Difficulty: Introductory | Prerequisites: Basic atomic structure, periodic table groups and periods.
This topic sits at the heart of general chemistry's bonding unit. Before you can predict molecular shapes (VSEPR), understand intermolecular forces, or make sense of reaction mechanisms, you need to know how atoms share or transfer electrons, and Lewis structures are the tool for mapping that out. Electronegativity is the property that tells you whether a bond's electrons sit evenly between two atoms or get pulled toward one side. If you are comfortable with valence electrons and the layout of the periodic table, you are ready for this material.
Electronegativity (EN) measures how strongly an atom pulls on shared electrons. The difference in EN between two bonded atoms determines whether a bond is nonpolar covalent, polar covalent, or ionic. Lewis dot structures are a step-by-step method for drawing how valence electrons are arranged in a molecule, showing bonds and lone pairs.
Electronegativity (EN)
The tendency of an atom to attract shared electrons toward itself in a chemical bond.
In simple terms, it is how "greedy" an atom is for electrons.
Valence electrons
The electrons in an atom's outermost energy level, which participate in bonding.
Think of them as the electrons that do the actual work when atoms connect.
Nonpolar covalent bond
A bond in which electrons are shared equally between two atoms, typically two identical nonmetal atoms.
In simple terms, neither atom hogs the electrons.
Polar covalent bond
A bond in which electrons are shared unequally, pulled more toward the more electronegative atom.
Think of it as a tug-of-war where one side is stronger.
Ionic bond
A bond formed when electrons are transferred (not shared) from a metal to a nonmetal.
In simple terms, one atom takes the electrons outright.
Delta notation (δ+ and δ−)
Symbols used to indicate partial charges on atoms in a polar bond. The more electronegative atom carries δ−; the less electronegative carries δ+.
Octet rule
The principle that most atoms bond in ways that give them eight electrons in their valence shell.
Think of eight as the "magic number" for stability (hydrogen only needs two).
Lewis dot structure (electron dot structure)
A diagram showing valence electrons as dots (lone pairs) and lines (bonding pairs) around atoms in a molecule.
It is your map of where every valence electron lives in a molecule.
Electronegativity is not a fixed charge; it is a relative scale describing how strongly an atom attracts bonding electrons.
Periodic table trends
EN increases as you move right across a row (period). Fluorine, in the top-right corner, is the most electronegative element.
EN decreases as you move down a column (group). Atoms get larger, and the nucleus holds outer electrons less tightly.
Bonding is a continuum, not three neat boxes. The difference in EN (ΔEN) between two atoms determines where a bond sits on that spectrum:
ΔEN < 0.4: nonpolar covalent (electrons shared equally). Example: N₂, O₂, Cl₂.
ΔEN ≥ 0.4: polar covalent (electrons shared unequally). Example: H–Cl, H–O.
Very large ΔEN (typically > 1.7): ionic (electrons effectively transferred). Example: NaCl.
Polarity and partial charges
In a polar bond, the more electronegative atom carries a partial negative charge (δ−).
The less electronegative atom carries a partial positive charge (δ+).
Polarity refers to the extent to which electrons are shifted from one atom to another during bonding.
Bond type by atom identity (quick rule of thumb)
Two identical nonmetal atoms: nonpolar covalent.
Two different nonmetal atoms: generally polar covalent.
A metal and a nonmetal: ionic.
Sum all valence electrons. Add up the valence electrons from every atom in the molecule. If the species carries a charge, add electrons for a negative charge or subtract for a positive charge.
Place the central atom. The least electronegative element (or the most metallic) goes in the centre. Hydrogen is never the central atom.
Draw single bonds from the central atom to each surrounding atom. Each single bond uses two electrons.
Distribute remaining electrons.
Complete octets on the outer atoms first (give each eight electrons, counting shared ones).
Place any leftover electrons on the central atom.
Check that every atom has an octet.
Octet exceptions to remember
Hydrogen needs only 2 electrons (a duet, not an octet).
Boron (B) and aluminium (Al) are stable with only 6 electrons.
Alkaline earth metals (Be, Mg) may be stable with only 4.
Third-period and heavier atoms can expand beyond 8 (e.g. SF₆ has 12 electrons around S).
Form multiple bonds if needed. If the central atom lacks an octet after step 4, convert lone pairs on outer atoms into double or triple bonds until the octet is satisfied. The nitrogen molecule (N₂) is a classic example: each N shares three electrons, forming a triple bond, so each N achieves an octet.
