English:Electron Configuration
Electron Configuration
Introduction
Electron configuration describes how the electrons of an atom or ion are distributed among shells, atomic orbitals, and subshells. At Grades 11–13 level, it is a bridge between the Periodic table, Quantum mechanics, Chemical bonding, periodic trends, ion formation, and magnetism. You will learn not only how to write configurations, but also how to explain why the notation works, where simple filling rules have exceptions, and how configurations help you predict chemical behavior.
By the end of this aiMOOC, you should be able to interpret the quantum numbers used to label orbitals, apply the Aufbau principle, the Pauli exclusion principle, and Hund's rule, write full and noble-gas configurations for atoms and common ions, recognize important transition-metal exceptions, connect configurations with the structure of the periodic table, and use unpaired electrons to reason about magnetic behavior.
The diagram above links the organization of the periodic table with electron configurations. As you work through the course, return to it and ask: Which subshell is being filled in each region of the table?
The Crash Course video gives a broad introduction to electrons, shells, orbitals, electron configurations, ionization, electron affinity, and the periodic table. Use it as an overview before studying the rules in detail.
From Shells to Orbitals
In an early shell model, electrons are grouped by their principal energy level. Modern atomic theory goes further: an electron in an atom is described by a quantum-mechanical state, and an atomic orbital is a spatial probability distribution associated with such a state. An orbital is not a small circular path around the nucleus. It is a mathematical description that tells you where an electron is likely to be found.
The principal quantum number n identifies the main shell and has positive integer values 1, 2, 3, and so on. Within a shell, orbitals are grouped into subshells labeled s, p, d, and f. These letters correspond to the orbital angular momentum quantum number l. For the chemistry normally studied at this level, s corresponds to l = 0, p to l = 1, d to l = 2, and f to l = 3.
The image shows probability-density clouds for representative s, p, d, and f orbitals. The shapes become more complex as the orbital angular momentum increases. These pictures are useful models, but remember that the electron is not a tiny object traveling along the surface of a colored lobe.
Quantum Numbers
Four quantum numbers are used to distinguish electron states in an atom. The principal quantum number n gives the shell. The orbital angular momentum quantum number l identifies the subshell. The magnetic quantum number ml identifies a particular orbital within a subshell and takes integer values from −l to +l. The spin quantum number ms can be +1/2 or −1/2.
These rules explain the number of orbitals in a subshell. An s subshell has one orbital, a p subshell has three, a d subshell has five, and an f subshell has seven. Since each orbital can contain at most two electrons, the corresponding electron capacities are 2, 6, 10, and 14.
| Subshell | Number of orbitals | Maximum electrons | Typical appearance in configurations |
|---|---|---|---|
| s | 1 | 2 | 1s², 2s², 3s² |
| p | 3 | 6 | 2p⁶, 3p⁶, 4p⁶ |
| d | 5 | 10 | 3d¹⁰, 4d¹⁰ |
| f | 7 | 14 | 4f¹⁴, 5f¹⁴ |
The s and p orbitals illustrate how one subshell can contain several orbitals with the same n and l values but different orientations or magnetic quantum numbers.
Shell Capacity and Energy Ordering
The maximum number of electrons allowed by the quantum numbers in shell n is 2n². This gives maximum shell capacities of 2, 8, 18, and 32 for the first four shells. However, maximum capacity is not the same as ground-state filling order. In many-electron atoms, subshell energies overlap: for example, 4s is occupied before 3d when building the neutral atoms potassium and calcium, even though 4s belongs to a higher principal shell.
This distinction matters. A shell-based picture is useful for counting electrons by principal level, but an electron configuration is usually written by subshell because chemistry depends strongly on the energies and occupancies of particular orbitals.
The shell diagram of titanium is a useful comparison model. It groups electrons by shell, while the more detailed ground-state configuration is [Ar] 3d² 4s².
