Periodic Table

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

Periodic Table — Key Facts for JEE Advanced

Modern Periodic Law: Properties of elements are periodic functions of their atomic number (not atomic mass, as Mendeleev originally proposed).

Classification:

• Groups: 18 vertical columns (IA to VIIIA, or 1 to 18)
• Periods: 7 horizontal rows ( Period 1 to Period 7)
• Blocks: s, p, d, f (based on filling of orbitals)

Key Trends (across period, left to right):

Property	Trend
Atomic size	Decreases (Z effective increases)
Ionization energy	Increases (generally, with exceptions)
Electronegativity	Increases
Metallic character	Decreases
Electron affinity	Increases (with irregularities at group 2, 15, 18)

Anomalies in Trends:

• Group 13: B has lower EA than Al (2p electron has poor shielding)
• Group 2: High EA (ns² fully filled, poor shielding)
• Group 15: Half-filled stability (p³), EA less negative than expected
• Group 18: Near zero EA (fully filled, no tendency to gain electrons)

⚡ Exam Tip: Atomic radius decreases across a period but ionic radius has a discontinuity — when cations form, the ionic radius drops sharply. Always compare ionic radii of isoelectronic species at the same charge to avoid confusion.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Periodic Table — Chemistry Study Guide

Effective Nuclear Charge (Z_eff):

Z_eff = Z − σ (screening constant) Slater’s rules for calculating σ:

• For ns or np electron: same group electrons: 0.35 (or 0.30 for 1s)
• For (n−1) shell: 0.85
• For (n−2) or lower: 1.00
• For ns, np electron in d or f group: electrons in same group: 0.35

Example: Nitrogen (1s² 2s² 2p³) Z = 7 σ = (1 × 0.30 for 1s²) + (5 × 0.35 for same n group) = 0.30 + 1.75 = 2.05 Z_eff = 7 − 2.05 = 4.95

Screening Order: (n−1) electrons > d and f electrons > np > ns (Inner shell shields better than same shell d/f due to penetration)

Atomic and Ionic Radii:

Covalent radius: Half the bond length in homonuclear diatomic molecules Van der Waals radius: Half the internuclear distance between nonbonded atoms Ionic radius: Based on ion distance in crystal lattices (isoelectronic series follow same pattern)

Isoelectronic series (same electron configuration): O²⁻ (8) > F⁻ (9) > Na⁺ (10) > Mg²⁺ (11) > Al³⁺ (12) Radii decrease with increasing Z (more protons pull electrons in) Order: O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺

Ionization Energy (IE):

Energy required to remove the most loosely held electron from a neutral gaseous atom. IE₁ < IE₂ < IE₃… (successive IEs, each higher than previous)

Factors affecting IE:

1. Z_eff (more protons = higher IE)
2. Atomic radius (larger = lower IE)
3. Electron configuration (half-filled and fully filled are stable, higher IE)
4. Penetration effect (s electrons penetrate more, harder to remove)

JEE Important Anomalies:

• B (IE₁ = 8.30 eV) < Be (IE₁ = 9.32 eV) — Be has stable 2s² configuration
• N (IE₁ = 14.5 eV) > O (IE₁ = 13.6 eV) — N has half-filled 2p³, very stable
• Al (IE₁ = 5.99 eV) < Mg (IE₁ = 7.64 eV) — Mg has stable 3s²
• P (IE₁ = 10.5 eV) > S (IE₁ = 10.4 eV) — P has half-filled 3p³

⚡ Exam Tip: When comparing IE values, check if the electron being removed is from a stable configuration (half-filled or fully filled subshell). An atom with a stable configuration will have anomalously high IE compared to its neighbors.

Electron Affinity (EA):

Energy released when an electron is added to a neutral gaseous atom. More negative EA = stronger electron acceptor = more reactive nonmetal.

Trends:

• Increases across period (more tendency to gain electrons)
• Maxima at Group 17 (halogens)
• Minima at Group 2 and 18 (stable configurations)

Exceptions:

• N has positive EA (+0.5 eV) despite being electronegative (half-filled stable, added electron goes to 2p orbital which is already half-filled — poor shielding)
• O has lower EA than S (due to inter-electron repulsion in small 2p orbital)
• Group 2 elements have very low EA (stable s²)

Values for halogens (for reference): F = −3.62 eV, Cl = −3.61 eV, Br = −3.36 eV, I = −3.06 eV Cl > F (odd-even effect — F is small, added electron suffers strong repulsion)

Electronegativity (EN):

Pauling scale: Relative measure of atom’s ability to attract bonding electrons.

χ = 0.18(eV) × [Δ EN pair] Δ = √(EA_diatomic − (EA₁ + EA₂)/2)

Pauline electronegativity values: F = 4.0, O = 3.5, N = 3.0, Cl = 3.2, Br = 3.0, S = 2.5, C = 2.5, H = 2.1, Si = 1.9, Na = 0.9

Allred-Rochow scale: Based on electrostatic force F = (Z_eff × e²)/(r²) More directly related to effective nuclear charge and covalent radius.

