Organic Chemistry: Acids and Bases

Organic Chemistry

Chapter 2: Acids and Bases

Learning Objectives

• 2.1 Polar covalent bonds: Electronegativity

• 2.2 Polar covalent bonds: Dipole moments

• 2.3 Acids and bases: The Brønsted–Lowry definition

• 2.4 Acid and base strength

• 2.5 Predicting acid–base reactions from pKa values

• 2.6 Organic acids and organic bases

Polar Covalent Bonds: Electronegativity

• Covalent bonds can exhibit ionic character.

• Polar covalent bonds: Bonding electrons are attracted more strongly by one atom than by the other, leading to an asymmetrical electron distribution between the atoms.

Electronegativity

• Definition: The intrinsic ability of an atom to attract shared electrons in a covalent bond.

• Effects of Electronegativity: Differences in electronegativity (EN) produce bond polarity.

  • Fluorine (F): Most electronegative element with an EN value of 4.0.

  • Cesium (Cs): Least electronegative with an EN value of 0.7.

  • Periodic Trend: Metals on the left side of the periodic table attract electrons weakly, while halogens (reactive nonmetals on the right) attract electrons strongly.

  • Electronegativity of Carbon (C): 2.5.

Figure 2.2 - Electronegativity Values and Trends

• Graphical representation of various elements' electronegativity values illustrating their relative strengths from Li (1.0) to F (4.0).

Bond Polarity and Inductive Effect

• Inductive effect: The shifting of electrons in a sigma (σ) bond in response to the electronegativity of nearby atoms.

• Polar Covalent Bonds: If the difference in EN of the bonded atoms is less than 2, the bond is categorized as polar. Examples include C–O and C–X bonds.

  • Carbon (C) acquires a partial positive charge (δ+) while the electronegative atom acquires a partial negative charge (δ-).

  • C–H bonds are considered relatively nonpolar.

• In contrast, if the difference in EN is greater than 2, the bond is classified as ionic (e.g., LiF, NaCl).

Electrostatic Potential Maps

• Example: Methanol (CH₃OH) possesses a polar covalent C–O bond, while methyllithium (CH₃Li) has a polar covalent C–Li bond.

• Electrostatic Potential Maps: Computer-generated representations that use color to indicate charge distributions in molecules, where red signifies electron-rich areas (δ-) and blue indicates electron-poor regions (δ+).

Worked Example

• Question: Identify the more electronegative element in the following pairs:
  (a) Na or Be
  (b) F or Br

• Solution:
  (a) Na (0.9) is less electronegative compared to Be (1.6).
  (b) F (4.0) is more electronegative compared to Br (2.8).

Polar Covalent Bonds: Dipole Moments

• Molecules exhibit polarity due to the vector summation of individual bond polarities along with contributions from lone pairs of electrons.

• Solubility Note: Strongly polar substances are generally soluble in polar solvents (e.g., water), whereas nonpolar substances are insoluble in water.

• Dipole Moment (μ): Defined as the net molecular polarity arising from charge differences across the molecule.

  • Formula: ext{μ} = Q imes r where

  • Q = magnitude of charge at the end of the molecular dipole,

  • r = distance between charges.

Additional Information on Dipole Moments

• The dipole moment measured in debyes (D) is an important parameter;

  • 1 D = 3.336 imes 10^{-30} coulomb meter.

  • A typical dipole moment for an average covalent bond is 1.60 × 10^{-29} C·m or approximately 4.80 D.

Absence of Dipole Moments

• In symmetrical molecules, the dipole moments from each bond may cancel each other due to opposite orientation.

Dipole Moments in Water and Ammonia

• Water (H₂O) and ammonia (NH₃) exhibit large dipole moments because the electronegativities of oxygen (O) and nitrogen (N) are greater than that of hydrogen (H).

• Both O and N have lone-pair electrons oriented away from all nuclei, intensifying the net dipole.

Worked Example

• Task: Draw a three-dimensional representation of H₂C═CH₂ and predict if it has a dipole moment.

• Solution: The dipole moment is zero, indicated by an arrow pointing from the less electronegative to the more electronegative element.

Acids and Bases: The Brønsted-Lowry Definition

• Traditional definitions of acids and bases (i.e., solutions containing H⁺ or OH⁻) are inadequate in organic chemistry.

• Brønsted-Lowry Theory: Defines acids and bases based on their role in proton (H⁺) transfer reactions.

  • Brønsted-Lowry Acid: Donates a hydrogen ion (H⁺).

  • Brønsted-Lowry Base: Accepts a hydrogen ion (H⁺).

• Importance of terminology: A proton is synonymous with H⁺, resulting from the loss of a valence electron from the hydrogen atom, leaving behind only its nucleus.

