Difficulty: Intermediate | Prerequisites: Chapter 2 (atomic structure, periodic table groups)
Chapters 7 and 8 explain why the periodic table is shaped the way it is. Electron configurations tell you how electrons are arranged in an atom, and those arrangements drive every periodic trend: atomic size, ion size, and the charges ions prefer. This material connects atomic structure (Chapter 2) to bonding (Chapter 9). If you understand electron configurations, predicting ion charges and comparing atomic radii becomes systematic rather than memorisation.
Electrons fill orbitals in a specific order (1s, 2s, 2p, 3s, ...). The outermost (valence) electrons determine an element's chemistry. Periodic trends in atomic radius, ionic radius, and effective nuclear charge follow predictable patterns across periods and down groups. Ions form by gaining or losing electrons to reach a noble-gas configuration, and species with the same electron count are isoelectronic.
Electron configuration
The complete listing of which orbitals an atom's electrons occupy and how many electrons are in each. Written as 1s² 2s² 2p⁶ 3s¹ for sodium, for example. Think of it as the atom's electron address book.
Abbreviated (noble-gas) electron configuration
A shorthand that replaces the inner-shell electrons with the symbol of the preceding noble gas in brackets. Sodium: [Ne] 3s¹. This highlights only the valence shell.
Aufbau (building-up) principle
Electrons fill the lowest-energy orbitals first before occupying higher-energy ones. The filling order does not simply follow the principal quantum number; 4s fills before 3d, for instance.
Valence electrons
The electrons in the outermost principal energy level. For main-group elements, the number of valence electrons equals the group number (using the 1-through-18 numbering, take the ones digit for groups 13 through 18). These are the electrons involved in chemical bonding.
Core electrons
All the electrons that are not valence electrons. They sit in filled inner shells and are generally not involved in bonding. In the abbreviated configuration, the noble-gas core represents these.
Effective nuclear charge (Z_eff)
The net positive charge felt by a valence electron after accounting for the shielding effect of inner-shell electrons. Z_eff = Z (atomic number) minus the shielding from core electrons. In simple terms, it is how strongly the nucleus "pulls" on the outer electrons.
Atomic radius
Half the distance between the nuclei of two bonded atoms of the same element. Increases going down a group (more electron shells). Decreases going across a period left to right (higher Z_eff pulls electrons closer).
Ionic radius
The radius of an ion. Cations are smaller than their parent atoms (fewer electrons, same nuclear charge). Anions are larger than their parent atoms (more electrons, same nuclear charge).
Isoelectronic
Species that have the same number of electrons and therefore the same electron configuration. For example, Na⁺, Ne, and F⁻ all have 10 electrons. Among isoelectronic species, the one with more protons is smaller (stronger nuclear pull).
Cation
A positive ion formed by losing electrons. Metals form cations. When writing the electron configuration of a cation, remove electrons from the highest principal energy level first (for transition metals, remove from the s orbital before the d).
Anion
A negative ion formed by gaining electrons. Nonmetals form anions. Electrons are added to the lowest-energy available orbital.
Orbital filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p.
Use the diagonal rule (draw an arrow chart) if you have not yet memorised this.
Complete configuration: list every subshell with its electron count. Example: Fe (Z = 26) is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶.
Abbreviated configuration: replace the core with the noble gas that precedes the element. Fe: [Ar] 4s² 3d⁶.
Valence electrons for main-group elements equal the group number (Group 1 = 1, Group 2 = 2, Group 13 = 3, Group 14 = 4, Group 15 = 5, Group 16 = 6, Group 17 = 7, Group 18 = 8).
Core electrons are everything inside the noble-gas bracket.
The "chemically important" electrons are the valence electrons: they determine bonding behaviour and chemical reactivity.
Ions form when atoms gain or lose electrons to achieve a noble-gas electron configuration (full outer shell).
Cation configurations: remove electrons from the highest n (principal quantum number) first.
For transition metals, remove s electrons before d electrons. Fe²⁺: [Ar] 3d⁶ (remove the two 4s electrons from [Ar] 4s² 3d⁶). Fe³⁺: [Ar] 3d⁵.
Anion configurations: add electrons to the next available orbital. Cl⁻: [Ne] 3s² 3p⁶ = [Ar].
Effective nuclear charge (Z_eff) increases across a period because protons are added to the nucleus while shielding from core electrons stays roughly constant. This stronger pull draws electrons closer.