Mulliken scale: Average of IE and EA (in eV) / 5.6 gives Pauling-like values

Metallic Character:

Metals: Low IE, low EN, form cations, basic oxides Nonmetals: High IE, high EN, form anions, acidic oxides Metalloids: Intermediate properties (B, Si, Ge, As, Sb, Te, At)

Trends:

• Metals on left and bottom of periodic table
• Nonmetals on right and top
• Metalloids at the border (step pattern)

Acid-base character of oxides: Across period: Basic → Amphoteric → Acidic Na₂O (basic), MgO (basic), Al₂O₃ (amphoteric), SiO₂ (acidic), P₄O₁₀ (acidic), SO₃ (acidic)

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Periodic Table — Comprehensive Chemistry Notes

Detailed Electronic Configurations and Anomalies:

Ground state electron configurations: Follow the Aufbau principle with exceptions:

Element	Expected	Actual	Reason
Cr	[Ar] 4s² 3d⁴	[Ar] 4s¹ 3d⁵	Half-filled d subshell is extra stable
Cu	[Ar] 4s² 3d⁹	[Ar] 4s¹ 3d¹⁰	Full d subshell provides extra stability
Mo	[Kr] 5s² 4d⁴	[Kr] 5s¹ 4d⁵	Half-filled d
Pd	[Kr] 5s² 4d⁸	[Kr] 4d¹⁰	Full d preferred
La	[Xe] 6s² 5d¹	[Xe] 5d¹ (but actually [Xe] 6s² 5d¹)	—
Gd	[Xe] 6s² 4f⁷ 5d¹	[Xe] 6s² 4f⁷ 5d¹	Half-filled f is stable
Pt	[Xe] 6s² 5d⁸	[Xe] 6s¹ 5d⁹	Relativistic effects, full d preferred

Relativistic Effects (Heavy Elements):

For very heavy elements (Z > 30, especially 6th period onward):

• s electrons experience higher Z_eff (contracted)
• d and f electrons are shielded (expand)
• This affects IE, atomic size, and chemistry

Example: Mercury (Hg) is liquid at room temperature despite being a post-transition metal — relativistic effects make 6s² stable (like 6s² in inert pair effect).

Inert Pair Effect:

Post-transition metals (Ga, In, Tl; Ge, Sn, Pb; Sb, Bi) show reluctance to form +3 oxidation state despite being in group 13-15.

• +1 oxidation state becomes more stable down the group
• Reason: ns² electrons become less reactive (poor shielding of inner electrons)
• Example: Pb²⁺ is more stable than Pb⁴⁺ in aqueous solution

Tendency order: In > Tl > Ga (for +1 stability in group 13) Sn²⁺ < Pb²⁺ (for group 14) Sb³⁺ < Bi³⁺ (for group 15)

Lanthanide Contraction:

4f electrons have poor shielding, so as we fill 4f, Z_eff increases This pulls the 6s and 5d electrons inward, reducing size of Hf relative to Zr.

Consequences:

• Zr (160 pm) ≈ Hf (159 pm) — similar atomic sizes
• 5d series has higher IE than expected
• Third row transition metals are pulled closer to fourth row
• Chemistry of 5d and 6d metals is more similar than expected

Trends in Oxidation States:

Variable oxidation states:

• s-block: +1 (group 1), +2 (group 2) — only stable oxidation states
• p-block: Varies from (group − 10) to (group) — e.g., C: +4 to −4, N: +5 to −3
• d-block: +2 to +7 (can use inner d electrons)
• f-block: +3 (lanthanides), actinides: +3 to +7

Oxoacids and common oxidation states:

Element	Common oxoacid	Max Ox. State
Cl	HClO₄	+7
Br	HBrO₄	+7
I	H₅IO₆	+7
S	H₂SO₄	+6
N	HNO₃	+5
P	H₃PO₄	+5
Cr	H₂CrO₄	+6
Mn	HMnO₄	+7

Diagonal Relationship:

Elements on the diagonal (Be-B, Al-Si, etc.) show similarities with elements below-right in the next period, not with their own group.

Reasons:

1. Similar Z_eff per unit charge (charge/radius ratio)
2. Similar polarising power of cation (Z/r)
3. Covalent character of compounds

Examples:

• BeCl₂ is covalent (like AlCl₃)
• Be(OH)₂ is amphoteric (like Al(OH)₃)
• Be forms complexes (like Al)
• B forms complexes (Si doesn’t but B₂O₃ and SiO₂ both acidic)

Hydride and Oxide Trends:

Hydrides (binary compounds with H):

• Covalent hydrides: Groups 13-17 (volatile, molecular)
• Ionic hydrides: Groups 1-2 (form H⁻ ions, crystalline)
• Metallic hydrides: d-block (hydrogen occupies interstitial sites)
• Complex hydrides: LiAlH₄, NaBH₄ (covalent with H⁻ character)

Oxides: Across period: Basic → Amphoteric → Acidic → Covalent Na₂O (basic), MgO (basic), Al₂O₃ (amphoteric), SiO₂ (acidic glass), P₄O₁₀ (acidic), SO₃ (acidic), Cl₂O₇ (acidic)

