Organic Test PDF

Ionic bonds Covalent bonds Dative bond electrostatic forces of electrostatic forces of must have a lone pair of e , attraction between 2 attraction between shared e- acceptor atom must have oppositely charged ions...

Ionic bonds Covalent bonds Dative bond electrostatic forces of electrostatic forces of must have a lone pair of e , attraction between 2 attraction between shared e- acceptor atom must have oppositely charged ions pair and +ve nuclei vacant low lying orbital Soluble in polar solvents Soluble with same IMF solvents (formation of strong ion ­dipole interactions release a lot of energy) Ionic with covalent character: electronegative cation polarises the anion→↑degree of covalency Covalent with ionic character: large electronegativity difference between bonding atoms → large dipole moment → ↑degree of ionic character VSEPR: molecular geometry of a species → e pairs are arranged as far as possible to minimise inter electronic repulsion - l.p − l.p repulsion > l.p − b.p repulsion > b.p − b.p repulsion - As electronegativity of central atom↑, ↑b.p- ­b.p repulsion, (nearer to each other) Bond strength= B.E Is dependent on: 1. No. of bonds (No. of shared e- → attraction→ bond strength) bonds 2. Degree of Eff overlap ( when size of orbital↑→ orbitals become more diffuse→ Eff overlap↓→↓bond strength) 3. ΔEN ( as ΔEN↑ → ↑polar→ resultant net dipoles cause additional electrostatic forces of attraction→↑bond strength) Instantaneous dipole- Permanent dipole- H­ydrogen Bonding induced dipole Permanent dipole (caused by momentary (polar molecules; forces of (a form of pd­-pd, strong movements of e- charge, attraction between -δ and forces of attraction between and is present in all +δ ) electron deficient H+ and particles) lone pair on F/O/N molecules) ↑e cloud →↑ I.M.F strength ↑δ moments→↑I.M.F strength ↑N H−bonds formed with other molecules per ↑surface area→↑ molecule → ↑I.M.F strength I.M.F strength (e.g branched hydrocarbons) Acid Base Eqm pg 40-42 ∆H= Hf products - Hf reactants ∆H= Hc reactants - Hc products ∆H= BB- BF ∆Hsolution = [Hhydration M+ - HhydrationX-] - LE (MX) Limitations: → Thermodynamic feasibility = spontaneity, whether it can occur → kinetic feasibility = whether the rate of rxn is observable (depends on E ) activation Organic Chemistry Key Concepts: 1. Steric hindrance: when presence of a substituent hinders the approach of an attacking rxt and prevents rxn/↓rxty of a particular site 2. Isomerism 3. Rotation of plane polarised light - Enantiomerism: optical isomerism - Racemic mixture: equimolar ratio of both enantiomers - No plane of symmetry - ≥ 1 chiral centre (sp hybridised) usually 4. Resonance: donation/ withdrawal of e through overlap of subst & benzene ring 5. Acidity of Carboxylic acid - Resonance structure of -COO- - -­ve charge dispersed/delocalised over two highly electronegative O atoms - stable conjugate base 6. Basicity of Amines - available lone pair on N atom for coordination with a proton - presence of e donating alkyl grps increase