JEE Main Inorganic Chemistry — Periodic Table & p-Block Revision Notes

This comprehensive revision resource covers the essential concepts of the Periodic Table and p-Block elements, crucial for excelling in JEE Main Inorganic Chemistry. We'll delve into fundamental principles, periodic trends, and the detailed chemistry of p-Block elements, providing a strong foundation for tackling a wide range of problems.

The Periodic Table and p-Block elements are high-weightage topics in JEE Main, often contributing 2-3 questions directly, and forming the basis for understanding other chapters like chemical bonding and coordination compounds. Mastering these concepts is vital for a good score.

Exam Weightage & Importance

In JEE Main, Inorganic Chemistry typically accounts for about 30-35% of the total Chemistry section. Within this, the Periodic Table and p-Block elements are consistently high-scoring areas. Expect 2-3 questions directly from these chapters. Questions range from straightforward recall of facts and trends to application-based problems involving reactions or properties. A solid understanding of periodic trends is also crucial for indirectly solving questions related to chemical bonding, d- & f-block elements, and even some aspects of coordination chemistry. Many qualitative analysis questions also indirectly rely on p-block chemistry.

Periodic Table: Fundamental Concepts & Trends

The periodic table is a systematic arrangement of elements based on their electronic configurations, which dictates their chemical and physical properties. Understanding its structure and the trends within it is fundamental.

Genesis and Evolution of the Periodic Table

• Dobereiner's Triads (1829): Grouped elements in threes (triads) where the atomic weight of the middle element was approximately the average of the other two. (e.g., Li, Na, K; Ca, Sr, Ba; Cl, Br, I). Limitations: Only a few triads could be identified.
• Newlands' Law of Octaves (1865): Arranged elements in increasing order of atomic weights, observing that every eighth element had properties similar to the first (like musical octaves). (e.g., Li, Na, K). Limitations: Applicable only up to Calcium, failed for elements beyond that.
• Mendeleev's Periodic Law (1869): "The physical and chemical properties of elements are a periodic function of their atomic weights."
  • Achievements:
    • Systematic study of elements.
    • Corrected atomic weights of some elements (e.g., Be, In, Au).
    • Predicted the properties of undiscovered elements and left gaps (e.g., Eka-Boron (Sc), Eka-Aluminium (Ga), Eka-Silicon (Ge)).
  • Limitations:
    • Position of hydrogen.
    • Anomalous pairs (e.g., Ar (39.9) before K (39.1); Co (58.9) before Ni (58.7)).
    • Position of isotopes.
    • No clear explanation for the cause of periodicity.
• Modern Periodic Law (Moseley, 1913): "The physical and chemical properties of elements are a periodic function of their atomic numbers."
  • This eliminated the anomalies of Mendeleev's table.
  • The modern long form of the periodic table is based on atomic number and electronic configuration.

Modern Periodic Table Structure

• Periods (Horizontal Rows):
  • 7 periods.
  • The period number signifies the highest principal quantum number ($n$) of the valence shell.
  • Number of elements in periods:
    • 1st (n=1): 2 elements (H, He) - Shortest period.
    • 2nd (n=2): 8 elements (Li-Ne) - Short period.
    • 3rd (n=3): 8 elements (Na-Ar) - Short period.
    • 4th (n=4): 18 elements (K-Kr) - Long period. Begins with filling of 4s, followed by 3d, then 4p.
    • 5th (n=5): 18 elements (Rb-Xe) - Long period. Begins with filling of 5s, followed by 4d, then 5p.
    • 6th (n=6): 32 elements (Cs-Rn) - Longest period. Begins with filling of 6s, then 4f (lanthanoids), 5d, then 6p.
    • 7th (n=7): 32 elements (Fr-Og) - Incomplete period. Begins with filling of 7s, then 5f (actinoids), 6d, then 7p.
• Groups (Vertical Columns):
  • 18 groups (IUPAC recommendation).
  • Elements in a group have similar outermost electronic configurations and thus similar chemical properties.
  • Classification based on valence electron filling:
    • s-block (Groups 1 & 2): Valence electrons enter s-orbital. General configuration: $ns^{1-2}$.
    • p-block (Groups 13-18): Valence electrons enter p-orbital. General configuration: $ns^2 np^{1-6}$.
    • d-block (Groups 3-12, Transition elements): Valence electrons enter $(n-1)d$ orbital. General configuration: $(n-1)d^{1-10} ns^{0-2}$ (exceptions like Pd: $4d^{10}5s^0$).
    • f-block (Lanthanoids & Actinoids, Inner Transition elements): Valence electrons enter $(n-2)f$ orbital. General configuration: $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$.

Periodic Properties and Trends

Understanding these trends is crucial for predicting chemical behavior.

1. Atomic Radius

• Distance from the center of the nucleus to the outermost electron shell.
• Types:
  • Covalent Radius: Half the distance between two identical atoms bonded covalently.
  • Metallic Radius: Half the inter-nuclear distance between two adjacent metal atoms in a metallic lattice.
  • Van der Waals Radius: Half the inter-nuclear distance between two non-bonded atoms of adjacent molecules. $\text{Van der Waals Radius} > \text{Metallic Radius} > \text{Covalent Radius}$.
• Trends:
  • Across a Period (L to R): Decreases. Due to increasing effective nuclear charge ($Z_{eff}$) and simultaneous addition of electrons into the same principal shell. The increased $Z_{eff}$ pulls the electrons closer to the nucleus.
  • Down a Group (Top to Bottom): Increases. Due to the addition of new electron shells, increasing the distance of valence electrons from the nucleus, despite increasing $Z_{eff}$.
• Exceptions/Anomalies:
  • For transition elements, the atomic radius decreases slightly from left to right, then becomes almost constant, and then increases slightly towards the end. This is due to the shielding effect of d-electrons. (e.g., Fe, Co, Ni have similar radii).
  • Lanthanoid Contraction: The steady decrease in atomic radii and ionic radii of the lanthanoids with increasing atomic number. This is due to the poor shielding effect of 4f electrons, leading to an increase in $Z_{eff}$.
    • Consequences of Lanthanoid Contraction:
      • Similar radii of 4d and 5d series elements (e.g., Zr (160 pm) and Hf (159 pm); Nb (147 pm) and Ta (147 pm)). This makes their chemical properties very similar and separation difficult.
      • Higher density of 5d series elements.
      • Similar basicity of lanthanoid hydroxides (decreases with increasing atomic number).

2. Ionic Radius

• Radius of an ion.
• Trends:
  • Cations: Smaller than their parent atoms (loss of electrons, increased $Z_{eff}$ per electron). The more positive the charge, the smaller the cation. (e.g., $\text{Fe} > \text{Fe}^{2+} > \text{Fe}^{3+}$).
  • Anions: Larger than their parent atoms (gain of electrons, decreased $Z_{eff}$ per electron, increased electron-electron repulsion). The more negative the charge, the larger the anion. (e.g., $\text{Cl} < \text{Cl}^{-}$).
  • Isoelectronic Species: Ions/atoms with the same number of electrons. For isoelectronic species, ionic radius decreases with an increase in atomic number (i.e., increase in nuclear charge).
    • Example: $\text{N}^{3-} > \text{O}^{2-} > \text{F}^{-} > \text{Na}^{+} > \text{Mg}^{2+} > \text{Al}^{3+}$ (all have 10 electrons, but nuclear charge increases from N to Al).