Down group: Acidic → Amphoteric → Basic N₂O₅ (acidic), P₄O₁₀ (acidic), As₂O₃ (amphoteric), Sb₂O₃ (amphoteric/basic), Bi₂O₃ (basic)

p-Block Chemistry Highlights:

Group 13 (Boron family):

• Boron: Only 3 valence electrons, forms covalent compounds
• Aluminum: Most abundant metal in Earth’s crust, amphoteric oxide
• Thallium: Shows +1 (Tl⁺ is stable, like alkali metal) and +3

Group 14 (Carbon family):

• Catenation: Ability to form chains (C >> Si > Ge > Sn > Pb)
• Allotropes: C has diamond, graphite, fullerenes; Sn has gray/white; Pb only metallic
• PbS still has +2 preference due to inert pair effect

Group 15:

• Allotropropy: N has N₂ (triple bond, diatomic gas)
• P has white (P₄, reactive) and red (polymeric) allotropes
• Bi forms only metallic allotrope (inert pair, too large for catenation)

Group 16 (Oxygen family):

• O₂ is a biradical (paramagnetic) in ground state (triplet state is stable)
• S has many allotropes: rhombic, monoclinic, plastic sulfur
• Se has metallic and nonmetallic forms
• Po is radioactive metal

Group 17 (Halogens):

• F₂: Most reactive element, oxidizing agent
• Cl₂, Br₂, I₂: Increasing reactivity down group (iodine is least reactive)
• Interhalogen compounds: XY, XY₃, XY₅, XY₇ (where X is less electronegative)
• Pseudohalogens: (CN)₂, SCN⁻, etc.

Group 18 (Noble gases):

• He: Lowest boiling point (−269°C), used in cryogenics
• Ne, Ar, Kr, Xe: Xe forms compounds (XeF₂, XeF₄, XeF₆, XeO₃, XeO₄)
• Rn is radioactive

d-Block (Transition Metal) Chemistry:

General characteristics:

• Metallic properties (luster, conductivity, malleability)
• Variable oxidation states
• Colored compounds (d-d transitions)
• Paramagnetism (unpaired electrons)
• Catalytic properties
• Form coordination compounds

First row trends:

• Atomic size decreases to Cr, then slight increase, then decrease to Cu
• IE₂ is very high (removing from stable d¹⁰ or d⁵ configurations)
• Magnetic moment: μ = √[n(n+2)] BM where n = unpaired electrons

Color in transition metal complexes:

• d-d transition: e_g → t_2g splitting
• Δ = hν = hc/λ
• Larger Δ → shorter wavelength (blue) → complementary color seen
• [Ti(H₂O)₆]³⁺: d¹, orange-red (λmax ~500 nm)

f-Block (Inner Transition Metals):

Lanthanides (Ce to Lu):

• 4f orbitals are buried (poor shielding, inner)
• All form +3 oxidation state (very similar chemistry)
• Ce and Eu also show +4 and +2 respectively (stable configurations: 4f⁰ for Ce⁴⁺ and 4f⁷ for Eu²⁺)
• Atomic radii decrease (lanthanide contraction)
• Magnetic properties: Most are paramagnetic

Actinides (Th to Lr):

• 5f orbitals participate in bonding (more extended than 4f)
• Variable oxidation states (+3 to +7)
• All radioactive (beyond uranium)
• Famous actinides: U, Pu, Am (nuclear fuels)

⚡ Exam Tip: In JEE, lanthanide contraction is frequently tested — it explains why Zr and Hf have similar size, and why second and third row d-block elements are similar. Also, the actinide series is less commonly tested than lanthanides.

Periodic Trends in Atomic Properties — Summary:

Property	Across Period (→)	Down Group (↓)
Atomic radius	↓	↑
Ionization energy	↑ (with exceptions)	↓
Electron affinity	↑ (with exceptions)	↓
Electronegativity	↑	↓
Metallic character	↓	↑
Nonmetallic character	↑	↓
Acidic oxide character	↑	↓

Some JEE Advanced Problem Types:

Type 1: Compare IE of elements in same period Given: Na (IE₁ = 5.14 eV), Mg (IE₁ = 7.64 eV), Al (IE₁ = 5.99 eV), Si (IE₁ = 8.15 eV) Why is Al’s IE less than Mg? Mg has stable 3s² configuration. Why is P’s IE greater than S? P has stable half-filled 3p³.

Type 2: Isoelectronic series radius ordering Arrange O²⁻, F⁻, Na⁺, Mg²⁺, Al³⁺ in order of increasing radius. All have 10 electrons. More protons = smaller radius: Al³⁺ (13p) < Mg²⁺ (12p) < Na⁺ (11p) < F⁻ (9p) < O²⁻ (8p)

Type 3: Oxidation state stability Why is Sn²⁺ stable but Pb²⁺ more stable than Pb⁴⁺? Inert pair effect: ns² electrons become inert down the group (due to poor shielding and relativistic effects). The heavier element (Pb) has more stable +2 state than +4.

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