Periodic Table

Periodic Table

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

The periodic table arranges elements by increasing atomic number (Z), the Modern Periodic Law stating that physical and chemical properties are periodic functions of Z. Seven horizontal periods correspond to the filling of principal quantum shells n = 1 through 7; eighteen vertical groups (1–18) collect elements with identical valence configurations.

Two high-yield formulas:

• Effective nuclear charge: Z_eff = Z − S, where S is the screening constant from inner electrons (Slater’s rules).
• Mulliken electronegativity: χ_M = (IE + EA)/5.6, with IE = ionization enthalpy and EA = electron gain enthalpy, both in eV.

Across a period (left → right): atomic radius decreases, IE increases, EA becomes more negative, electronegativity rises, metallic character falls. Down a group the trends reverse. Memorise the anomalous positions of Li, Be, B (diagonal relationships with Mg, Al, Si) and remember that d-block and lanthanoid contraction shrink post-f-block radii.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Classification of Elements into Blocks

The block of an element is fixed by which sub-shell receives the differentiating electron:

• s-block: Groups 1, 2 — configuration [noble gas] ns¹⁻² (alkali and alkaline-earth metals).
• p-block: Groups 13–18 — configuration ends in np¹⁻⁶ (metals, metalloids, non-metals, halogens, noble gases).
• d-block: Groups 3–12 — (n−1)d¹⁻¹⁰ ns⁰⁻²; transition metals of the 4th, 5th, and 6th periods.
• f-block: Lanthanoids (4f) and actinoids (5f), placed separately below the main table.

Periodic Trends — Across a Period and Down a Group

Property	Across period (→)	Down group (↓)
Atomic radius	Decreases (Z_eff ↑)	Increases (new shell added)
Ionization enthalpy (IE)	Increases	Decreases
Electron gain enthalpy (EA)	More negative	Less negative
Electronegativity (χ)	Increases	Decreases
Metallic character	Decreases	Increases

Anomalous Behaviour and Diagonal Relationship

The first member of each p-block group (Li, Be, B, C, N, O, F) differs markedly from its congeners: very small radius, high charge density, no available d-orbitals, and high electronegativity. Diagonal neighbours share similar charge/radius ratios, producing close resemblance: Li ↔ Mg (both form nitrides, hydroxides are weak bases), Be ↔ Al (amphoteric oxides), B ↔ Si (semiconductors, formation of oxo-acids).

Contractions

• d-block contraction: poor shielding by 3d electrons shrinks the radii of 4p elements (Ga, Ge) relative to expectation, raising their IE.
• Lanthanoid contraction: 4f electrons shield poorly, pulling post-lanthanoid 5d/6p elements (Hf, W, Au) inward; this is why Zr ≈ Hf and Nb ≈ Mo in size.

Common JEE Question Patterns

• Identify period/group from a given configuration (e.g., [Ar] 3d¹⁰ 4s² 4p⁵ → Period 4, Group 17).
• Order elements by IE, EA, atomic radius, or electronegativity; watch for half-filled (N, P) and fully-filled (Be, Mg) IE anomalies.
• Isoelectronic species comparison using r ∝ 1/Z_eff.

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Effective Nuclear Charge in Depth

Slater’s rules quantify S: electrons in the same (ns, np) group shield each other by 0.35 (0.30 for 1s); an electron in (n−1) shell shields by 0.85; electrons in (n−2) or lower shield fully (1.00). For a 2p electron of fluorine (Z = 9, configuration 1s² 2s² 2p⁵): S = 2(0.85) + 6(0.35) = 1.70 + 2.10 = 3.80, giving Z_eff ≈ 5.20 — the high value that drives fluorine’s extreme electronegativity (χ ≈ 4.0).

Ionization Enthalpy Anomalies — Why They Exist

IE order across Period 2 is not monotonic: IE₁(Be) > IE₁(B) because Be’s 2s² is fully filled (extra stability), and IE₁(N) > IE₁(O) because N’s 2p³ is half-filled (symmetric, low electron–electron repulsion). Removing an electron from a more stable configuration costs more energy. The same anomaly reappears at P > S and Mg > Al.

Electron Gain Enthalpy — Subtle Behaviour

EA becomes progressively more negative from B to Cl, but three exceptions break the trend: EA(N) ≈ 0 or slightly positive (added electron enters an already half-filled 2p³, forcing pairing); EA(O) less negative than S? No — EA(O) is in fact less negative than EA(S) because O’s small 2p orbitals concentrate the new electron’s repulsion, while S’s larger 3p accommodates it. EA(Cl) is the most negative single-atom value (−349 kJ mol⁻¹).

Isoelectronic Series

For species with identical electron counts (N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺), radius shrinks as nuclear charge rises because Z_eff rises while S stays nearly constant: r ∝ 1/Z_eff. This single relation explains why Al³⁺ < Mg²⁺ < Na⁺ < Ne < F⁻ < O²⁻ < N³⁻ in size.

Connection to Bonding

Electronegativity difference Δχ between bonded atoms predicts bond polarity; Pauling’s Δχ = 0.1017√(Δ in kJ mol⁻¹) converts bond-energy differences into electronegativity values. High IE + highly negative EA = strong non-metal; low IE + low |EA| = metal. These periodic inputs flow directly into Chemical Bonding and Coordination Chemistry problems.

Common Mistakes

• Confusing period number with principal quantum number n of the valence shell — for d-block elements the period number equals n of the outermost s-electron shell.
• Assuming radius always decreases across a transition series; in d-block the radii decrease only slightly (Sc → Cu) because added d-electrons poorly shield.
• Treating EA of F as more negative than Cl; Cl actually has the most negative EA in the periodic table.

Practice Prompts

1. Arrange N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺, Al³⁺ in increasing ionic radius, justifying with Z_eff reasoning.
2. Explain why IE₁ of nitrogen is greater than that of oxygen using electronic configuration and exchange energy.

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