BS

Chapter 3: Acids and Bases

  • Intro to Bronsted-Lowry Acids and Bases

    • based on the transfer of a proton (H+)

    • acids are ‘proton donors’

    • bases are ‘proton acceptors’

    • acid + base → conj. acid + conj. base

  • Flow of Electron Density: Curved-Arrow Notation

    • all reactions happen via movement of electrons → including acid-base

    • illustrated with curved arrows, beginning on the base’s lone pair and ending on the acid’s proton, which has a secondary arrow that shows the bond dissolving into a new lone pair on the acid

    • the reaction mechanism

    • involves electrons from a base ‘deprotonating’ an acid → acids can’t lose protons without a participating base

  • Bronsted-Lowry Acidity: Comparing pK_a values

    • using pK_a values to compare acidity

      • K_a and pK_a refer to general acidity. (pK_a is log of K_a)

      • acidity is important when considering the equilibrium of an acid-base reaction (more accurately, the stability of the acid vs conj acid)

      • equilibrium will favor the side of the relatively weaker acid (more stable/less reactive)

      • stability is influenced by: resonance, inductive effect, electronegativity/atomic size, location of formal charges

    • using pK_a values to compare basicity

      • compounds can be only acidic or basic in a specific scenario

      • therefore, the stronger parent/conj. acid has the weakest base

    • using pK_a Values to predict the position of equilibrium

      • equilibrium will always favor formation of the weaker acid (the higher pK_a)

      • indicated by drawing a longer arrow pointed to the favored side, and a shorter arrow pointing to the least favorable side

  • 3.4 Bronsted-Lowry Acidity: Factors Affecting the Stability of Anions

    • Conjugate base stability

      • look to the conj bases when comparing acids with no net charge

      • the one with the most stable/weakest conj. base is the strongest acid

    • Factors affecting the stability of anionic conjugate bases

      • which atom bears the charge: the most electronegative atom prefers to carry negative charges, and the least prefers to carry the positive charges

      • Resonance: molecules are more stable if a charge/lone pair is delocalized via resonance, so molecules with resonance will be more stable than those without

      • Induction: compounds carrying a very electronegative atom/substituent such as a halogen will be more stable since the electronegative atoms draw electron density towards themselves.

      • Orbitals: in cases where two protons exhibit the same affects (such as a negative charge on a carbon atom in the conjugate base), considering what orbitals the lone pairs are in can clarify. orbitals hold electrons closer as they become less hybridized (sp3→sp) Therefore, negative charges on less hybridized (sp→ sp3) are more stable.

    • Ranking the factors that affect the stability of anions

      • order of priority is: atom characteristics (electronegativity/size) → resonance → induction → orbital

      • remember: ARIO

  • Bronsted-Lowry Acidity: Assessing the Relative Acidity of Cationic Acids

    • rather than comparing conjugate bases, simply compare stability of positive charge on each acid

    • uses the same factors as with anions→ focusing on atom characteristics since the others are unpredictably effective

  • Positions of Equilibrium and Choice of Reagents

    • equilibrium can be predicted from relative stability of ions rather than pK_a if those values are unavailable

    • in reactions involving anions: the most stable anion is favored

    • in reactions involving cations: the most stable cation is favored

    • position of either situations is determined by finding the relative stability of each ion

    • when the answer is unclear for cations: evaluate the conjugate bases (remember that the answer is the inverse→ weakest base=strongest acid)

  • Leveling Effect

    • bases stronger than hydroxide can’t be used in a water solvent

    • This is because, while the base is strong enough to deprotonate water molecules, the resulting hydroxide is more stable and is favored over the base to the point that the base is completely used up

    • the leveling effect says that the base cannot be stronger than the conjugate base of the solvent. → if the solvent is water, hydroxide is the strongest possible base

    • this is also true for acids

  • Solvating Effects

    • solvent effects can explain small differences in pKa values.

    • bulkier conjugate bases experience ‘steric hindrance’ when interacting with the solvent

    • this means that molecule size also affects the relative strength of acids and bases

  • Counterions

    • negative charged bases are accompanied by a cation like Li, Na, or K always

    • They do not participate in the reaction and thus are spectator ions

  • Lewis Acids and Bases

    • defined in broader terms: transfer of electron pairs rather than protons

    • lewis bases are electron-pair donors (donating a lone pair to the acid)

    • lewis acids are electron-pair acceptors

    • expands acid/base classification because it includes compounds that otherwise wouldn’t qualify as either (i.e. BF3)

    • Involves only one curved-arrow when drawn out in that notation rather than two