Organic Biochemistry Lecture Notes

ACIDS, BASES, AND SALTS

  • General Characteristics:

    • Acids: Sour in taste; contains H+H^+ and a nonmetal (e.g., HClHCl) or Hydrogen and a negative radical (polyatomic ion) (e.g., HNO3HNO_3).
    • Bases: Bitter in taste; contains OH−OH^-, a metal, and hydroxide (e.g., KOHKOH), or a positively charged radical and hydroxide (e.g., NH4OHNH_4OH).
    • Salts: Salty in taste; contains a Metal and a Nonmetal or a Negative Radical (Polyatomic Ion) (e.g., Na2SO4Na_2SO_4).

  • Arrhenius Acid-Base Theory:

    • Arrhenius Acid: A hydrogen-containing compound that produces H+H^+ ions in solution upon the addition of water.
      • Example: HNO3→H++NO3−HNO_3 \rightarrow H^+ + NO_3^-.
      • Example (forming hydronium): Nitric Acid added to water produces Hydronium Ion (H3O+H_3O^+) and Nitrate Ion (NO3−NO_3^-).
    • Arrhenius Base: A hydroxide-containing compound that produces OH−OH^- (Hydroxide) ions in solution.
      • Example: NaOH→Na++OH−NaOH \rightarrow Na^+ + OH^-.

  • Ionization vs. Dissociation:

    • Ionization: The process where individual positive and negative ions are produced from a molecular compound dissolved in water (e.g., Arrhenius acids like HCl→H++Cl−HCl \rightarrow H^+ + Cl^-).
    • Dissociation: The process where individual positive and negative ions are released from an ionic compound dissolved in solution (e.g., Arrhenius bases where Metal + OH−OH^- separates, such as NaOH→Na++OH−NaOH \rightarrow Na^+ + OH^-).

  • Bronsted-Lowry Acid-Base Theory:

    • Bronsted-Lowry Acid: A substance that can donate a proton (H+H^+ ion) to another substance; defined as a proton donor.
    • Bronsted-Lowry Base: A substance that can accept a proton (H+H^+ ion) from another substance; defined as a proton acceptor.
    • Example: HCl+H2O→Cl−+H3O+HCl + H_2O \rightarrow Cl^- + H_3O^+ (HClHCl is the acid donor; H2OH_2O is the base acceptor).

  • Acids in Water Behavior:

    • Safety Rule: Always add Acid to Water, never water to acid. Adding water to acid can cause a violent reaction, especially with strong acids. This should be performed under a fume hood.
    • General Equation: HA(aq)+H2O(l)⇌H3O(aq)++A(aq)−HA_{(aq)} + H_2O_{(l)} \rightleftharpoons H_3O^+_{(aq)} + A^-_{(aq)}.
    • Conjugate Acid: The species formed after the base accepts a proton.
    • Conjugate Base: The species left behind (ion) after the acid has donated its proton.
    • Conjugate Pairs: HA/A−HA/A^- (Acid and Conjugate Base) and H2O/H3O+H_2O/H_3O^+ (Base and Conjugate Acid).

  • Acid Ionization Equilibrium:

    • Defined as the state where the rate of reactant consumption equals the rate of product formation.
    • Reactant ⇌\rightleftharpoons Product (Rate of Ion++Ion−Ion^+ + Ion^- formation is equal).

  • Exercises: Conjugate Identification:

    • Conjugate base of HSO4−HSO_4^-: −H+→SO42−-H^+ \rightarrow SO_4^{2-}.
    • Conjugate acid of NO3−NO_3^-: +H+→HNO3+H^+ \rightarrow HNO_3.
    • Conjugate base of H3PO4H_3PO_4: −H+→H2PO4−-H^+ \rightarrow H_2PO_4^-.
    • Conjugate acid of HC2O4−HC_2O_4^-: +H+→H2C2O4+H^+ \rightarrow H_2C_2O_4.

  • Amphiprotic Substances:

    • A substance that can either lose or accept a proton, functioning as either an acid or a base (e.g., H2OH_2O can become H3O+H_3O^+ or OH−OH^-).

  • Proticity of Acids:

    • Monoprotic Acid: Supplies one proton (H+H^+) per molecule (e.g., HA+H2O⇌A−+H3O+HA + H_2O \rightleftharpoons A^- + H_3O^+).
    • Diprotic Acid: Supplies two protons per molecule, occurring in two steps (e.g., H2A→HA−→A2−H_2A \rightarrow HA^- \rightarrow A^{2-}).
    • Triprotic Acid: Supplies three protons per molecule, occurring in three steps (e.g., H3A→H2A−→HA2−→A3−H_3A \rightarrow H_2A^- \rightarrow HA^{2-} \rightarrow A^{3-}).
    • Polyprotic Acid: Includes any acid supplying two or more protons.