3. Ionization Enthalpy (IE) or Ionization Energy (IP)

• The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
• First Ionization Enthalpy ($\text{IE}_1$): $\text{M}(g) \rightarrow \text{M}^{+}(g) + e^{-}$
• Second Ionization Enthalpy ($\text{IE}_2$): $\text{M}^{+}(g) \rightarrow \text{M}^{2+}(g) + e^{-}$
• Always $\text{IE}_1 < \text{IE}_2 < \text{IE}_3...$ (Successive ionization enthalpies increase because it becomes harder to remove an electron from an already positive ion).
• Factors Affecting IE:
  • Atomic Size: Inversely proportional. Smaller atoms have higher IE.
  • Effective Nuclear Charge ($Z_{eff}$): Directly proportional. Higher $Z_{eff}$ means tighter hold on electrons, higher IE.
  • Shielding/Screening Effect: Inversely proportional. Greater shielding by inner electrons reduces $Z_{eff}$ on valence electrons, lower IE.
  • Penetration Effect of Orbitals: For a given shell, the penetrating power (and thus attraction to nucleus) order is $s > p > d > f$. An electron in an s-orbital is harder to remove than one in a p-orbital of the same shell.
  • Electronic Configuration Stability: Half-filled ($p^3, d^5, f^7$) and completely filled ($p^6, d^{10}, f^{14}$) subshells are exceptionally stable and require more energy to remove an electron.
• Trends:
  • Across a Period (L to R): Generally Increases. Due to increasing $Z_{eff}$ and decreasing atomic size.
  • Down a Group (Top to Bottom): Generally Decreases. Due to increasing atomic size and increasing shielding effect, despite increasing $Z_{eff}$.
• Exceptions/Anomalies in Period 2 & 3:
  • $\text{IE}_1(\text{Be}) > \text{IE}_1(\text{B})$: ($2s^2$ is stable, fully filled configuration).
  • $\text{IE}_1(\text{N}) > \text{IE}_1(\text{O})$: ($2p^3$ is stable, half-filled configuration).
  • Similar exceptions are observed in the 3rd period: $\text{IE}_1(\text{Mg}) > \text{IE}_1(\text{Al})$ and $\text{IE}_1(\text{P}) > \text{IE}_1(\text{S})$.

4. Electron Gain Enthalpy ($\Delta H_{eg}$) or Electron Affinity (EA)

• The energy released when an electron is added to an isolated gaseous atom to form an anion.
• $\text{X}(g) + e^{-} \rightarrow \text{X}^{-}(g)$
• Conventionally, energy released is negative, so $\Delta H_{eg}$ is usually negative.
• If energy is absorbed (e.g., for noble gases or elements like Be, Mg, N), $\Delta H_{eg}$ is positive.
• Factors Affecting $\Delta H_{eg}$:
  • Atomic Size: Inversely proportional. Smaller atoms generally have more negative (favorable) $\Delta H_{eg}$ because the added electron is closer to the nucleus.
  • Effective Nuclear Charge ($Z_{eff}$): Directly proportional. Higher $Z_{eff}$ means a stronger attraction for the incoming electron, more negative $\Delta H_{eg}$.
  • Electronic Configuration Stability: Atoms with stable configurations ($s^2, p^6$, half-filled $p^3$) have less negative or even positive $\Delta H_{eg}$.
• Trends:
  • Across a Period (L to R): Generally Becomes More Negative. Due to increasing $Z_{eff}$ and decreasing atomic size.
  • Down a Group (Top to Bottom): Generally Becomes Less Negative. Due to increasing atomic size and increasing shielding, the attraction for the incoming electron decreases.
• Exceptions/Anomalies:
  • $\text{Cl} > \text{F}$: Even though F is smaller than Cl, its electron gain enthalpy is less negative. This is due to the very small size of F, leading to strong inter-electronic repulsion among existing 2p electrons and the incoming electron. The same applies to P vs N ($\text{S} > \text{O}$).
  • Noble gases have positive $\Delta H_{eg}$ (highly unstable to add an electron).
  • Group 2 elements ($ns^2$) and Group 15 elements ($np^3$) have positive or slightly positive $\Delta H_{eg}$ due to stable configurations.

5. Electronegativity (EN)

• The tendency of an atom in a chemical bond to attract the shared pair of electrons towards itself. (It's a relative property, not an absolute energy value).
• Different scales: Pauling scale (most common), Mulliken scale, Allred-Rochow scale.
• Mulliken Electronegativity: $\text{EN} = \frac{\text{IE} + \text{EA}}{2}$ (where IE and EA are in eV).
• Factors Affecting EN:
  • Atomic Size: Inversely proportional. Smaller atoms have higher EN.
  • Effective Nuclear Charge ($Z_{eff}$): Directly proportional. Higher $Z_{eff}$ means tighter hold on electrons, higher EN.
  • Hybridization: Electronegativity increases with increasing s-character ($sp > sp^2 > sp^3$).
  • Oxidation State: Atoms in higher oxidation states are more electronegative.
• Trends:
  • Across a Period (L to R): Increases. Due to increasing $Z_{eff}$ and decreasing atomic size. F (4.0) is the most electronegative element.
  • Down a Group (Top to Bottom): Decreases. Due to increasing atomic size and increasing shielding. Cesium (0.7) is the least electronegative element (excluding Fr).
• No units for Electronegativity.

Summary of Trends (General)

Common Mistakes & Mnemonics

• IE vs. EA: Remember IE is energy required (always +ve values for the atom), EA is energy released (usually -ve for atom) when adding an electron.
• F vs. Cl EA: Fluorine due to its tiny size experiences greater electron-electron repulsion than chlorine, making its EA less negative than Cl. This is a very common trap.
• Lanthanoid Contraction effect: Remember its effect on 4d and 5d elements and their properties (size, density, chemical similarity).
• Mnemonics for Electronegativity order: FONCl BrIS CH (Fluorine, Oxygen, Nitrogen, Chlorine, Bromine, Iodine, Sulfur, Carbon, Hydrogen) for rough EN scale.

p-Block Elements

The p-block elements are characterized by the filling of p-orbitals in their valence shell. They exhibit a wide range of properties, from non-metals to metalloids to metals, and show diverse oxidation states.

General electronic configuration: $ns^2 np^{1-6}$ (except He which is $1s^2$). Oxidation states: The most common oxidation state is (Group number - 10) for groups 13-16, and (Group number - 18) for groups 17 and 18. They also show variable oxidation states due to the inert pair effect. Inert Pair Effect: The reluctance of the $ns^2$ electrons to participate in bonding, especially in heavier p-block elements, leading to a more stable lower oxidation state (Group Number - 10 - 2, or Group Number - 18 - 2). This effect increases down the group.

Group 13 Elements: The Boron Family ($ns^2 np^1$)

Elements: Boron (B), Aluminium (Al), Gallium (Ga), Indium (In), Thallium (Tl).

