Electron Configuration Rules

Why This Matters

Electron configuration isn't just about memorizing orbital filling orders—it's the foundation for understanding why atoms behave the way they do. When you're tested on atomic physics, you're being asked to connect the arrangement of electrons to everything from periodic trends, chemical bonding, spectral lines, and magnetic properties. The rules governing electron configuration explain why certain elements are reactive while others are inert, why transition metals have unusual properties, and why atoms absorb or emit specific wavelengths of light.

Here's the key insight: these rules exist because nature minimizes energy and avoids electron crowding. Every principle you'll learn—Aufbau, Pauli, Hund's—reflects electrons finding the most stable arrangement possible. Don't just memorize which orbital fills first; understand why it fills first and what happens when electrons break the rules. That conceptual understanding is what separates a 3 from a 5 on exam day.

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Energy Minimization Principles

Electrons arrange themselves to achieve the lowest possible energy state. This drive toward stability governs the entire filling process and explains most of what you need to know about ground-state configurations.

Aufbau Principle

• Electrons fill orbitals from lowest to highest energy—this creates the ground state configuration you'll use for most problems
• Energy determines filling order, not distance from nucleus—the 4s orbital fills before 3d because it has lower energy, despite being farther out
• Predicts ground state configurations for most elements, though some transition metals are exceptions you'll need to memorize

Ground State vs. Excited State Configurations

• Ground state is the lowest energy arrangement—this is the default configuration you write for any element
• Excited states occur when electrons absorb energy and jump to higher orbitals, which is directly testable in questions about atomic spectra
• Electrons in excited states are unstable and will emit photons when returning to ground state—connecting configuration to emission spectra

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Electron Exclusion and Distribution

These rules explain how electrons avoid each other within the same atom. Electrons repel each other due to their negative charge, so nature has built-in rules to keep them separated.

Pauli Exclusion Principle

• No two electrons can share the same four quantum numbers—this is a fundamental law, not a tendency
• Each orbital holds maximum two electrons with opposite spins—represented as $\uparrow\downarrow$ in orbital diagrams
• Explains why electrons spread across orbitals rather than all piling into the lowest energy level

Hund's Rule

• Electrons occupy degenerate orbitals singly before pairing—think of it like passengers on a bus taking empty seats before sitting next to someone
• Single electrons maintain parallel spins (all $\uparrow$ or all $\downarrow$), which minimizes repulsion and creates magnetic properties
• Half-filled subshells are surprisingly stable—this explains exceptions like chromium's configuration

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Orbital Structure and Capacity

Understanding orbital types and their electron capacities lets you quickly determine configurations without memorizing every element. The capacity of each orbital type comes from quantum mechanics and the allowed values of quantum numbers.

Electron Capacity of Orbitals

• s = 2, p = 6, d = 10, f = 14 electrons maximum—memorize these; they appear constantly in configuration problems
• Capacity follows the formula $2(2l + 1)$ where $l$ is the angular momentum quantum number (s = 0, p = 1, d = 2, f = 3)
• Each subshell contains $(2l + 1)$ orbitals, and each orbital holds 2 electrons—this connects orbital shape to electron capacity

Orbital Filling Order

• Standard sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d...—the diagonal rule or an energy diagram helps you remember this
• 4s fills before 3d because 4s has slightly lower energy in multi-electron atoms—this trips up many students
• After ionization, 4s electrons leave before 3d—the filling order reverses when removing electrons from transition metals

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Electron Classification and Notation

These concepts help you quickly identify which electrons matter for chemistry and how to write configurations efficiently. Distinguishing between electron types is essential for predicting bonding and reactivity.

Valence Electrons and Core Electrons

• Valence electrons occupy the outermost shell and determine chemical bonding, reactivity, and periodic group placement
• Core electrons are in filled inner shells and shield valence electrons from the full nuclear charge
• Group number often equals valence electron count for main group elements—a quick check for your configurations

Noble Gas Configuration Shorthand

• Use the previous noble gas in brackets to replace core electrons—$[Ar]$ replaces $1s^2 2s^2 2p^6 3s^2 3p^6$
• Only valence and outer electrons are written explicitly—this highlights the chemically relevant electrons
• Example: Calcium is $[Ar]4s^2$ instead of writing all 20 electrons—faster and cleaner for exam work

Shorthand Notation for Electron Configuration

• Condenses full configurations using noble gas cores—essential for elements beyond the second period
• Superscripts indicate electron count in each subshell: $3d^{10}$ means 10 electrons in the 3d subshell
• Practice converting between full and shorthand—exams test both directions

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Exceptions and Special Cases

Not every element follows the standard rules. These exceptions occur because half-filled and fully-filled d subshells provide extra stability through exchange energy.

Exceptions to the Aufbau Principle

• Chromium is $[Ar]3d^5 4s^1$, not $[Ar]3d^4 4s^2$—the half-filled d subshell provides extra stability
• Copper is $[Ar]3d^{10} 4s^1$, not $[Ar]3d^9 4s^2$—the fully-filled d subshell is more stable than expected
• Other exceptions include Mo, Ag, Au, and Pt—transition metals in groups 6 and 11 are most likely to deviate

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Quick Reference Table

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Self-Check Questions

1. Which two principles both serve to minimize electron-electron repulsion, and how do they accomplish this differently?

2. An element has the configuration $[Kr]4d^{10}5s^2 5p^4$. How many valence electrons does it have, and which electrons would be lost first during ionization?

3. Compare and contrast the electron configurations of Cr and Mn—why does chromium break the expected pattern while manganese does not?

4. If an atom absorbs a photon and an electron moves from 3p to 4s, is the resulting configuration a ground state or excited state? What happens when the atom returns to its lowest energy arrangement?

5. Using Hund's Rule, draw the orbital diagram for nitrogen's 2p subshell. Why do the three electrons remain unpaired rather than filling one orbital completely?