Rules for Ground-State Electron Configurations
For many atoms, you can construct a useful ground-state electron configuration by applying three central ideas: the Aufbau principle, the Pauli exclusion principle, and Hund's rule. These rules are powerful, but they are models of the observed energy ordering rather than a complete solution of the many-electron Schrödinger equation.
Aufbau Principle and the Madelung Order
Aufbau means building up. In the simplest school-level application, electrons are added to the available orbitals in order of increasing energy. A commonly used order is:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p.
A compact way to remember much of this sequence is the Madelung rule: subshells are generally filled in order of increasing n + l, and when two subshells have the same value of n + l, the one with smaller n is usually filled first. This rule predicts many ground-state configurations well, but important exceptions exist, especially among transition and inner-transition elements.
The Khan Academy video introduces Aufbau filling, orbital notation, and the Pauli exclusion principle using the first-period elements.
Pauli Exclusion Principle
The Pauli exclusion principle states that no two electrons in the same atom can have the same set of four quantum numbers. In practical orbital-box notation, this means that one orbital can hold at most two electrons, and if two electrons occupy the same orbital, they must have opposite spins.
For example, a filled 2s orbital is represented by two opposite arrows in one box. Writing two parallel arrows in the same orbital would violate the Pauli exclusion principle.
Hund's Rule
Hund's rule applies when several orbitals of the same subshell have equal energy in an isolated atom. Electrons occupy these degenerate orbitals singly with parallel spins before pairing occurs. For carbon, the two 2p electrons therefore occupy two different 2p orbitals rather than pairing in one orbital.
This arrangement is associated with a lower-energy state for the isolated atom because of the combined effects of electron-electron interactions and exchange. Hund's rule is especially useful when you count unpaired electrons and predict whether an atom or ion is paramagnetic.
Writing Electron Configurations
A full electron configuration lists occupied subshells and shows the number of electrons in each subshell as a superscript. For example, nitrogen has atomic number 7, so a neutral nitrogen atom has seven electrons:
N: 1s² 2s² 2p³
The superscripts must add up to the total number of electrons. This simple check catches many mistakes.
Orbital-box diagrams add information about spin and unpaired electrons. In a box diagram, each orbital is shown separately and each electron is shown as an arrow. This representation makes the Pauli and Hund rules visible rather than merely implied.
Worked Examples for Main-Group Atoms
Oxygen, Z = 8: 1s² 2s² 2p⁴. The 2p subshell contains four electrons. Hund's rule gives two unpaired electrons in the isolated ground-state atom.
Sodium, Z = 11: 1s² 2s² 2p⁶ 3s¹. The outer 3s electron is the valence electron most readily removed when Na⁺ forms.
Chlorine, Z = 17: 1s² 2s² 2p⁶ 3s² 3p⁵. The valence-shell pattern 3s² 3p⁵ helps explain why chlorine commonly gains one electron to reach a noble-gas configuration.
Calcium, Z = 20: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². Both 4s electrons are typically lost when Ca²⁺ forms.
Noble-Gas Shorthand
For larger atoms, you can replace the filled inner-shell configuration by the symbol of the preceding noble gas in square brackets. For example:
Na: [Ne] 3s¹
Cl: [Ne] 3s² 3p⁵
Ca: [Ar] 4s²
Fe: [Ar] 3d⁶ 4s²
Noble-gas shorthand emphasizes the valence and near-valence electrons that are often most relevant to chemical behavior. It is a notation shortcut; it does not mean that the inner electrons have disappeared.
The Periodic Table as an Electron-Configuration Map
The structure of the Periodic table reflects recurring electron configurations. The table can be divided into s, p, d, and f blocks according to the type of subshell being filled.
The s block includes Groups 1 and 2 and helium by configuration. The p block includes Groups 13–18 except helium. The d block contains the transition elements, where a d subshell is being filled. The f block contains the lanthanides and actinides, where f subshells are being filled.
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The periodic table also helps you reconstruct configurations without memorizing the entire filling sequence. Moving across a block tells you how many electrons are being added to the corresponding subshell.