• Electronic Configuration: $ns^2 np^1$.
• Oxidation States: +3 (most common), +1 (due to inert pair effect, more prominent for In, Tl). Tl is stable in +1 state.
• Atomic Radii: B < Ga < Al < In < Tl. (Anomaly: Ga is smaller than Al due to poor shielding of 3d electrons in Ga, leading to increased $Z_{eff}$).
• Ionization Enthalpy: Decreases irregularly (B > Tl > Ga > Al > In). Anomalies due to penetration effect and d-electron shielding.
• Electronegativity: B > Tl > In > Ga > Al.
• Nature: Boron is a non-metal (shows allotropy). Aluminum is a metal. Ga, In, Tl are metals.
• Density: Increases down the group.
• M.P./B.P.: B has very high M.P./B.P. (covalent network). Ga has an unusually low M.P. (29.8 °C) and very high B.P. (2400 °C), making it useful in high-temp thermometers.

Important Compounds of Group 13

1. Boron:

   • Diborane ($\text{B}_2\text{H}_6$):
     • Structure: Two $\text{BH}_2$ units joined by two bridging hydrogen atoms. It's an electron-deficient compound (6 BH bonds using 12 electrons, but need 14 electrons).
     • It has 4 terminal B-H bonds (2c-2e) and 2 bridging B-H-B bonds (3c-2e, "banana bonds").
     • Highly reactive towards oxygen.
     • Reacts with water to form boric acid. $\text{B}_2\text{H}_6 + 6\text{H}_2\text{O} \rightarrow 2\text{H}_3\text{BO}_3 + 6\text{H}_2$.
     • Forms addition compounds with Lewis bases (e.g., $\text{B}_2\text{H}_6 + 2\text{N}\text{Me}_3 \rightarrow 2\text{BH}_3\cdot\text{N}\text{Me}_3$).
   • Borax ($\text{Na}_2\text{B}_4\text{O}_7\cdot10\text{H}_2\text{O}$):
     • Contains tetrahedral $\text{BO}_4$ and planar $\text{BO}_3$ units. Structure is $[\text{B}_4\text{O}_5(\text{OH})_4]^{2-}$ in solution.
     • On heating, it swells up and forms a transparent glass bead (borax bead test). $\text{Na}_2\text{B}_4\text{O}_7\cdot10\text{H}_2\text{O} \xrightarrow{\text{heat}} \text{Na}_2\text{B}_4\text{O}_7 \xrightarrow{\text{heat}} 2\text{NaBO}_2 + \text{B}_2\text{O}_3$. $\text{B}_2\text{O}_3$ combines with metal oxides to form colored metaborates.
     • Aqueous solution is alkaline due to hydrolysis: $\text{Na_2B_4O_7} + 7\text{H}_2\text{O} \rightleftharpoons 2\text{NaOH} + 4\text{H}_3\text{BO}_3$.
   • Boric Acid ($\text{H}_3\text{BO}_3$ or $\text{B(OH)}_3$):
     • Weak mono-basic Lewis acid, not a Brønsted acid. It accepts a hydroxyl ion from water.
     • $\text{B(OH)}_3 + \text{H}_2\text{O} \rightleftharpoons [\text{B(OH)}_4]^{-} + \text{H}^+$.
     • Planar, sheet-like structure with hydrogen bonding.
     • On heating: $\text{H}_3\text{BO}_3 \xrightarrow{100^\circ\text{C}} \text{HBO}_2 (\text{metaboric acid}) \xrightarrow{160^\circ\text{C}} \text{H}_2\text{B}_4\text{O}_7 (\text{tetraboric acid}) \xrightarrow{\text{Red hot}} \text{B}_2\text{O}_3 (\text{boric anhydride})$.
   • Boron Nitride (BN): "Inorganic graphite," has a layered structure similar to graphite. Also exists in an extremely hard diamond-like form called "borazon."

2. Aluminium:

   • Amphoteric metal. Reacts with acids ($\text{2Al} + 6\text{HCl} \rightarrow 2\text{AlCl}_3 + 3\text{H}_2$) and bases ($\text{2Al} + 2\text{NaOH} + 6\text{H}_2\text{O} \rightarrow 2\text{Na}[\text{Al(OH)}_4] + 3\text{H}_2$).
   • Does not react with conc. $\text{HNO}_3$ due to the formation of a passive oxide layer.
   • $\text{AlCl}_3$ exists as a dimer ($\text{Al}_2\text{Cl}_6$) in the vapor phase. It's a Lewis acid.

Group 14 Elements: The Carbon Family ($ns^2 np^2$)

Elements: Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), Lead (Pb).

• Electronic Configuration: $ns^2 np^2$.
• Oxidation States: +4 and +2.
  • Stability of +4 state decreases down the group (C > Si > Ge > Sn > Pb).
  • Stability of +2 state increases down the group (C, Si don't show +2 easily; Ge < Sn < Pb). Pb is stable in +2 state due to inert pair effect.
• Catenation: Ability to form bonds with other atoms of the same element.
  • $\text{C} >> \text{Si} > \text{Ge} \approx \text{Sn} >> \text{Pb}$. Carbon shows maximum catenation.
• Atomic Size/Radii: Increases down the group.
• Ionization Enthalpy: Decreases down the group.
• Electronegativity: Fairly constant (Si to Pb). C > Si > Ge = Sn = Pb.
• Nature: C is a non-metal. Si, Ge are metalloids. Sn, Pb are metals.
• Melting Point: Decreases down the group (C > Si > Ge > Sn > Pb).

Important Compounds of Group 14

1. Carbon:

   • Allotropes:
     • Diamond: Hardest natural substance, covalent network solid, $sp^3$ hybridized, insulator ($\text{C-C}$ bond length 154 pm).
     • Graphite: Soft, lustrous, good conductor of electricity, layered structure ($sp^2$ hybridized, $\pi$-electron system), hexagonal rings ( $\text{C-C}$ bond length 141.5 pm), layers held by Van der Waals forces.
     • Fullerenes ($\text{C}_{60}$): "Buckyball," cage-like structure, $sp^2$ hybridized.
     • Carbon Nanotubes: Graphene sheets rolled into tubes.
   • Carbon Monoxide (CO):
     • Colorless, odorless, highly poisonous (forms carboxyhemoglobin).
     • Reducing agent (e.g., in metallurgy, $\text{Fe}_2\text{O}_3 + 3\text{CO} \rightarrow 2\text{Fe} + 3\text{CO}_2$).
     • Neutral oxide.
   • Carbon Dioxide ($\text{CO}_2$):
     • Linear molecule, non-polar.
     • Acidic oxide (reacts with bases, dissolves in water to form carbonic acid $\text{H}_2\text{CO}_3$).
     • Used as a refrigerant ("dry ice").
   • Carbides: Ionic ($\text{e.g.,} \text{CaC}_2$), Covalent ($\text{e.g.,} \text{SiC}$), Interstitial ($\text{e.g.,} \text{TiC}$). $\text{CaC}_2$ on hydrolysis gives acetylene.