  • Strength of Acids and Bases:

    • Strong Acids: Transfer approximately 100%100\% of protons to water. Equilibrium lies far to the right, yielding a weak conjugate base.
    • Weak Acids: Transfer only a small percentage of protons. Equilibrium lies far to the left. The weaker the acid, the stronger its conjugate base.
    • Strong Bases: Typically hydroxides of Groups IA and IIA metals.

  • Ionization Constants (KaK_a and KbK_b):

    • Acid Ionization Constant (KaK_a): The equilibrium constant for a weak acid reacting with water. Ka=[H3O+][A−][HA]K_a = \frac{[H_3O^+][A^-]}{[HA]}.
    • Base Ionization Constant (KbK_b): The equilibrium constant for a weak base reacting with water. Kb=[BH+][OH−][B]K_b = \frac{[BH^+][OH^-]}{[B]}.
    • Relationships: Acid strength increases as percent ionization increases and as the magnitude of KaK_a increases.

  • Salts and Neutralization:

    • Salts are ionic compounds containing a metal or polyatomic ion as the positive ion and a nonmetal or polyatomic ion (except hydroxide) as the negative ion.
    • Neutralization Reaction: Acid+Base→Salt+WaterAcid + Base \rightarrow Salt + Water.
      • Example: HCl+NaOH→NaCl+H2OHCl + NaOH \rightarrow NaCl + H_2O.
      • Example: H2SO4+2KOH→K2SO4+2H2OH_2SO_4 + 2KOH \rightarrow K_2SO_4 + 2H_2O.

  • Ion Product Constant for Water (KwK_w):

    • At 24∘C24^\circ C (and 25∘C25^\circ C): Kw=[H3O+][OH−]=1.00×10−14K_w = [H_3O^+][OH^-] = 1.00 \times 10^{-14}.
    • pKw=−log⁡(Kw)=14pK_w = -\log(K_w) = 14.
    • pH+pOH=14pH + pOH = 14.
    • pKa+pKb=pKw=14pK_a + pK_b = pK_w = 14.

  • pH and concentration Calculations:

    • pH=−log⁡[H3O+]pH = -\log[H_3O^+].
    • pOH=−log⁡[OH−]pOH = -\log[OH^-].
    • If pH=5pH = 5, then [H+]=10−5 M[H^+] = 10^{-5}\,M. Antilog calculation: shift+log⁡(−pH)\text{shift} + \log (-pH).
    • Solution States:
      • Neutral: [H3O+]=[OH−][H_3O^+] = [OH^-].
      • Acidic: [H3O+]>[OH−][H_3O^+] > [OH^-].
      • Basic: [H3O+]<[OH−][H_3O^+] < [OH^-].
    • Exercise: If [OH−]=1.0×10−4 M[OH^-] = 1.0 \times 10^{-4}\,M, then pOH=4pOH = 4, pH=10pH = 10, [H3O+]=1.0×10−10 M[H_3O^+] = 1.0 \times 10^{-10}\,M (Basic).
    • Exercise: If pH=5.85pH = 5.85, then [H3O+]=1.4×10−6 M[H_3O^+] = 1.4 \times 10^{-6}\,M.

  • Salt Hydrolysis Types:

    1. Strong Acid + Strong Base: No hydrolysis; solution is neutral (e.g., KCl,NaNO3KCl, NaNO_3).
    2. Strong Acid + Weak Base: Hydrolyzes to produce an acidic solution (e.g., NH4ClNH_4Cl).
    3. Weak Acid + Strong Base: Hydrolyzes to produce a basic solution (e.g., NaF,KC2H3O2NaF, KC_2H_3O_2).
    4. Weak Acid + Weak Base: Result depends on the relative weakness of components.

  • Buffers:

    • An aqueous solution that prevents major changes in pH when small amounts of acid/base are added.
    • Composed of a weak acid and its conjugate base (salt).
    • Mechanism:
      • Added OH−OH^- reacts with H3O+H_3O^+; equilibrium shifts right to replenish H3O+H_3O^+.
      • Added H3O+H_3O^+ shifts equilibrium left to consume the excess.
    • Henderson-Hasselbalch Equation: pH=pKa+log⁡[A−][HA]pH = pK_a + \log\frac{[A^-]}{[HA]}.
    • Exercise: Buffer with 0.45 M0.45\,M acetic acid and 0.85 M0.85\,M sodium acetate (Ka=1.8×10−5K_a = 1.8 \times 10^{-5}). pKa=4.74pK_a = 4.74. pH=4.74+log⁡0.850.45=5.02pH = 4.74 + \log \frac{0.85}{0.45} = 5.02.

