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Introduction to p-Block Elements The p-block elements are located from Group 13 to Group 18 of the periodic table. Their distinguishing feature is that the last electron enters the outermost p-orbital. The general outer electronic configuration is $ns^2 np^{1-6}$ (except for Helium, which has $1s^2$). Key characteristics: Include metals, non-metals, and metalloids. Exhibit a wide variation in properties. Oxidation states vary significantly across the groups. Tendency to show inert pair effect down the group (stability of lower oxidation state increases). Group 13: Boron Family ($ns^2 np^1$) Elements: B, Al, Ga, In, Tl Boron (B): Non-metal, very hard, high melting point, forms covalent compounds. Aluminum (Al): Metal, most abundant metal in Earth's crust, forms ionic compounds. Gallium (Ga), Indium (In), Thallium (Tl): Metals. General Trends: Atomic Radii: Increases down the group. Anomaly: Ga has smaller radius than Al due to poor shielding by d-electrons in Ga. Ionization Enthalpy: Decreases from B to Al, then increases for Ga (d-electrons), decreases for In, then increases for Tl (f- and d-electrons). Electronegativity: Decreases from B to Al, then increases slightly. Oxidation States: +3 is common. +1 becomes more stable down the group (inert pair effect, e.g., $Tl^+$ is more stable than $Tl^{3+}$). Important Compounds: Boron Hydrides (Boranes): e.g., Diborane ($B_2H_6$). Structure: Two $B-H-B$ bridge bonds (3-center-2-electron bonds) and four terminal $B-H$ bonds (2-center-2-electron bonds). Preparation: $2NaBH_4 + I_2 \rightarrow B_2H_6 + 2NaI + H_2$ Boron Trihalides ($BX_3$): Lewis acids (electron deficient). Acidity order: $BF_3 Boric Acid ($H_3BO_3$): Weak monobasic Lewis acid. $B(OH)_3 + 2H_2O \rightleftharpoons [B(OH)_4]^- + H_3O^+$. Borax ($Na_2B_4O_7 \cdot 10H_2O$): Used in borax bead test for colored metal salts. Alum ($K_2SO_4 \cdot Al_2(SO_4)_3 \cdot 24H_2O$): Double salt, used as a coagulant. Group 14: Carbon Family ($ns^2 np^2$) Elements: C, Si, Ge, Sn, Pb Carbon (C): Non-metal, forms extensive covalent bonds, allotropy (diamond, graphite, fullerenes). Silicon (Si): Metalloid, semiconductor. Germanium (Ge): Metalloid, semiconductor. Tin (Sn): Metal, allotropes (white, grey, rhombic). Lead (Pb): Metal, soft, heavy. General Trends: Atomic Radii: Increases down the group. Ionization Enthalpy: Decreases down the group. Electronegativity: Decreases from C to Si, then remains almost constant. Oxidation States: +4 and +2. +2 becomes more stable down the group (inert pair effect, e.g., $Pb^{2+}$ is more stable than $Pb^{4+}$). Catenation: Tendency to form bonds with identical atoms. Order: $C >> Si > Ge \approx Sn$. Pb shows very little catenation. Important Compounds: Carbon Monoxide (CO): Neutral oxide, highly toxic, reducing agent. Carbon Dioxide ($CO_2$): Acidic oxide, greenhouse gas. Silicon Dioxide ($SiO_2$): Covalent network solid (quartz). Acidic. Silicates: Basic structural unit is $SiO_4^{4-}$ tetrahedron. e.g., feldspar, mica, asbestos. Silicones: Organosilicon polymers containing $R_2SiO$ units. Water repellent, high thermal stability. Carbides: Ionic ($CaC_2$), covalent ($SiC$), interstitial ($TiC$). Group 15: Nitrogen Family ($ns^2 np^3$) Elements: N, P, As, Sb, Bi Nitrogen (N): Non-metal, diatomic gas ($N_2$). Phosphorus (P): Non-metal, exists as $P_4$ (white, red, black). Arsenic (As), Antimony (Sb): Metalloids. Bismuth (Bi): Metal. General Trends: Atomic Radii: Increases down the group. Ionization Enthalpy: Decreases down the group. Generally higher than Group 14 due to stable half-filled p-orbitals. Electronegativity: Decreases down the group. Oxidation States: -3, +3, +5. +3 becomes more stable down the group (inert pair effect, e.g., $Bi^{3+}$ is more stable than $Bi^{5+}$). Nitrogen shows various oxidation states from -3 to +5. Allotropy: P, As, Sb show allotropy. N and Bi do not. Important Compounds: Ammonia ($NH_3$): Basic, Lewis base. Haber process: $N_2 + 3H_2 \rightleftharpoons 2NH_3$. Nitric Acid ($HNO_3$): Strong oxidizing acid. Ostwald process for industrial production. Oxides of Nitrogen: $N_2O$ (nitrous oxide), $NO$ (nitric oxide), $N_2O_3$, $NO_2$, $N_2O_4$, $N_2O_5$. Phosphorus Allotropes: White P (reactive, tetrahedral $P_4$), Red P (polymeric, less reactive), Black P. Phosphine ($PH_3$): Poisonous gas, weaker base than $NH_3$. Phosphorus Halides: $PX_3$ and $PX_5$ (e.g., $PCl_3$, $PCl_5$). $PCl_5$ has trigonal bipyramidal geometry in gas phase. Oxoacids of Phosphorus: e.g., $H_3PO_3$ (phosphorous acid), $H_3PO_4$ (phosphoric acid). Group 16: Oxygen Family (Chalcogens, $ns^2 