Atomic radius trends:
Decreases across a period (left to right): rising Z_eff contracts the electron cloud.
Increases down a group: each new period adds an electron shell, pushing valence electrons farther from the nucleus.
Ionic radius trends:
Cations are smaller than their parent atoms (lost electrons, same nuclear charge).
Anions are larger than their parent atoms (gained electrons, same nuclear charge).
Within an isoelectronic series, the species with more protons is smaller (e.g. Al³⁺ < Mg²⁺ < Na⁺ < Ne < F⁻ < O²⁻).
Isoelectronic species share the same electron count and configuration. To compare their sizes, look at the nuclear charge: more protons = smaller radius.
Orbital filling order (diagonal rule)
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p
Maximum electrons per subshell
Subshell | Max electrons |
|---|---|
s | 2 |
p | 6 |
d | 10 |
f | 14 |
Abbreviated configuration examples
Element | Full configuration | Abbreviated |
|---|---|---|
Na (Z=11) | 1s² 2s² 2p⁶ 3s¹ | [Ne] 3s¹ |
Cl (Z=17) | 1s² 2s² 2p⁶ 3s² 3p⁵ | [Ne] 3s² 3p⁵ |
Fe (Z=26) | 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ | [Ar] 4s² 3d⁶ |
Fe²⁺ | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ | [Ar] 3d⁶ |
Electron configurations explain why copper is a better conductor than iron (its single 4s electron is easily delocalised). The periodic trend in atomic radius is why lithium-ion batteries use lithium: it is the smallest, lightest metal, so it packs the most energy per unit mass.
Students often remove d electrons first when writing transition-metal cation configurations. You remove the electrons from the highest principal quantum number first: for Fe²⁺, remove the two 4s electrons, not the 3d electrons.
Students often assume that atomic radius increases across a period because more electrons are being added. The radius actually shrinks because the rising nuclear charge pulls the electron cloud tighter.
Students often confuse "isoelectronic" with "isotope." Isoelectronic species have the same electron count. Isotopes have the same proton count but different neutron counts.
Students often forget that cations are smaller and anions are larger than the parent atom. Losing an electron removes electron-electron repulsion and lets the nucleus pull the remaining electrons closer.
⚠️ Writing complete and abbreviated electron configurations for any atom or monatomic ion is a core exam skill. Practise both directions: configuration from element, and element from configuration.
⚠️ Transition-metal cation configurations (removing s before d) are a favourite trick question.
⚠️ Ranking atoms and ions by size using periodic trends and isoelectronic comparisons appears regularly.
⚠️ Counting valence electrons from a configuration or from the group number is the bridge to Lewis structures in Chapter 9.
True or False: The abbreviated electron configuration of potassium (K, Z=19) is [Ar] 4s¹. (True)
Fill in the blank: A neutral atom of oxygen has ______ valence electrons. (6)
True or False: Atomic radius increases from left to right across a period. (False; it decreases)
Fill in the blank: Na⁺, Ne, and F⁻ are examples of ______ species. (isoelectronic)
True or False: To write the electron configuration of Fe³⁺, you remove three 3d electrons from Fe. (False; remove the two 4s electrons first, then one 3d electron)
Q: Write the complete and abbreviated electron configurations for sulfur (S, Z=16).
A: Complete: 1s² 2s² 2p⁶ 3s² 3p⁴. Abbreviated: [Ne] 3s² 3p⁴.
Q: How many valence electrons does phosphorus (Group 15) have?
A: Five.
Q: Write the electron configuration for Cr³⁺ (Z=24). Note: Cr is an exception; its ground state is [Ar] 4s¹ 3d⁵.
A: Remove the 4s electron first, then two 3d electrons: [Ar] 3d³.
Q: Rank these in order of increasing atomic radius: Na, Mg, Al.
A: Al < Mg < Na. All are in Period 3; moving left increases the radius because Z_eff decreases.
Q: Rank these isoelectronic species in order of increasing ionic radius: O²⁻, F⁻, Na⁺, Mg²⁺.
A: Mg²⁺ < Na⁺ < F⁻ < O²⁻. All have 10 electrons. More protons = stronger pull = smaller radius.
Electron configurations feed directly into Chapter 9: the number of valence electrons determines how many bonds an atom forms, how many lone pairs it carries, and whether bonding is ionic or covalent. Periodic trends in electronegativity (also Chapter 9) follow the same logic as atomic radius trends. The ion charges you predict here are exactly what you need for naming ionic compounds (Chapter 2).
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