P-Block Elements Essentials

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P-Block Elements Overview Definition: Elements where the last electron enters the outermost p-orbital. Groups: 13 (Boron family) to 18 (Noble gases). General Electronic Configuration: $ns^2np^{1-6}$ (except Helium: $1s^2$). Key Characteristics: Non-metals, metalloids, and metals are all present. Properties vary significantly within a group, especially from the first member. Exhibit a wide range of oxidation states. General Trends in Properties Atomic and Ionic Radii: Generally increase down the group, but anomalies exist due to poor shielding by $d$ and $f$ electrons (e.g., Ga vs Al, Ge vs Si, Sn vs Pb). Ionization Enthalpy: Decreases down the group, but not smoothly. Poor shielding by $d$ and $f$ electrons can lead to increases (e.g., Tl > In). Electronegativity: Generally decreases down the group. Metallic Character: Increases down the group. The heaviest element in each p-block group is typically the most metallic. Oxidation States: Maximum Oxidation State: Equal to the group number (sum of $s$ and $p$ electrons). Inert Pair Effect: The stability of oxidation states two units less than the group oxidation state increases for heavier elements down the group due to the reluctance of $ns^2$ electrons to participate in bonding. Covalence: First member of each group (e.g., B, C, N, O, F) cannot expand its octet due to the absence of d-orbitals. Subsequent members can expand their covalence. Nature of Oxides: Non-metal oxides are generally acidic or neutral. Metalloid oxides are amphoteric. Metal oxides are basic. Acidity of oxides decreases down a group and increases across a period. Group 13 Elements: The Boron Family ($ns^2np^1$) Members: Boron (B), Aluminium (Al), Gallium (Ga), Indium (In), Thallium (Tl). Nature: Boron (non-metal), Aluminium (metal, but forms covalent compounds), Ga, In, Tl (metallic). Oxidation States: +3 is common. +1 becomes more stable down the group (In, Tl) due to inert pair effect. Anomalous Behaviour of Boron: Small size, high ionization enthalpy, high electronegativity. Absence of d-orbitals (max covalence 4). Forms covalent compounds exclusively. Forms $p\pi-p\pi$ multiple bonds in some compounds. Reactivity of Group 13 Elements Towards Air (Oxygen): Boron: Amorphous B forms $B_2O_3$ on heating. Crystalline B unreactive below $700^\circ C$. Aluminium: Forms a thin, protective oxide layer ($Al_2O_3$) that prevents further reaction. All elements form $E_2O_3$ at high temperatures. Towards Dinitrogen: Form nitrides ($EN$) at high temperatures ($2E + N_2 \rightarrow 2EN$). Towards Acids: Boron: Unreactive with non-oxidizing acids. Oxidizing acids (conc. $HNO_3, H_2SO_4$) oxidize it to boric acid ($H_3BO_3$). Aluminium: Amphoteric. Reacts with dilute acids (e.g., HCl) to liberate $H_2$. Passive with conc. $HNO_3$. Ga, In, Tl: React with acids. Ga is amphoteric. Towards Alkalies: Boron: Reacts with fused NaOH to form sodium borate. Aluminium: Amphoteric. $2Al + 2NaOH + 6H_2O \rightarrow 2Na[Al(OH)_4] + 3H_2$. Gallium: Amphoteric. Towards Halogens: Form trihalides $EX_3$ (e.g., $BX_3, AlX_3$). $2E + 3X_2 \rightarrow 2EX_3$. Trihalides are Lewis acids due to electron deficiency (e.g., $BF_3$ forms adducts like $BF_4^-$). $AlCl_3$ exists as a dimer ($Al_2Cl_6$) in vapour phase. Important Compounds of Boron 1. Borax ($Na_2B_4O_7 \cdot 10H_2O$) Common Name: Sodium tetraborate decahydrate. Preparation: From colemanite ($Ca_2B_6O_{11}$): $Ca_2B_6O_{11} + 2Na_2CO_3 \rightarrow 2CaCO_3 + Na_2B_4O_7 + 2NaBO_2$. Sodium metaborate is then converted to borax. Properties: White crystalline solid, dissolves in water to give alkaline solution due to hydrolysis: $Na_2B_4O_7 + 7H_2O \rightleftharpoons 2NaOH + 4H_3BO_3$. Borax Bead Test: On heating, borax loses water, swells, then forms a transparent glassy bead ($NaBO_2 + B_2O_3$). This bead reacts with transition metal oxides to form characteristic coloured metaborates. $Na_2B_4O_7 \cdot 10H_2O \xrightarrow{\Delta} Na_2B_4O_7 \xrightarrow{\Delta} 2NaBO_2 + B_2O_3$. $CuO + B_2O_3 \rightarrow Cu(BO_2)_2$ (blue bead). 2. Orthoboric Acid ($H_3BO_3$ or $B(OH)_3$) Common Name: Boric acid. Preparation: From borax: $Na_2B_4O_7 + 2HCl + 5H_2O \rightarrow 2NaCl + 4H_3BO_3$. From colemanite: $Ca_2B_6O_{11} + SO_2 + H_2O \rightarrow Ca(HSO_3)_2 + H_3BO_3$. Properties: White crystalline solid, slippery to touch. Weak monobasic Lewis acid (accepts $OH^-$ from water): $B(OH)_3 + H_2O \rightleftharpoons [B(OH)_4]^- + H^+$. Reaction on heating: $H_3BO_3 \xrightarrow{370K} HBO_2$ (metaboric acid). $HBO_2 \xrightarrow{>370K} H_2B_4O_7$ (tetraboric acid). $H_2B_4O_7 \xrightarrow{\text{red hot}} B_2O_3$ (boric anhydride). Uses: Mild antiseptic, eye wash. 3. Diborane ($B_2H_6$) Preparation: Lab: $2NaBH_4 + I_2 \rightarrow B_2H_6 + 2NaI + H_2$. Lab: $2BF_3 + 6LiH \rightarrow B_2H_6 + 6LiF$. Industrial: $2BF_3 + 6NaH \xrightarrow{450K} B_2H_6 + 6NaF$. Properties: Colourless, highly toxic gas, pyrophoric (ignites spontaneously). Structure: Two $BH_2$ units linked by two bridging H atoms (banana bonds). Four terminal B-H bonds are 2e-2c bonds. Two bridging B-H-B bonds are 2e-3c bonds. Reactions: Hydrolysis: $B_2H_6 + 6H_2O \rightarrow 2H_3BO_3 + 6H_2$. Combustion: $B_2H_6 + 3O_2 \rightarrow B_2O_3 + 3H_2O$. With Lewis bases (L): Forms adducts $BH_3 \cdot L$ (e.g., $B_2H_6 + 2NMe_3 \rightarrow 2H_3B \cdot NMe_3$). With Ammonia: At low temp: $[BH_2(NH_3)_2]^+[BH_4]^-$. At high temp: Borazine ($B_3N_3H_6$, "inorganic benzene"). Group 14 Elements: The Carbon Family ($ns^2np^2$) Members: Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), Lead (Pb). Nature: C, Si (non-metals), Ge (metalloid), Sn, Pb (metals). Oxidation States: +4 (common). +2 (stability increases down the group due to inert pair effect, especially for Pb). Carbon also shows negative oxidation states. Catenation: Strong tendency to form bonds with itself (C >> Si > Ge $\approx$ Sn > Pb). Allotropy: All members except Pb exhibit allotropy. Reactivity of Group 14 Elements Towards Oxygen: Form monoxides (MO) and dioxides ($MO_2$). $CO_2, SiO_2, GeO_2$: Acidic. $SnO_2, PbO_2$: Amphoteric. $CO$: Neutral. $GeO$: Acidic. $SnO, PbO$: Amphoteric. Towards