Chemistry Essentials

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1. Periodic Properties Definitions & Trends Atomic Size: Distance from nucleus to outermost electron shell. Period: Decreases (L to R) due to increased nuclear charge pulling electrons closer. Group: Increases (Top to Bottom) due to addition of new electron shells. Metallic Character: Tendency to lose electrons. Period: Decreases (L to R). Group: Increases (Top to Bottom). Non-metallic Character: Tendency to gain electrons. Period: Increases (L to R). Group: Decreases (Top to Bottom). Ionisation Potential (IE): Energy required to remove an electron from an isolated gaseous atom. Period: Increases (L to R) due to increased nuclear charge and smaller size. Group: Decreases (Top to Bottom) due to increased size and shielding effect. Electron Affinity (EA): Energy released when an electron is added to an isolated gaseous atom. Period: Increases (L to R) due to increased nuclear charge and smaller size. Group: Decreases (Top to Bottom). Electronegativity (EN): Tendency of an atom to attract shared electron pairs in a covalent bond. Period: Increases (L to R). Group: Decreases (Top to Bottom). 2. Chemical Bonding Electrovalent (Ionic) Bonding Formed by complete transfer of electrons between a metal and a non-metal. Electron Dot Structures: NaCl: $Na \cdot + \cdot \ddot{Cl}: \rightarrow Na^+ [: \ddot{Cl}: ]^-$ MgCl$_2$: $Mg \cdot + 2 \cdot \ddot{Cl}: \rightarrow Mg^{2+} 2[: \ddot{Cl}: ]^-$ CaO: $Ca \cdot + \cdot \ddot{O}: \rightarrow Ca^{2+} [: \ddot{O}: ]^{2-}$ Characteristics: Solid state, high melting/boiling points (strong electrostatic forces). Conduct electricity in molten state or aqueous solution (free ions). Soluble in polar solvents (e.g., water). Covalent Bonding Formed by sharing of electrons between non-metal atoms. Electron Dot Structures (Duplet/Octet): H$_2$: $H \cdot + \cdot H \rightarrow H:H$ Cl$_2$: $: \ddot{Cl} \cdot + \cdot \ddot{Cl}: \rightarrow : \ddot{Cl}: \ddot{Cl}: $ N$_2$: $:N \equiv N:$ NH$_3$: $H-\underset{\overset{H}{|}}{\ddot{N}}-H$ CH$_4$: $\underset{\overset{H}{|}}{H}-\underset{\overset{H}{|}}{C}-H$ CCl$_4$: $\underset{\overset{Cl}{|}}{Cl}-\underset{\overset{Cl}{|}}{C}-Cl$ Polar Covalent Compounds: Unequal sharing of electrons due to electronegativity difference. HCl: $H^{\delta+}-Cl^{\delta-}$ H$_2$O: $H^{\delta+}-\ddot{O}^{\delta-}-H^{\delta+}$ Characteristics: Usually gases, liquids, or soft solids. Lower melting/boiling points. Generally poor conductors of electricity (no free ions). Soluble in non-polar solvents. Coordinate (Dative) Bonding A type of covalent bond where both shared electrons are donated by one atom. Lone Pair: A pair of valence electrons not involved in bonding. Formation of H$_3$O$^+$: $H_2\ddot{O}: + H^+ \rightarrow [H-\ddot{O}-H]^{\text{+}}$ Formation of NH$_4^+$: $H_3\ddot{N}: + H^+ \rightarrow [H-\underset{\overset{H}{|}}{N}-H]^{\text{+}}$ 3. Acids, Bases and Salts Definitions & Properties Acids: Produce $H^+$ (or $H_3O^+$) ions in water. Turn blue litmus red. E.g., $HCl \rightarrow H^+ + Cl^-$, $H_2SO_4 \rightarrow 2H^+ + SO_4^{2-}$ Bases/Alkalis: Produce $OH^-$ ions in water. Turn red litmus blue. E.g., $NaOH \rightarrow Na^+ + OH^-$, $Ca(OH)_2 \rightarrow Ca^{2+} + 2OH^-$ Salts: Formed by replacement of $H^+$ of an acid by a metal or ammonium ion. pH Scale: Measures acidity/alkalinity. $pH $pH = 7$: Neutral $pH > 7$: Basic/Alkaline Types of Salts Normal Salt: Complete replacement of $H^+$ ions (e.g., $NaCl, K_2SO_4$). Acid Salt: Partial replacement of $H^+$ ions (e.g., $NaHCO_3, NaHSO_4$). Basic Salt: Partial replacement of $OH^-$ ions (e.g., $Pb(OH)Cl$). Action of Dilute Acids on Salts Carbonates ($CO_3^{2-}$): Acid + Carbonate $\rightarrow$ Salt + Water + $CO_2$ $CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$ Hydrogen Carbonates ($HCO_3^-$): Acid + Hydrogen Carbonate $\rightarrow$ Salt + Water + $CO_2$ $NaHCO_3(s) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l) + CO_2(g)$ Sulphites ($SO_3^{2-}$): Acid + Sulphite $\rightarrow$ Salt + Water + $SO_2$ $Na_2SO_3(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2O(l) + SO_2(g)$ Sulphides ($S^{2-}$): Acid + Sulphide $\rightarrow$ Salt + $H_2S$ $FeS(s) + 2HCl(aq) \rightarrow FeCl_2(aq) + H_2S(g)$ Preparation of Normal Salts Direct Combination: $2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$ Displacement: $Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)$ Precipitation (Double Decomposition): $AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$ Neutralisation of Insoluble Base: $CuO(s) + 2HCl(aq) \rightarrow CuCl_2(aq) + H_2O(l)$ Neutralisation of Alkali (Titration): $NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)$ Action of Dilute Acids on Carbonates/Bicarbonates: (See above) 4. Analytical Chemistry Action of Ammonium Hydroxide and Sodium Hydroxide on Salt Solutions Salt Solution Color of Solution NaOH (drop by drop) NaOH (in excess) NH$_4$OH (drop by drop) NH$_4$OH (in excess) Ca$^{2+}$ (e.g., $CaCl_2$) Colorless White ppt. ($Ca(OH)_2$) Insoluble No ppt. No ppt. Fe$^{2+}$ (e.g., $FeSO_4$) Pale Green Dirty Green ppt. ($Fe(OH)_2$) Insoluble Dirty Green ppt. ($Fe(OH)_2$) Insoluble Fe$^{3+}$ (e.g., $FeCl_3$) Yellowish-Brown Reddish-Brown ppt. ($Fe(OH)_3$) Insoluble Reddish-Brown ppt. ($Fe(OH)_3$) Insoluble Cu$^{2+}$ (e.g., $CuSO_4$) Blue Pale Blue ppt. ($Cu(OH)_2$) Insoluble Pale Blue ppt. ($Cu(OH)_2$) Deep Blue Solution ($[Cu(NH_3)_4](OH)_2$) Zn$^{2+}$ (e.g., $ZnSO_4$) Colorless White gelatinous ppt. ($Zn(OH)_2$) Soluble (forms $Na_2ZnO_2$) White gelatinous ppt. ($Zn(OH)_2$) Soluble (forms $[Zn(NH_3)_4](OH)_2$) Pb$^{2+}$ (e.g., $Pb(NO_3)_2$) Colorless Chalky White ppt. ($Pb(OH)_2$) Soluble (forms $Na_2PbO_2$) Chalky White ppt. ($Pb(OH)_2$) Insoluble Special Action: $CuSO_4(aq) + 4NH_4OH(aq) \rightarrow [Cu(NH_3)_4]SO_4(aq) + 4H_2O(l)$ (Deep blue solution) $NH_4Cl(aq) + NaOH(aq) \xrightarrow{\Delta} NaCl(aq) + H_2O(l) + NH_3(g)$ (Ammonia gas evolved) Amphoteric Nature Metals/Oxides/Hydroxides that react with both acids and strong bases. Examples: Al, Zn, Pb, $Al_2O_3, ZnO, PbO, Al(OH)_3, Zn(OH)_2, Pb(OH)_2$ $ZnO(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2O(l)$ $ZnO(s) + 2NaOH(aq) \rightarrow Na_2ZnO_2(aq) + H_2O(l)$ (Sodium zincate) 5. Mole Concept and Stoichiometry Gay-Lussac’s Law of Combining Volumes: When gases react, they do so in volumes that bear a simple whole number ratio to one another and to the volumes