Source: PSO Worksheet 1, CHM 25500, Purdue University
Difficulty: Introductory to Intermediate | Prerequisites: General chemistry electron configuration, periodic table trends in electronegativity.
This material sits right at the start of organic chemistry and builds on everything you learnt about atoms, valence electrons, and bonding in general chemistry. The ability to draw Lewis structures, assign lone pairs, calculate formal charges, and read bond polarity is foundational: every reaction mechanism you encounter later depends on knowing where the electrons are and which atoms are electron-rich or electron-poor. If you can do this fluently, the rest of organic chemistry becomes far more intuitive.
Lewis structures show how valence electrons are distributed across a molecule, using dots for lone pairs and lines for bonds. Formal charge tells you whether an atom "owns" more or fewer electrons than it brought to the party. Bond polarity, marked with δ+ and δ–, shows the direction electrons are pulled in a covalent bond based on electronegativity differences.
Lewis structure (electron-dot structure)
A diagram showing all valence electrons in a molecule: bonding pairs as lines and non-bonding (lone) pairs as dots. Every atom should satisfy the octet rule (or duet for hydrogen).
In simple terms, it is a map of where every electron lives in the molecule.
Lone pair (non-bonding pair)
A pair of valence electrons on an atom that is not shared in a bond. These are drawn as two dots.
Think of it as electrons an atom keeps to itself rather than sharing with a neighbour.
Bonding pair
A pair of electrons shared between two atoms, forming a covalent bond. Represented as a line in structural formulae.
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: valence electrons – lone pair electrons – (1/2 × bonding electrons). It is a bookkeeping tool, not a true physical charge.
In simple terms, formal charge asks: does this atom have more or fewer electrons than a neutral, unbonded version of itself?
Octet rule
Most atoms in the second period (C, N, O, F) are most stable when surrounded by eight electrons (four pairs). Hydrogen is stable with two (a duet).
Electronegativity
A measure of how strongly an atom attracts shared electrons in a covalent bond. Increases going up and to the right on the periodic table (fluorine is the most electronegative element).
Bond polarity (δ+ and δ– notation)
When two bonded atoms have different electronegativities, the electron density shifts toward the more electronegative atom. That atom carries a partial negative charge (δ–) and the other carries a partial positive charge (δ+).
Think of it as a tug-of-war for electrons: the stronger puller gets δ–.
Line notation (skeletal structure / bond-line structure)
A shorthand where carbon atoms are implied at each vertex and line endpoint, hydrogen atoms on carbon are omitted, and heteroatoms (O, N, S, etc.) are written explicitly.
In simple terms, it is a stripped-down drawing that shows only the skeleton and the interesting atoms.
Start by counting total valence electrons for the molecule. For ions, add one electron per negative charge or subtract one per positive charge.
Draw single bonds between atoms first, then distribute remaining electrons as lone pairs to satisfy the octet rule on each atom.
Oxygen has two bonds and two lone pairs in its neutral, stable state. Nitrogen has three bonds and one lone pair. Carbon has four bonds and no lone pairs.
Halogens (F, Cl, Br, I) typically form one bond and carry three lone pairs.
Hydrogen never has lone pairs; it needs only two electrons (a duet).
Sulphur can expand its octet in some structures because it has accessible d-orbitals, but in organic chemistry it usually follows the octet rule.
When looking at a complex molecule like amoxicillin, work atom by atom: count the bonds already drawn, then fill in lone pairs until each atom has a full octet.
A covalent bond between two atoms of different electronegativity is polar. The more electronegative atom pulls electron density toward itself (δ–), leaving the other atom slightly electron-poor (δ+).
Common electronegativity ranking to memorise: F > O > N > Cl > Br > S > C > H. Carbon–hydrogen bonds are usually treated as essentially nonpolar.
A C–O bond is polar: oxygen is δ–, carbon is δ+. A C–N bond is polar: nitrogen is δ–, carbon is δ+. An N–H bond is polar: nitrogen is δ–, hydrogen is δ+.
If both atoms in a bond have the same electronegativity (e.g. C–C), the bond is nonpolar and no δ symbols are assigned.
Direction of polarity matters: the arrow or δ symbols point from δ+ toward δ–. In arrow notation, the arrowhead points toward the more electronegative atom.
Formal charge = (valence electrons of the free atom) – (lone pair electrons) – (half of bonding electrons).
The best Lewis structure minimises formal charges. If charges are unavoidable, negative formal charges should sit on the more electronegative atom.
Worked examples from the worksheet:
Ethyl acetate (C₄H₈O₂): the carbonyl oxygen is double-bonded to carbon (two lone pairs, formal charge 0). The ester oxygen has two single bonds and two lone pairs (formal charge 0).
Triethyl amine ((CH₃CH₂)₃N): nitrogen forms three single bonds and carries one lone pair. Formal charge on N is 0.
Ammonium ion (NH₄⁺): nitrogen forms four bonds, no lone pairs. Formal charge on N is +1. The overall ion charge is +1.
Acetonitrile (CH₃CN): the carbon–nitrogen bond is a triple bond. Nitrogen has one lone pair. Both C and N carry formal charge 0.
