Other Important Acid-Base Concepts - CHM 241 Chapter 3 Lecture Part 3

Predicting the Position of Equilibrium in Acid-Base Reactions

• When considering any acid-base reaction, there are two distinct methodologies utilized to predict which side of the chemical equation is favored at equilibrium:

  • pKa Values: By comparing the $pKa$ values of the two acids present in the reaction ($H-A$ and the conjugate acid $H-B$), the side containing the acid with the higher $pKa$ value (the weaker acid) will be favored at equilibrium.

  • Relative Stability of Bases: By analyzing the relative stability of the bases ($A^-$ and $B^-$), one can determine the equilibrium position. The reaction will favor the side with the more stable (weaker) conjugate base.

• Refer to Skillbuilder 3.10 for practice in applying these predictive methods.

Choosing an Appropriate Reagent for Acid-Base Reactions

• An essential skill in organic chemistry is the selection of a suitable reagent to facilitate a specific acid-base transformation.

• To effectively protonate a specific molecule, one must choose an acid from a standardized list (such as Table 3.1 in the Klein textbook) that possesses a sufficiently low $pKa$ to drive the reaction toward the desired products.

The Leveling Effect of Solvents

• Solvents play a critical role in acid-base chemistry by surrounding the reactants and facilitating the molecular collisions necessary for a reaction to occur. A key requirement is that the solvent must not react with the reagents themselves.

• Definition of the Leveling Effect: Because water ($H_2O$) can function as either an acid or a base, it exerts a "leveling effect" on strong acids and bases.

  • Strong Acids: Acids that are stronger than the hydronium ion ($H_3O^+$) cannot be used in water. For example, if sulfuric acid ($H_2SO_4$) is added to water, it will react immediately to produce $H_3O^+$. Consequently, no actual sulfuric acid remains in the solution to react with other intended reagents.

  • Strong Bases: Bases that are stronger than the hydroxide ion ($OH^-$) cannot be used in water. Such bases would spontaneously deprotonate water to form $OH^-$, thereby effectively replacing the stronger base with hydroxide.

• Appropriate Use of Water: Water is a suitable solvent only when the base used is not stronger than hydroxide ($pKa$ of conjugate acid of hydroxide/water = $15.7$).

  • Example Case: Reaction of the acetate ion ($CH_3CO_2^-$). When water is the solvent, the acetate ion might react with water, but because the $pKa$ of the resulting acetic acid ($4.75$) is much lower than the $pKa$ of water ($15.7$), the equilibrium heavily favors the left side (the reactants). Thus, water remains an appropriate solvent because it does not significantly interfere with the reagent.

Solvating Effects and Acidity

• While the ARIO (Atom, Resonance, Induction, Orbital) framework is usually sufficient to explain acidity, it cannot explain the difference in $pKa$ between certain similar molecules, such as ethanol and tert-butanol.

• Relative Stability of Conjugate Bases: As with all acids, differences in acidity are rooted in the stability of their conjugate bases. In cases where ARIO factors are identical or very similar, the ability of the solvent to stabilize the conjugate base becomes the deciding factor.

• Ion-Dipole Attractions: To stabilize a formal negative charge, the solvent must form ion-dipole attractions around the conjugate base.

• Steric Hindrance:

  • Ethoxide: This base is relatively unhindered, allowing solvent molecules to surround and stabilize the negative charge efficiently.

  • tert-Butoxide: This base is sterically hindered (bulky). Due to its size and shape, solvent molecules cannot get as close to the negative charge to stabilize it.

• Conclusion: Because tert-butoxide is less effectively solvated and therefore less stable than ethoxide, tert-butanol is less acidic than ethanol.

Counter Ions (Spectator Ions)

• Counter ions are also referred to as spectator ions. They are always present in a solution to maintain electrical neutrality and balance the overall charge.

• Role in Reactions: Although they are physically present, they typically do not participate in the chemical transformation.

• Notation: In full chemical equations, counter ions (such as $Na^+$ or $Li^+$) are included. However, in organic chemistry mechanisms and shorthand reaction equations, they are frequently omitted for clarity, focusing only on the reactive species.

Lewis Acids and Bases

• The Lewis definition of acids and bases focuses on electron pairs rather than protons:

  • Lewis Acid: A substance that accepts and shares a pair of electrons.

  • Lewis Base: A substance that donates and shares a pair of electrons.

• Relationship with Brønsted-Lowry Definition:

  • All acids categorized under the Brønsted-Lowry definition (proton donors) also qualify as Lewis acids because the proton acts as the electron pair acceptor.

  • All bases categorized under the Brønsted-Lowry definition (proton acceptors) also qualify as Lewis bases because they use an electron pair to bond with the proton.

• Distinct Differences: Some reactions can be classified as Lewis acid-base reactions but cannot be classified as Brønsted-Lowry reactions. This occurs when an electron pair is donated to an atom other than a hydrogen proton (for example, donation to a boron or metal center).

• Practice identifying these reactions is available in SkillBuilder 3.12.