Electron configuration and energy levels explain how electrons are arranged around the nucleus and why elements behave the way they do. Understanding these arrangements is the foundation for predicting how atoms bond and react with each other.

Electron Arrangement Principles

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Fundamental Principles of Electron Configuration

Three rules govern how electrons fill orbitals in an atom:

The Aufbau Principle says electrons fill orbitals starting from the lowest energy and working up. The filling order is:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

Notice that 4s fills before 3d. That's because 4s is actually lower in energy than 3d due to how electrons interact with each other and the nucleus. This kind of overlap between energy levels is a common source of confusion.

The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers. In practice, this means each orbital holds a maximum of two electrons, and those two must have opposite spins (one "up," one "down").

Hund's Rule requires electrons to spread out among orbitals of equal energy before any pairing occurs. Think of it like passengers on a bus: everyone takes their own seat before anyone sits next to a stranger. This maximizes the number of unpaired electrons in a sublevel.

Electron configuration notation uses numbers, letters, and superscripts to show where all the electrons are. For example, carbon (6 electrons) is written as $1s^2 \, 2s^2 \, 2p^2$ . The number is the energy level, the letter is the sublevel, and the superscript is how many electrons are in that sublevel.

Applications and Exceptions

Noble gas shorthand simplifies long configurations by replacing the inner-shell electrons with the symbol of the preceding noble gas in brackets. Instead of writing out all 18 inner electrons for calcium, you write $[Ar] \, 4s^2$ .

A few transition metals break the Aufbau pattern because half-filled and fully filled d sublevels are unusually stable:

Chromium (Cr): Expected $[Ar] \, 4s^2 \, 3d^4$ , but actual is $[Ar] \, 4s^1 \, 3d^5$ (half-filled d sublevel)
Copper (Cu): Expected $[Ar] \, 4s^2 \, 3d^9$ , but actual is $[Ar] \, 4s^1 \, 3d^{10}$ (fully filled d sublevel)

Orbital diagrams give a visual version of electron configurations. Each orbital is drawn as a box, and electrons are shown as arrows (↑ for spin-up, ↓ for spin-down). These diagrams make it easy to see unpaired electrons and verify that Hund's rule is being followed.

Electron Energy Levels

Structure of Atomic Energy Levels

Energy levels (also called shells) are numbered outward from the nucleus: $n = 1, 2, 3, ...$ The principal quantum number $n$ determines the main energy level. Higher $n$ means higher energy and greater average distance from the nucleus. For naturally occurring elements, $n$ ranges from 1 to 7.

Each energy level contains one or more sublevels, and each sublevel contains a specific number of orbitals:

s sublevel: 1 orbital (holds 2 electrons)
p sublevel: 3 orbitals (holds 6 electrons)
d sublevel: 5 orbitals (holds 10 electrons)
f sublevel: 7 orbitals (holds 14 electrons)

The maximum number of electrons in any energy level is $2n^2$ . So the first shell holds 2, the second holds 8, the third holds 18, and so on.

Orbitals describe the region of space where you're most likely to find an electron. They have distinct shapes: s orbitals are spherical, p orbitals are dumbbell-shaped, and d orbitals have more complex shapes (often described as cloverleaf). These aren't fixed paths; the quantum mechanical model treats electrons as probability clouds rather than particles orbiting in neat circles.

Fundamental Principles of Electron Configuration, Electronic Structure of Atoms (Electron Configurations) | Chemistry

Electron Behavior and Energy States

Normally, electrons occupy the lowest available energy levels. This is called the ground state. When an atom absorbs energy (from heat, light, or electricity), an electron can jump to a higher energy level, creating an excited state. That excited electron is unstable, so it quickly drops back down and releases the extra energy as a photon of light. This is exactly why heated elements glow with characteristic colors.

The energy of orbitals generally increases as both $n$ and the sublevel type increase. The overlap between levels (like 4s filling before 3d) happens because electron-electron repulsion shifts the relative energies of certain sublevels.

Electron Types

Valence Electrons and Chemical Behavior

Valence electrons are the electrons in the outermost energy level. They're the ones that participate in chemical bonding and determine an element's reactivity.

You can figure out how many valence electrons a main-group element has by looking at its group number on the periodic table:

Group 1 (alkali metals): 1 valence electron
Group 2 (alkaline earth metals): 2 valence electrons
Groups 13–18: the number of valence electrons equals the group number minus 10 (so Group 15 has 5)
Group 18 (noble gases): 8 valence electrons (except helium, which has 2)

Elements tend to gain, lose, or share electrons to reach a full outer shell, mimicking the stable electron configuration of the nearest noble gas. This drive toward 8 valence electrons is often called the octet rule.

Electron dot diagrams (Lewis dot structures) are a quick way to represent valence electrons. You write the element's symbol and place dots around it for each valence electron. For example, nitrogen has 5 valence electrons, so its dot diagram shows 5 dots arranged around the N symbol.

Core Electrons and Atomic Properties

Core electrons are all the electrons in the inner shells, below the valence level. They don't typically participate in bonding, but they play a major role in determining atomic properties.

Core electrons create a shielding effect: they sit between the nucleus and the valence electrons, partially blocking the nuclear charge. Because of this shielding, valence electrons don't feel the full positive pull of the nucleus. What they do feel is called the effective nuclear charge ( $Z_{eff}$ ).

These two ideas explain important periodic trends:

Across a period (left to right): electrons are added to the same shell, so shielding stays roughly constant, but the nucleus gains protons. Effective nuclear charge increases, pulling valence electrons closer. Result: smaller atomic radius and higher ionization energy.
Down a group (top to bottom): each new row adds another shell of core electrons, increasing shielding. Valence electrons are farther from the nucleus and held less tightly. Result: larger atomic radius and lower ionization energy.

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