Periodic Table and Periodicity — JEE Main & Advanced Notes

Turn the periodic table into predictable trends: radius, ionization enthalpy, electron affinity, electronegativity and chemical character.

Priority: Must Do. Unit: Inorganic Chemistry. Level: Foundation.

How the uploaded material was used: Mapped from periodic classification, periodic trend and exception-based inorganic sheets. The final student-facing notes and questions are original, rewritten and copyright-safe.

These are the ideas that decide most correct answers in Periodic Table and Periodicity.

• Effective nuclear charge and shielding — the root cause of almost every trend: The actual charge felt by a valence electron is not the full nuclear charge Z but the effective nuclear charge Zeff = Z − σ. Inner electrons (core electrons) partially shield the valence electrons from the nucleus. Across a period (left to right), Z increases while electrons are added to the same shell (shielding increases only slightly) → Zeff increases → stronger pull on valence electrons. Down a group, n increases (new shell) → electrons are farther from the nucleus and more shielded → Zeff on the outer electrons changes slowly, but the increased distance dominates.
• Atomic radius — across a period and down a group: Across a period (left to right), atomic radius decreases because Zeff increases while the valence shell stays the same — the nucleus pulls electrons closer. Down a group, atomic radius increases because a new principal quantum shell is added, placing the outermost electrons farther from the nucleus despite the higher Z. The shielding by inner shells is very effective, so atomic size grows. For transition metals, atomic radii barely change across a period because d electrons shield poorly and Zeff increases but the d orbital is also filling.
• Ionisation enthalpy (IE) — trends and exceptions: IE₁ generally increases across a period (more Zeff → harder to remove electron). IE₁ decreases down a group (higher n → electron farther and more shielded). Critical exceptions: (1) IE₁(B) < IE₁(Be): removing a 2p electron (B) is easier than removing a 2s electron (Be) because 2p is higher energy and partially shielded by the 2s. (2) IE₁(O) < IE₁(N): nitrogen has a half-filled 2p subshell (3 unpaired electrons) which is extra stable; oxygen's fourth 2p electron is paired and repelled → easier to remove. Large jumps in successive IEs indicate a jump from valence to core electrons — used to identify the group of an element.
• Electron affinity (EA) and electronegativity (EN) — related but distinct: EA is the energy released when an electron is added to a neutral gaseous atom (negative means energy released, exothermic). EN is the tendency of an atom in a molecule to attract the shared electron pair. Both generally increase across a period and decrease down a group. Key exceptions: EA(F) < EA(Cl) because F is small and the added electron enters a very compact 2p orbital with high electron–electron repulsion; EA(N) ≈ 0 because N's 2p is half-filled (stable) and electron addition disturbs this. EA₂ (second electron affinity) is always positive (endothermic) because adding an electron to an anion requires overcoming repulsion.
• Isoelectronic species — size is determined solely by Z: Isoelectronic species have the same number of electrons (same electron–electron repulsion) but different nuclear charges. The one with more protons (higher Z) attracts electrons more strongly → smaller size. Example: O²⁻, F⁻, Ne, Na⁺, Mg²⁺ all have 10 electrons. Size order: O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ (Z = 8 < 9 < 10 < 11 < 12). This is the key rule for comparing ionic sizes and neutral atoms in isoelectronic series.
• Anomalous properties of second-period elements — the diagonal relationship: The first element of each group differs from the rest because it is uniquely small (no d orbitals, maximum electronegative character, smallest size). This leads to: Li behaving more like Mg than Na; Be behaving more like Al than Mg; B behaving more like Si than Al. The diagonal similarity arises because diagonally placed elements have similar charge density (charge/volume ratio). Second-period elements also have a maximum of four bonds (octet rule strict) while heavier elements can expand their octet using d orbitals.
• Metallic and non-metallic character: Metallic character (ease of losing electrons) increases down a group and decreases across a period (left to right). Non-metallic character (tendency to gain electrons) is the opposite. The line dividing metals from non-metals in the periodic table runs diagonally — elements near this line (B, Si, Ge, As, Sb, Te) are metalloids with intermediate properties. Oxides of metals are basic (or amphoteric); oxides of non-metals are acidic.
• Periodicity of oxides and hydrides: Across period 3: Na₂O (strongly basic) → MgO (basic) → Al₂O₃ (amphoteric) → SiO₂ (weakly acidic) → P₄O₁₀ (acidic) → SO₃ (very acidic) → Cl₂O₇ (most acidic). This progression directly mirrors the increase in non-metallic character across the period. Hydrides: group 1 hydrides are ionic (NaH → Na⁺H⁻); group 4 hydrides are covalent (CH₄, SiH₄); group 15–17 hydrides show hydrogen bonding (NH₃, H₂O, HF) and have anomalously high boiling points.