Water: C, Si, Ge: Unaffected. Sn: Reacts with steam to form $SnO_2$ and $H_2$. Pb: Unaffected (forms protective oxide layer). Towards Halogens: Form tetrahalides ($MX_4$) and dihalides ($MX_2$). $MX_4$ are generally covalent (except $SnF_4, PbF_4$). Tetrahedral geometry. $PbI_4$ does not exist (oxidizing power of $Pb^{4+}$ and reducing power of $I^-$). Stability of $MX_2$ increases down the group ($GeX_2 Hydrolysis: $CCl_4$ is resistant to hydrolysis. Other $MX_4$ are easily hydrolysed (e.g., $SiCl_4 + 4H_2O \rightarrow Si(OH)_4 + 4HCl$). Allotropes of Carbon 1. Diamond Structure: Each C atom $sp^3$ hybridized, covalently bonded to four other C atoms in a tetrahedral arrangement. Rigid 3D network. Properties: Hardest known natural substance, high melting point, electrical insulator, transparent. Uses: Abrasive, cutting tools, jewellery. 2. Graphite Structure: Layers of hexagonal rings. Each C atom $sp^2$ hybridized, bonded to three others. Layers held by weak van der Waals forces. Delocalized $\pi$ electrons within layers. Properties: Soft, greasy (lubricant), good electrical conductor, opaque. Uses: Pencil lead, lubricants, electrodes, in nuclear reactors as moderator. 3. Fullerenes (e.g., $C_{60}$) Common Name: Buckminsterfullerene. Preparation: Heating graphite in an electric arc in presence of inert gases (He, Ar). Structure: Cage-like molecule, 20 six-membered rings and 12 five-membered rings. All C atoms are $sp^2$ hybridized. Properties: Aromatic, soluble in organic solvents. Uses: Superconductors, catalysts, medicinal applications (drug delivery). 4. Carbon Nanotubes Structure: Rolled-up sheets of graphite (cylindrical fullerenes). Single-walled (SWNT) or multi-walled (MWNT). Properties: Extremely strong, excellent electrical and thermal conductors. Uses: Composites, electronics, supercapacitors, nanodevices. Important Compounds of Carbon and Silicon 1. Carbon Monoxide (CO) Preparation: Limited oxygen combustion of carbon: $2C(s) + O_2(g) \xrightarrow{\Delta} 2CO(g)$. Dehydration of formic acid: $HCOOH \xrightarrow{conc. H_2SO_4} CO + H_2O$. Water gas (CO + $H_2$): $C(s) + H_2O(g) \xrightarrow{1000^\circ C} CO(g) + H_2(g)$. Producer gas (CO + $N_2$): $2C(s) + O_2(g) + 4N_2(g) \xrightarrow{1273K} 2CO(g) + 4N_2(g)$. Properties: Colourless, odourless, highly poisonous gas (forms carboxyhemoglobin). Reducing agent. Structure: Polar molecule, triple bond between C and O (:C$\equiv$O:). 2. Carbon Dioxide ($CO_2$) Preparation: Complete combustion: $C(s) + O_2(g) \rightarrow CO_2(g)$. Lab: $CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$. Industrial: Calcination of limestone: $CaCO_3(s) \xrightarrow{\Delta} CaO(s) + CO_2(g)$. Properties: Colourless, odourless gas. Acidic oxide (forms carbonic acid, $H_2CO_3$). Solid $CO_2$ is dry ice. Structure: Linear molecule, $sp$ hybridization on C. Uses: Refrigerant (dry ice), fire extinguishers, carbonated beverages, photosynthesis. 