availability of lone pair for donation - aromatic amines: the lone pair delocalises into benzene ring = unavailable for donation Oxidation (✖ 3°‒OH, Ketone, ‒COOH) Common products: H₂O, CO₂ Exceptions: methanoic/ ethanedioic acid w KMnO₄, H₂SO₄ (aq), Heat Alkenes 1°‒OH 2°‒OH Aldehyde K₂Cr₂O₇ (aq) H₂SO₄ (aq) H₂SO₄ (aq) H₂SO₄ (aq) Heat Heat Heat w immediate distillation Ketone Acid Aldehyde OR  H₂SO₄ (aq) Heat under reflux Acid KMnO₄ (aq) (mild) H₂SO₄/ H₂SO₄ (aq) H₂SO₄ (aq) H₂SO₄ (aq) H₂SO₄/ NaOH (aq) Heat under Heat Heat NaOH (aq) Cold reflux Cold Ketone Acid Benzoic Acid Diol Acid OR (oxidative cleavage) H₂SO₄ (aq) Heat Reduction H₂ (g) Reactant Alkenes, Aldehydes, Ketones, Nitriles (CN) Arene (Ni/Pt high temp high pressure) Ni,heat / Pt, room Product Alkane, 1° Alcohol, 2° Alcohol, 1° Amine temp Cyclohexane NaBH₄ Reactant Aldehydes, Ketones alkaline methanol Product 1° Alcohol, 2° Alcohol LiAlH₄ Reactant Aldehydes, Ketones, Carboxylic Acid, Nitriles, Ester in dry ether room temp stronger Product 1° Alcohol, 2° Alcohol, 1° Alcohol, 1°Amine, Acid+ Alcohol than NaBH4 because more polar 1. Sn(s), Reactant ⚠︎ Nitrobenzene conc. HCl, heat Avoid LiAlH₄ to prevent reduction of other fn. groups / formation of 2. NaOH side products (aq), room temp Product Phenylamine Na (Redox) NaOH (Acid-Base) NaCO𖾔 (Acid-Base) ‒OH ✔ ✔ ✔ ‒COOH ✔ ✔ ✔ NaOH (gives H₂SO₄ H₂O (ℓ) H/OH) Hydrolysis With heat, With heat, At rtp. 1. Ester 1. Ester 1. -COCl Alcohol/Acid Alcohol/Acid Acid 2. ‒CN Acid 2. ‒CN Acid 3. Amide Acid 3. Amide Acid Acid-Base At rtp. At rtp. 1. Phenol 1. Amine 2. ‒COOH 2. Amino Acid 3. Amino Acid - zwitterions Electrophilic Addition Alkenes 1. Alkenes 1. Conc. H₂SO₄ , - (Lab) Heat cold - (Industrial) steam, 2. Water heat conc. H𖾔PO₄, high temp, high Pa  1° Alcohol 1° Alcohol Condensation Conc. , heat Alcohol + Acid = Ester Elimination Or KOH in 1. Alcohol - conc. , heat ethanol, heat Alkene RX Alkene Halogen X₂, UV light, HX (g) X₂ (aq) PX₅ (s) / high temp rtp. rtp. PX₃(l)/ SOX₂(l) Anhydrous rtp. FRS Alkane Side-chain methyl benzene Mono-addition Alkene (X/H) Alkene (X/OH) Multi-addition Alkene CCl₄, dark Mono-sub Alcohol Arene Alcohol Anhydrous Dry HX(g)/ FeX₃/AlX₃/ Fe(s) Carboxylic acid Conc. HCl, ZnCl₂/ Phenol PX₃ most Conc. HBr CCl₄ effective Multi-sub Phenol Phenylamine Acyl Chlorides CH₃CO+ Cl- CH₃+ Cl-  Arene Anhydrous FeX₃/ AlX₃ rtp. Anhydrous FeX₃/ AlX₃ rtp. E. sub Alcohol Anhydrous Condensation forms esters rtp. (RCOCl + R’OH =RCOROH) vs RCOCl +NaOH = RCOOH Phenol NaOH(aq) Condensation rtp. (from sidechain)/ heat (from ring) Amine NH₃ in ethanol Condensation Heat with excess CH₃COCl in RNHR’COCl- sealed tube !! R can replace all H with prolonged reaction Electrophilic Sub. NO₂ -OR -CN Arene Conc. HNO₃, Conc. H₂SO₄ Heat in water bath Phenol (mono) dilute HNO₃ rtp. (multi) conc. HNO₃ rtp RX NaOR in methanal , NaCN / KCN in heat ethanol, heat Other Reactions 1. Thermal/Catalytic Cracking - 900℃ / 500℃ with catalyst 2. Nucleo. Add.