Difficulty: Introductory | Prerequisites: General chemistry (atomic structure, electron configuration, electronegativity)
This topic is the foundation of everything in organic chemistry. Before you can understand reactions, mechanisms, or molecular behaviour, you need to be comfortable drawing Lewis structures, placing lone pairs correctly, and recognising where partial charges sit. If you skipped Gen Chem or feel shaky on electronegativity and valence electrons, revisit those first.
Lewis structures show how atoms share and hold electrons in a molecule. Lone pairs are non-bonding electron pairs that sit on an atom and affect its geometry and reactivity. Partial charges (noted with the Greek delta symbol) arise when electrons are shared unequally between atoms of different electronegativity.
Lewis structure (electron-dot structure)
A diagram showing every valence electron in a molecule as either a bonding pair (a line between two atoms) or a lone pair (dots on a single atom). Every atom must satisfy its valence, and the total electron count must match the sum of each atom's valence electrons.
In simple terms, it is the blueprint that tells you which atoms are connected, how many bonds they share, and where the leftover electrons sit.
Lone pair (non-bonding pair)
A pair of valence electrons on an atom that is not involved in bonding. Shown as two dots. Lone pairs still occupy space around the atom, affecting molecular shape and reactivity.
Think of it as electrons that "belong" to one atom rather than being shared between two.
Bonding pair
A pair of electrons shared between two atoms, forming a covalent bond. Drawn as a line between the atoms.
Partial charge (dipole, delta notation)
When two bonded atoms have different electronegativities, the more electronegative atom pulls electron density toward itself, gaining a slight negative charge (noted as a lowercase delta with a minus sign). The less electronegative atom carries a slight positive charge (delta plus). The bond is polar.
Think of it as a tug-of-war for electrons: the stronger atom wins a larger share, but neither atom fully gains or loses an electron.
Electronegativity
A measure of how strongly an atom attracts shared electrons toward itself in a covalent bond. Increases across a period (left to right) and up a group on the periodic table. Fluorine is the most electronegative element.
Valence electrons
The outermost-shell electrons of an atom, which participate in bonding. Carbon has 4, nitrogen has 5, oxygen has 6, halogens have 7.
Count the total number of valence electrons for all atoms in the molecule.
Identify the central atom (usually the least electronegative atom, never hydrogen).
Connect outer atoms to the central atom with single bonds (each bond uses 2 electrons).
Distribute remaining electrons as lone pairs on outer atoms first, then on the central atom, to satisfy the octet rule (duet rule for hydrogen).
If the central atom lacks an octet, convert lone pairs on outer atoms into double or triple bonds.
The worksheet uses Tylenol (acetaminophen) to practise reading a larger organic Lewis structure. Key observations:
The carbonyl group (C=O) carries a partial negative charge on the oxygen (delta minus) because oxygen is more electronegative than carbon.
The O-H bond in the hydroxyl group is polarised: oxygen is delta minus, hydrogen is delta plus.
The nitrogen atom carries lone pairs that are important for the molecule's chemistry.
Dichloromethane, CH2Cl2
Carbon is the central atom with 4 valence electrons.
Two C-H bonds and two C-Cl bonds, all single.
Each chlorine carries 3 lone pairs (6 non-bonding electrons).
Total valence electrons: 4 + 2(1) + 2(7) = 20.
Hydrazine, N2H4
Two nitrogen atoms bonded to each other, each also bonded to two hydrogens.
Each nitrogen carries one lone pair.
Total valence electrons: 2(5) + 4(1) = 14.
Methanol, CH3OH
Carbon bonded to three hydrogens and one oxygen; oxygen bonded to one hydrogen.
Oxygen carries two lone pairs.
Total valence electrons: 4 + 4(1) + 6 = 14.
Formal charge = (valence electrons of the free atom) - (non-bonding electrons) - (1/2 bonding electrons)
Total valence electron count = sum of all atoms' valence electrons, adjusted for any overall charge on the molecule or ion.
Octet rule: most atoms in Period 2 (C, N, O, F) require 8 electrons in their valence shell. Hydrogen requires 2 (the duet rule).
Partial charges on bonds are exactly what makes water a good solvent: the polar O-H bonds let water molecules interact with ions and other polar molecules. Drug design relies heavily on understanding lone pairs and partial charges, because a drug molecule like acetaminophen (Tylenol) must fit into an enzyme's active site through specific electrostatic interactions.
Students often forget lone pairs on atoms like chlorine and nitrogen. Every Lewis structure must account for all valence electrons, including lone pairs.
Hydrogen is never the central atom. It can only form one bond.
A polar bond does not mean the molecule is ionic. Delta plus and delta minus indicate partial charges from unequal sharing, not full electron transfer.
Students sometimes confuse the number of bonds an atom forms with its number of valence electrons. Carbon has 4 valence electrons and forms 4 bonds, but oxygen has 6 valence electrons and typically forms only 2 bonds (with 2 lone pairs).
Drawing correct Lewis structures is tested on virtually every organic chemistry exam. Expect to be asked to draw structures from molecular formulas, identify lone pairs, and mark partial charges on polar bonds.
Knowing where lone pairs sit is essential for predicting molecular geometry (VSEPR) and understanding reaction mechanisms later in the course.
True or False: Oxygen in a C=O bond carries a partial positive charge.
Fill in the blank: Nitrogen has ___ valence electrons and typically forms ___ bonds with ___ lone pair(s).
True or False: A lone pair on chlorine in CH2Cl2 counts toward chlorine's octet.
Fill in the blank: In a polar bond, the more electronegative atom carries a partial ___ charge.
True or False: Hydrogen can serve as the central atom in a Lewis structure.
Q: Draw the Lewis structure for CH2Cl2. How many lone pairs does each chlorine atom have?
A: Carbon is the central atom with two C-H bonds and two C-Cl bonds. Each chlorine has 3 lone pairs (6 non-bonding electrons). Total valence electrons = 20.
Q: In the Tylenol molecule, which atoms carry partial negative charges and why?
A: The oxygen atoms carry partial negative charges because oxygen is more electronegative than carbon and hydrogen. Oxygen pulls electron density toward itself in the C=O and O-H bonds.
Q: Draw the Lewis structure for hydrazine (N2H4). How many lone pairs are on each nitrogen?
A: Each nitrogen is bonded to two hydrogens and to the other nitrogen (single bond). Each nitrogen carries one lone pair. Total valence electrons = 14.
Q: What is the difference between a lone pair and a bonding pair?
A: A bonding pair is shared between two atoms to form a covalent bond. A lone pair sits on a single atom and is not shared. Both types count toward the atom's octet.
Q: Explain why the O-H bond in methanol is polar.
A: Oxygen (electronegativity about 3.5) is substantially more electronegative than hydrogen (about 2.1). Oxygen pulls the shared electrons closer to itself, creating a partial negative charge on oxygen and a partial positive charge on hydrogen.
Lewis structures connect directly to VSEPR theory (Week 1, Question 6), where lone pairs determine molecular geometry and bond angles. They also feed into resonance structures (covered in the next set of notes), because you need to draw correct Lewis structures before you can identify resonance contributors. Later in the course, lone pairs drive nucleophilic attack in reaction mechanisms.
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