Organic Chemistry | Purdue University | Lecture 2
Difficulty: Introductory Prerequisites: Electronic configuration, the octet rule (see Part 1 of these notes).
Covalent bonding is the mechanism by which most organic molecules hold together. Where the first part of this lecture covered how electrons are arranged and why atoms seek a full valence shell, this part covers what happens when atoms share electrons rather than transferring them. You will also learn how to assign formal charges, which is critical for drawing correct Lewis structures and understanding reaction mechanisms later in the course. If you are not yet comfortable with electron configurations and the octet rule, review those first.
A covalent bond forms when two atoms share one or more pairs of electrons so both can satisfy the octet rule. Bonds come in single, double, and triple varieties depending on how many pairs are shared. When an atom has more or fewer bonds than the valency rules predict, it carries a formal charge, which you calculate by comparing its bonded electron count to its neutral-atom electron count.
Covalent bond
A bond formed when one or more pairs of electrons are shared between two atomic nuclei. The shared electron pair counts toward the valence shell of both atoms. In simple terms, two atoms hold hands by sharing electrons instead of one stealing from the other.
Lewis dot structure (electron dot structure)
A diagram showing the valence electrons of an atom as dots arranged around the element symbol. Gilbert Lewis introduced this notation in 1916. Think of it as a map of where every valence electron lives on an atom or molecule.
Lone pair (non-bonding electron pair)
A pair of valence electrons on an atom that is not shared in a bond. Lone pairs sit on one atom only and are shown as two dots in a Lewis structure. In simple terms, these are electrons that belong to one atom and are not involved in holding atoms together.
Single bond
One shared pair of electrons between two atoms (two valence electrons involved). Drawn as a single line (e.g. H–H).
Double bond
Two shared pairs of electrons between two atoms (four valence electrons involved). Drawn as two lines (e.g. O=O).
Triple bond
Three shared pairs of electrons between two atoms (six valence electrons involved). Drawn as three lines (e.g. N≡N).
Valency (valence)
The typical number of bonds an atom forms. Halogens form 1, hydrogen forms 1, oxygen forms 2, nitrogen forms 3, carbon forms 4. Think of it as the number of "hands" an atom has available for bonding.
Formal charge
The charge assigned to an atom in a molecule by assuming that all bonding electrons are shared equally (split 50/50). Calculated as: (valence electrons of the neutral atom) minus (lone pair electrons) minus (half the bonding electrons). In simple terms, it is a bookkeeping charge that tells you whether an atom has more or fewer electrons than it would as a free, neutral atom.
Atoms share electron pairs so that each atom can achieve a full valence shell (octet, or duet for hydrogen).
The shared electron pair counts toward the valence shell of both atoms simultaneously.
Example: two hydrogen atoms (each 1s¹) share their single electrons to form H:H (or H–H). Each hydrogen now "sees" two electrons, satisfying its duet.
Gilbert Lewis described this model of bonding in 1916.
A single bond shares one pair (2 electrons). Example: H–H.
A double bond shares two pairs (4 electrons). Example: O=O.
A triple bond shares three pairs (6 electrons). Example: N≡N.
In the nitrogen example, each N atom has five valence electrons (1s² 2s² 2p³). Two nitrogen atoms share three pairs to form a triple bond (:N≡N:), leaving one lone pair on each nitrogen.
These rules tell you how many bonds each common atom typically forms:
Halogens (F, Cl, Br, I): 1 bond
Hydrogen (H): 1 bond
Oxygen (O): 2 bonds
Nitrogen (N): 3 bonds
Carbon (C): 4 bonds
An atom with more or fewer electrons than these rules predict will carry a charge (it becomes an ion).
Special cases (expanded octets):
Phosphorus (P) can form 3 or 5 bonds.
Sulphur (S) can form 2, 4, or 6 bonds.
These elements sit in period 3 or below and have access to d-orbitals, which allows them to exceed eight electrons.
Methane (CH₄): carbon forms four single bonds, one to each hydrogen.
In the Lewis structure, carbon sits in the centre bonded to four hydrogen atoms.
Each H achieves a duet (2 valence electrons). Carbon achieves an octet (8 valence electrons, all in bonding pairs).
Formal charge arises when an atom has more or fewer bonds than the standard valency rules predict.
To calculate formal charge, use the bookkeeping method:
Split each covalent bond in half, giving one electron from the bond to each atom.
Add the lone pair electrons (these stay with their atom).
Compare this count to the number of valence electrons the neutral atom has.
Formal charge = (valence electrons of neutral atom) – (lone pair electrons) – (½ bonding electrons).
Worked example: the ammonium ion (NH₄⁺)
Nitrogen forms four bonds to hydrogen (one more bond than nitrogen's usual three).
