Difficulty: Introductory | Prerequisites: Basic periodic table familiarity, electron configuration basics
Atoms bond by sharing or transferring electrons. In covalent bonds, atoms share electron pairs to reach a stable electron count (usually eight, per the octet rule). Lewis structures are flat diagrams that map out which electrons are shared (bonding pairs) and which sit unused (lone pairs) on each atom. Understanding these diagrams is the first step to predicting molecular shape, polarity, and reactivity.
Covalent bond
A bond formed when two atoms share one or more pairs of electrons. Typically occurs between two non-metal atoms with similar electronegativities. In simple terms, both atoms contribute electrons to a shared pool so each can reach a stable count.
Ionic bond
A bond formed when the electronegativity difference between two atoms is large enough (generally greater than 1.7) that electrons are transferred rather than shared. Think of it as one atom taking the electron outright instead of sharing.
Lewis structure
A 2D diagram showing only the valence electrons of each atom in a molecule. Bonds appear as dashes (lines) and non-bonding electrons appear as dot pairs. Only valence electrons are drawn, never core electrons.
Valence electrons
The electrons in an atom's outermost shell. For main-group elements, the group number on the periodic table tells you how many valence electrons the atom has. These are the electrons that participate in bonding.
Core electrons
All electrons that are not valence electrons. They sit in inner shells and do not participate in bonding.
Octet rule
The principle that atoms tend to gain, lose, or share electrons until they have eight valence electrons, achieving maximum stability. In simple terms, atoms "want" a full outer shell of eight.
Duet rule
Hydrogen and helium need only two electrons to fill their outer shell, not eight. This is why hydrogen forms just one bond.
Lone pair
A pair of valence electrons on an atom that is not involved in bonding. These are drawn as two dots sitting on the atom in a Lewis structure. They still count toward the atom's octet.
Bonding pair
A pair of electrons shared between two atoms in a covalent bond. Drawn as a dash (line) in a Lewis structure.
Bond order
The number of electron pairs shared between two atoms. Single bond = 1, double bond = 2, triple bond = 3.
Electronegativity
A measure of how strongly an atom attracts shared electrons toward itself in a bond. Higher electronegativity means the atom pulls electrons closer.
Polar bond
A covalent bond in which electrons are shared unequally because the two atoms have different electronegativities. The electrons spend more time near the more electronegative atom.
Non-polar bond
A covalent bond in which electrons are shared roughly equally, because the atoms have similar electronegativities.
Covalent compounds form discrete molecules. They can be gases, liquids, or non-brittle solids at room temperature. They have relatively lower melting points and higher vapour pressures.
Ionic compounds are typically crystalline solids with high melting points. Many dissolve readily in water to produce electrically conducting solutions.
The dividing line: when the electronegativity difference between two bonded atoms exceeds roughly 1.7, the bond is considered ionic rather than covalent.
For any main-group element, the group number equals the number of valence electrons.
Carbon (Group IV): 4 valence electrons
Nitrogen (Group V): 5 valence electrons
Oxygen (Group VI): 6 valence electrons
Fluorine/Chlorine (Group VII): 7 valence electrons
Noble gases (Group VIII) already have 8 valence electrons, which is why they rarely form compounds.
Atoms form bonds to reach 8 valence electrons (or 2 for hydrogen and helium).
Carbon has 4 valence electrons and forms 4 bonds to complete its octet.
Nitrogen has 5 valence electrons, forms 3 bonds, and keeps 1 lone pair (3 x 2 + 2 = 8).
Oxygen has 6 valence electrons, forms 2 bonds, and keeps 2 lone pairs (2 x 2 + 4 = 8).
Hydrogen has 1 valence electron and forms 1 bond (duet rule).
Write the atomic symbol for each atom.
Only valence electrons appear in the diagram.
Place bonds (dashes) and lone pairs (dot pairs) in the north, south, east, west positions first, then use additional positions if needed.
A single dash represents a single bond (2 shared electrons).
Two dashes (or two lines) represent a double bond (4 shared electrons).
Three dashes represent a triple bond (6 shared electrons).
Single bond: 2 electrons shared, bond order = 1. Example: each C–H bond in methane (CH₄).
Double bond: 4 electrons shared, bond order = 2. Example: the C=O bonds in carbon dioxide (CO₂).
Triple bond: 6 electrons shared, bond order = 3. Example: the N≡N bond in molecular nitrogen (N₂).
As bond order increases, bond length decreases and bond strength increases.
When two bonded atoms have nearly identical electronegativities, the bond is non-polar (electrons shared equally).
When electronegativities differ significantly (but not enough for an ionic bond), the bond is polar covalent. The electrons sit closer to the more electronegative atom.
A useful rough threshold: electronegativity difference less than about 0.4 is non-polar, 0.4 to 1.7 is polar covalent, greater than 1.7 is ionic.