2. Silicon:

   • Silica ($\text{SiO}_2$):
     • Covalent network solid, $sp^3$ hybridized silicon. Each Si is tetrahedrally bonded to 4 oxygen atoms, and each oxygen is bonded to 2 silicon atoms.
     • Many forms (Quartz, Cristobalite, Tridymite).
     • Highly inert, reacts with HF ($\text{SiO}_2 + 4\text{HF} \rightarrow \text{SiF}_4 + 2\text{H}_2\text{O}$) and strong bases ($\text{SiO}_2 + 2\text{NaOH} \rightarrow \text{Na}_2\text{SiO}_3 + \text{H}_2\text{O}$).
   • Silicates: Basic structural unit is $\text{SiO}_4^{4-}$ tetrahedron. Can be chain, ring, sheet, or 3D. (e.g., Feldspar, Zeolites, Asbestos, Mica).
   • Silicones (R2SiO)n: Organo-silicon polymers containing $\text{R}_2\text{SiO}$ repeating units. Stable towards heat, resistant to chemicals, water repellent. Used as lubricants, sealants, etc.
     • Preparation: Hydrolysis of $\text{R}_2\text{SiCl}_2$. $(\text{CH}_3)_2\text{SiCl}_2 + 2\text{H}_2\text{O} \rightarrow (\text{CH}_3)_2\text{Si(OH)}_2 + 2\text{HCl}$. Then condensation polymerization: $n(\text{CH}_3)_2\text{Si(OH)}_2 \rightarrow -[\text{O}-\text{Si}(\text{CH}_3)_2]_n- + n\text{H}_2\text{O}$.

Group 15 Elements: The Nitrogen Family ($ns^2 np^3$)

Elements: Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi).

• Electronic Configuration: $ns^2 np^3$.
• Oxidation States: -3, +3, +5.
  • -3: Decreases stability down the group (N, P > As > Sb > Bi).
  • +3: Increases stability down the group (N, P < As < Sb < Bi) due to inert pair effect.
  • +5: Decreases stability down the group (N, P > As > Sb > Bi). Bi(V) is a strong oxidizing agent.
• Atomic Radii: Increases down the group.
• Ionization Enthalpy: Decreases down the group. Generally higher than Group 14 due to half-filled p-orbitals.
• Electronegativity: Decreases down the group.
• Nature: N, P are non-metals. As, Sb are metalloids. Bi is a metal.
• Bonding: N forms p$\pi$-p$\pi$ multiple bonds (e.g., $\text{N}_2$). Heavier elements do not due to larger atomic size and diffuse orbitals.

Important Compounds of Group 15

1. Nitrogen:

   • Dinitrogen ($\text{N}_2$):
     • Inert at room temperature (high bond dissociation enthalpy of triple bond).
     • Used in Haber process for ammonia synthesis ($\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$).
   • Ammonia ($\text{NH}_3$):
     • Pyramidal shape, $sp^3$ hybridized, has a lone pair, hence Lewis base.
     • Forms hydrogen bonds, highly soluble in water.
     • Forms complex compounds with transition metal ions.
   • Nitric Acid ($\text{HNO}_3$):
     • Prepared by Ostwald process.
     • Strong oxidizing agent. Reacts with metals, reducing itself to various products depending on conditions (conc. $\rightarrow \text{NO}_2$; dil. $\rightarrow \text{NO}$; very dil. $\rightarrow \text{N}_2\text{O}$ or $\text{NH}_4^+$).
     • Passivates Fe, Cr, Al.
   • Oxides of Nitrogen:
     • $\text{N}_2\text{O}$ (Nitrous oxide) - Neutral, "laughing gas".
     • $\text{NO}$ (Nitric oxide) - Neutral, paramagnetic.
     • $\text{N}_2\text{O}_3$ (Dinitrogen trioxide) - Acidic, blue liquid.
     • $\text{NO}_2$ (Nitrogen dioxide) - Acidic, brown gas, paramagnetic. Dimerizes to $\text{N}_2\text{O}_4$ (colorless, diamagnetic).
     • $\text{N}_2\text{O}_4$ (Dinitrogen tetroxide) - Acidic.
     • $\text{N}_2\text{O}_5$ (Dinitrogen pentoxide) - Acidic.

2. Phosphorus:

   • Allotropes:
     • White Phosphorus ($\text{P}_4$): Discrete tetrahedral molecules, highly reactive, pyrophoric (ignites in air), soluble in $\text{CS}_2$, poisonous.
     • Red Phosphorus: Polymeric, less reactive, insoluble in $\text{CS}_2$, non-poisonous. Formed by heating white P at 573 K in inert atmosphere.
     • Black Phosphorus: Most stable, made under high pressure.
   • Phosphine ($\text{PH}_3$):
     • Less basic than $\text{NH}_3$.
     • Highly poisonous, used in Holme's signals.
     • Spontaneously flammable (due to $\text{P}_2\text{H}_4$ impurities).
   • Phosphorus Halides:
     • $\text{PCl}_3$: Pyramidal, $sp^3$ hybridized P. Reacts with organic compounds having -OH groups.
     • $\text{PCl}_5$: Trigonal bipyramidal in gaseous/liquid state ($sp^3d$). In solid state, exists as $[\text{PCl}_4]^+[\text{PCl}_6]^-$.
   • Oxoacids of Phosphorus:
     • Hypophosphorous acid ($\text{H}_3\text{PO}_2$): Strong reducing agent, monobasic (1 P-OH, 2 P-H).
     • Phosphorous acid ($\text{H}_3\text{PO}_3$): Dibasic (2 P-OH, 1 P-H), reducing agent.
     • Orthophosphoric acid ($\text{H}_3\text{PO}_4$): Tribasic (3 P-OH), used in fertilizers.
     • Pyrophosphoric acid ($\text{H}_4\text{P}_2\text{O}_7$).
     • Hypophosphoric acid ($\text{H}_4\text{P}_2\text{O}_6$).
     • Key Concept: Only H atoms attached to O are acidic. H atoms directly attached to P are not acidic but contribute to reducing properties.

Group 16 Elements: The Oxygen Family (Chalcogens) ($ns^2 np^4$)

Elements: Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), Polonium (Po).

• Electronic Configuration: $ns^2 np^4$.
• Oxidation States: -2, +2, +4, +6.
  • -2: Oxygen generally shows -2 (except in $\text{OF}_2$ (+2), peroxides (-1), superoxides (-1/2)). Stability decreases down the group.
  • +4, +6: More stable for heavier elements due to inert pair effect. Po shows +2 more prominently.
• Atomic Radii: Increases down the group.
• Ionization Enthalpy: Decreases down the group. Lower than Group 15 due to less stable $np^4$ configuration.
• Electronegativity: Decreases down the group. Oxygen is the second most electronegative element (3.5).
• Nature: O, S are non-metals. Se, Te are metalloids. Po is a metal (radioactive).
• Allotropy: All elements show allotropy (O: $\text{O}_2, \text{O}_3$; S: rhombic, monoclinic, plastic etc.).
• Hydrides ($\text{H}_2\text{E}$): $\text{H}_2\text{O}, \text{H}_2\text{S}, \text{H}_2\text{Se}, \text{H}_2\text{Te}$.
  • Boiling Point: $\text{H}_2\text{O} > \text{H}_2\text{Te} > \text{H}_2\text{Se} > \text{H}_2\text{S}$. (Anomaly: $\text{H}_2\text{O}$ has highest due to H-bonding).
  • Acidic Strength: $\text{H}_2\text{O} < \text{H}_2\text{S} < \text{H}_2\text{Se} < \text{H}_2\text{Te}$. (Increases down group due to decreasing bond dissociation enthalpy).
  • Reducing Property: $\text{H}_2\text{O} < \text{H}_2\text{S} < \text{H}_2\text{Se} < \text{H}_2\text{Te}$. (Increases down group due to decreasing stability of bond).