of the gaseous products, provided all volumes are measured at the same temperature and pressure. Avogadro’s Law: Equal volumes of all gases, at the same temperature and pressure, contain an equal number of molecules. Mole: A unit representing $6.022 \times 10^{23}$ particles (Avogadro's Number, $N_A$). Molar Volume: 1 mole of any gas at S.T.P. (0$^\circ$C, 1 atm) occupies 22.4 L. Molecular Mass = 2 $\times$ Vapour Density (VD) Calculations: Moles = Mass / Molar Mass Moles = Volume (at STP) / 22.4 L Moles = Number of Particles / $N_A$ Empirical Formula: Simplest whole number ratio of atoms in a compound. Molecular Formula: Actual number of atoms of each element in a molecule. Molecular Formula = n $\times$ Empirical Formula, where $n = \frac{\text{Molecular Mass}}{\text{Empirical Formula Mass}}$ 6. Electrolysis Electrolytes & Non-Electrolytes Electrolytes: Substances that conduct electricity in molten state or aqueous solution due to presence of free ions. (e.g., acids, bases, salts). Strong Electrolytes: Almost completely dissociate (e.g., NaCl, HCl, NaOH). Weak Electrolytes: Partially dissociate (e.g., $CH_3COOH, NH_4OH$). Non-Electrolytes: Do not conduct electricity (e.g., glucose, urea, alcohol). Contain only molecules. Terms Electrolysis: Decomposition of a compound by passage of electric current. Electrolyte: Substance undergoing electrolysis. Electrodes: Conductors through which current enters/leaves the electrolyte. Anode: Positive electrode (oxidation occurs, anions move towards it). Cathode: Negative electrode (reduction occurs, cations move towards it). Anion: Negatively charged ion. Cation: Positively charged ion. Oxidation: Loss of electrons (at anode). Reduction: Gain of electrons (at cathode). Selective Discharge: Ions with lower discharge potential are preferentially discharged. (Activity Series helps predict). Examples of Electrolysis Molten Lead Bromide ($PbBr_2$): Cathode (Reduction): $Pb^{2+} + 2e^- \rightarrow Pb(l)$ Anode (Oxidation): $2Br^- \rightarrow Br_2(g) + 2e^-$ Acidified Water ($H_2O$ with $H_2SO_4$): Cathode: $2H^+ + 2e^- \rightarrow H_2(g)$ Anode: $4OH^- \rightarrow O_2(g) + 2H_2O(l) + 4e^-$ (or $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$) Aqueous Copper(II) Sulphate ($CuSO_4$) with Copper Electrodes: Cathode: $Cu^{2+} + 2e^- \rightarrow Cu(s)$ (Copper deposits) Anode: $Cu(s) \rightarrow Cu^{2+} + 2e^-$ (Anode dissolves) Applications of Electrolysis Electroplating: Coating a metal with a thin layer of another metal. Object to be plated: Cathode. Plating metal: Anode. Electrolyte: Salt solution of the plating metal. Example (Silver plating): Anode: Ag, Cathode: Object, Electrolyte: $K[Ag(CN)_2]$ Cathode: $Ag^+ + e^- \rightarrow Ag(s)$ Anode: $Ag(s) \rightarrow Ag^+ + e^-$ Electrorefining of Copper: Impure Copper: Anode. Pure Copper: Cathode. Electrolyte: $CuSO_4$ solution with dilute $H_2SO_4$. Cathode: $Cu^{2+} + 2e^- \rightarrow Cu(s)$ Anode: $Cu(s) \rightarrow Cu^{2+} + 2e^-$ (more reactive impurities oxidize, less reactive fall as anode mud) 7. Metallurgy Occurrence of Metals Mineral: Naturally occurring compound of a metal. Ore: Mineral from which metal can be profitably extracted. Common Ores: Iron: Hematite ($Fe_2O_3$), Magnetite ($Fe_3O_4$) Aluminium: Bauxite ($Al_2O_3 \cdot 2H_2O$) Zinc: Zinc Blende ($ZnS$), Calamine ($ZnCO_3$) Stages of Metal Extraction Dressing of Ore (Concentration): Removing impurities (gangue). Hydrolytic Method: For heavier ores (e.g., tin stone). Magnetic Separation: For magnetic ores (e.g., magnetite). Froth Flotation: For sulphide ores (e.g., zinc blende). Conversion of Concentrated Ore to Oxide: Roasting: Heating sulphide ore in air. $2ZnS(s) + 3O_2(g) \xrightarrow{\Delta} 2ZnO(s) + 2SO_2(g)$ Calcination: Heating carbonate/hydroxide ore in absence of air. $ZnCO_3(s) \xrightarrow{\Delta} ZnO(s) + CO_2(g)$ Reduction of Metallic Oxides: By Carbon/CO/H$_2$: For less reactive metals (e.g., Fe, Pb, Cu, Zn). $Fe_2O_3(s) + 3CO(g) \xrightarrow{\Delta} 2Fe(s) + 3CO_2(g)$ $CuO(s) + H_2(g) \xrightarrow{\Delta} Cu(s) + H_2O(l)$ By Electrolysis: For highly reactive metals (e.g., Na, K, Ca, Al). Refining: Further purification of crude metal. Extraction of Aluminium (Hall-Héroult Process) Purification of Bauxite (Bayer's Process): Removes $Fe_2O_3, SiO_2$. $Al_2O_3 \cdot 2H_2O + 2NaOH \rightarrow 2Na[Al(OH)_4]$ $Na[Al(OH)_4] + 2H_2O \rightarrow Al(OH)_3 \downarrow + NaOH$ $2Al(OH)_3 \xrightarrow{1000^\circ C} Al_2O_3 + 3H_2O$ Electrolytic Reduction: Electrolyte: Molten alumina ($Al_2O_3$) dissolved in cryolite ($Na_3AlF_6$) and fluorspar ($CaF_2$). Cathode: Carbon lining of the steel tank. Anode: Graphite rods. Reactions: Cathode: $Al^{3+} + 3e^- \rightarrow Al(l)$ Anode: $2O^{2-} \rightarrow O_2(g) + 4e^-$; $C(s) + O_2(g) \rightarrow CO_2(g)$ (anodes burn away) Alloys (Composition & Uses) Stainless Steel: Fe, Cr, Ni (Corrosion resistance, utensils) Duralumin: Al, Cu, Mg, Mn (Aircraft parts, pressure cookers) Brass: Cu, Zn (Utensils, decorative articles) Bronze: Cu, Sn (Statues, coins) Fuse Metal/Solder: Pb, Sn (Joining electrical wires) 8. Study of Compounds A. Hydrogen Chloride (HCl) Preparation: $NaCl(s) + H_2SO_4(conc.) \xrightarrow{ Properties: Colorless gas, pungent smell. Denser than air. Highly soluble in water (fountain experiment). Forms hydrochloric acid. Reacts with ammonia: $HCl(g) + NH_3(g) \rightarrow NH_4Cl(s)$ (dense white fumes). Acidic Properties (in solution): Reacts with active metals: $Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$ Reacts with metal oxides/hydroxides: $CuO(s) + 2HCl(aq) \rightarrow CuCl_2(aq) + H_2O(l)$ Reacts with carbonates/bicarbonates: $Na_2CO_3(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2O(l) + CO_2(g)$ Precipitation Reactions: With $AgNO_3$: $AgNO_3(aq) + HCl(aq) \rightarrow AgCl(s) \downarrow + HNO_3(aq)$ (White ppt., soluble in $NH_4OH$) With $Pb(NO_3)_2$: $Pb(NO_3)_2(aq) + 2HCl(aq) \rightarrow PbCl_2(s) \downarrow + 2HNO_3(aq)$ (White ppt., soluble in hot water) B. Ammonia (NH$_3$) Lab Preparation: $2NH_4Cl(s) + Ca(OH)_2(s) \xrightarrow{\Delta} CaCl_2(s) + 2H_2O(l) + 2NH_3(g)$ (Collected by downward displacement of air). From Nitrides: $Mg_3N_2(s) + 6H_2O(l) \xrightarrow{\Delta} 3Mg(OH)_2(s) + 2NH_3(g)$ Haber's Process (Manufacture): $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$ (Conditions: Fe catalyst, Mo promoter, 450-500$^\circ$C, 200-900 atm). Properties: Colorless gas, pungent smell. Lighter than