3. Silicon Dioxide ($SiO_2$) Common Name: Silica. Occurrence: Quartz, sand, cristobalite, tridymite. Structure: Covalent 3D network. Each Si is tetrahedrally bonded to four O atoms, and each O is bonded to two Si atoms. Properties: Very stable, high melting point. Acidic oxide. Reactions: Reacts with strong bases and HF. $SiO_2 + 2NaOH \rightarrow Na_2SiO_3 + H_2O$. $SiO_2 + 4HF \rightarrow SiF_4 + 2H_2O$. Uses: Piezoelectric material, manufacturing glass, cement, ceramics. 4. Silicones Definition: Organosilicon polymers containing $R_2SiO$ repeating units. Preparation: Hydrolysis of alkyl or aryl substituted chlorosilanes (e.g., $R_2SiCl_2$) followed by polymerization. $R_2SiCl_2 + 2H_2O \rightarrow R_2Si(OH)_2 + 2HCl$. $n R_2Si(OH)_2 \rightarrow (-R_2SiO-)_n + nH_2O$. Properties: Water repellent, chemically inert, high thermal stability. Uses: Sealants, lubricants, electrical insulators, water-proofing fabrics, surgical implants. 5. Silicates Definition: Compounds containing silicon and oxygen, often with other metals. Basic unit is $SiO_4^{4-}$ tetrahedron. Types: Orthosilicates, pyrosilicates, cyclic silicates, chain silicates, sheet silicates, 3D silicates. Examples: Zeolites, feldspar, mica, asbestos, talc. Uses: Building materials (cement, glass), ceramics. Group 15 Elements: The Nitrogen Family ($ns^2np^3$) Members: Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi). Nature: N, P (non-metals), As, Sb (metalloids), Bi (metal). Oxidation States: -3, +3, +5. Nitrogen: Many states (-3 to +5). Phosphorus: -3, +3, +5. As, Sb: +3, +5 (stability of +3 increases down the group). Bi: +3 is more stable than +5. Anomalous Behaviour of Nitrogen: Small size, high electronegativity, high ionization enthalpy. Absence of d-orbitals (max covalence 4). Forms $p\pi-p\pi$ multiple bonds (N$\equiv$N, N=O). Does not form pentahalides. Reactivity of Group 15 Elements Towards Hydrogen: Form hydrides $EH_3$ (e.g., $NH_3, PH_3$). Stability decreases from $NH_3$ to $BiH_3$. Reducing character increases down the group. Basic character decreases down the group. Towards Oxygen: Form oxides $E_2O_3, E_2O_5$. Acidic character decreases down the group. $N_2O_3, P_2O_3$ (acidic), $As_2O_3, Sb_2O_3$ (amphoteric), $Bi_2O_3$ (basic). Towards Halogens: Form trihalides ($EX_3$) and pentahalides ($EX_5$). All form $EX_3$ (except $NI_3$). $NCl_3$ is explosive. All elements except N form $EX_5$ (N lacks d-orbitals). $PX_5$ are more covalent than $PX_3$. Towards Metals: React with metals to form binary compounds (e.g., $Ca_3N_2, Ca_3P_2$). Important Compounds of Nitrogen and Phosphorus 1. Dinitrogen ($N_2$) Preparation: Lab: $NH_4Cl(aq) + NaNO_2(aq) \xrightarrow{\Delta} N_2(g) + 2H_2O(l) + NaCl(aq)$. Thermal decomposition of ammonium dichromate: $(NH_4)_2Cr_2O_7 \xrightarrow{\Delta} N_2 + Cr_2O_3 + 4H_2O$. Pure $N_2$: Thermal decomposition of sodium or barium azide: $Ba(N_3)_2 \xrightarrow{\Delta} Ba + 3N_2$. Industrial: Fractional distillation of liquid air. Properties: Colourless, odourless, tasteless, non-toxic gas. Very unreactive due to high bond enthalpy of N$\equiv$N. Reactions: With metals: Forms nitrides (e.g., $6Li + N_2 \rightarrow 2Li_3N$). With $H_2$ (Haber process): $N_2 + 3H_2 \rightleftharpoons 2NH_3$. With $O_2$: $N_2 + O_2 \xrightarrow{2000K} 2NO$. 