Important Compounds of Group 16

1. Oxygen:

   • $\text{O}_2$ (Dioxygen): Paramagnetic.
   • $\text{O}_3$ (Ozone):
     • Prepared by silent electric discharge through dry oxygen.
     • Powerful oxidizing agent. Decomposes to $\text{O}_2$ (exothermic).
     • Bent structure ($sp^2$ hybridized).
     • Liberates iodine from iodide solution: $\text{2I}^- + \text{H}_2\text{O} + \text{O}_3 \rightarrow \text{2OH}^- + \text{I}_2 + \text{O}_2$.

2. Sulfur:

   • Allotropes:
     • Rhombic Sulfur ($\text{S}_\alpha$): Most stable form at room temperature, yellow, octahedral crystals. Contains $\text{S}_8$ puckered ring structure.
     • Monoclinic Sulfur ($\text{S}_\beta$): Stable above 369 K, dull yellow, needle-like crystals. Contains $\text{S}_8$ rings.
   • Sulfur Dioxide ($\text{SO}_2$):
     • Colorless gas, pungent smell. Bent structure.
     • Reducing agent (e.g., reduces $\text{Fe}^{3+}$ to $\text{Fe}^{2+}$, decolorizes $\text{KMnO}_4$).
     • Bleaching agent (temporary, by reduction: $\text{dye} + \text{SO}_2 + 2\text{H}_2\text{O} \rightarrow \text{dyeH}_2\text{SO}_4$).
     • Forms $\text{SO}_3^{2-}$ in water (sulfurous acid).
   • Sulfuric Acid ($\text{H}_2\text{SO}_4$):
     • Prepared by Contact Process.
     • Strong acid, dehydrating agent, oxidizing agent.
     • $\text{C} + 2\text{H}_2\text{SO}_4(\text{conc}.) \rightarrow \text{CO}_2 + 2\text{SO}_2 + 2\text{H}_2\text{O}$.
     • $\text{S} + 2\text{H}_2\text{SO}_4(\text{conc}.) \rightarrow 3\text{SO}_2 + 2\text{H}_2\text{O}$.
     • $\text{P}_4 + 10\text{H}_2\text{SO}_4(\text{conc}.) \rightarrow 4\text{H}_3\text{PO}_4 + 10\text{SO}_2 + 4\text{H}_2\text{O}$.
     • Has high viscosity and density due to hydrogen bonding.

Group 17 Elements: The Halogen Family ($ns^2 np^5$)

Elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At).

• Electronic Configuration: $ns^2 np^5$.
• Oxidation States: -1 (most common), +1, +3, +5, +7 (except F, which only shows -1).
  • -1: Stability decreases down the group.
  • Higher positive oxidation states are shown by Cl, Br, I when bonded to more electronegative oxygen or fluorine.
• Atomic and Ionic Radii: Increases down the group.
• Ionization Enthalpy: Decreases down the group. Very high due to high $Z_{eff}$.
• Electron Gain Enthalpy: Most negative in the periodic table. $\text{Cl} > \text{F} > \text{Br} > \text{I}$. (Anomaly: F is less negative than Cl).
• Electronegativity: Decreases down the group. F is the most electronegative element.
• Nature: All are non-metals. F, Cl are gases, Br is liquid, I is solid.
• Bond Dissociation Enthalpy: $\text{Cl}_2 > \text{Br}_2 > \text{F}_2 > \text{I}_2$. (Anomaly: $\text{F}_2$ is unexpectedly low due to strong inter-electronic repulsion in small F-F bond and the resulting weaker bond).
• Oxidizing Power: $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$. Reactivity decreases down the group. Halogens oxidize halide ions of higher atomic number. (e.g., $\text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2$).

Important Compounds of Group 17

1. Fluorine:

   • Most reactive of all halogens.
   • Reacts with water: $2\text{F}_2 + 2\text{H}_2\text{O} \rightarrow 4\text{HF} + \text{O}_2$.
   • Hydrofluoric acid (HF) is a weak acid but attacks glass. It is stored in plastic bottles.

2. Chlorine:

   • Prepared by Deacon's process ($\text{HCl} + \text{O}_2 \rightarrow \text{Cl}_2 + \text{H}_2\text{O}$) or Electrolysis of brine.
   • Oxidizing agent, bleaching agent (permanent, by oxidation: $\text{Cl}_2 + \text{H}_2\text{O} \rightarrow \text{HCl} + \text{HOCl}$; $\text{HOCl} \rightarrow \text{HCl} + [\text{O}] $).
   • Reacts with cold, dilute $\text{NaOH}$ to form $\text{NaCl}$ and $\text{NaClO}$. $\text{Cl}_2 + 2\text{NaOH}(\text{cold, dil}) \rightarrow \text{NaCl} + \text{NaClO} + \text{H}_2\text{O}$.
   • Reacts with hot, concentrated $\text{NaOH}$ to form $\text{NaCl}$ and $\text{NaClO}_3$. $3\text{Cl}_2 + 6\text{NaOH}(\text{hot, conc}) \rightarrow 5\text{NaCl} + \text{NaClO}_3 + 3\text{H}_2\text{O}$.

3. Hydrohalic Acids (HX):

   • Acidic Strength: $\text{HF} < \text{HCl} < \text{HBr} < \text{HI}$. (HF is weak due to strong H-bonding).
   • Reducing Power: $\text{HF} < \text{HCl} < \text{HBr} < \text{HI}$.

4. Oxoacids of Halogens:

   • Can't be isolated in pure form (except $\text{HClO}_4$).
   • Acidic strength increases with increasing oxidation state of halogen: $\text{HClO} < \text{HClO}_2 < \text{HClO}_3 < \text{HClO}_4$.
   • Acidic strength of similar oxoacids: $\text{HClO}_4 > \text{HBrO}_4 > \text{HIO}_4$. (Due to decreasing electronegativity, less ability to stabilize conjugate base).

5. Interhalogen Compounds:

   • Formed between two different halogens (e.g., $\text{ClF}, \text{ClF}_3, \text{ClF}_5, \text{BrF}_3, \text{IF}_7$).
   • General formula: $\text{AX}_n$ (where A is a larger halogen, X is a smaller halogen, n = 1, 3, 5, 7).
   • More reactive than halogens (except $\text{F}_2$) because A-X bond is weaker than X-X bond in halogens.
   • Hydrolyze to form halide and oxyhalide acids.

Group 18 Elements: The Noble Gases ($ns^2 np^6$)

Elements: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn), Oganesson (Og).

• Electronic Configuration: $ns^2 np^6$ (He is $1s^2$).
• Oxidation States: Generally 0. Xe forms compounds with highly electronegative elements like F and O (e.g., $\text{XeF}_2, \text{XeF}_4, \text{XeF}_6, \text{XeO}_3, \text{XeOF}_4$).
• Ionization Enthalpy: Very high, decreases down the group.
• Electron Gain Enthalpy: Positive (unstable to add an electron).
• Atomic Radii: Largest in their respective periods (Van der Waals radii). Increases down the group.
• Boiling Point: Very low, increases down the group (due to increasing Van der Waals forces).
• Nature: Monatomic gases, unreactive due to stable octet.