Bookkeeping: N gets 4 electrons (half of 8 bonding electrons) and 0 lone pairs = 4 electrons total.
Neutral nitrogen has 5 valence electrons.
Formal charge on N = 5 – 0 – 4 = +1.
Each H gets 1 electron from its bond, matching the 1 valence electron of neutral H, so each H has a formal charge of 0.
The +1 formal charge on nitrogen accounts for the overall +1 charge of the ammonium ion.
Important nuance: formal charge is a bookkeeping tool. In reality, the positive charge in NH₄⁺ is not sitting entirely on nitrogen. Because nitrogen is more electronegative than hydrogen, the electron density is pulled toward nitrogen. The real charge distribution is more spread out than the formal charge suggests.
Charged species are typically less stable than neutral ones.
Molecules will rearrange, if possible, to minimise formal charges.
When drawing Lewis structures, the structure with the fewest and smallest formal charges is generally preferred.
Students often confuse lone pairs with bonding pairs. Lone pairs belong to one atom only and are not shared between nuclei. They still count toward that atom's octet.
A common mistake is forgetting to count lone pairs when calculating formal charge. Both lone pair electrons and half the bonding electrons contribute to the count.
Some students believe that a formal charge of +1 on nitrogen in NH₄⁺ means nitrogen has literally lost an electron. It has not. Formal charge is a bookkeeping convention, not a description of where the electrons physically are.
Students sometimes assume that carbon can form five bonds because they see expanded octets on phosphorus and sulphur. Carbon is a second-period element with no accessible d-orbitals, so it is strictly limited to four bonds.
⚠️ You will draw Lewis structures constantly throughout organic chemistry. Knowing the valency rules cold (C = 4, N = 3, O = 2, halogens = 1, H = 1) saves time on every problem.
⚠️ Formal charge calculation is tested directly and also appears embedded in larger problems (resonance structures, reaction mechanisms). Practise the bookkeeping method until it is automatic.
⚠️ Recognising that charged species are less stable than neutral ones is key for predicting which resonance structure is the major contributor.
⚠️ Know the distinction between formal charge (a bookkeeping tool) and actual charge distribution (determined by electronegativity). This distinction becomes important when you study polar bonds and dipole moments.
True or False: A double bond involves the sharing of four electrons.
True.
Fill in the blank: Carbon typically forms ____ covalent bonds.
Four.
True or False: A lone pair counts toward an atom's octet.
True.
Fill in the blank: The formal charge on nitrogen in NH₄⁺ is ____.
+1.
True or False: Formal charge reflects the exact physical location of charge in a molecule.
False. It is a bookkeeping convention. Electronegativity determines the real charge distribution.
Q: How many valence electrons does each atom in methane (CH₄) "see," and does each satisfy the octet rule?
A: Each hydrogen sees 2 electrons (its bonding pair), satisfying the duet rule. Carbon sees 8 electrons (four bonding pairs), satisfying the octet rule.
Q: Draw the Lewis structure of N₂ and identify the bond type and any lone pairs.
A: Each nitrogen has five valence electrons. They share three pairs to form a triple bond (:N≡N:). Each nitrogen also retains one lone pair. Total electrons shown: 10 (matching 2 × 5 valence electrons).
Q: Calculate the formal charge on nitrogen in NH₄⁺ using the bookkeeping method.
A: Nitrogen has 4 bonds and 0 lone pairs. Half of 8 bonding electrons = 4. Neutral nitrogen has 5 valence electrons. Formal charge = 5 – 0 – 4 = +1.
Q: Sulphur can form 2, 4, or 6 bonds. Why can sulphur exceed the octet rule while oxygen cannot?
A: Sulphur is in period 3 and has empty 3d orbitals available to accommodate additional electrons beyond eight. Oxygen is in period 2 and has no accessible d-orbitals, so it is limited to eight electrons.
Q: A molecule has two possible Lewis structures. One places a formal charge of +1 on oxygen and –1 on carbon. The other has no formal charges. Which is preferred and why?
A: The structure with no formal charges is preferred. Structures with fewer and smaller formal charges are more stable. Placing a positive formal charge on oxygen (the more electronegative atom) is especially unfavourable.
Lewis structures and formal charge feed directly into resonance (upcoming lectures), where you draw multiple valid Lewis structures for the same molecule and evaluate which contributes most to the real structure. Formal charge is also central to understanding nucleophiles and electrophiles: atoms with negative formal charges or lone pairs tend to act as nucleophiles, while atoms with positive formal charges tend to be electrophilic sites. The valency rules here are the foundation for functional group chemistry, which is the organising principle of the rest of the course.
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