Classification of Elements & Periodicity in Properties

1 · The history ladder matching favourite

• Dobereiner (1829) — Triads: middle element's atomic weight ≈ average of the other two (Li-Na-K, Ca-Sr-Ba, Cl-Br-I). Dismissed as coincidence.
• A.E.B. de Chancourtois (1862) — French geologist; cylindrical table by increasing atomic weight; ignored.
• Newlands (1865) — Law of Octaves: every 8th element resembles the first; worked only up to calcium; Davy Medal 1887.
• Lothar Meyer (1868) — plotted atomic volume/m.p./b.p. vs atomic weight; table resembling the modern one, published after Mendeleev.
• Mendeleev (1869) — Periodic Law: properties are a periodic function of atomic weights. Left gaps (Eka-aluminium = gallium, Eka-silicon = germanium) and predicted their properties; swapped iodine/tellurium by properties.
• Moseley (1913) — √ν of characteristic X-rays vs Z gives a straight line → Modern Periodic Law: properties are periodic functions of atomic numbers.

2 · Long form architecture

Horizontal rows = periods (7 of them); vertical columns = groups (1–18, IUPAC 1984, replacing IA…VIIA/VIII/IB…VIIB/0). Period number = highest principal quantum number n. Elements per period: 2, 8, 8, 18, 18, 32; the 7th is incomplete (theoretical max 32, ends at Z = 118, a noble gas). The 14-member lanthanoid (4f) and actinoid (5f) series sit in separate bottom panels — a layout owed to Seaborg (element 106 = Seaborgium; Nobel 1951).

3 · Why 2, 8, 18, 32?

Elements per period = 2 × (number of orbitals filled in that energy level). 5th period: orbitals 5s, 4d, 5p = 9 orbitals → 18 electrons → 18 elements. Same logic gives 2 (1s), 8 (2s2p), 32 (6s 4f 5d 6p = 16 orbitals).

4 · The four blocks

• s-block (Gp 1–2): ns¹⁻². Alkali & alkaline earth metals; reactive, mostly ionic compounds.
• p-block (Gp 13–18): ns²np¹⁻⁶. s + p = representative/main-group elements. Halogens (17) & chalcogens (16) have highly negative ΔegH.
• d-block (Gp 3–12): (n−1)d¹⁻¹⁰ns⁰⁻². Transition elements; all metals, coloured ions, variable valence. Exception: Pd = 4d¹⁰5s⁰.
• f-block: (n−2)f¹⁻¹⁴(n−1)d⁰⁻¹ns². Inner-transition — lanthanoids & actinoids (all radioactive; after uranium = transuranium).

Placement quirks: Helium is s-block by config (1s²) but sits in Group 18 — filled valence shell, noble-gas behaviour. Hydrogen (1s¹) fits Group 1 AND Group 17 logic — so it is placed separately at the top.

5 · Metals, non-metals, metalloids

Metals = >78% of elements, left side, solids at room temp (Hg liquid; Ga 303 K and Cs 302 K melt near room temp), malleable, ductile, conductors. Non-metals: top right, low m.p./b.p. (boron & carbon are exceptions). Along the zig-zag: Si, Ge, As, Sb, Te = metalloids/semi-metals. Metallic character ↑ down a group, ↓ across a period.

6 · Atomic & ionic radius

Radius ↓ across a period (nuclear charge outruns shielding), ↑ down a group (new shells). Cation < parent atom (fewer electrons, same charge: Na 186 pm → Na⁺ 95 pm); anion > parent (F 64 pm → F⁻ 136 pm). Isoelectronic species (same electron count, e.g., O²⁻, F⁻, Na⁺, Mg²⁺ — all 10 e⁻): more protons = smaller. So O²⁻ > F⁻ > Na⁺ > Mg²⁺.

7 · Ionization enthalpy (ΔiH) and its two anomalies most tested

X(g) → X⁺(g) + e⁻; always positive; IE₂ > IE₁. Trend: ↑ across (shielding can't compensate rising nuclear charge), ↓ down (distance + shielding win). Graph: maxima at noble gases, minima at alkali metals.

• Be > B: boron loses a 2p electron — less penetrating, better shielded — than beryllium's 2s.
• N > O: nitrogen's 2p³ is exactly half-filled (Hund); oxygen's paired 2p electron suffers extra repulsion and leaves easily.

8 · Electron gain enthalpy (ΔegH)

X(g) + e⁻ → X⁻(g); can be exothermic (halogens, very negative) or endothermic (noble gases, LARGE POSITIVE — the new electron must enter the next shell). Becomes more negative across a period, less negative down a group — BUT O and F are less negative than S and Cl: in the compact n = 2 shell, electron-electron repulsion resists the newcomer. Cl has the most negative ΔegH of all.