Important Compounds of Noble Gases (Xenon)

• $\text{XeF}_2$: Linear. Formed by $\text{Xe} + \text{F}_2$ (1:5 ratio, 673K, 1 bar).
• $\text{XeF}_4$: Square planar. Formed by $\text{Xe} + \text{F}_2$ (1:5 ratio, 873K, 7 bar).
• $\text{XeF}_6$: Distorted octahedral. Formed by $\text{Xe} + \text{F}_2$ (1:20 ratio, 573K, 60-70 bar) or $\text{XeF}_4 + \text{O}_2\text{F}_2$.
• Hydrolysis of Xenon Fluorides:
  • $\text{2XeF}_2 + 2\text{H}_2\text{O} \rightarrow 2\text{Xe} + 4\text{HF} + \text{O}_2$.
  • $\text{6XeF}_4 + 12\text{H}_2\text{O} \rightarrow 4\text{Xe} + 2\text{XeO}_3 + 24\text{HF} + 3\text{O}_2$.
  • $\text{XeF}_6 + 3\text{H}_2\text{O} \rightarrow \text{XeO}_3 + 6\text{HF}$. ($\text{XeO}_3$ is explosive).
  • Partial hydrolysis of $\text{XeF}_6$ gives $\text{XeOF}_4$ and $\text{XeO}_2\text{F}_2$.
• Bonding in Xe compounds: Xe uses its empty d-orbitals for expansion of octet. Example: In $\text{XeF}_2$, Xe is $sp^3d$ hybridized with 3 lone pairs. In $\text{XeF}_4$, Xe is $sp^3d^2$ hybridized with 2 lone pairs.

Worked Examples

Example 1: Periodic Trends Among the elements B, Al, C, Si, N, the correct order of decreasing first ionization enthalpy is: (A) $\text{N} > \text{Al} > \text{B} > \text{C} > \text{Si}$ (B) $\text{N} > \text{C} > \text{B} > \text{Al} > \text{Si}$ (C) $\text{N} > \text{C} > \text{Si} > \text{Al} > \text{B}$ (D) $\text{B} > \text{Al} > \text{C} > \text{Si} > \text{N}$

Solution:

1. Identify elements and their positions:

   • Period 2: B (Group 13), C (Group 14), N (Group 15)
   • Period 3: Al (Group 13), Si (Group 14)

2. General trend: IE increases across a period and decreases down a group.

   • For Period 2: $\text{B} < \text{C} < \text{N}$ (expected, with some anomalies).
   • For Group 13: $\text{B} > \text{Al}$.
   • For Group 14: $\text{C} > \text{Si}$.

3. Specific comparisons and anomalies:

   • Across Period 2: $\text{IE}_1(\text{C}) > \text{IE}_1(\text{B})$. $\text{IE}_1(\text{N}) > \text{IE}_1(\text{C})$ because N has a stable half-filled p-orbital configuration ($2p^3$). So, $\text{N} > \text{C} > \text{B}$.
   • Compare Group 13 elements: $\text{B} > \text{Al}$.
   • Compare Group 14 elements: $\text{C} > \text{Si}$.
   • Now, relative position for Al and Si: Si is to the right of Al in the same period (Period 3). So, $\text{IE}_1(\text{Si}) > \text{IE}_1(\text{Al})$.
   • The main "crossing over" happens when comparing elements from different groups/periods.
   • Comparing B and Al: B is smaller, so $\text{IE}_1(\text{B}) > \text{IE}_1(\text{Al})$.
   • Comparing C and Si: C is smaller, so $\text{IE}_1(\text{C}) > \text{IE}_1(\text{Si})$.
   • Now, place Al and Si relative to B, C, N. Both Al and Si are significantly larger than B, C, N and thus will have lower IE.
   • So, we have a general order of $\text{N} > \text{C} > \text{B} > \text{Si} > \text{Al}$ (considering the general trend and the N anomaly).
   • Let's check the values (approx. in kJ/mol):
     • N: 1402
     • C: 1086
     • B: 801
     • Si: 786
     • Al: 577
   • The correct order is indeed $\text{N} > \text{C} > \text{B} > \text{Si} > \text{Al}$.

Final Answer: (B) $\text{N} > \text{C} > \text{B} > \text{Al} > \text{Si}$

Example 2: p-Block Reactions An aqueous solution of borax is: (A) acidic (B) basic (C) neutral (D) amphoteric

Solution:

1. Recall the nature of borax: Borax is $\text{Na}_2\text{B}_4\text{O}_7\cdot10\text{H}_2\text{O}$.
2. Hydrolysis: When borax dissolves in water, it undergoes hydrolysis. $\text{Na}_2\text{B}_4\text{O}_7 + 7\text{H}_2\text{O} \rightleftharpoons 2\text{NaOH} + 4\text{H}_3\text{BO}_3$.
3. Strength of products:
   • $\text{NaOH}$ is a strong base.
   • $\text{H}_3\text{BO}_3$ (boric acid) is a very weak acid.
4. Conclusion: Since a strong base and a weak acid are formed, the resulting aqueous solution will be basic.

Final Answer: (B) basic

Example 3: Hybridization and Shape The hybridization of boron in diborane ($\text{B}_2\text{H}_6$) is: (A) $sp$ (B) $sp^2$ (C) $sp^3$ (D) $dsp^2$

Solution:

1. Diborane structure: Diborane is $\text{B}_2\text{H}_6$. It has a unique structure containing two boron atoms and six hydrogen atoms.
2. Terminal and bridging hydrogens: There are 4 terminal B-H bonds and 2 bridging B-H-B bonds.
3. Coordination of Boron: Each boron atom is bonded to two terminal hydrogen atoms and two bridging hydrogen atoms. Therefore, each boron atom is surrounded by 4 groups (2 terminal H, 2 bridging H).
4. Hybridization: For 4 groups (4 areas of electron density), the hybridization is $sp^3$.
5. Nature of bonds: Even though the bridging bonds are 3-center-2-electron (banana bonds), the boron atom itself still forms four effective bonds (two 2c-2e and two 3c-2e), implying a tetrahedral geometry around each boron and thus $sp^3$ hybridization.

Final Answer: (C) $sp^3$

Example 4: Relative Reactivity Which of the following halogens has the highest bond dissociation enthalpy? (A) $\text{F}_2$ (B) $\text{Cl}_2$ (C) $\text{Br}_2$ (D) $\text{I}_2$

Solution:

1. General trend: Down a group, bond dissociation enthalpy generally decreases due to increasing atomic size and bond length. So, $\text{X}_2$ bond length increases from $\text{F}_2$ to $\text{I}_2$.
2. Expected order: Based on size, one would expect $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$.
3. Anomaly: The bond dissociation enthalpy of $\text{F}_2$ is unexpectedly low. This is because fluorine is very small, and the lone pairs on the two fluorine atoms in $\text{F}_2$ experience significant inter-electronic repulsion, weakening the F-F bond.
4. Actual order: Therefore, it is $\text{Cl}_2 > \text{Br}_2 > \text{F}_2 > \text{I}_2$.
5. Conclusion: Chlorine has the highest bond dissociation enthalpy among the halogens.