air. Highly soluble in water (fountain experiment). Forms $NH_4OH$ (weak base). Reactions: Acidic gas: $NH_3(g) + HCl(g) \rightarrow NH_4Cl(s)$ With hot $CuO$: $3CuO(s) + 2NH_3(g) \xrightarrow{\Delta} 3Cu(s) + 3H_2O(l) + N_2(g)$ With $Cl_2$: Excess $NH_3$: $8NH_3(g) + 3Cl_2(g) \rightarrow 6NH_4Cl(s) + N_2(g)$ Excess $Cl_2$: $NH_3(g) + 3Cl_2(g) \rightarrow NCl_3(l) + 3HCl(g)$ Burning in oxygen: $4NH_3(g) + 3O_2(g) \xrightarrow{\Delta} 2N_2(g) + 6H_2O(l)$ Catalytic Oxidation: $4NH_3(g) + 5O_2(g) \xrightarrow{Pt/Rh, 800^\circ C} 4NO(g) + 6H_2O(g)$ Aqueous solution reactions (gives hydroxide ppts.): $FeCl_3(aq) + 3NH_4OH(aq) \rightarrow Fe(OH)_3(s) \downarrow + 3NH_4Cl(aq)$ (Reddish-brown ppt.) $CuSO_4(aq) + 2NH_4OH(aq) \rightarrow Cu(OH)_2(s) \downarrow + (NH_4)_2SO_4(aq)$ (Pale blue ppt., forms deep blue solution in excess $NH_4OH$) C. Nitric Acid (HNO$_3$) Lab Preparation: $KNO_3(s) + H_2SO_4(conc.) \xrightarrow{ Manufacture (Ostwald's Process): Catalytic oxidation of $NH_3$: $4NH_3 + 5O_2 \xrightarrow{Pt/Rh, 800^\circ C} 4NO + 6H_2O$ Oxidation of NO: $2NO + O_2 \rightarrow 2NO_2$ Absorption of $NO_2$ in water: $3NO_2 + H_2O \rightarrow 2HNO_3 + NO$ As an Oxidizing Agent: With Copper: Dilute $HNO_3$: $3Cu + 8HNO_3(dilute) \rightarrow 3Cu(NO_3)_2 + 4H_2O + 2NO$ Conc. $HNO_3$: $Cu + 4HNO_3(conc.) \rightarrow Cu(NO_3)_2 + 2H_2O + 2NO_2$ With Carbon: $C + 4HNO_3(conc.) \rightarrow CO_2 + 2H_2O + 4NO_2$ With Sulphur: $S + 6HNO_3(conc.) \rightarrow H_2SO_4 + 2H_2O + 6NO_2$ D. Sulphuric Acid (H$_2$SO$_4$) Manufacture (Contact Process): Burning Sulphur: $S + O_2 \rightarrow SO_2$ Catalytic Oxidation: $2SO_2 + O_2 \xrightarrow{V_2O_5, 450^\circ C} 2SO_3$ Absorption: $SO_3 + H_2SO_4(conc.) \rightarrow H_2S_2O_7$ (Oleum) Dilution: $H_2S_2O_7 + H_2O \rightarrow 2H_2SO_4$ As an Acid (Dilute): Reacts with metals, oxides, hydroxides, carbonates, sulphites, sulphides (like HCl). $Zn + H_2SO_4(dilute) \rightarrow ZnSO_4 + H_2$ As an Oxidizing Agent (Concentrated): With Carbon: $C + 2H_2SO_4(conc.) \rightarrow CO_2 + 2H_2O + 2SO_2$ With Sulphur: $S + 2H_2SO_4(conc.) \rightarrow 3SO_2 + 2H_2O$ As a Dehydrating Agent (Concentrated): Dehydration of sugar: $C_{12}H_{22}O_{11} \xrightarrow{conc. H_2SO_4} 12C + 11H_2O$ (Black charring) Dehydration of $CuSO_4 \cdot 5H_2O$: $CuSO_4 \cdot 5H_2O \xrightarrow{conc. H_2SO_4} CuSO_4 + 5H_2O$ (Blue to white) Non-volatile Nature: Used to prepare more volatile acids. $2NaCl + H_2SO_4(conc.) \xrightarrow{\Delta} Na_2SO_4 + 2HCl$ 9. Organic Chemistry Introduction Unique Nature of Carbon: Tetravalency: Forms 4 bonds. Catenation: Forms long chains, branches, and rings with other carbon atoms. Forms single, double, and triple bonds. Structures: Straight chain, branched chain, cyclic (e.g., Benzene). Structure and Isomerism Structural Formulae: Alkanes (C-C single bond): $C_nH_{2n+2}$ Methane ($CH_4$), Ethane ($CH_3CH_3$), Propane ($CH_3CH_2CH_3$) Alkenes (C=C double bond): $C_nH_{2n}$ Ethene ($CH_2=CH_2$), Propene ($CH_3CH=CH_2$) Alkynes (C$\equiv$C triple bond): $C_nH_{2n-2}$ Ethyne ($CH \equiv CH$), Propyne ($CH_3C \equiv CH$) Isomerism: Compounds with the same molecular formula but different structural formulae. Chain Isomerism: Different carbon chain arrangements (e.g., n-butane, isobutane). Position Isomerism: Different positions of functional group/substituent (e.g., 1-propanol, 2-propanol). Homologous Series Series of organic compounds with similar chemical properties, same general formula, and successive members differ by a $CH_2$ group. Characteristics: Same general formula. Gradual change in physical properties (e.g., boiling point increases with molecular mass). Similar chemical properties. Can be prepared by general methods. Examples: Alkanes, Alkenes, Alkynes, Alcohols, Carboxylic acids. Nomenclature (IUPAC) Longest Chain Rule: Select the longest continuous carbon chain. Smallest Number Rule: Number the chain to give the functional group/substituent the lowest possible number. Functional Groups: -OH (Alcohol): -ol -CHO (Aldehyde): -al -COOH (Carboxylic Acid): -oic acid C=C (Alkene): -ene C$\equiv$C (Alkyne): -yne Hydrocarbons Alkanes: Saturated (single bonds). Preparation: Wurtz Reaction (from alkyl halides): $2CH_3I + 2Na \rightarrow CH_3-CH_3 + 2NaI$ Decarboxylation (from sodium salts of fatty acids): $CH_3COONa + NaOH \xrightarrow{CaO, \Delta} CH_4 + Na_2CO_3$ Combustion: $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$ Substitution (with $Cl_2$ in UV light): $CH_4 + Cl_2 \xrightarrow{UV} CH_3Cl + HCl$ Alkenes (Ethene): Unsaturated (double bond). Preparation: Dehydrohalogenation: $CH_3CH_2Br + KOH(alc.) \rightarrow CH_2=CH_2 + KBr + H_2O$ Dehydration of Ethanol: $CH_3CH_2OH \xrightarrow{conc. H_2SO_4, 170^\circ C} CH_2=CH_2 + H_2O$ Addition Reactions: Hydrogenation: $CH_2=CH_2 + H_2 \xrightarrow{Ni, \Delta} CH_3-CH_3$ Halogenation: $CH_2=CH_2 + Br_2 \rightarrow CH_2Br-CH_2Br$ (Decolorizes bromine water) Alkynes (Ethyne): Unsaturated (triple bond). Preparation: From Calcium Carbide: $CaC_2 + 2H_2O \rightarrow Ca(OH)_2 + C_2H_2$ From 1,2-dibromoethane: $CH_2Br-CH_2Br + 2KOH(alc.) \rightarrow CH \equiv CH + 2KBr + 2H_2O$ Addition Reactions: Similar to alkenes, but can add twice. Alcohols (Ethanol, $CH_3CH_2OH$) Preparation: Hydrolysis of alkyl halide: $CH_3CH_2Br + KOH(aq) \rightarrow CH_3CH_2OH + KBr$ Properties: Physical: Colorless liquid, characteristic smell, soluble in water, boiling point higher than corresponding alkanes (H-bonding). Chemical: Combustion: $C_2H_5OH + 3O_2 \rightarrow 2CO_2 + 3H_2O$ Action with Na: $2CH_3CH_2OH + 2Na \rightarrow 2CH_3CH_2ONa + H_2$ Esterification (with acetic acid): $CH_3COOH + CH_3CH_2OH \rightleftharpoons CH_3COOCH_2CH_3 + H_2O$ (Sweet smell) Dehydration (to ethene): $CH_3CH_2OH \xrightarrow{conc. H_2SO_4, 170^\circ C} CH_2=CH_2 + H_2O$ Denatured Alcohol: Ethanol made unfit for drinking by adding poisonous substances (e.g., methanol). Uses: Solvent, fuel, antiseptic, in beverages. Carboxylic Acids (Acetic Acid, $CH_3COOH$) Structure: $\underset{\overset{O}{||}}{CH_3-C-OH}$ Properties: Physical: Pungent smell (vinegar), glacial acetic acid freezes below 17$^\circ$C to ice-like crystals. Chemical: Acidic (turns blue litmus red). Reacts with alkalis: $CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O$ Esterification (with ethanol): $CH_3COOH + C_2H_5OH \rightleftharpoons CH_3COOC_2H_5 + H_2O$ Uses: Flavoring agent, solvent, in preparing esters, vinegar (5-8% solution).