2. Ammonia ($NH_3$) Preparation: Lab: Heating ammonium salts with bases: $2NH_4Cl + Ca(OH)_2 \xrightarrow{\Delta} CaCl_2 + 2NH_3 + 2H_2O$. Haber's process (industrial): $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$ ($\Delta H = -46.1 \text{ kJ mol}^{-1}$). Conditions: High pressure ($200 \text{ atm}$), optimal temperature ($700 \text{ K}$), catalyst ($Fe_2O_3$ with $K_2O, Al_2O_3$). Properties: Colourless gas, pungent smell. Basic nature (lone pair on N), forms ammonium salts. $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$. Precipitates metal hydroxides: $FeCl_3 + 3NH_4OH \rightarrow Fe(OH)_3 \downarrow + 3NH_4Cl$. Forms complex compounds with transition metal ions: $Cu^{2+} + 4NH_3 \rightarrow [Cu(NH_3)_4]^{2+}$ (deep blue). Structure: Pyramidal shape, $sp^3$ hybridization with one lone pair. 3. Nitric Acid ($HNO_3$) Preparation: Lab: $NaNO_3 + H_2SO_4 \xrightarrow{\Delta} NaHSO_4 + HNO_3$. Ostwald's process (industrial): Catalytic oxidation of $NH_3$: $4NH_3 + 5O_2 \xrightarrow{Pt/Rh \text{ gauze}, 500K, 9 \text{ bar}} 4NO + 6H_2O$. Oxidation of NO: $2NO + O_2 \rightarrow 2NO_2$. Absorption of $NO_2$ in water: $3NO_2 + H_2O \rightarrow 2HNO_3 + NO$. (NO recycled). Properties: Strong oxidizing agent, strong acid. Reaction with metals: Conc. $HNO_3$: $Cu + 4HNO_3(conc.) \rightarrow Cu(NO_3)_2 + 2NO_2 + 2H_2O$. Dilute $HNO_3$: $3Cu + 8HNO_3(dilute) \rightarrow 3Cu(NO_3)_2 + 2NO + 4H_2O$. Very dilute $HNO_3$ with active metals (Mg, Mn): Forms $N_2O$. Does not react with noble metals (Au, Pt) but forms aqua regia (1:3 conc. $HNO_3$:conc. $HCl$). Non-metals: Oxidizes C, S, I to their respective oxyacids. Brown Ring Test: For nitrates. $NO_3^- + 3Fe^{2+} + 4H^+ \rightarrow NO + 3Fe^{3+} + 2H_2O$. $Fe^{2+} + NO \rightarrow [Fe(H_2O)_5(NO)]^{2+}$ (brown ring complex). 4. Oxides of Nitrogen Oxide Common Name Oxidation State Nature $N_2O$ Nitrous oxide (Laughing gas) +1 Neutral $NO$ Nitric oxide +2 Neutral $N_2O_3$ Dinitrogen trioxide +3 Acidic $NO_2$ Nitrogen dioxide +4 Acidic $N_2O_4$ Dinitrogen tetroxide +4 Acidic $N_2O_5$ Dinitrogen pentoxide +5 Acidic 5. Allotropes of Phosphorus White Phosphorus ($P_4$): Preparation: Heating phosphate rock ($Ca_3(PO_4)_2$) with coke and sand in an electric furnace. Properties: Translucent white waxy solid, poisonous, glows in dark (chemiluminescence), highly reactive, ignites in air. Insoluble in water, soluble in $CS_2$. Structure: Tetrahedral $P_4$ unit, angle $60^\circ$ (high strain). Red Phosphorus: Preparation: Heating white P at $573K$ in an inert atmosphere. Properties: Polymeric, less reactive, non-poisonous, insoluble in $CS_2$. Black Phosphorus: Preparation: Heating white P at high pressure. Properties: Two forms ($\alpha$ and $\beta$), most stable (metallic lustre). 6. Phosphine ($PH_3$) Preparation: Lab: $Ca_3P_2 + 6HCl \rightarrow 3CaCl_2 + 2PH_3$. (Impure, spontaneously combustible). From white phosphorus: $P_4 + 3NaOH + 3H_2O \rightarrow PH_3 + 3NaH_2PO_2$. Properties: Colourless, poisonous gas, rotten fish smell. Weakly basic. Uses: 'Holme's signals' (containers with $CaC_2$ and $Ca_3P_2$ for generating $PH_3$ and $C_2H_2$ on contact with water). 