Final Answer: (B) $\text{Cl}_2$

Example 5: Oxoacids of Phosphorus Among the oxoacids of phosphorus, the one which is a strong reducing agent and monobasic is: (A) $\text{H}_3\text{PO}_4$ (B) $\text{H}_3\text{PO}_3$ (C) $\text{H}_3\text{PO}_2$ (D) $\text{H}_4\text{P}_2\text{O}_7$

Solution:

1. Monobasic: An acid is monobasic if it has only one dissociable H atom (i.e., one P-OH bond).
2. Reducing agent: The reducing character usually relates to the presence of P-H bonds.
3. Structures and properties:
   • (A) $\text{H}_3\text{PO}_4$ (Orthophosphoric acid): 3 P-OH bonds, no P-H bonds. Tribasic. Not a strong reducing agent.
   • (B) $\text{H}_3\text{PO}_3$ (Phosphorous acid): 2 P-OH bonds, 1 P-H bond. Dibasic. Reducing agent.
   • (C) $\text{H}_3\text{PO}_2$ (Hypophosphorous acid): 1 P-OH bond, 2 P-H bonds. Monobasic. Strong reducing agent.
   • (D) $\text{H}_4\text{P}_2\text{O}_7$ (Pyrophosphoric acid): 4 P-OH bonds, no P-H bonds. Tetrabasic. Not a reducing agent.
4. Conclusion: $\text{H}_3\text{PO}_2$ fits both criteria: it is monobasic and a strong reducing agent.

Final Answer: (C) $\text{H}_3\text{PO}_2$

Example 6: Hybridization in Xenon Fluorides What is the hybridization of Xenon in $\text{XeF}_4$? (A) $sp^3$ (B) $sp^3d$ (C) $sp^3d^2$ (D) $sp^2$

Solution:

1. Count valence electrons: Xenon (central atom) has 8 valence electrons. Each fluorine atom contributes 1 electron to form a single bond. There are 4 fluorine atoms.
2. Calculate steric number (SN):
   • Number of bond pairs = 4 (Xe-F bonds).
   • Number of lone pairs = (Total valence electrons on central atom - Electrons used in bonding) / 2 = $(8 - 4 \times 1) / 2 = 4 / 2 = 2$ lone pairs.
   • Steric Number (SN) = Number of bond pairs + Number of lone pairs = $4 + 2 = 6$.
3. Determine hybridization: A steric number of 6 corresponds to $sp^3d^2$ hybridization.
4. Predict geometry (VSEPR): Shape due to 6 electron pairs is octahedral. With 4 bond pairs and 2 lone pairs, the actual molecular geometry is square planar (lone pairs occupy opposite positions to minimize repulsion).

Final Answer: (C) $sp^3d^2$

Common Mistakes and How to Avoid Them

1. IE/EA/EN Trends:

   • Mistake: Blindly applying general trends without considering exceptions.
   • Correction: Always remember the exceptions mentioned:
     • $\text{IE}_1(\text{Be}) > \text{IE}_1(\text{B})$, $\text{IE}_1(\text{Mg}) > \text{IE}_1(\text{Al})$ (due to $s^2$ stability).
     • $\text{IE}_1(\text{N}) > \text{IE}_1(\text{O})$, $\text{IE}_1(\text{P}) > \text{IE}_1(\text{S})$ (due to half-filled $p^3$ stability).
     • $\Delta H_{eg}(\text{Cl}) > \Delta H_{eg}(\text{F})$ (F is smaller, higher electron-electron repulsion). Similar for 2nd period vs 3rd period for other groups.
     • Bond dissociation enthalpy: $\text{Cl}_2 > \text{Br}_2 > \text{F}_2 > \text{I}_2$ ($\text{F}_2$ anomaly due to small size and lone pair repulsion).

2. Lanthanoid Contraction:

   • Mistake: Forgetting its implications.
   • Correction: Remember it causes 4d and 5d series elements (especially from Group 4 onwards) to have similar atomic radii and chemical properties (e.g., Zr/Hf, Nb/Ta). Also leads to higher density of 5d elements.

3. Nature of Oxides:

   • Mistake: Assuming all non-metal oxides are acidic and all metal oxides are basic.
   • Correction:
     • Some non-metal oxides are neutral ($\text{CO, NO, N}_2\text{O}$).
     • Some metal oxides are amphoteric ($\text{Al}_2\text{O}_3, \text{Ga}_2\text{O}_3, \text{SnO}_2, \text{PbO}_2$). Group 13 oxides show an increase in basicity down the group, with $\text{Al}_2\text{O}_3$ amphoteric and $\text{Tl}_2\text{O}_3$ basic.

4. $\text{H}_3\text{BO}_3$ as Acid:

   • Mistake: Considering it a Brønsted acid (proton donor).
   • Correction: Boric acid is a weak mono-basic Lewis acid. It accepts a hydroxyl ion from water, not donate a proton. $\text{B(OH)}_3 + 2\text{H}_2\text{O} \rightleftharpoons [\text{B(OH)}_4]^- + \text{H}_3\text{O}^+$.

5. Basicity/Acidity of Hydrides/Oxoacids:

   • Mistake: Confusing trends across period vs. down group.
   • Correction:
     • Hydrides of group 15: $\text{NH}_3 > \text{PH}_3 > \text{AsH}_3 > \text{SbH}_3 > \text{BiH}_3$ (basic strength decreases).
     • Hydrides of group 16: $\text{H}_2\text{O} < \text{H}_2\text{S} < \text{H}_2\text{Se} < \text{H}_2\text{Te}$ (acidic strength increases).
     • Oxoacids with same central atom (e.g., Cl): $\text{HClO} < \text{HClO}_2 < \text{HClO}_3 < \text{HClO}_4$ (acidity increases with oxidation state).
     • Oxoacids with different central atom in same group (e.g., $\text{HXO}_4$): $\text{HIO}_4 < \text{HBrO}_4 < \text{HClO}_4$ (acidity increases with EN of halogen).

6. Shape and Hybridization:

   • Mistake: Not accounting for lone pairs correctly in VSEPR theory.
   • Correction: Always calculate steric number (bond pairs + lone pairs) to determine hybridization. Then use VSEPR rules to predict the molecular geometry by minimizing repulsion. For instance, $\text{XeF}_4$ ($sp^3d^2$ and square planar, not octahedral). $\text{NH}_3$ ($sp^3$ and pyramidal, not tetrahedral).

7. Inert Pair Effect:

   • Mistake: Forgetting that it promotes lower oxidation states for heavier p-block elements.
   • Correction: Remember that for Groups 13, 14, 15, the stability of (Group number - 10 - 2) or (Group number - 18 - 2) oxidation state increases down the group. (e.g., $\text{Pb}^{2+}$ is more stable than $\text{Pb}^{4+}$).

Memory Tricks/Mnemonics

• Order of Ionization Energy Exceptions (Period 2): Be Boasts No Ordinary (Be > B, N > O).
• Electron Affinity Order of Halogens: Cute Funny Boys Ignore (Cl > F > Br > I).
• Acidic oxides (Group 14): Carbon Dioxide $\text{CO}_2$ is Acidic. Carbon Monoxide $\text{CO}$ is Neutral (think neutral - mono).
• Reducing Nature of Hydrides of Group 15: 'No PHDs Here' (NH3, PH3, AsH3, SbH3, BiH3). Basicity decreases, reducing nature increases.
• Order of stability of Hydrides (thermal stability): 'What Se Te' (H2O > H2S > H2Se > H2Te). Thermal stability decreases, acidic nature increases.
• Hybridization and lone pairs in Xe compounds:
  • $\text{XeF}_2$: 2 bonds, 3 lone pairs (Linear, $sp^3d$)
  • $\text{XeF}_4$: 4 bonds, 2 lone pairs (Square planar, $sp^3d^2$)
  • $\text{XeF}_6$: 6 bonds, 1 lone pair (Distorted Octahedral, $sp^3d^3$)
  • Count total bonds + lone pairs for steric number.

