Module 2-2

• Polar Covalent Bonds

  • The bonding electrons are attracted more strongly by one atom than the other, resulting in an unsymmetrical electron distribution between the atoms.
• Bond polarity is due to differences in electronegativity (EN), the intrinsic ability of an atom to attract the shared electrons in a covalent bond.
• Arrhenius Acid - Increase the concentration of H+ ions when dissolved in water
• Arrhenius Base - Increase the concentration of OHions when dissolved in water
• Brønsted–Lowry Acid - Proton donor
• Brønsted–Lowry Base - Proton acceptor
• Lewis Acid - electron acceptor
• Lewis Base - electron donor
• The Brønsted–Lowry Definition

  • Structural effects on acidity and basicity of organic molecules.
• The Lewis Definition

  • The use of curved arrows to show simple proton-transfer reactions
• Acidity – measure of the tendency of a compound to give up a proton
• Basicity – measure of a compound’s affinity for a proton
• Acid Strength: pH and pKa

  • The tendency of an acid to donate a proton.
  • The more readily the compound donates a proton, the stronger the acid.
• Do not confuse pH and pKa:

  • The pH scale is used to describe the acidity (concentration of positively charged hydrogen ions) of a solution
  • The pKa is characteristic of a particular compound, which indicates the tendency of the compound to give up its proton.
• The acid dissociation constant, Ka
• Equilibrium always favors formation of the weaker acid and base
• An acid can be deprotonated by the conjugate base of any acid having a higher pKa.
• Remember, we said that the stronger the acid, the weaker its conjugate base. Therefore, a weaker conjugate base means a more stable conjugate base because if it was not as stable as it is, it would have reacted with the proton and shifted the reaction backward, forming the acid.
• pKa and Acid Strength – What Factors Affect the Acidity?

  The acidity relies on the stability of the conjugate base. There are four main factors, in the following priority order, that affect the stability of the conjugate base:

  1. Atom
  2. Resonance
  3. Induction
  4. Orbital
• Atom Effects

  • Consider an atom that is connected to the hydrogen: From left to right of the periodic table, acidity increases as the electronegativity increases.
  • The better it stabilizes the negative charge, the more stable the conjugate base is.
  • Down the periodic table, the atomic size determines the acidity and not the electronegativity.
• Resonance Effect

  • Delocalized electrons are more stable.
• Inductive Effects

  • The stronger the electronegativity, the stronger the inductive effect
• Inductive Effects

  • Another factor is the distance of the electronegative element from the negative charge. The closer it is, the better it helps to stabilize the negative charge
• Orbital (or Hybridization) Effects

  • The only difference is these carbons is their hybridization state. Remember, alkanes are sp3 , alkenes are sp2 and alkynes are sp-hybridized.
  • s orbitals are more electronegative than p orbital, and the more s character the hybrid orbital has, the better it stabilizes the negative charge.
• The Lewis Definition

  • Lewis definition of acids and bases is broader and more encompassing than the Bronsted-Lowry definition because it is not limited to substances that donate or accept protons only
• Lewis bases are structurally the same as Brønsted–Lowry bases. Both have an available electron pair—a lone pair or an electron pair in a π bond.
• A Brønsted–Lowry base always donates this electron pair to a proton, but a Lewis base donates this electron pair to anything that is electron deficient.
• Any species that is electron deficient and capable of accepting an electron pair is also a Lewis acid.
• All Brønsted–Lowry acids are also Lewis acids, but the reverse is not necessarily true.
• A Lewis acid is also called an “electrophile” (electron-deficient)
• When a Lewis base reacts with an electrophile other than a proton, the Lewis base is also called a “nucleophile” (electron-rich).