This Khan Academy video develops the periodic-table method for electron configurations and noble-gas shorthand.
Valence-Electron Patterns
Main-group elements show especially clear patterns. Group 1 elements generally end in ns¹, Group 2 in ns², Group 17 in ns² np⁵, and Group 18 in ns² np⁶, with helium as the special case 1s².
These patterns help explain similarities within groups. Elements with similar valence configurations often form similar ions and types of bonds. Electron configuration therefore provides a microscopic explanation for periodic repetition in chemical properties.
Transition Metals, Exceptions, and Ions
The simple filling diagram is a starting point, not an infallible law. In transition metals, the relative energies of ns and n−1 d orbitals are close and change as nuclear charge and electron occupancy change. As a result, some observed ground states differ from the simplest Aufbau prediction.
Two important examples are:
Chromium: [Ar] 3d⁵ 4s¹ rather than the simple prediction [Ar] 3d⁴ 4s².
Copper: [Ar] 3d¹⁰ 4s¹ rather than the simple prediction [Ar] 3d⁹ 4s².
These configurations reflect the actual energy balance in many-electron atoms. Avoid explaining them only by saying that half-filled or filled subshells are always automatically stable; electron-electron interactions and orbital energies are part of the full explanation.
The tutorial above provides additional practice, including transition-metal ions.
Forming Cations and Anions
For a main-group cation, remove electrons from the orbitals with the largest principal quantum number first. Sodium therefore changes from [Ne] 3s¹ to Na⁺: [Ne]. Calcium changes from [Ar] 4s² to Ca²⁺: [Ar].
For anions, add electrons according to the available orbitals. Oxygen is 1s² 2s² 2p⁴, so oxide is O²⁻: 1s² 2s² 2p⁶ = [Ne].
Transition-metal ions need special care. Although 4s is filled before 3d in the neutral atoms potassium and calcium and is written in many Aufbau sequences before 3d, 4s electrons are removed before 3d electrons when common transition-metal cations form. Thus:
Fe: [Ar] 3d⁶ 4s²
Fe²⁺: [Ar] 3d⁶
Fe³⁺: [Ar] 3d⁵
Cu: [Ar] 3d¹⁰ 4s¹
Cu⁺: [Ar] 3d¹⁰
Cu²⁺: [Ar] 3d⁹
The key lesson is that orbital-energy ordering can change with electron occupancy and ionization. Use experimentally established configurations for exceptions rather than forcing every atom into one rigid filling diagram.
Electron Configuration and Magnetism
An atom or ion with one or more unpaired electrons is usually paramagnetic: it is attracted into an applied magnetic field. A species in which all electrons are paired is diamagnetic: it is weakly repelled by an applied magnetic field.
Orbital-box diagrams make this easy to predict. For example, N with 2p³ has three unpaired electrons. Ne with 2p⁶ has no unpaired electrons. Fe³⁺ with 3d⁵ has five unpaired electrons in the free-ion high-spin picture, although the behavior of transition-metal ions in compounds can also depend on ligand-field effects.
This connection is important because it turns electron configuration into an experimentally testable model. Magnetic measurements can reveal information about electron pairing.
Electron Configuration, Spectra, and Chemical Behavior
Electron configurations are not only bookkeeping. They help explain why atoms absorb and emit particular energies, why elements in the same group often react similarly, why ionization energies show periodic patterns, and why transition metals can have several common oxidation states.
When an atom absorbs energy, an electron can be promoted to a higher-energy state. When the atom returns to a lower-energy state, it may emit a photon. The energy difference between states determines the photon energy. In real atoms, spectra involve more detail than a single configuration because electron-electron interactions, spin-orbit coupling, and selection rules also matter, but configuration language remains a foundation for interpreting atomic structure.
The most useful habit is to connect three representations: periodic-table position, electron-configuration notation, and orbital occupancy. If these three agree, your reasoning is usually on solid ground.