Previous Year Question Patterns

JEE Main questions from the Periodic Table and p-Block can be broadly categorized as:

1. Direct Trend-based Questions:

   • Comparing atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity for a given set of elements.
   • Often involves exceptions to general trends. (e.g., $\text{F}/\text{Cl}$ EA, $\text{Be}/\text{B}$ IE).
   • Example: Order of acidic strength, reducing power of hydrides, oxides, etc.

2. Structural Questions:

   • Hybridization, shape, bond angles, number of lone pairs in p-block compounds (e.g., Diborane, $\text{PCl}_5, \text{XeF}_4$, oxoacids).
   • Isostructural or isoelectronic species identification.

3. Reaction-based Questions:

   • Reactions of Borax, Boric acid, Ammonia, Phosphine, Halogens with bases/water, $\text{SO}_2$, $\text{H}_2\text{SO}_4$, Xenon fluorides.
   • Identifying products, reactants, or conditions.
   • Properties (oxidizing/reducing agent, dehydrating agent).

4. Properties & Uses Questions:

   • Allotropes of Carbon (diamond, graphite properties), Phosphorus, Sulphur.
   • Unique properties like paramagnetism ($\text{O}_2, \text{NO}, \text{NO}_2$), inert pair effect consequences.
   • Uses of compounds (e.g., Ga in thermometers, Xe compounds as powerful oxidizing agents, Silicones).

5. Identification of Substances:

   • Via flame test, borax bead test (though less common in Main, more in Advanced).
   • Reactions involving color changes.

Recent Trends:

• More concept-based questions rather than rote memorization of obscure facts.
• Emphasis on exceptions to periodic trends.
• Hybridization and VSEPR theory applications remain crucial for p-block compounds.
• Reactions of Group 13, 15, 16, 17, and Noble gases are important, particularly those with water, acids, and bases.

Revision Checklist

• Periodic Table:
  • Modern Periodic Law & its basis.
  • Classification of elements (s, p, d, f block).
  • Periodicity and definitions of atomic radius, ionic radius, IE, EA, EN.
  • General trends across periods and down groups.
  • All major exceptions to trends and their reasons (F/Cl EA, N/O IE, Be/B IE, Lanthanoid Contraction effect, $\text{F}_2$ bond dissociation enthalpy).
  • Isoelectronic species.
  • Amphoteric nature (Al, Ga, Sn, Pb).
• p-Block Elements (Group-wise):
  • General electronic configuration.
  • Oxidation states and inert pair effect.
  • General trends in physical and chemical properties within each group.
  • Group 13 (Boron Family):
    • Anomaly in atomic radii (Al-Ga).
    • Diborane: Structure, bonding, reactions ($\text{H}_2\text{O}$, Lewis bases).
    • Borax: Hydrolysis, borax bead test.
    • Boric acid: Structure, nature as Lewis acid, heating effects.
    • Aluminium: Amphoteric nature, passivity with $\text{HNO}_3$, $\text{AlCl}_3$ dimer.
  • Group 14 (Carbon Family):
    • Allotropes of Carbon (properties, hybridization).
    • Hydrides (catenation).
    • Oxides of Carbon ($\text{CO}, \text{CO}_2$ - nature, properties).
    • Silica ($\text{SiO}_2$): Structure, reactions ($\text{HF}$, bases).
    • Silicates: Basic unit.
    • Silicones: Preparation, structure, properties.
    • Stability of +2 and +4 oxidation states.
  • Group 15 (Nitrogen Family):
    • Allotropes of Phosphorus.
    • $\text{N}_2$: Inertness, uses.
    • $\text{NH}_3$: Lewis basicity, H-bonding, complexes.
    • $\text{HNO}_3$: Preparation (Ostwald), oxidizing nature, passivity.
    • Oxides of Nitrogen: Nature (acidic/neutral), oxidation states, paramagnetism.
    • Phosphine ($\text{PH}_3$): Basicity, toxicity.
    • Phosphorus Halides ($\text{PCl}_3, \text{PCl}_5$): Structure, reactions, solid state of $\text{PCl}_5$.
    • Oxoacids of Phosphorus: Structure, basicity, reducing character (P-H bonds).
  • Group 16 (Oxygen Family):
    • Hydrides ($\text{H}_2\text{X}$): Boiling point anomaly ($\text{H}_2\text{O}$), acidic strength, reducing power.
    • Oxygen: Paramagnetism of $\text{O}_2$.
    • Ozone ($\text{O}_3$): Preparation, oxidizing nature, structure.
    • Allotropes of Sulfur (Rhombic, Monoclinic).
    • Sulfur Dioxide ($\text{SO}_2$): Structure, reducing and bleaching action ($\text{KMnO}_4$, dyes).
    • Sulfuric Acid ($\text{H}_2\text{SO}_4$): Preparation (Contact), oxidizing, dehydrating properties.
  • Group 17 (Halogen Family):
    • Electron gain enthalpy order.
    • Bond dissociation enthalpy order.
    • Oxidizing power.
    • Reactions of $\text{Cl}_2$ with cold/hot $\text{NaOH}$.
    • Hydrohalic acids (HX): Acidic strength, reducing power.
    • Oxoacids of Halogens: Acidity trend with oxidation state and electronegativity.
    • Interhalogen Compounds: Formula, reactivity, hydrolysis.
  • Group 18 (Noble Gases):
    • Inertness, boiling points.
    • Xenon compounds ($\text{XeF}_2, \text{XeF}_4, \text{XeF}_6, \text{XeO}_3, \text{XeOF}_4$): Preparation, shapes, hybridization, hydrolysis reactions.

Last-Day Strategy

1. Prioritize Exceptions and Anomalies: These are favorite areas for JEE Main questions. Quickly review the list of exceptions for IE, EA, EN, bond dissociation enthalpy, atomic radii (Ga-Al, Zr-Hf), and check you know the reasons.
2. Functional Structures: For p-block, quickly recall the structures, hybridization, and number of lone pairs for compounds like Diborane, $\text{PCl}_5$, oxoacids of Phosphorus, $\text{SO}_2$, $\text{Cl F}_3$, $\text{XeF}_4, \text{XeO}_3$.
3. Key Reactions: Focus on hydrolysis reactions (Borax, Diborane, $\text{XeF}_x$, hydrides), reactions of halogens with alkalies, oxidizing/reducing nature (e.g., $\text{HNO}_3$, $\text{H}_2\text{SO}_4$, $\text{SO}_2$, $\text{O}_3$, $\text{H}_3\text{PO}_2$).
4. Trends in Acidic/Basic Properties: Review how acidity/basicity changes for oxides, hydroxides, and hydrides across periods and down groups.
5. Don't Start New Topics: Stick to quick revision of what you've already studied. Avoid diving into new concepts or complex derivations.
6. Practice Previous Year Questions: Solve a few last-minute PYQs to get into the exam mindset and identify any lingering weak spots. Prioritize questions that combine multiple concepts.
7. Positive Mindset: Inorganic Chemistry often relies on recall. Trust your preparation and stay calm. A clear mind helps retrieve information efficiently.
8. Formula & Concept Sheet: If you've maintained a concise revision sheet, go through it one last time. Highlight the most crucial points.