9 · Electronegativity

Not measurable, only relative (Pauling scale most common): tendency to attract shared electrons in a compound. F = 4.0 (highest), Cs ≈ 0.7. ↑ across, ↓ down; directly tied to non-metallic character. (ΔegH is for an isolated gaseous atom; EN is inside a bond — the classic distinction question.)

10 · Valence, oxides & reactivity extremes

Valence = outer electrons OR 8 − outer electrons (Gp 1→18: 1, 2, 3, 4, 3/5, 2/6, 1/7, 0/8). Oxidation state from electronegativity: OF₂ → O is +2; Na₂O → O is −2. Reactivity peaks at both ends of a period (alkali metals lose e⁻; halogens gain e⁻) and is lowest in the middle. Oxides: extreme left = most basic (Na₂O), extreme right = most acidic (Cl₂O₇), centre = amphoteric (Al₂O₃, As₂O₃) or neutral (CO, NO, N₂O).

11 · Anomalous second period & diagonal relationship

First members (Li → F) differ from their groups: small size, large charge/radius ratio, high EN, and only 4 valence orbitals (2s + 2p) → max covalency 4 (B forms only BF₄⁻; Al can form AlF₆³⁻ with covalency 6). They also form pπ–pπ multiple bonds (C=C, C≡N, N=O) that heavier members don't. Diagonal relationship: Li ~ Mg, Be ~ Al — first element resembles the second element of the NEXT group.

Modern Periodic Law

"The physical and chemical properties of the elements are periodic functions of their atomic numbers." (Mendeleev's original used atomic weights.)

Ionization enthalpy

Energy required to remove an electron from an isolated gaseous atom in its ground state: X(g) → X⁺(g) + e⁻ (kJ mol⁻¹, always positive).

Electron gain enthalpy

Enthalpy change when an electron is added to a neutral gaseous atom: X(g) + e⁻ → X⁻(g). Negative = exothermic (halogens), positive = endothermic (noble gases).

Electronegativity

Qualitative ability of an atom in a chemical compound to attract shared electrons; no single measurable value; Pauling scale most common.

Isoelectronic species

Atoms/ions with the same number of electrons, e.g., O²⁻, F⁻, Na⁺, Mg²⁺ (10 e⁻ each) — radius shrinks as nuclear charge grows.

Shielding / effective nuclear charge

Core electrons screen the valence electron; net charge felt < actual nuclear charge. Shielding is most effective when inner shells are completely filled.

Oxidation state

Charge acquired by an atom in a compound on the basis of electronegativity of the other atoms (modern usage of "valence").

P1 · Largest/smallest among Mg, Mg²⁺, Al, Al³⁺

Radius ↓ across; cations smaller than parents; among isoelectronic ions more charge = smaller. Largest Mg, smallest Al³⁺.

P2 · Predict ΔiH of Al given Na 496, Mg 737, Si 786

Closer to 575 than 760 — Al's 3p electron is shielded by 3s², so its ΔiH is LOWER than Mg's despite higher Z.

P3 · Place undiscovered Z = 117 and Z = 120

117 → halogen family, Group 17, [Rn]5f¹⁴6d¹⁰7s²7p⁵. 120 → alkaline earth, Group 2, [Uuo]8s².

P4 · Formula prediction from group valence

Si (Gp 14, valence 4) + Br (valence 1) → SiBr₄. Al (Gp 13, valence 3) + S (Gp 16, valence 2) → Al₂S₃.

P5 · Metallic character order: Si, Be, Mg, Na, P

↑ down, ↓ across → P < Si < Be < Mg < Na.

P6 · Justify 18 elements in the 5th period

n = 5 fills 5s, 4d, 5p = 9 orbitals = 18 electrons = 18 elements.

P7 · Oxidation state vs covalency in [AlCl(H₂O)₅]²⁺

Not the same: oxidation state of Al = +3, covalency = 6 (six ligands).

G1 · ΔiH vs Z (Fig 3.5)

Maxima at noble gases (closed shells), minima at alkali metals (single ns electron, well shielded) — the sawtooth every NEET aspirant must sketch from memory.

G2 · Moseley's plot (why the law changed)

√ν of characteristic X-rays vs Z = straight line; vs atomic mass = not. Atomic number is the more fundamental property.

G3 · Atomic radius across period 2 vs down Group 1/17

Across Li → F: falling curve. Down alkali metals Li → Cs and halogens F → I: rising curves. Cs is the largest common atom (262 pm).

G4 · Lothar Meyer's original idea

Atomic volume / m.p. / b.p. plotted vs atomic weight repeat periodically — the graph that anticipated the table.