Carbonate ion (CO₃²⁻): one C=O double bond and two C–O single bonds (with resonance). The two singly bonded oxygens each carry formal charge –1, giving an overall charge of –2.
In line (skeletal) notation, carbon atoms are at every vertex and at the end of every line. Hydrogens on carbon are implicit: fill them in so each carbon has exactly four bonds.
Heteroatoms (O, N, S, Cl, etc.) and their attached hydrogens are always written explicitly.
To convert a full structure to line notation: remove all C and H labels, draw the carbon skeleton as a zigzag, and write only heteroatoms.
To convert line notation to a full structure: place a C at every vertex and endpoint, then add H atoms until each carbon has four bonds. Count all atoms to write the molecular formula.
For a molecule like amoxicillin, trace the full expanded structure carefully, then redraw it using the zigzag shorthand, keeping every heteroatom visible.
Formal charge formula:
Formal charge = V – L – B/2
Where V = number of valence electrons for the free atom, L = number of lone pair electrons on the atom, B = number of bonding electrons (count each bond line as 2).
Total valence electron count for a molecule:
Sum the valence electrons from each atom. For anions, add electrons equal to the charge. For cations, subtract electrons equal to the charge.
Electronegativity ranking (memorise):
F (4.0) > O (3.5) > N (3.0) > Cl (3.0) > Br (2.8) > S (2.5) > C (2.5) > H (2.1)
Bond polarity is why water dissolves salt but oil does not: the polar O–H bonds in water interact favourably with ions, while nonpolar C–H bonds in oil do not. Understanding which bonds are polar is the foundation for predicting solubility, reactivity, and intermolecular forces in drug design, materials science, and biochemistry.
Formal charge analysis helps chemists evaluate the stability of reactive intermediates, which matters when designing synthetic routes in pharmaceutical manufacturing.
Students often forget lone pairs on heteroatoms in Lewis structures. Every oxygen, nitrogen, and halogen needs its lone pairs drawn, not just the bonds.
Students sometimes confuse formal charge with oxidation state. Formal charge splits bonding electrons equally between the two atoms; oxidation state gives them entirely to the more electronegative atom. They are different tools for different purposes.
Students often assume that a C–H bond is polar. The electronegativity difference is small (about 0.4), so C–H bonds are treated as essentially nonpolar in organic chemistry.
Students sometimes draw five bonds to carbon or forget that hydrogen can only ever form one bond. Carbon always has exactly four bonds. Hydrogen always has exactly one.
⚠️ Drawing Lewis structures with correct lone pairs and formal charges is tested on nearly every organic chemistry exam. Expect to do this under time pressure.
⚠️ Bond polarity questions often appear as "assign δ+ and δ– to the following bonds." Know your electronegativity rankings cold.
⚠️ Converting between full structures and line notation is a bread-and-butter skill. You will need it to read problems for the rest of the course.
⚠️ Formal charge calculation shows up in resonance structure analysis (covered in the next set of notes) and in evaluating reaction intermediates.
True or False: Oxygen in a neutral molecule typically has two bonds and two lone pairs. (True)
Fill in the blank: The formal charge on nitrogen in NH₄⁺ is ____. (+1)
True or False: A C–O bond is nonpolar. (False, oxygen is more electronegative so the bond is polar)
Fill in the blank: In line notation, carbon atoms are found at every ____ and ____. (vertex, endpoint)
True or False: Formal charge and oxidation state are the same thing. (False)
Q: Calculate the formal charge on the nitrogen atom in triethyl amine, (CH₃CH₂)₃N. Nitrogen has three bonds to carbon and one lone pair.
A: Formal charge = 5 (valence e–) – 2 (lone pair e–) – 3 (half of 6 bonding e–) = 0. Nitrogen is neutral.
Q: In the ammonium ion NH₄⁺, nitrogen has four N–H bonds and no lone pairs. What is its formal charge?
A: Formal charge = 5 – 0 – 4 = +1.
Q: Assign δ+ and δ– to a C=O bond.
A: Oxygen is more electronegative than carbon. Oxygen is δ–, carbon is δ+.
Q: How many lone pairs does each oxygen in the carbonate ion CO₃²⁻ carry (in the Lewis structure with one C=O double bond and two C–O single bonds)?
A: The doubly bonded oxygen has two lone pairs. Each singly bonded oxygen has three lone pairs and a formal charge of –1.
Q: Convert the line structure of a molecule that appears as a zigzag with four vertices and an OH group at the end. How many carbon atoms does it have and what is its molecular formula?
A: Four vertices plus the endpoint gives five carbons (assuming a straight chain). With an OH at the terminus, the formula is C₅H₁₂O (pentan-1-ol), after filling hydrogens so each carbon has four bonds.
Formal charge and lone pair placement feed directly into resonance structure analysis (PSO Worksheet 1, Q4), where you move lone pairs into bonds and evaluate which contributor is the most stable. If you cannot reliably draw lone pairs and assign formal charges, resonance will not make sense.
Bond polarity connects to molecular polarity and dipole moments (later in the course), and ultimately to intermolecular forces, solubility, and physical properties.
Line notation fluency is assumed for every subsequent topic. Reaction mechanisms, spectroscopy problems, and synthesis questions all use skeletal structures.
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