Interactive Tasks
Quiz: Test Your Knowledge
What does an electron configuration describe? (The distribution of electrons among atomic orbitals) (!The number of neutrons in the nucleus) (!The path of an electron around the nucleus) (!The mass of an isotope)
What is the maximum number of electrons in a p subshell? (Six) (!Two) (!Ten) (!Fourteen)
Which is the ground-state electron configuration of carbon? (1s2 2s2 2p2) (!1s2 2s2 2p6) (!1s2 2s1 2p3) (!1s2 2p4)
What does the Pauli exclusion principle require for two electrons in the same orbital? (They must have opposite spins) (!They must have parallel spins) (!They must be in different shells) (!They must have the same four quantum numbers)
What does Hund's rule predict for three electrons entering three equal-energy p orbitals? (They occupy separate orbitals with parallel spins) (!They pair in the first orbital before using the others) (!They all enter the same orbital) (!They must occupy different principal shells)
Which element has the ground-state configuration [Ne] 3s² 3p⁵? (Chlorine) (!Sodium) (!Magnesium) (!Argon)
Which configuration is the observed ground state of chromium? (Argon core 3d5 4s1) (!Argon core 3d4 4s2) (!Argon core 3d6) (!Neon core 3s2 3p6)
What is the ground-state configuration of Fe²⁺? (Argon core 3d6) (!Argon core 3d4 4s2) (!Argon core 3d6 4s2) (!Argon core 3d8)
Which subshell is being filled across the first transition series from scandium to zinc? (3d) (!3p) (!4p) (!4f)
What makes an isolated atom or ion paramagnetic? (The presence of one or more unpaired electrons) (!A completely filled valence shell only) (!An even atomic number) (!A positive ionic charge)
Memory Game
| Aufbau principle | Electrons are placed into available orbitals in a generally increasing energy order |
| Pauli exclusion principle | No two electrons in one atom can have the same four quantum numbers |
| Hund's rule | Degenerate orbitals are singly occupied with parallel spins before pairing |
| Noble-gas shorthand | Core electrons are represented by the symbol of a preceding noble gas |
| Valence electron | Electron in the outer or chemically active part of an atom |
| Paramagnetism | Attraction associated with the presence of unpaired electrons |
| Subshell | Set of orbitals sharing the same principal and angular momentum quantum numbers |
| Orbital | Quantum state associated with a spatial probability distribution for an electron |
Drag and Drop
| Match the correct terms. | Topic |
|---|---|
| Opposite spins in one orbital | Pauli exclusion principle |
| Single occupancy before pairing | Hund's rule |
| Use of a noble-gas core symbol | Condensed configuration |
| Removal of 4s electrons before 3d electrons | Transition-metal cations |
| Unpaired electrons cause attraction to a field | Paramagnetism |
...
Crossword Puzzle
| Orbital | What quantum-mechanical region can hold up to two electrons? |
| Aufbau | What building-up principle gives a common orbital filling sequence? |
| Hund | Which rule says degenerate orbitals are singly occupied before pairing? |
| Pauli | Which exclusion principle limits an orbital to two electrons with opposite spins? |
| Valence | What word describes electrons most directly involved in many chemical reactions? |
| Paramagnetic | What describes a species with unpaired electrons that is attracted to a magnetic field? |
LearningApps
Cloze Text
Open-Ended Tasks
Easy
- Electron configuration chart: Create a one-page visual guide showing the s, p, d, and f subshell capacities and the common filling order; include at least three example atoms and explain your design choices.
- Orbital box poster: Draw orbital-box diagrams for carbon, nitrogen, oxygen, and neon, then annotate where the Pauli exclusion principle and Hund's rule are visible.
- Periodic table color map: Produce an image or digital graphic that marks the s, p, d, and f blocks of the periodic table and explains what each block means for electron configuration.
- Configuration interview: Interview a classmate about how they decide which subshell fills next, record the explanation in clear English, and write a short reflection on one idea you would improve.