• Effective nuclear charge: Zeff = Z − σ where σ = total shielding constant (Slater's rules)
• Slater shielding: electrons in same subshell shield 0.35 each; electrons in (n−1) shell shield 0.85 each; electrons in n−2 and below shield 1.00 each
• Covalent atomic radius = half the bond length in a homodiatomic molecule (e.g., C−C in diamond/2)
• Van der Waals radius > covalent radius (van der Waals is the non-bonded contact radius)
• Ionic radius: cations < parent atom (electron removed → less electron–electron repulsion); anions > parent atom
• Isoelectronic species: same number of electrons, different Z. Radius order: higher Z → smaller size
• IE₁ < IE₂ < IE₃ … for successive ionisation energies of the same element (large jump when core electron is removed)

Derivation / logic hint: Do not plug values blindly. Start from conservation of mass/charge, equilibrium definition, energy balance, electron movement, structure-property relation, or stability of the product/intermediate.

Most negative marks in this chapter come from condition errors, not lack of memory.

• Memorising trends without understanding the exception logic: IE₁(B)
• Confusing electron affinity sign conventions: a large negative EA (e.g., F: −328 kJ/mol) means a lot of energy is released — highly favourable. A positive EA means energy must be supplied. EA₂ is always positive. Do not confuse 'large EA' with 'high energy released' without checking the sign convention used.
• Comparing isoelectronic sizes incorrectly: bigger Z → smaller ion (not bigger) among isoelectronic species. Students who apply the group trend (bigger Z → more electrons → bigger atom) make this mistake because in isoelectronic series, the number of electrons is fixed, so only Z matters.
• Assuming second-period elements always follow the same trends as heavier elements in their group: F's EA is less than Cl's; N's IE is greater than O's; these are not random exceptions but systematic consequences of the compact 2p shell. Second-period elements often break the trends seen from the 3rd period onward.
• Forgetting that successive ionisation energies always increase, with a large jump at the point where a full shell is breached: IE₁ < IE₂ < IE₃ always. The jump identifies the valence shell size. If there's a big jump between IE₂ and IE₃, the element is in Group 2.

For JEE Main, prioritise direct formula use, NCERT-aligned facts, named-reaction recognition, trend comparison and quick elimination. Target 60–90 seconds per question.

• Effective nuclear charge and shielding
• Atomic and ionic radius trends
• Ionisation enthalpy and exceptions
• Electron affinity and sign conventions

For JEE Advanced, combine ideas. Expect assertion-reason, integer, multiple-correct, paragraph-style and hidden-condition problems. Before finalising, ask which assumption the question is testing.

• Ionisation enthalpy and exceptions
• Electron affinity and sign conventions
• Electronegativity (Pauling scale)
• Isoelectronic species and size comparison
• Anomalous properties of 2nd period elements
• Diagonal relationship

Use this block in the final 24–48 hours before a mock.

• Atomic radius: decreases across a period (↑Zeff, same shell); increases down a group (new shell, increased shielding). Exception: transition metals (d filling) barely change across.
• IE₁ exceptions: B < Be (2p > 2s energy); O < N (half-filled 2p stability). IE jumps reveal group number.
• EA exceptions: F < Cl (crowded 2p); N ≈ 0 (half-filled stability).
• Isoelectronic size: higher Z → smaller size. Compare O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺.
• Diagonal relationship: Li↔Mg, Be↔Al, B↔Si behave similarly due to similar charge density.