Electronic Configuration & Filling Rules Why electronic configuration matters Electronic configuration tells you where electrons live in an atom — which shells, subshells, and orbitals. It explains reactivity, bonding, colors, and magnetism. For NEET, you must write ground‑state configurations fast (especially up to Z = 30), know the filling order, understand exceptions (Cr, Cu), and handle cations (which electrons are removed). Periodic table blocks: s (left), p (right), d (middle), f (bottom). Block identity follows the subshell where the last electron enters. Real‑life link: Unpaired electrons from specific configurations make many transition‑metal ions paramagnetic and colored — exploited in MRI contrast agents (e.g., Gd³⁺ with several unpaired 4f electrons) and in industrial catalysis. remember Quantum numbers → shells, subshells, orbitals (quick refresher) Think of an atom like a building: floors (shells, n), rooms on each floor (subshells, l), seats in each room (orbitals, m l ), and two students per seat with opposite caps (spins, m s ). This picture helps remember capacities and filling rules. Shell (n) Principal energy level. n = 1, 2, 3, ... Capacity given by 2n 2. Larger n means generally higher energy and larger size. Shape label within a shell. l = 0(s), 1(p), 2(d), 3(f). Each subshell holds a fixed number of orbitals. Subshell (l) Orbital ( m l ) 3D region where an electron is most likely found. Each orbital holds max 2 electrons with opposite spins. Spin ( m s ) Intrinsic property of an electron. Two allowed values, represented as up ( ) or down ( ) arrows in box diagrams. Degenerate orbitals Orbitals of the same subshell (same n and l) that have equal energy, e.g., the three 2p orbitals. Max electrons per shell A shell with principal quantum number n can accommodate up to 2n 2 electrons. Max electrons per subshell A subshell with azimuthal quantum number l has (2l+1) orbitals, each holding 2 electrons. Subshell Number of orbitals (2l+1) Max electrons (2(2l+1)) Type Subshells, orbitals, and capacities 10 14 Aufbau principle and the (n+l) rule Aufbau principle: In the ground state, electrons occupy orbitals in order of increasing energy. Practically, we use the (n+l) rule to rank subshell energies: the subshell with lower (n+l) fills first; if (n+l) is equal, the subshell with lower n fills first. This generates the familiar order used in NEET. (n+l) selection rule Energy ordering heuristic for ground‑state filling. Standard filling order to remember (through 7p) 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p Aufbau diagonal‑arrow guide: follow arrows to get 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p ... Order Subshell n+l (n+l) values for the main subshells (1s to 7p) 1s 2s 2p 3s 3p 4s 3d 4p 5s 10 4d 11 5p 12 6s 13 4f 14 5d 15 6p 16 7s 17 5f 18 6d 19 7p Pauli Exclusion Principle No two electrons in the same atom can have all four quantum numbers identical. Practically, an orbital can hold at most 2 electrons with opposite spins (one and one ). This restricts how we fill box diagrams and explains why capacities are what they are. remember Pauli + capacities: s holds 2, p holds 6, d holds 10, f holds 14 — derived from the number of degenerate orbitals times 2 spins each. Hund's Rule: spread out before pairing Hund's Rule of Maximum Multiplicity says: electrons occupy all degenerate orbitals singly with parallel spins before pairing. Analogy: empty bus seats — students take separate seats (same cap direction) before sharing a seat. This minimizes electron–electron repulsion and slightly stabilizes the atom. Box diagram for nitrogen (Z = 7): 2p shows (not pairing prematurely). Labels highlight Pauli and Hund. Hund = Hunt for empty seats first: fill one arrow (same direction) in each box of a subshell before any pairing. 3d fills before 4s for all species. For neutral atoms, 4s fills before 3d by the (n+l) rule. But when forming cations, 4s electrons are removed before 3d because 4s becomes higher in energy once 3d starts filling. Incorrect. In a given subshell (e.g., 2p), electrons occupy different orbitals singly with parallel spins first; only then does pairing occur. Hund's Rule can be ignored; you may pair in degenerate orbitals whenever. Write configurations fast: method and NEET traps Find Z (atomic number). Walk through the standard filling order (use diagonal arrows or (n+l)). Respect capacities: s2, p6, d10, f14; apply Pauli and Hund in box diagrams. Use noble‑gas shorthand to compress inner shells. Check for known exceptions (Cr, Cu; later Mo, Ag, Au, Pt, Pd). For cations, remove electrons from the highest n orbital first (4s before 3d for first‑row transition metals). Stepwise method High‑yield traps: (1) 4s fills before 3d in neutrals but is removed first in cations. (2) Cr is [Ar] 3d 5 4s 1 ; Cu is [Ar] 3d 10 4s 1 . (3) Don’t pair in p/d boxes until each is singly occupied (Hund). neet-alert Worked ground‑state configurations: Z = 1 to 30 Hydrogen 1s 1 1s 1 Helium He 1s 2 1s 2 Lithium Li [He] 2s 1 2s 1 Beryllium Be [He] 2s 2 2s 2 Boron [He] 2s 2 2p 1 2p 1 Carbon [He] 2s 2 2p 2 2p 2 Nitrogen [He] 2s 2 2p 3 2p 3 Oxygen [He] 2s 2 2p 4 2p 4 Fluorine [He] 2s 2 2p 5 2p 5 10 Neon Ne [He] 2s 2 2p 6 2p 6 11 Sodium Na [Ne] 3s 1 3s 1 12 Magnesium Mg [Ne] 3s 2 3s 2 13 Aluminium Al [Ne] 3s 2 3p 1 3p 1 14 Silicon Si [Ne] 3s 2 3p 2 3p 2 15 Phosphorus [Ne] 3s 2 3p 3 3p 3 16 Sulfur [Ne] 3s 2 3p 4 3p 4 17 Chlorine Cl [Ne] 3s 2 3p 5 3p 5 18 Argon Ar [Ne] 3s 2 3p 6 3p 6 19 Potassium [Ar] 4s 1 4s 1 20 Calcium Ca [Ar] 4s 2 4s 2 21 Scandium Sc [Ar] 3d 1 4s 2 3d 1 4s 2 22 Titanium Ti [Ar] 3d 2 4s 2 3d 2 4s 2 23 Vanadium [Ar] 3d 3 4s 2 3d 3 4s 2 24 Chromium Cr [Ar] 3d 5 4s 1 3d 5 4s 1 (exception) 25 Manganese Mn [Ar] 3d 5 4s 2 3d 5 4s 2 26 Iron Fe [Ar] 3d 6 4s 2 3d 6 4s 2 27 Cobalt Co [Ar] 3d 7 4s 2 3d 7 4s 2 28 Nickel Ni [Ar] 3d 8 4s 2 3d 8 4s 2 29 Copper Cu [Ar] 3d 10 4s 1 3d 10 4s 1 (exception) 30 Zinc Zn [Ar] 3d 10 4s 2 3d 10 4s 2 Element Symbol Configuration (shorthand) Valence subshell occupancy Electronic configurations (noble‑gas shorthand) up to Z = 30 Why exceptions occur: extra stability of half‑filled and fully‑filled d Observed configurations sometimes deviate from the naive Aufbau order because half‑filled ( d 5 ) and fully‑filled ( d 10 $) subshells gain stability (exchange energy, symmetry, and reduced electron–electron repulsion). The famous first‑row cases are Cr and Cu. Cr (24) [Ar] 3d 4 4s 2 [Ar] 3d 5 4s 1 Half‑filled d 5 stability Cu (29) [Ar] 3d 9 4s 2 [Ar] 3d 10 4s 1 Fully‑filled d 10 stability