Formal charge formula
Formal charge = (number of valence electrons) – (number of bonds) – (number of unshared electrons)
This is used to evaluate which Lewis structure is the "best" representation of a molecule.
Typical bond counts by element
Hydrogen: 1 bond
Carbon: 4 bonds
Nitrogen: 3 bonds (plus 1 lone pair)
Oxygen: 2 bonds (plus 2 lone pairs)
Halogens (F, Cl, Br, I): 1 bond (plus 3 lone pairs)
Electron configuration of carbon
Carbon (Z = 6): [He] 2s² 2p², giving 4 valence electrons. The group number (IV) matches the valence electron count for all main-group elements.
"Polar bonds are ionic bonds." They are not. A polar covalent bond still involves sharing electrons, just unequally. Only when the electronegativity difference is large enough (roughly above 1.7) does the bond become ionic.
"Double bonds are just two single bonds side by side." A double bond is shorter and stronger than a single bond between the same two atoms. The second pair of shared electrons sits in a different region of space (a pi bond), not simply a duplicate of the first.
"Every atom needs exactly 8 electrons." Hydrogen and helium follow the duet rule (2 electrons). Some elements in Period 3 and below can have expanded octets (more than 8), which you will see with molecules like PCl₅ and SF₄.
"Lone pairs do not matter." Lone pairs occupy space around the central atom and directly affect molecular shape through VSEPR theory. Ignoring them leads to wrong geometry predictions.
⚠️ You will be expected to draw correct Lewis structures from a molecular formula. Practise counting valence electrons and distributing them as bonds and lone pairs until it is automatic.
⚠️ Know the typical bond counts: H = 1, C = 4, N = 3, O = 2. If your Lewis structure gives an atom more or fewer bonds than expected, check your work.
⚠️ Be able to distinguish ionic from covalent from polar covalent using electronegativity differences.
⚠️ Formal charge calculations appear on exams. The formula (valence electrons – bonds – unshared electrons) must be memorised.
⚠️ Understand that lone pairs count toward the octet but are not visible in a molecular model kit.
True or false: A covalent bond involves the transfer of electrons from one atom to another.
Fill in the blank: Carbon has ___ valence electrons and typically forms ___ bonds.
True or false: A polar covalent bond is the same thing as an ionic bond.
Fill in the blank: The octet rule states that atoms tend to have ___ valence electrons for maximum stability.
True or false: Lone pairs are drawn in a Lewis structure but do not affect molecular shape.
(Answers: 1. False, sharing not transfer. 2. Four, four. 3. False, polar covalent still shares electrons. 4. Eight. 5. False, they strongly affect shape.)
Q: Consider the diatomic molecules H₂, O₂, N₂, F₂, and HCl. What type of bond (single, double, or triple) does each have?
A: H₂ has a single bond (each H contributes 1 electron). O₂ has a double bond (each O needs 2 more electrons). N₂ has a triple bond (each N needs 3 more electrons). F₂ has a single bond (each F needs 1 more electron). HCl has a single bond (H needs 1, Cl needs 1).
Q: Draw the Lewis structure of methane (CH₄) and explain why carbon forms exactly four bonds.
A: Carbon has 4 valence electrons (Group IV). To satisfy the octet rule, it needs 4 more electrons, which it obtains by forming 4 single bonds with 4 hydrogen atoms. Each hydrogen achieves a duet. The Lewis structure shows C in the centre with a single dash to each H.
Q: What is the difference between a bonding pair and a lone pair?
A: A bonding pair is shared between two atoms and forms a covalent bond (drawn as a dash). A lone pair sits on a single atom and is not shared (drawn as two dots). Both count toward the atom's octet.
Q: A bond has an electronegativity difference of 1.2. Is it ionic, polar covalent, or non-polar covalent?
A: Polar covalent. The difference is large enough for unequal sharing but below the roughly 1.7 threshold for ionic character.
Q: Calculate the formal charge on nitrogen in NH₃. Nitrogen has 5 valence electrons, 3 bonds, and 2 unshared electrons (one lone pair).
A: Formal charge = 5 – 3 – 2 = 0. Nitrogen has a formal charge of zero in ammonia.
Lewis structures are the foundation for VSEPR theory (covered in Part 2 of these notes), which predicts molecular geometry from the number of bonding and lone pairs around a central atom. Bond polarity feeds directly into determining whether an entire molecule is polar or non-polar, which in turn governs physical properties such as solubility and boiling point. Formal charges become essential when you study resonance structures, where the "best" Lewis structure is the one that minimises formal charges and places negative charges on more electronegative atoms.
Lewis dot structure, electron dot diagram, covalent bonding, ionic bonding, valence shell, octet rule, duet rule, bond order, single bond, double bond, triple bond, electronegativity, polar covalent, non-polar covalent, formal charge, bonding pair, lone pair, valence electrons, core electrons, G. N. Lewis, general chemistry, molecular modelling lab