INTRODUCTION TO ORGANIC CHEMISTRY

  • Scope: Study of carbon compounds. Living things consist of organic chemicals like proteins (hair), DNA (genetics), foods, and medicines.

  • Composition: Carbon plus Hydrogen, Oxygen, and Nitrogen. Sometimes contains Sulfur, Phosphorus, and Halogens (F,Cl,Br,IF, Cl, Br, I).

  • Exclusions: CO2,CO,Na2CO3,CO32−,CO_2, CO, Na_2CO_3, CO_3^{2-}, and CN−CN^- are considered inorganic.

  • Carbon Criticality: Over 90%90\% of 30 million compounds contain carbon. Carbon (Group 4A) shares 4 valence electrons via 4 covalent bonds (1s22s22p21s^2 2s^2 2p^2).

  • History: Vital Force Theory suggested living organisms were needed to produce organic compounds. Wohler (1828) debunked this via experiment.

  • VSEPR Theory: Common bond angles are 109.5∘109.5^\circ (tetrahedral), 120∘120^\circ (trigonal), and 180∘180^\circ (linear).

  • Electronegativity and Bond Polarity:

    • Electronegativity (EN): Ability to attract shared electrons. F=4.0F = 4.0 (highest), Cs=0.7Cs = 0.7 (lowest), C=2.5C = 2.5.
    • Nonpolar Covalent: EN difference ≤0.4\le 0.4.
    • Polar Covalent: EN difference >0.4> 0.4 to 2.02.0.
    • Ionic: EN difference >2.0> 2.0.
    • Dipole Moment (μ\mu): Net molecular polarity. Symmetrical molecules may have local dipoles that cancel out.

  • Organic Structure Bonding Rules:

    • Carbon: 4 bonds, 0 unshared pairs.
    • Hydrogen: 1 bond, 0 unshared pairs.
    • Nitrogen: 3 bonds, 1 unshared pair.
    • Oxygen: 2 bonds, 2 unshared pairs.
    • Halogen: 1 bond, 3 unshared pairs.

  • Kinds of Organic Reactions:

    1. Addition: Two molecules combine.
    2. Elimination: One molecule splits into two.
    3. Substitution: Parts of two molecules exchange.
    4. Rearrangement: Molecule undergoes structural connection changes.

  • Reaction Mechanisms:

    • Describes the step-by-step transformation from reactant to product.
    • Concerted: Several steps occurring simultaneously.
    • Homolytic (Radical): Symmetrical bond breaking; indicated by "fish-hook" arrows.
    • Heterolytic (Polar): Unsymmetrical bond breaking; indicated by full-head arrows.

  • Hybridization and Orbitals:

    • Hybridization: Mixing atomic orbitals from a single atom to produce identical hybrid orbitals.
    • sp3sp^3: Tetrahedral; 1s + 3p overlap; 109.5∘109.5^\circ; 4 sigma bonds (e.g., Methane).
    • sp2sp^2: Trigonal Planar; 1s + 2p overlap; 120∘120^\circ; 1 sigma and 1 pi bond (e.g., Ethylene).
    • spsp: Linear; 1s + 1p overlap; 180∘180^\circ; 1 sigma and 2 pi bonds (e.g., Acetylene).
    • Sigma ($\sigma$) bond: Head-to-head overlap with maximum electron density on the internuclear axis.
    • Pi ($\pi$) bond: Side-to-side overlap of unhybridized parallel p orbitals.

FUNCTIONAL GROUPS OVERVIEW

  • Alkane: C−CC-C single bond; formula CnH2n+2C_nH_{2n+2}; Suffix: -ane.
  • Alkene: C=CC=C double bond; formula CnH2nC_nH_{2n}; Suffix: -ene.
  • Alkyne: C≡CC \equiv C triple bond; formula CnH2n−2C_nH_{2n-2}; Suffix: -yne.
  • Alcohol: −OH-OH (Hydroxyl); formula CnH2n+1OHC_nH_{2n+1}OH; Suffix: -ol.
  • Ether: R−O−R′R-O-R' (Alkoxy group); Suffix: -yl -oxy.
  • Aldehyde: Carbonyl group (C=OC=O) with a terminal hydrogen; Suffix: -al.
  • Ketone: Carbonyl group bonded to two alkyl groups; Suffix: -one.
  • Carboxylic Acid: Carbonyl + Alcohol (−COOH-COOH); Suffix: -oic acid.
  • Ester: Carbonyl + OR′OR'; Suffix: -oate.
  • Amine: Contains Nitrogen (1∘,2∘,3∘1^\circ, 2^\circ, 3^\circ); Suffix: -amine.
  • Amide: Carbonyl group attached to Nitrogen; Suffix: -amide.
  • Thiol: −SH-SH (Sulfhydryl); Suffix: -thiol or mercaptan.
  • Aromatic: Benzene-like rings (e.g., C6H6C_6H_6).