Standard
- Ion configuration investigation: Build a comparison table for Na, Na⁺, O, O²⁻, Fe, Fe²⁺, Fe³⁺, Cu, Cu⁺, and Cu²⁺, then explain the different electron-removal rules for main-group and transition-metal ions.
- Aufbau demonstration: Design a physical or digital model that lets users place electrons into orbital boxes and receive feedback when they violate Pauli or Hund rules; test it with another learner.
- Magnetism prediction project: Choose six atoms or monatomic ions, predict the number of unpaired electrons from orbital diagrams, classify each as paramagnetic or diamagnetic, and explain how a laboratory magnetic test could check the prediction.
- Chemistry explainer video: Produce a three-to-five-minute video that teaches noble-gas shorthand and includes one main-group example, one transition-metal example, and one ion.
Advanced
- Chromium and copper exception study: Research why chromium and copper differ from the simplest Aufbau predictions, compare at least two reliable chemistry sources, and present a reasoned explanation that avoids the claim that half-filled shells are always automatically stable.
- Spectroscopy connection: Investigate how electron transitions produce atomic emission or absorption spectra, select one element with an accessible spectrum, and connect observed spectral lines to changes in electronic energy.
- Transition-metal field study: Visit a school laboratory, university laboratory, science museum, or virtual laboratory resource and document examples of transition-metal compounds; relate observed colors or magnetic behavior to electron configuration while clearly separating atomic configurations from ligand-field effects.
- Configuration data project: Create a small dataset of ground-state configurations for at least thirty elements, identify exceptions to the simplest Madelung order, visualize where the exceptions occur in the periodic table, and propose a rule-based algorithm with clearly stated limits.
Learning Assessment
- Configuration reasoning: Given an unfamiliar element or ion, derive its electron count, write a justified ground-state configuration, and explain each decision using the periodic table and filling rules.
- Error diagnosis: Analyze a set of deliberately incorrect configurations, identify whether each error concerns electron count, orbital capacity, Pauli exclusion, Hund's rule, ionization order, or an Aufbau exception, and correct it.
- Representation transfer: Convert the same species among full configuration, noble-gas shorthand, orbital-box notation, and periodic-table position, then explain what information is gained or lost in each representation.
- Magnetic prediction: Use configurations and orbital diagrams to predict relative numbers of unpaired electrons for several atoms or ions and justify which should be paramagnetic or diamagnetic.
- Periodic explanation: Explain a selected periodic trend by connecting valence configurations with position in the periodic table, while noting at least one limitation of a simple configuration-only explanation.
- Model evaluation: Compare the Aufbau diagram with the more nuanced energy behavior of transition-metal atoms and ions, then argue where the model is useful and where experimental configurations must take priority.
Evidence of Learning
Knowledge: You can explain shells, subshells, orbitals, quantum numbers, electron capacities, Aufbau ordering, Pauli exclusion, Hund's rule, noble-gas shorthand, common exceptions, ion configurations, and the link between unpaired electrons and magnetism.
Skills: You can count electrons in atoms and ions, write and check configurations, construct orbital-box diagrams, use the periodic table as a configuration map, identify unpaired electrons, compare representations, diagnose errors, and communicate reasoning clearly.
Products: Strong evidence may include annotated configuration charts, orbital diagrams, a periodic-table map, a digital model, a short explainer video, a spectroscopy report, a magnetism investigation, or a documented dataset of configurations.
Transfer achievements: You can apply configuration ideas to unfamiliar elements and ions, relate microscopic electron arrangements to periodic behavior, use magnetic or spectroscopic evidence to test predictions, and recognize when a simplified filling rule is no longer sufficient.
OERs on the Topic
The English Wikipedia article below provides a broad reference on shells, notation, the Aufbau principle, the periodic table, exceptions, ionization, and applications. Use it as a reference source and compare its explanations with your course notes and other reliable chemistry materials.
Linked Learning Areas
Electron configuration connects atomic structure with chemical patterns. Understanding it supports later work on bonding, spectroscopy, transition-metal chemistry, periodic trends, and introductory quantum chemistry.
