Atomic Structure & Bonding

Atomic Orbitals & Electron Configuration

Part 1 of 7

Every reaction you will draw in Organic Chemistry is, underneath the curved arrows, a story about electrons moving between orbitals. Before you can predict where a nucleophile attacks or why a carbocation rearranges, you need a physically honest picture of where electrons actually live around an atom. That picture is the atomic orbital.

An orbital is not a tiny planetary orbit. It is a region of space — described by a wavefunction $\psi$ — where there is roughly a 90% probability of finding an electron. The square of the wavefunction, $\psi^2$, gives the electron probability density. This matters for orgo because bonding is orbital overlap: a $\sigma$ bond forms where two orbitals share electron density between two nuclei, and the shape of the orbital dictates the direction a bond can point.

The orbitals you will use constantly are:

• $s$ orbitals — spherical, no directional preference. One per shell.
• $p$ orbitals — two-lobed (dumbbell-shaped), aligned along the $x$, $y$, and $z$ axes. Three per shell from $n = 2$ upward.

For carbon and its neighbors (N, O, F), the $1s$, $2s$, and three $2p$ orbitals are the entire toolkit. Everything else in this course is built from how those five orbitals are filled, mixed, and overlapped.

Quantum Numbers: The Orbital's Address

Each orbital is specified by a set of quantum numbers, and three rules govern how electrons fill them. You do not need to solve the Schrödinger equation, but you must be fluent with the bookkeeping, because it tells you the number of valence electrons — and that number is the single most important fact about an atom in orgo.

Worked Example: The Ground-State Configuration of Carbon

Carbon has 6 electrons. Applying Aufbau, Pauli, and Hund in order:

• $1s^2$ — the first two electrons pair up in the lowest orbital (these are core electrons, chemically inert).
• $2s^2$ — the next two fill and pair the $2s$.
• $$ — the final two electrons enter the set. By they go into orbitals with parallel spin, not into the same one.

Checkpoint — Orbitals & Filling Rules

The Octet Rule and Why It Works

Second-row elements (C, N, O, F) are most stable when surrounded by eight valence electrons — a filled $2s$ plus filled $2p$ set, the configuration of neon. This is the octet rule, and it is the reason Lewis structures (Part 2) are drawn the way they are.

The octet rule has a hard physical ceiling for the second row: there are only four valence orbitals available ($2s$ + three $2p$), and each holds two electrons, so $4 \times 2 = 8$ is the maximum. This is why carbon, nitrogen, oxygen, and fluorine exceed an octet — they have no orbitals to expand into. Memorize this: an "expanded octet" on carbon is always a mistake.

Checkpoint — The Octet Rule

Exit Ticket — Part 1 Synthesis

Lewis Structures & Formal Charge

Part 2 of 7

A Lewis structure is the working diagram of organic chemistry. It shows every bond, every lone pair, and — once you add formal charges — exactly where a molecule is electron-rich (nucleophilic) and electron-poor (electrophilic). You will draw thousands of these. Drawing them correctly and quickly is a non-negotiable skill, because a wrong Lewis structure produces a wrong reaction mechanism every single time.

A reliable five-step procedure:

1. Count total valence electrons. Sum the valence electrons of every atom. For an anion, add one electron per negative charge; for a cation, subtract one per positive charge.
2. Place the skeleton. Carbon is almost always central; hydrogen and halogens are always terminal (one bond each). The least electronegative atom (other than H) tends to be central.
3. Connect with single bonds, then distribute remaining electrons as lone pairs, completing octets on the outer atoms first.
4. Form multiple bonds by converting lone pairs into bonding pairs wherever a central atom is short of an octet.
5. Assign formal charges to check that you have the best structure.

Steps 1 and 5 are where students lose points, so we will drill the electron count and formal charge hardest.

Formal Charge: The Bookkeeping That Reveals Reactivity

Formal charge answers the question: relative to a free, neutral atom, did this atom gain or lose ownership of electrons when it joined the molecule? The rule for counting "owned" electrons is the heart of it: an atom owns all of its lone-pair electrons but only half of each bonding pair (the other half belongs to its bonding partner).

Hybridization

Part 3 of 7

In Part 1 we hit a contradiction: ground-state carbon ($1s^2\,2s^2\,2p^2$) has only two unpaired electrons, yet methane has four identical bonds in a perfect tetrahedron. Hybridization is the model that resolves this — and it is arguably the single most useful concept in all of orgo, because hybridization at a carbon predicts its geometry, its bond angles, its bond strengths, and even its acidity.

Sigma & Pi Bonds: Orbital Overlap and Molecular Orbitals

Part 4 of 7

Hybridization (Part 3) told us the shape of a carbon's bonding orbitals. This part tells us what happens when those orbitals overlap to actually form a bond — and why the two flavors of covalent bond, $\sigma$ (sigma) and $\pi$ (pi), behave so differently. The distinction between $\sigma$ and $\pi$ is the engine behind almost every reaction of alkenes, alkynes, and carbonyls.

Molecular orbital (MO) theory is the underlying framework. When two atomic orbitals overlap, they combine into two molecular orbitals (orbitals are conserved — two in, two out):

Electronegativity, Bond Polarity & Dipole Moments

Part 5 of 7

A covalent bond shares electrons — but rarely equally. Electronegativity is an atom's pull on the shared electrons of a bond. When two bonded atoms differ in electronegativity, the electron density shifts toward the greedier atom, creating a polar bond: a separation of partial charge, written $\delta^-$ (partial negative, on the more electronegative atom) and $\delta^+$ (partial positive). This polarity is the origin of almost every intermolecular force and every "where does the nucleophile attack" question in orgo.

The trend on the periodic table, which you should know without a chart:

• Electronegativity increases left right across a period (more nuclear charge pulling on the same shell).

Problem-Solving Workshop

Part 6 of 7

You now have all the tools: Lewis structures and formal charge (Part 2), hybridization and s-character (Part 3), $\sigma$/$\pi$ bonds and bond order (Part 4), and electronegativity, dipoles, and VSEPR geometry (Part 5). Real exam questions rarely test these in isolation — they hand you a structure and ask you to chain the ideas together. This workshop drills that integrated workflow on real organic molecules.

A master procedure for analyzing any structure handed to you:

1. Draw / verify the Lewis structure — correct electron count, octets satisfied.
2. Assign formal charges — locate any charged atoms; these flag reactive sites.
3. Count regions of electron density per central atom ($\sigma$ bonds + lone pairs) to assign hybridization and VSEPR geometry.
4. Identify vs bonds — singles are ; doubles add one ; triples add two .

Resonance, Delocalization & Capstone Review

Part 7 of 7

This final part does two things. First, it gives resonance — touched on in Parts 2 and 6 — the full treatment it deserves, because resonance is the concept students most often misunderstand and the one orgo leans on most heavily (acidity, stability, aromaticity, mechanism all depend on it). Second, it ties the whole unit together: structure determines properties, from the orbital all the way up to boiling point and reactivity.

The thesis of the unit: a molecule's behavior is dictated, in a clean chain, by

$\text{electron configuration} \rightarrow \text{Lewis structure / formal charge} \rightarrow \text{hybridization} \rightarrow \text{geometry} \rightarrow \text{polarity} \rightarrow \text{properties and reactivity}$