7. Phosphorus Halides ($PX_3, PX_5$) Phosphorus Trichloride ($PCl_3$): Preparation: $P_4 + 6Cl_2 \rightarrow 4PCl_3$. Hydrolysis: $PCl_3 + 3H_2O \rightarrow H_3PO_3 + 3HCl$. Phosphorus Pentachloride ($PCl_5$): Preparation: $P_4 + 10Cl_2 \rightarrow 4PCl_5$. Hydrolysis: $PCl_5 + H_2O \rightarrow POCl_3 + 2HCl$ (partial). $POCl_3 + 3H_2O \rightarrow H_3PO_4 + 3HCl$ (complete). Reactions: $C_2H_5OH + PCl_5 \rightarrow C_2H_5Cl + POCl_3 + HCl$. Structure: Trigonal bipyramidal in gas/liquid state. Ionic $[PCl_4]^+\text{ (tetrahedral)} [PCl_6]^-\text{ (octahedral)}$ in solid state. 8. Oxoacids of Phosphorus Hypophosphorous Acid ($H_3PO_2$): +1 oxidation state. Monobasic. Strong reducing agent. Phosphorous Acid ($H_3PO_3$): +3 oxidation state. Dibasic. Reducing agent. Orthophosphoric Acid ($H_3PO_4$): +5 oxidation state. Tribasic. Pyrophosphoric Acid ($H_4P_2O_7$): +5 oxidation state. Tetrabasic. Key: P-H bonds are not ionizable. P-OH bonds are ionizable. Group 16 Elements: The Oxygen Family (Chalcogens) ($ns^2np^4$) Members: Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), Polonium (Po). Nature: O, S (non-metals), Se, Te (metalloids), Po (metal, radioactive). Oxidation States: -2, +2, +4, +6. Oxygen: -2 (common), -1 (peroxides), -1/2 (superoxides), +2 ($OF_2$). Sulfur, Se, Te: -2 (less stable down group), +2, +4, +6. Stability of +4 increases down group, +6 decreases (inert pair effect). Anomalous Behaviour of Oxygen: Small size, high electronegativity, absence of d-orbitals (max covalence 2). Forms $p\pi-p\pi$ multiple bonds ($O_2$). Hydrogen bonding. Reactivity of Group 16 Elements Towards Hydrogen: Form hydrides $H_2E$ (e.g., $H_2O, H_2S$). Acidity increases down the group ($H_2O Thermal stability decreases down the group. Reducing character increases down the group. Towards Oxygen: All form oxides. $SO_2, SO_3$ are common. Towards Halogens: Form halides $EX_2, EX_4, EX_6$. Oxygen forms $OF_2, O_2F_2$. Sulfur forms $SF_6$ (stable), $SF_4, S_2F_2$. Stability of hexahalides decreases down the group. Dihalides are generally unstable. Important Compounds of Oxygen and Sulfur 1. Dioxygen ($O_2$) Preparation: Lab: Decomposition of oxygen-containing salts: $2KClO_3 \xrightarrow{MnO_2, \Delta} 2KCl + 3O_2$. From hydrogen peroxide: $2H_2O_2 \xrightarrow{\text{finely divided metals/MnO_2}} 2H_2O + O_2$. Industrial: Electrolysis of water, fractional distillation of liquid air. Properties: Colourless, odourless gas. Paramagnetic (due to two unpaired electrons in antibonding $\pi$ orbitals). Reactions: Supports combustion. Reacts with metals and non-metals to form oxides. 2. Ozone ($O_3$) Preparation: Silent electric discharge through dry oxygen: $3O_2 \rightleftharpoons 2O_3$ ($\Delta H = +142 \text{ kJ mol}^{-1}$). Properties: Pale blue gas, pungent smell. Powerful oxidizing agent. Oxidizes $KI$ to $I_2$: $2KI + H_2O + O_3 \rightarrow 2KOH + I_2 + O_2$. Oxidizes PbS to $PbSO_4$: $PbS + 4O_3 \rightarrow PbSO_4 + 4O_2$. Structure: Angular molecule, $sp^2$ hybridization. Resonance structures. Uses: Germicide, disinfectant, bleaching oils, water purification. 