Acids and Bases; Functional Groups

Polarity of Bonds and Molecules

Bond polarity describes the distribution of electron density between atoms in a chemical bond. It ranges from nonpolar covalent, through polar covalent, to ionic bonds, depending on the difference in electronegativity between the bonded atoms.

• Bond Dipole Moment (μ): A measure of the separation of positive and negative charges in a bond, expressed in debye (D).

• Polarity Examples:

  • H3C–CH3 (non-polar)

  • H3C–NH2 (δ+ δ−)

  • H3C–OH (δ+ δ−)

  • H3C–Cl (ionic)

• Bond Dipole Moments in Organic Compounds:

  • H–C: 0.3 D

  • H–O: 1.53 D

  • C–N: 0.22 D

  • C–O: 0.86 D

  • C=O: 1.56 D

  • C≡N: 3.6 D

• Lone Pairs: Lone pairs of electrons contribute to molecular polarity by creating regions of partial positive and negative charge.

Intermolecular Forces and Boiling Points

Intermolecular forces are non-covalent interactions between molecules that affect physical properties such as boiling points.

• Types of Intermolecular Forces:

  • London Dispersion Forces: Temporary attractive forces due to induced dipoles. Strength increases with molecular surface area.

  • Dipole-Dipole Forces: Attractions between permanent dipoles in polar molecules. Polar compounds have higher boiling points than non-polar ones.

  • Hydrogen Bonding: Strong dipole-dipole attraction involving O–H or N–H bonds. Not a true bond, but significantly increases boiling points.

• Boiling Point Trends:

  • Longer carbon chains → higher boiling point

  • More branching → lower boiling point

• Examples:

  • n-butane: b.p. 0°C

  • n-pentane: b.p. 36°C

  • isopentane: b.p. 28°C

  • neopentane: b.p. 10°C

• Hydrogen Bonding Energy: ~29 kJ/mol (7 kcal/mol)

• Hydrogen Bonding Boiling Points:

  • H3C–CH2–OH: 78°C

  • H3C–CH2–NH2: 17°C

  • H3C–CH2–Cl: 12°C

Functional Groups

Functional groups are specific groups of atoms within molecules that are responsible for characteristic chemical reactions.

• Alkane: Only C–C and C–H single bonds

• Alkene: Contains C=C double bond

• Alkyne: Contains C≡C triple bond

• Benzene: Aromatic hydrocarbon

• Alkyl Halide: R–X (X = halogen)

• Alcohol: R–OH

• Thiol: R–SH

• Ether: R–O–R'

• Aldehyde: R–CHO

• Ketone: RCOR'

• Carboxylic Acid: RCOOH or RCO2H

• Ester: RCOOR' or RCO2R'

• Amide: RCONH2, RCONHR', RCONR'2

• Amine: R–NH2, R–NHR', R–NR'2

• Examples: Structures of propanal (aldehyde), propanol (alcohol), propanamide (amide), and ethyl acetate (ester).

Acids and Bases

Acids and bases are defined in organic chemistry by the Brønsted-Lowry and Lewis definitions.

• Brønsted-Lowry Acids: Species that donate a proton (H+).

• Brønsted-Lowry Bases: Species that accept a proton.

• Lewis Acids: Species that accept a pair of electrons.

• Lewis Bases: Species with non-bonding electrons that can be donated.

• Acid Strength: Expressed by the extent of ionization in water.

• Acid Dissociation Constant ():

• pKa: Convenient logarithmic measure of acid strength.

  • The smaller the pKa, the stronger the acid.

• Examples:

  • Ethanol: ,

  • Acetic acid: ,

Acid and Conjugate Base Table

Factors Affecting Acidity

The strength of an acid is influenced by the stability of its conjugate base. Several factors contribute to this stability:

• Electronegativity: More electronegative atoms stabilize negative charge better, increasing acid strength.

  • Stability order: F− > O− > N− > C−

  • Acidity order: HF > H2O > NH3 > CH4

• Size of Anions: Larger anions spread negative charge over a greater volume, stabilizing the conjugate base.

  • Size order: F− < Cl− < Br− < I−

  • Acidity order: HF > HCl > HBr > HI

• Resonance Stabilization: Delocalization of negative charge via resonance increases stability of the conjugate base.

  • Example: Acetic acid is more acidic than ethanol because its conjugate base is resonance stabilized.

• Inductive Effect: Electron-withdrawing groups stabilize the conjugate base by pulling electron density away, spreading out the negative charge.

  • Example: Chloroacetic acid () is more acidic than acetic acid () due to the inductive effect of Cl.

• Hybridization: Non-bonding electrons in orbitals with more s-character are more stable.

  • Stability order: > >

  • Example: Acetylene () > Ethylene () > Ethane ()

Summary Table: Factors Affecting Acidity

Key Equations

Examples and Applications

• Comparing Acid Strength: Acetic acid () is stronger than ethanol ().

• Conjugate Base Stability: The more stable the conjugate base, the stronger the acid.

• Acid-Base Equilibria: Reactions favor formation of the weaker acid and weaker base.

Additional info: These notes summarize the foundational concepts of acids, bases, functional groups, and factors affecting acidity in organic chemistry, suitable for college-level study and exam preparation.