np^4$) Elements: O, S, Se, Te, Po Oxygen (O): Non-metal, diatomic gas ($O_2$), highly reactive. Sulfur (S): Non-metal, exists as $S_8$ (rhombic, monoclinic). Selenium (Se), Tellurium (Te): Metalloids. Polonium (Po): Metal, radioactive. General Trends: Atomic Radii: Increases down the group. Ionization Enthalpy: Decreases down the group. Generally lower than Group 15 due to less stable $np^4$ configuration. Electronegativity: Decreases down the group. Oxygen is the second most electronegative element. Oxidation States: -2, +2, +4, +6. Oxygen mainly shows -2 (except in $OF_2$ and peroxides). +4 and +6 are common for S, Se, Te. +4 becomes more stable down the group. Allotropy: O (ozone), S (rhombic, monoclinic), Se, Te show allotropy. Important Compounds: Dioxygen ($O_2$): Paramagnetic. Ozone ($O_3$): Allotrope of oxygen, strong oxidizing agent. Structure: bent. Hydrides ($H_2E$): $H_2O$ (liquid), $H_2S$, $H_2Se$, $H_2Te$ (gases). Volatility order: $H_2O Sulfur Dioxide ($SO_2$): Acidic gas, reducing agent. Used in bleaching. Sulfuric Acid ($H_2SO_4$): King of chemicals. Strong oxidizing, dehydrating acid. Contact process for industrial production. Halides: e.g., $SF_6$ (octahedral), $SCl_2$. Group 17: Halogen Family ($ns^2 np^5$) Elements: F, Cl, Br, I, At Fluorine (F): Non-metal, most electronegative element, diatomic gas ($F_2$). Chlorine (Cl): Non-metal, diatomic gas ($Cl_2$). Bromine (Br): Non-metal, diatomic liquid ($Br_2$). Iodine (I): Non-metal, diatomic solid ($I_2$), sublimes. Astatine (At): Metalloid, radioactive. General Trends: Atomic Radii: Increases down the group. Ionization Enthalpy: Decreases down the group. Halogens have very high ionization enthalpies. Electron Gain Enthalpy: Highly negative. Order: $Cl > F > Br > I$. (F is smaller, so electron-electron repulsion is higher). Electronegativity: Decreases down the group. Fluorine is highest. Oxidation States: -1 is most common. Cl, Br, I show +1, +3, +5, +7 (due to d-orbitals). F always shows -1. Bond Dissociation Enthalpy: Order: $Cl_2 > Br_2 > F_2 > I_2$. (Anomalously low for $F_2$ due to small size and lone pair repulsion). Oxidizing Power: Decreases down the group. $F_2 > Cl_2 > Br_2 > I_2$. Important Compounds: Hydrogen Halides ($HX$): $HF, HCl, HBr, HI$. Acidic strength: $HF Oxoacids of Halogens: e.g., Hypochlorous acid ($HClO$), Chlorous acid ($HClO_2$), Chloric acid ($HClO_3$), Perchloric acid ($HClO_4$). Acidic strength increases with oxidation state of halogen. Interhalogen Compounds: $XX'$, $XX'_3$, $XX'_5$, $XX'_7$. Formed between two different halogens. e.g., $ClF_3$ (T-shaped), $IF_7$ (pentagonal bipyramidal). Bleaching Powder ($CaOCl_2$): Contains $Ca^{2+}$, $Cl^-$, and $OCl^-$ ions. Group 18: Noble Gases ($ns^2 np^6$) Elements: He, Ne, Ar, Kr, Xe, Rn All are gases at room temperature. Monoatomic, chemically unreactive (inert). Exist in very small quantities in the atmosphere (except Argon). General Trends: Atomic Radii: Increases down the group. Ionization Enthalpy: Very high, decreases down the group. Electron Gain Enthalpy: Very positive (tendency to accept electrons is negligible). Melting/Boiling Points: Very low, increases down the group due to increasing London dispersion forces. Important Compounds: For a long time, considered non-reactive. Neil Bartlett prepared the first compound, $XePtF_6$. Xenon Compounds: Xe forms compounds with highly electronegative elements like F and O. Fluorides: $XeF_2$ (linear), $XeF_4$ (square planar), $XeF_6$ (distorted octahedral). Oxyfluorides: $XeOF_4$ (square pyramidal), $XeO_2F_2$. Oxides: $XeO_3$ (pyramidal), $XeO_4$ (tetrahedral). Uses: He: Non-flammable, lighter-than-air (balloons, cryogenics, diving), MRI. Ne: Neon signs. Ar: Inert atmosphere for welding, light bulbs. Kr, Xe: Specialized lighting. Rn: Radioactive, used in cancer therapy. Anomalous Behavior of First Element in Each Group The first element of each p-block group (B, C, N, O, F) differs significantly from its congeners due to: Small Size: Leads to high ionization enthalpy and electronegativity. Absence of d-orbitals: Limits coordination number (max 4). Cannot expand octet. Ability to form $p\pi-p\pi$ multiple bonds: Especially true for C, N, O, F (e.g., $C=C, C \equiv C, N \equiv N, O=O$). Heavier elements form $p\pi-d\pi$ bonds, but these are weaker. Examples: Boron forms only covalent compounds, while Al forms ionic. Carbon shows extensive catenation, but Si less so. Nitrogen exists as $N_2$ (triple bond), while P exists as $P_4$ (single bonds). Oxygen is a gas, Sulfur is a solid. Fluorine shows only -1 oxidation state, other halogens show positive oxidation states.