Additional Definitions in Acids and Bases

• Conjugate Base: The species formed by the deprotonation of a Brønsted-Lowry acid.

• Conjugate Acid: The species formed by the protonation of a Brønsted-Lowry base.

Acid Base Strength

• Acidity Constant (Ka): Measures the strength of an acid in reaction with water to produce hydronium ions (H₃O⁺):
  Ka = rac{[H3O^+][A^-]}{[HA]}

• In this equation, brackets [$ ext{ }$] denote concentrations measured in moles per liter.

Measures of Acid Strength

• Acid strengths are commonly represented using pKa values.

  • pKa Definition: The negative logarithm of the acidity constant (Ka).

  • The relationship: pKa = - ext{log} Ka

  • Strong acids correspond to lower pKa values, while weaker acids correspond to higher pKa values.

Table 2.3: Relative Strengths of Some Common Acids and Their Conjugate Bases

Acid and Base Strength

• Water is both an acid and base solvent, characterized by its ion product:
  K_w = [H₃O^+][OH⁻]

• The molar concentration of pure water at 25°C is: [H₂O] = 55.4 ext{M}

  • This interaction can be measured as K = rac{1.0 imes 10^{-14}}{55.4} leading to pKa = 15.74.

Worked Example

• Problem: Compare the acid strengths of phenylalanine (pKa = 1.83) and tryptophan (pKa = 2.83).

• Solution: The stronger acid has the lower pKa value. Therefore, phenylalanine is stronger than tryptophan based on their pKa values.

Predicting Acid-Base Reactions from pKa Values

• The pKa values indicate the logarithmic relation to the equilibrium constants, helping predict the likelihood of a given acid-base reaction occurring.

• The difference in pKa values offers a method to calculate the transfer extent in protonation or deprotonation reactions.

Example: pKa Values of Acetic Acid and Water

• Acetic acid has a pKa of 4.76, while water has a pKa of 15.74.

  • Which is the stronger acid?

  • Does hydroxide ion (
    e.g., OH⁻) react significantly with acetylene?

Organic Acids

• Characterized typically by the presence of a positively polarized hydrogen atom (H) covalently bonded to an electronegative atom, often oxygen.

• Examples of Organic Acids:

  • Methanol (pKa = 15.54)

  • Acetic acid (pKa = 4.76)

  • Acetone (pKa = 19.3)

Anion Stabilization

• Two main types of organic acids can stabilize their respective anions:

  • Those with a hydrogen atom bonded to an electronegative oxygen atom (O–H).

  • Others have a hydrogen atom linked to a carbon adjacent to a C═O bond (O═C─C─H).

Anion Stabilization in Detail

• An anion can be stabilized through resonance and by possessing a negative charge on a highly electronegative atom, which significantly impacts acidity.

Organic Bases

• Characterization: Organic bases contain atoms like nitrogen or oxygen with lone pairs capable of bonding to H⁺ ions.

• Common Organic Bases: Compounds derived from ammonia are prevalent organic bases.

• Organic bases may also react as acids with strong bases.

Zwitterion Example: Alanine

• Uncharged Form: H₂N, O=C, CH₃ (Alanine)

• Zwitterion Form: H₃N⁺, O=, CH₃ (Zwitterionic structure of alanine)

Acids and Bases: The Lewis Definition

• Lewis Acid: Defined as an electron pair acceptor.

• Lewis Base: Defined as an electron pair donor.

• Characterized by filled orbitals (Lewis base) and vacant orbitals (Lewis acid).

  • Diagram visualizing the bond formation with a Lewis base donating electrons to a Lewis acid indicating electron movement with curved arrows.

Some Examples of Lewis Acids and Bases

• Lewis Acids Include:

  • Neutral proton donors, e.g., H₂O, HCl, HBr, HNO₃.

  • Cations such as Li⁺, Mg²⁺, and other metal compounds (AlCl₃, TiCl₄, FeCl₃).

• Lewis Bases Include:

  • Oxygen and nitrogen-containing organic compounds with transferable electron pairs.

  • Alcohols, ethers, aldehydes, ketones, carboxylic acids, amines, etc., showcase versatility as both acids and bases.

Acid-Base Complexes

• An acid-base complex forms when a Lewis acid accepts an electron pair from a Lewis base.

• Example Process: Interaction between BF₃ (Lewis acid) and NH₃ (Lewis base), illustrated with curved arrows showing electron pair movement.

• Visual representations can include structures of complexes formed between Lewis acids and bases showing their interactions and resultant structures.

Worked Example of Acetaldehyde

• Task: Illustrate how acetaldehyde (CH₃CHO) acts as a Lewis base.

• Structure Interaction: Acetaldehyde demonstrates the ability to bond with protons (H⁺) through its carbonyl oxygen while facilitating a reaction with a Lewis acid represented as dashed lines in diagrams.