Organic Chemistry · Structure and Bonding

Atomic Structure: Electron Configurations

7 min read

Configurations given are standard ground-state configurations for neutral atoms; the aufbau order applies through argon without exception.

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In 30 seconds

Now that orbitals exist, the question becomes: how do an atom's electrons fill them? The answer — the atom's — is a compact notation that lists which orbitals are occupied and how many electrons each holds. Carbon, for instance, has the configuration 1s² 2s² 2p².

Three rules govern the filling. The says electrons occupy the lowest-energy orbitals first. The caps each orbital at two electrons. says electrons spread out one per orbital within a subshell before pairing up. Together, these rules predict everything from an atom's magnetism to how many bonds it can form — and the configuration of carbon is the single most important one in organic chemistry.

Why this matters

Electron configurations are the bridge between the periodic table and chemical behavior:

decide reactivity. The electrons in the outermost shell — the valence electrons — are the ones that form bonds. Carbon's four valence electrons (2s² 2p²) are why carbon forms four bonds, the backbone of all organic molecules.
Hund's rule explains magnetism. An atom with unpaired electrons is (attracted into a magnetic field); one with all electrons paired is . The unpaired electrons of oxygen (1s² 2s² 2p⁴) are why liquid oxygen is attracted to a magnet.
Periodic trends: Configurations explain why elements in the same column of the periodic table behave alike — they have the same number of valence electrons.
Bonding theory depends on it: Knowing that carbon's 2s and 2p electrons differ in energy sets up the hybridization story of later topics in this chapter.

The college version

Core Concepts

The aufbau order: filling lowest energy first

Electrons fill orbitals in order of increasing energy. For the elements most important to organic chemistry (H through Ne), the order is:

$1s \to 2s \to 2p \to 3s \to 3p$

In this range, energy increases with n, and within a shell, s is lower than p: 1s < 2s < 2p < 3s < 3p. (For heavier elements, the 4s orbital fills before 3d — a famous exception that matters for transition metals but rarely for carbon chemistry.)

Pauli exclusion principle: two per orbital

No two electrons in an atom may have the same four quantum numbers. Since the first three (n, ℓ, m_ℓ) are fixed by the orbital, the only remaining freedom is spin: one electron with m_s = +12 (spin up, ↑) and one with m_s = -12 (spin down, ↓). Result: each orbital holds a maximum of two electrons, and they must have opposite spins.

Hund's rule: spread out before pairing

Within a subshell (the three 2p orbitals, for example), electrons first occupy each orbital singly with parallel spins before any orbital receives a second electron. The reason is repulsion: electrons would rather sit in separate orbitals (farther apart) than share one. So nitrogen, with three 2p electrons, has one electron in each of 2p_x, 2p_y, and 2p_z — all unpaired.

Writing configurations: the notation

An electron configuration lists the occupied orbitals with superscripts for electron counts:

$carbon: 1s 2 2s 2 2p 2 nitrogen: 1s 2 2s 2 2p 3 oxygen: 1s 2 2s 2 2p 4$

Check the counts: carbon has 2 + 2 + 2 = 6 electrons, matching its atomic number. A shorthand uses the preceding noble gas in brackets:

$carbon: [He] 2s 2 2p 2$

The valence electrons are those in the highest principal shell (n = 2 for carbon: 2s² 2p² = four valence electrons); the are the rest.

Magnetism from unpaired electrons

Paramagnetic: has unpaired electrons; attracted into a magnetic field (e.g., oxygen, with two unpaired 2p electrons).
Diamagnetic: all electrons paired; weakly repelled by a magnetic field (e.g., neon, 1s² 2s² 2p⁶).

How It Works / Step-by-Step Process

To write the electron configuration of an element:

Find the atomic number Z — that is the number of electrons to place.
Fill orbitals in aufbau order (1s, 2s, 2p, 3s, 3p, …), two electrons per orbital.
Within a subshell, apply Hund's rule: one electron per orbital first, then pair.
Stop when all Z electrons are placed.
Identify the valence shell (highest n) and count its electrons.
Optionally compress the core with the preceding noble gas in brackets.

Common Confusions

Common Confusion	Correct Understanding
"The 2p subshell can hold 8 electrons."	The 2p subshell has three orbitals × 2 electrons = 6; the whole second shell (2s + 2p) holds 8.
"Electrons pair up in an orbital before filling others."	Hund's rule says the opposite: fill each degenerate orbital singly first.
"All six electrons of carbon are valence electrons."	Only the four in the outermost shell (n = 2) are valence; the 1s pair is core.
"A full subshell is required for stability."	Stability comes from filled or half-filled subshells (e.g., nitrogen's half-filled 2p); unfilled subshells still form stable atoms.
"The 3d orbitals fill before 4s."	For potassium and calcium, 4s fills before 3d — but for elements in the organic core (through neon), the simple order 1s, 2s, 2p, 3s, 3p applies.
"Paramagnetic means magnetic like an iron bar."	Paramagnetism is a weak attraction into a magnetic field due to unpaired electrons — not ferromagnetism.

