionisation energy and readily loses an electron; chlorine has a high electron affinity and readily gains one. The electronegativity difference is well above 1.7, so the electron is transferred rather than shared.
Q: Rank single, double, and triple bonds in order of increasing bond energy and decreasing bond length.
A: Bond energy (lowest to highest): single < double < triple. Bond length (longest to shortest): single > double > triple. More shared electron pairs means a stronger, shorter bond.
Q: A molecule has an electronegativity difference of 1.2 between its two atoms. Classify the bond and describe the charge distribution.
A: Polar covalent bond. The more electronegative atom carries a partial negative charge (δ−) and the less electronegative atom carries a partial positive charge (δ+). Electrons are shared, but unequally.
Q: Why can phosphorus form five bonds in PCl₅, while nitrogen cannot form five bonds?
A: Phosphorus is in period 3 and has access to empty 3d orbitals, allowing an expanded octet beyond eight electrons. Nitrogen is in period 2 and has no accessible d orbitals, so it is limited to the octet rule.
Q: What is a coordinate covalent bond, and how does it differ from a standard covalent bond?
A: In a coordinate covalent bond, both electrons in the shared pair come from the same atom (the Lewis base donates a lone pair to the Lewis acid). Once formed, the bond is identical to any other covalent bond in terms of strength and character.
This material connects directly to Lewis structures and formal charge (covered in Part 2 of these notes), because drawing correct structures requires you to know which atoms share, which transfer, and how many electrons each atom needs. It also feeds into VSEPR theory and molecular geometry (Part 3), where the number and type of bonds around a central atom determine the molecule's three-dimensional shape.
More broadly, understanding bond polarity is the foundation for intermolecular forces (hydrogen bonds, dipole-dipole, London dispersion), which in turn explain physical properties like boiling points and solubility.
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