ALKANES AND CYCLOALKANES

  • Physical Properties:

    • Almost complete lack of polarity.
    • Interaction via weak London dispersion forces.
    • Solubility: Insoluble in water ("like dissolves like"); soluble in nonpolar solvents (toluene, ether).
    • Density: Less dense than water (<1.0 g/mL< 1.0\,g/mL); floats.

  • Isomerism:

    • Constitutional Isomers: Same molecular formula, different connectivity.
    • Number of isomers increases with carbon count (C4H10C_4H_{10} has 2; C6H14C_6H_{14} has 5).

  • IUPAC Nomenclature for Alkanes:

    • Prefix indicates carbon count (1-meth, 2-eth, 3-prop, 4-but, 5-pent, 6-hex, 7-hept, 8-oct, 9-non, 10-dec).
    • 11-undec, 12-dodec, 13-tridec, 14-tetradec, 15-pentadec, 16-hexadec, 17-heptadec, 18-octadec, 19-nonadec, 20-eicos.
    • Rules:
      1. Find longest parent chain.
      2. Number from the end closest to a substituent.
      3. Use prefixes di-, tri-, tetra- for multiple identical substituents.
      4. Alphabetize different substituents (ignore di-, tert-, sec- in alphabetization, except iso).

  • Cycloalkanes:

    • Saturated cyclic hydrocarbons (cyclopentane,cyclohexanecyclopentane, cyclohexane).
    • Cis-Trans Isomerism: Restricted rotation by the ring allows stereoisomers.
      • Cis: Substituents on the same side.
      • Trans: Substituents on opposite sides.

OXYGEN AND NITROGEN DERIVATIVES

  • Alcohols:

    • Compounds with −OH-OH on a tetrahedral carbon.
    • IUPAC: Parent -ane changes to -ol. Carbon-1 is the one bearing the hydroxy group in cyclic versions.
    • Glycols: Hydroxyl groups on adjacent carbons.

  • Aldehydes and Ketones:

    • Physical Properties: Polar molecules due to C=OC=O bond. Lower boiling points than alcohols (no hydrogen bonding). Small molecules are soluble in water.
    • Oxidation: Aldehydes oxidize to carboxylic acids. Ketones resist oxidation. Tollens' Reagent: Specific for aldehydes, forms a "silver mirror."
    • Reduction: Aldehydes $\rightarrow$ Primary Alcohol; Ketones $\rightarrow$ Secondary Alcohol. Laboratory reagent: NaBH4NaBH_4. Biological agent: NADHNADH.
    • Hemiacetals and Acetals: Formed by addition of alcohols. Cyclic hemiacetals (5-6 membered rings) are very stable.

  • Carboxylic Acids and Derivatives:

    • Dicarboxylic Acids: Suffix -anedioic acid (e.g., Ethanedioic acid/Oxalic acid).
    • Fischer Esterification: Acid+Alcohol+Acid Catalyst→Ester+WaterAcid + Alcohol + Acid\,Catalyst \rightarrow Ester + Water.
    • Saponification: Base-catalyzed hydrolysis of an ester using hot aqueous base (NaOHNaOH).
    • Anhydrides: Two carbonyls shared by one oxygen. React with alcohols to give ester + acid.
    • Amides: Carbonyl bonded to Nitrogen. Cyclic amides are called lactams (e.g., Penicillin).

  • Amines:

    • Basic compounds. Aliphatic amines are stronger bases than ammonia.
    • Amine Salts: Formed by reacting amines with strong acids (water-soluble).
    • Heterocyclic Amine: Nitrogen is part of a ring.

POLYMERS

  • Nylon-66: First purely synthetic fiber; a polyamide made from two six-carbon monomers.
  • Kevlar: Polyaromatic amide from aromatic dicarboxylic acid and diamine.
  • Lexan: Common polycarbonate formed from the disodium salt of bisphenol A and phosgene.
  • Polyesters: Involve polymerization of diesters and diols.