Eli explains

The same idea, in plain words

Explain it like I’m 10

Filling an atom with electrons is like seating guests in a movie theater. Everyone picks the best (lowest) seats first, no seat holds more than two people, and until the theater is half full, people sit apart rather than squeezing in together. Carbon's six electrons take the front row's two seats and four of the second row's seats — leaving four "seats" (the valence electrons) ready to hold hands with other atoms.

Worked example

Example 1: Electron configuration of oxygen

Oxygen has Z = 8. Fill in aufbau order, applying Hund's rule in the 2p subshell:

$1s 2 \to 2s 2 \to 2p 4$

Check the total: 2 + 2 + 4 = 8 ✓. In the 2p subshell, Hund's rule puts one electron in each of the three orbitals first (3 electrons), then the fourth pairs up in one of them. Result: two unpaired electrons → oxygen is paramagnetic. Its six valence electrons (2s² 2p⁴) mean oxygen needs two more electrons to complete an octet — that is why water is H₂O (two O–H bonds) and why oxygen forms a double bond in O₂.

Example 2: Carbon's valence electrons and bonding capacity

Carbon has Z = 6:

$1s 2 2s 2 2p 2 or [He] 2s 2 2p 2$

The valence shell is n = 2, holding 2 + 2 = 4 electrons. Carbon therefore has four valence electrons and can form four covalent bonds (as in CH₄, where each hydrogen contributes one electron to a shared pair). This tetravalence is the single most important fact in organic chemistry: it lets carbon chain together into molecules of unlimited size — proteins, DNA, fuels, and drugs.

Example 3: Paramagnetic or diamagnetic? Nitrogen vs. neon

Nitrogen, 1s² 2s² 2p³: the three 2p electrons occupy three different orbitals singly (Hund), so nitrogen has three unpaired electrons → paramagnetic.

Neon, 1s² 2s² 2p⁶: the 2p subshell is full (two electrons in each orbital, all paired) → diamagnetic. Its filled octet explains neon's total lack of reactivity.

Key takeaways

Filling order for organic-relevant elements: 1s < 2s < 2p < 3s < 3p.
Pauli: two electrons per orbital, opposite spins.
Hund: fill degenerate orbitals singly (parallel spins) before pairing.
Carbon: 1s² 2s² 2p² → four valence electrons → four bonds (tetravalence).
Valence electrons = outermost shell; they alone participate in bonding.
Unpaired electrons → paramagnetic; all paired → diamagnetic.
Elements in the same periodic-table column share valence-electron count and therefore similar chemistry.
¹3C NMR spectroscopy (seen later) is sensitive to the electronic environment of each carbon, which depends on its configuration and bonding.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

Write the full electron configuration of nitrogen and count its unpaired electrons.
Show answer
1s² 2s² 2p³; three unpaired electrons (one in each 2p orbital, per Hund's rule).
State the three rules for filling orbitals, in your own words.
Show answer
Aufbau: fill lowest-energy orbitals first. Pauli: two electrons per orbital, opposite spins. Hund: occupy degenerate orbitals singly before pairing.
How many valence electrons does carbon have, and what does that imply about its bonding?
Show answer
Four valence electrons (2s² 2p²); carbon can form four covalent bonds, enabling the vast chain-and-ring structures of organic chemistry.
Is an atom with configuration 1s² 2s² 2p⁶ paramagnetic or diamagnetic? Why?
Show answer
Diamagnetic — the 2p subshell is full, so all electrons are paired (this is neon, a noble gas).
Which element has the configuration [He] 2s² 2p³?
Show answer
Nitrogen (Z = 7).
Why do fluorine and chlorine, in the same column, behave similarly?
Show answer
They have the same number of valence electrons (seven: ns² np⁵); valence-electron count, not total electrons, governs chemical behavior.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

electron configuration: A notation listing occupied orbitals and their electron counts (e.g., 1s² 2s² 2p²)
aufbau principle: Rule that electrons fill the lowest-energy orbitals first
Pauli exclusion principle: No two electrons can share all four quantum numbers
Hund's rule: Electrons occupy degenerate orbitals singly before pairing
valence electrons: Electrons in the outermost occupied shell
core electrons: Electrons in filled inner shells
paramagnetic: Having one or more unpaired electrons; attracted into a magnetic field
diamagnetic: Having all electrons paired; weakly repelled by a magnetic field

Sources & references

openstax.org — Organic Chemistry

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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In 30 seconds

Why this matters

The college version

Core Concepts

The aufbau order: filling lowest energy first

Pauli exclusion principle: two per orbital

Hund's rule: spread out before pairing

Writing configurations: the notation

Magnetism from unpaired electrons

How It Works / Step-by-Step Process

Common Confusions

The same idea, in plain words

Worked example

Example 1: Electron configuration of oxygen

Example 2: Carbon's valence electrons and bonding capacity

Example 3: Paramagnetic or diamagnetic? Nitrogen vs. neon

Key takeaways

Check yourself

Keep learning

Key vocabulary

Related lessons

Sources & references