-CN with Aldehyde - HCN, trace NaOH/NaCN - Cold Distinguishing Tests Alkene Br2 in CCl4, rtp Orange-red Br2 decolourises  Halogen 1. Heat with NaOH(aq) and cool Cl: white ppt. 2. Excess HNO3, then AgNO3 Br: cream ppt. 3. Excess NH3 I: yellow ppt. Alcohol PCl5(s) / SOCl2(l), rtp White fumes HCl(g) Phenol Neutral FeCl3(aq), rtp Violet colouration Carbonyl 2,4- DNPH, rtp Orange ppt. Aldehyde Tollen’s [Ag(NH3)2]+ (aq) Silver mirror Fehling’s [Alkaline Cu2+ Reddish-brown Cu2O ppt. complex] , heat Carboxylic Carbonate Effervescence. CO2 gas gives white ppt. In limewater. Acid Chlorides AgNO3 (aq), rtp White ppt. AgCl, White fumes HCl(g) Ester 1. Dilute H2SO4, heat (breaks ester linkages) with immediate distillation 2. Test for hydrolysed products (-OH) Primary Amide NaOH (aq), heat Effervescence. NH3 gas turns damp red litmus paper blue. Phenylamine 1. Br2(aq), rtp Orange solution decolourises, white ppt. 2. Neutral FeCl3(aq), rtp forms. No further changes in step 2. Transition Metals Cu(H2O)6]2+ MnO4- Cr2O72- [Fe(H2O)6]2+ NaOH Cu(OH)2 Mn(OH)2 ppt Cr(OH)3 ppt Fe(OH)2 ppt [Cr(OH)6]3+ Fe(OH)3 ppt NH3 Cu(OH)2 Mn(OH)2 ppt Cr(OH)3 ppt Fe(OH)2 ppt [Cu(NH3)4(H2O)]2+ KI / acidified Fe2+ CuI ppt brown Cr(H2O)6]3+ solution Na2S2O3 CuI ppt clear  [Cu(S2O3)2]3- Heat CuO Na2CO3 CuCO3 ppt MnCO3 ppt Cr(OH)3 ppt Fe(OH)3 ppt FeCO3 ppt HCl [CuCl4]2- Acid medium Mn(H2O)6]2+ Basic medium MnO2 ppt H2O2/ KMnO4/ Mn(OH)3 ppt CrO42- Fe(OH)3 ppt air MnO2 ppt Fe(H2O)6]3+ NH4SCN Fe(SCN)(H2O)5]2+ K4[Fe(CN)6] Fe4[Fe(CN)6]3 ppt “prussian” Periodic Table Period 3 Element Na Mg Al Si P S  Max. +1 +2 +3 +4 +5 +6 Oxidation State *Explain the variation in the highest oxidation number of the elements in oxides and chlorides Oxides Na2O MgO Al2O3 SiO2 P4O10 SO3 Solid Liquid Soluble Insoluble Soluble Conductor of electricity Poor N/A Ionic Giant Simple covalent Covalent pH 12-14 pH 8-9 pH 7 pH 2-3 pH 0-1 Chlorides NaCl MgCl2 AlCl3 SiCl4 PCl5 N/A Solid Liquid Solid Soluble Reacts to form acid Conductor Poor N/A Ionic Simple molecular Electrostatic forces id­id pH 7 pH 6-7 pH 3-4 pH 1-2 pH 1-2 *Explain the variation in bonding in oxides and chlorides in terms of electronegativity Ionic: Significant difference in electronegativity Ionic with covalent character: Intermediate electronegativity Covalent: Small difference in electronegativity Group 2 (Reducing Agents) Be Mg Ca Sr Ba Ra  θ 𝐸𝑐𝑒𝑙𝑙 More negative θ θ θ 𝐸𝑐𝑒𝑙𝑙 = (𝐸𝑟𝑒𝑑 ⇓) - 𝐸𝑜𝑥𝑖 Carbonate BeCO3 MgCO3 CaCO3 SrCO3 BaCO3 RaCO3 Density/ Polarisabili Decreasing ty Stability Increasing Why does thermal stability↑ down the grp? ↓charge density → ability to distort e cloud and polarise C − O bond↓ ­Therefore more heat req to break covalent C­O bond, causing↑decomp temp Group 17 (Oxidising Agents) F Cl Br I At θ 𝐸𝑐𝑒𝑙𝑙 Less positive θ θ θ 𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑟𝑒𝑑 - (𝐸𝑜𝑥𝑖 ⇓) Hydrides HF HCl HBr HI Bond energy Decreasing Stability Decreasing (i) electronic configuration (ii) atomic radius and ionic radius (iii) ionisation energy (iv) electronegativity (v) melting point (vi) electrical conductivity Enthalpy H chem energetics