Polarity threshold: ΔEN < 0.4 = nonpolar covalent; ΔEN ≥ 0.4 = polar covalent.
Valence electron count: Group 1A = 1, Group 2A = 2, Group 3A = 3, Group 4A = 4, Group 5A = 5, Group 6A = 6, Group 7A = 7, Group 8A = 8.
Formal charge (useful for checking structures): FC = (valence electrons) − (lone pair electrons) − (½ bonding electrons).
The polarity of water (O is far more electronegative than H) is the reason water is such a good solvent: its polar bonds let it pull apart ionic compounds and interact with other polar molecules. Understanding bond polarity is also how chemists predict whether a drug molecule will dissolve in blood (polar) or pass through a cell membrane (nonpolar).
"Ionic and covalent are completely different things." They are not. Bonding is a continuum. A bond with ΔEN of 1.6 is not magically different from one at 1.8; it is just further along the polar spectrum.
"Hydrogen can be a central atom." It cannot. Hydrogen forms only one bond, so it always sits on the outside of a Lewis structure.
"Every atom must have exactly eight electrons." Not true. Hydrogen needs two. Boron and aluminium are stable with six. Elements in period 3 and beyond can exceed eight.
"A polar bond means the whole molecule is polar." Not necessarily. If the geometry is symmetrical (e.g. CO₂), polar bonds can cancel out and produce a nonpolar molecule overall. Molecular polarity depends on both bond polarity and shape.
⚠️ You will almost certainly be asked to draw a Lewis structure from scratch. Practise the four-step method until it is automatic.
⚠️ Expect at least one question asking you to classify a bond as nonpolar covalent, polar covalent, or ionic based on ΔEN.
⚠️ Know the periodic table trends for electronegativity cold: increases going right, decreases going down.
⚠️ Octet exceptions (H, B, Al, expanded octets) are a favourite source of exam traps. If a structure looks wrong, check whether the central atom is one of these exceptions.
True or false: Electronegativity increases as you move down a group. (False, it decreases.)
Fill in the blank: A bond with ΔEN < 0.4 is classified as ______. (Nonpolar covalent.)
True or false: Hydrogen needs eight electrons to satisfy its bonding requirements. (False, it needs two.)
Fill in the blank: In a polar bond, the more electronegative atom carries a partial ______ charge. (Negative, written δ−.)
True or false: Bonding character is a continuum, not a set of discrete categories. (True.)
Q: Draw the Lewis structure for water (H₂O). How many lone pairs are on the oxygen atom?
A: Oxygen has 6 valence electrons. Two are used in bonds to the two hydrogens (one bond each). The remaining four electrons form two lone pairs on oxygen.
Q: Is the bond between H and Cl polar or nonpolar? Which atom carries the partial negative charge?
A: Polar covalent. Chlorine is more electronegative than hydrogen, so Cl carries δ− and H carries δ+.
Q: Why does nitrogen gas (N₂) contain a triple bond?
A: Each nitrogen has 5 valence electrons. To achieve an octet, each N must share 3 electrons with the other, resulting in a triple bond (plus one lone pair on each N).
Q: A student places hydrogen in the centre of a Lewis structure for HCN. What is wrong with this?
A: Hydrogen can only form one bond, so it cannot be a central atom. Carbon (the least electronegative of C and N) should be the central atom.
Q: Classify the bond in NaCl as ionic, polar covalent, or nonpolar covalent, and explain why.
A: Ionic. Sodium is a metal and chlorine is a nonmetal, giving a very large ΔEN. The electron is effectively transferred from Na to Cl rather than shared.
This material connects directly to VSEPR theory, which uses Lewis structures as its starting point to predict three-dimensional molecular shapes. Bond polarity feeds into intermolecular forces (dipole-dipole, hydrogen bonding), which in turn explain physical properties like boiling point and solubility. Understanding ionic vs. covalent character also lays the groundwork for predicting chemical reactivity and the properties of ionic vs. molecular compounds.
Lewis structures, electron dot diagrams, electronegativity, EN, bond polarity, polar covalent, nonpolar covalent, ionic bond, delta plus, delta minus, partial charge, octet rule, valence electrons, bonding continuum, VSEPR, molecular geometry, lone pairs, bonding pairs, formal charge, expanded octet, duet rule, Pauling electronegativity scale, general chemistry, Purdue, CHM 115, CHM 116