3. Sulfur Dioxide ($SO_2$) Preparation: Burning sulfur: $S + O_2 \rightarrow SO_2$. Roasting sulfide ores: $4FeS_2 + 11O_2 \rightarrow 2Fe_2O_3 + 8SO_2$. Lab: $Na_2SO_3 + H_2SO_4 \rightarrow Na_2SO_4 + H_2O + SO_2$. Properties: Colourless gas, pungent smell. Acidic oxide. Reducing agent (e.g., reduces $Fe^{3+}$ to $Fe^{2+}$). Bleaching agent (temporary, due to reduction). Structure: Bent shape, $sp^2$ hybridization. Uses: Refining petroleum, bleaching wool/silk, preservative, raw material for $H_2SO_4$. 4. Sulfuric Acid ($H_2SO_4$) Common Name: King of Chemicals. Preparation: Contact Process (industrial): Burning sulfur or sulfide ores: $S + O_2 \rightarrow SO_2$. Catalytic oxidation of $SO_2$: $2SO_2 + O_2 \xrightarrow{V_2O_5, 720K} 2SO_3$. Absorption of $SO_3$ in $H_2SO_4$: $SO_3 + H_2SO_4 \rightarrow H_2S_2O_7$ (oleum). Dilution of oleum: $H_2S_2O_7 + H_2O \rightarrow 2H_2SO_4$. Properties: Dense, oily, corrosive liquid. Strong acid, oxidizing agent, dehydrating agent. As acid: $H_2SO_4 + H_2O \rightarrow H_3O^+ + HSO_4^-$. As oxidizing agent: $C + 2H_2SO_4 \rightarrow CO_2 + 2SO_2 + 2H_2O$. As dehydrating agent: $C_{12}H_{22}O_{11} \xrightarrow{conc. H_2SO_4} 12C + 11H_2O$. Uses: Fertilizers, detergents, pigments, petroleum refining, metallurgy. 5. Oxoacids of Sulfur Sulfurous Acid ($H_2SO_3$): +4 oxidation state. Sulfuric Acid ($H_2SO_4$): +6 oxidation state. Peroxomonosulfuric Acid ($H_2SO_5$): Caro's acid. Peroxodisulfuric Acid ($H_2S_2O_8$): Marshall's acid. Pyrosulfuric Acid ($H_2S_2O_7$): Oleum. 6. Allotropes of Sulfur Rhombic Sulfur ($\alpha$-sulfur): Yellow, most stable at room temperature. $S_8$ rings. Monoclinic Sulfur ($\beta$-sulfur): Colourless needles. Stable above $369K$. $S_8$ rings. Plastic Sulfur: Formed by pouring molten sulfur into cold water. Non-crystalline. Group 17 Elements: The Halogens ($ns^2np^5$) Members: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At). Nature: All are non-metals. Oxidation States: -1 (common). Cl, Br, I also show +1, +3, +5, +7. Anomalous Behaviour of Fluorine: Smallest size, highest electronegativity, highest ionization enthalpy. Absence of d-orbitals (max covalence 1). Forms only one oxoacid ($HFO$). Forms strongest hydrogen bonds. Reactivity of Group 17 Elements Towards Hydrogen: Form hydrogen halides $HX$ (e.g., $HF, HCl$). Acidity increases down the group ($HF Thermal stability decreases down the group. Reducing character increases down the group. Towards Oxygen: Form various oxides (e.g., $Cl_2O, ClO_2, Cl_2O_7$). Fluorine forms $OF_2, O_2F_2$. Towards Metals: React to form metal halides (e.g., $2Na + Cl_2 \rightarrow 2NaCl$). Towards Other Halogens (Interhalogen Compounds): Form $XX', XX'_3, XX'_5, XX'_7$. $X$ is larger, $X'$ is smaller. (e.g., $ClF, BrF_3, IF_5, IF_7$). All are covalent. More reactive than halogens (except $F_2$). Important Compounds of Halogens 1. Chlorine ($Cl_2$) Preparation: Lab: $MnO_2 + 4HCl \rightarrow MnCl_2 + Cl_2 + 2H_2O$. Deacon's process (industrial): $4HCl + O_2 \xrightarrow{CuCl_2 \text{ catalyst}} 2Cl_2 + 2H_2O$. Electrolytic process: Electrolysis of