Electron Configuration Generator

Electron configuration is like the “address” of every electron in an atom. It tells us which energy level (n), which subshell (s, p, d, f) and how many electrons are present in each. From this single line of information we can deduce: (i) the element's position in the periodic table, (ii) its valence electrons and so its valency, (iii) the type of bonding it prefers, (iv) magnetic properties (paramagnetic vs diamagnetic), and (v) the chemical reactivity. It is the foundation of inorganic chemistry.

Three rules guide us. (1) Aufbau principle — fill subshells in increasing order of energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s … (2) Pauli exclusion principle — no two electrons can have all four quantum numbers the same; an orbital holds at most two electrons of opposite spin. (3) Hund's rule — electrons occupy degenerate orbitals singly (parallel spins) before pairing up. Apply these in order, and the configuration writes itself. Filling order: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 …

Step-by-step: (1) Find the atomic number Z, which equals the number of electrons in a neutral atom. (2) Use the diagonal “Madelung” diagram or the Aufbau order. (3) Distribute the Z electrons into subshells, putting at most 2 in s, 6 in p, 10 in d, 14 in f. (4) Verify the total adds up to Z. (5) For ions, add electrons (anions) or remove from the highest n shell first (cations). Always double-check exceptions like Cr (3d5 4s1) and Cu (3d10 4s1) which prefer half-filled or fully-filled stability.

Ground state means the lowest possible energy arrangement of an atom's electrons — every electron occupies the lowest available orbital obeying the Aufbau, Pauli and Hund rules. Any other arrangement (an electron promoted to a higher orbital) is called an excited state. For example, ground-state hydrogen is 1s1; if its electron jumps to 2s, the atom is in 2s1 excited state and will quickly fall back to 1s1, emitting a photon. Ground state is always the most stable.

Potassium is in Group 1, Period 4. Distributing 19 electrons in increasing energy order — 1s2, 2s2, 2p6, 3s2, 3p6, 4s1 — confirms it is an alkali metal with only one valence electron in the 4s subshell, which it readily loses to form K+. Note that the 4s fills before 3d, even though n = 3 is lower, because 4s has slightly lower energy in this region. Shorthand: [Ar] 4s1. K (Z=19) : 1s2 2s2 2p6 3s2 3p6 4s1 = [Ar] 4s1 K Electron shells: 2, 8, 8, 1

Calcium is the next element after potassium in Period 4. We simply add one more electron to the 4s orbital, giving 4s2 which is now full. Calcium therefore belongs to Group 2 (alkaline earth metals) and easily loses both 4s electrons to form Ca2+, which is the form found in bones, teeth and milk. Shorthand notation [Ar] 4s2 summarises this nicely. Ca (Z=20) : 1s2 2s2 2p6 3s2 3p6 4s2 = [Ar] 4s2

Sulfur has 16 electrons. Filling: 1s2 2s2 2p6 3s2 3p4. Note the 3p subshell holds 4 electrons — by Hund's rule, three of them occupy the three 3p orbitals singly with parallel spin, and the fourth pairs up in one of them. Sulfur belongs to Group 16, has 6 valence electrons (3s2 3p4), and typically gains 2 electrons to form sulphide S2− or shares electrons in compounds like H2S, SO2, SO3. S (Z=16) : 1s2 2s2 2p6 3s2 3p4 = [Ne] 3s2 3p4

Bromine has 35 electrons. Filling step by step: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5. Notice 4s fills before 3d, then 3d completes, then 4p starts. Bromine has 7 valence electrons (4s2 4p5) and easily gains one electron to form bromide Br−, completing the octet. It is in Group 17 (halogens), Period 4. Shorthand: [Ar] 3d10 4s2 4p5. Br (Z=35) : [Ar] 3d10 4s2 4p5

Carbon has just 6 electrons: 1s2 2s2 2p2. The two electrons in 2p go into separate p orbitals (Hund's rule), so carbon has two unpaired electrons. However, in most compounds the 2s and 2p orbitals hybridise (sp3, sp2, sp) to give 4 equivalent bonds. This explains why carbon always shows valency 4 and forms millions of organic compounds — the entire field of organic chemistry rests on this configuration. C (Z=6) : 1s2 2s2 2p2 = [He] 2s2 2p2 C Electron shells: 2, 4

Sodium has 11 electrons. Filling order: 1s2 2s2 2p6 3s1. The single 3s electron is loosely held and easily lost, so sodium readily forms Na+ with the stable electron configuration of neon. That is why sodium is highly reactive — kept in kerosene, as we discussed earlier. Shorthand: [Ne] 3s1. Na (Z=11) : 1s2 2s2 2p6 3s1 = [Ne] 3s1

Oxygen has 8 electrons: 1s2 2s2 2p4. By Hund's rule, the four 2p electrons occupy the three 2p orbitals as 2,1,1 — meaning oxygen has two unpaired electrons. This explains why O2 is paramagnetic (a fact predicted nicely by molecular orbital theory). Oxygen needs 2 more electrons to complete its octet, hence its valency of 2 and the formation of O2− in oxides. O (Z=8) : 1s2 2s2 2p4 = [He] 2s2 2p4

Magnesium has 12 electrons: 1s2 2s2 2p6 3s2. It has two valence electrons in the 3s subshell, which it loses to form Mg2+, achieving the stable neon configuration. This makes Mg an active alkaline-earth metal, important biologically (centre of chlorophyll!) and industrially (light alloys for aircraft). Shorthand: [Ne] 3s2. Mg (Z=12) : [Ne] 3s2

Iron has 26 electrons. Configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d6, or in shorthand [Ar] 3d6 4s2. Note that 4s fills before 3d in neutral atoms, but for the iron ions electrons are removed from 4s first: Fe2+ is [Ar] 3d6 and Fe3+ is [Ar] 3d5 (half-filled, extra stable — that is why Fe3+ is so common). The unpaired d-electrons explain why iron is ferromagnetic and forms coloured compounds. Fe (Z=26) : [Ar] 3d6 4s2 ; Fe3+ : [Ar] 3d5