Mo (42) [Kr] 4d 4 5s 2 [Kr] 4d 5 5s 1 Half‑filled d 5 Ag (47) [Kr] 4d 9 5s 2 [Kr] 4d 10 5s 1 Fully‑filled d 10 Pd (46) [Kr] 4d 8 5s 2 [Kr] 4d 10 5s 0 Completely filled 4d Pt (78) [Xe] 4f 14 5d 8 6s 2 [Xe] 4f 14 5d 9 6s 1 Stabilized 5d 9 Au (79) [Xe] 4f 14 5d 9 6s 2 [Xe] 4f 14 5d 10 6s 1 Fully‑filled 5d 10 Element Predicted (Aufbau) Observed (ground state) Note Common exceptions (predicted vs observed) NEET‑friendly explanation: Say “extra stability for half‑filled and fully‑filled d” — do not dive into advanced relativistic or many‑electron corrections in the exam. tip Cations: remove 4s before 3d Once 3d starts filling, the 3d subshell becomes lower in energy than 4s in most transition‑metal atoms/ions. So, when forming cations, electrons are removed from 4s before 3d. This is heavily tested. Species Neutral atom config Ion formation Ion config Common cation configurations (ground state of ions) Fe²⁺ [Ar] 3d 6 4s 2 Remove 2 from 4s [Ar] 3d 6 Fe³⁺ [Ar] 3d 6 4s 2 Remove 2 from 4s, 1 from 3d [Ar] 3d 5 Cu²⁺ [Ar] 3d 10 4s 1 Remove 1 from 4s, 1 from 3d [Ar] 3d 9 Cu⁺ [Ar] 3d 10 4s 1 Remove 1 from 4s [Ar] 3d 10 Sc³⁺ [Ar] 3d 1 4s 2 Remove 2 from 4s, 1 from 3d [Ar] Zn²⁺ [Ar] 3d 10 4s 2 Remove 2 from 4s [Ar] 3d 10 neet-alert Don’t write Fe²⁺ as [Ar] 3d 4 4s 2 . The 4s electrons are lost first; Fe²⁺ is [Ar] 3d 6 . Degenerate orbitals, pairing, and magnetism Unpaired electrons in degenerate orbitals cause paramagnetism (attracted to a magnetic field); all electrons paired causes diamagnetism (weakly repelled). Quick check: count unpaired electrons from the box diagram based on Hund’s rule. Order these by filling using (n+l): 4p, 5s, 3d. Write configurations for Na, P, Ca, Sc, Fe, Cu, Zn. Explain why Cr is [Ar] 3d 5 4s 1 and not [Ar] 3d 4 4s 2 . Write Fe²⁺ and Fe³⁺ configurations from Fe. State which subshells are degenerate in the n = 3 shell and how many orbitals each has. NEET practice prompts Semiconductor doping (B or P in Si) tunes the number of valence electrons (p‑type vs n‑type), directly tied to electronic configuration. Spin‑based memory (MRAM) and quantum dots also rely on electron arrangements and spin states. remember Aufbau principle Ground‑state electrons occupy the lowest‑energy orbitals first, following the (n+l) order. Rank subshells by (n+l). Lower (n+l) fills first; if equal, lower n fills first. (n+l) rule Pauli Exclusion Principle No two electrons in an atom share all four quantum numbers; an orbital holds max two electrons with opposite spins. Hund's Rule of Maximum Multiplicity Electrons singly occupy all degenerate orbitals with parallel spins before any pairing. Valence electrons Electrons in the outermost shell (highest n) of an atom; chiefly responsible for bonding. Core electrons Inner electrons not involved directly in bonding; represented by the noble‑gas core in shorthand. Noble‑gas core Compressed notation using the previous noble gas to represent fully filled inner shells. Ground state Lowest‑energy electronic configuration of an atom or ion. Any configuration where one or more electrons are promoted to higher‑energy orbitals. Excited state Half‑filled/fully‑filled stability Extra stability seen for exactly half‑filled (e.g., d 5 ) or fully‑filled (e.g., d 10 $) subshells. Key terms at a glance