brine (aqueous NaCl): $2NaCl(aq) + 2H_2O(l) \xrightarrow{\text{electrolysis}} 2NaOH(aq) + Cl_2(g) + H_2(g)$. Properties: Greenish-yellow gas, pungent smell. Oxidizing agent, bleaching agent (permanent, due to oxidation). Reaction with $H_2S$: $H_2S + Cl_2 \rightarrow 2HCl + S$. Reaction with $NH_3$: Excess $NH_3$: $8NH_3 + 3Cl_2 \rightarrow 6NH_4Cl + N_2$. Excess $Cl_2$: $NH_3 + 3Cl_2 \rightarrow NCl_3 + 3HCl$. Reaction with NaOH: Cold, dilute: $2NaOH + Cl_2 \rightarrow NaCl + NaOCl + H_2O$. Hot, concentrated: $6NaOH + 3Cl_2 \rightarrow 5NaCl + NaClO_3 + 3H_2O$. Uses: Bleaching cotton/wood pulp, water purification, manufacturing PVC, $DDT$, chloroform. 2. Hydrogen Chloride ($HCl$) Preparation: Lab: $NaCl + H_2SO_4 \xrightarrow{420K} NaHSO_4 + HCl$. Industrial: By-product of chlor-alkali process, direct combination of $H_2$ and $Cl_2$. Properties: Colourless, pungent gas. Highly soluble in water (forms hydrochloric acid). Strong acid. Reacts with active metals, metal oxides, hydroxides, carbonates. Aqua Regia: 3 parts conc. $HCl$ + 1 part conc. $HNO_3$. Dissolves noble metals ($Au, Pt$). 3. Oxoacids of Halogens Hypohalous acids ($HOX$): $HClO, HBrO, HIO$. +1 oxidation state. Halous acids ($HOXO$): $HClO_2$. +3 oxidation state. Halic acids ($HOXO_2$): $HClO_3, HBrO_3, HIO_3$. +5 oxidation state. Perhalic acids ($HOXO_3$): $HClO_4, HBrO_4, HIO_4$. +7 oxidation state. Acid strength increases with increasing oxidation state of halogen (e.g., $HClO Group 18 Elements: The Noble Gases ($ns^2np^6$) Members: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn). Nature: All are monatomic gases. Electronic Configuration: $ns^2np^6$ (except He: $1s^2$). Fully filled valence shells, hence very stable. Ionization Enthalpy: Very high. Electron Gain Enthalpy: Positive (no tendency to accept electrons). Atomic Radii: Largest in their respective periods (van der Waals radii). Melting/Boiling Points: Very low due to weak dispersion forces. Chemical Reactivity: Generally unreactive. First compound of noble gas: $XePtF_6$ by Bartlett (1962). Xenon forms compounds with F and O (e.g., $XeF_2, XeF_4, XeF_6, XeO_3, XeOF_4$). Krypton forms $KrF_2$. Radon forms $RnF_2$. Important Compounds of Xenon 1. Xenon Fluorides ($XeF_2, XeF_4, XeF_6$) Preparation: Direct reaction of Xe and $F_2$ under specific conditions. $Xe + F_2 \xrightarrow{673K, 1 \text{ bar}} XeF_2$. $Xe + 2F_2 \xrightarrow{873K, 7 \text{ bar}} XeF_4$. $Xe + 3F_2 \xrightarrow{573K, 60-70 \text{ bar}} XeF_6$. Properties: White crystalline solids, powerful fluorinating agents. Hydrolysis: $2XeF_2 + 2H_2O \rightarrow 2Xe + 4HF + O_2$. $XeF_4 + 12H_2O \rightarrow Xe + 2XeO_3 + 24HF + 3O_2$. $XeF_6 + 3H_2O \rightarrow XeO_3 + 6HF$. Structures: $XeF_2$: Linear ($sp^3d$, 3 lone pairs equatorial). $XeF_4$: Square planar ($sp^3d^2$, 2 lone pairs axial). $XeF_6$: Distorted octahedral ($sp^3d^3$, 1 lone pair). 2. Xenon Oxides and Oxyfluorides ($XeO_3, XeOF_4, XeO_2F_2$) $XeO_3$: Prepared by hydrolysis of $XeF_4$ or $XeF_6$. Explosive solid. Pyramidal structure. $XeOF_4$: Prepared by partial hydrolysis of $XeF_6$. Square pyramidal structure.