Predict the position of equilibrium, and calculate the equilibrium constant, \(K_{\text {cq, }}\), for each acid-base reaction. (a) \(\mathrm{CH}_{3} \mathrm{NH}_{2}+\mathrm{CH}_{3} \mathrm{COOH} \rightleftharpoons \mathrm{CH}_{3} \mathrm{NH}_{3}^{+}+\mathrm{CH}_{3} \mathrm{COO}^{-}\) \(\begin{array}{ccc}\text { Methylamine Aceticacid } & \begin{array}{c}\text { Methylammonium } \\ \text { ion }\end{array} & \begin{array}{c}\text { Acetate } \\ \text { ion }\end{array}\end{array}\) (b) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{O}^{-}+\mathrm{NH}_{3} \rightleftharpoons \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}+\mathrm{NH}_{2}{ }^{-}\) Ethoxide ion Ammonia Ethanol Amide ion

Understanding acid-base reactions is crucial for predicting the behavior of substances in a chemical reaction. An acid is a substance that donates a proton (H⁺), while a base is a substance that accepts a proton. When an acid reacts with a base, it forms a conjugate base and a conjugate acid, as seen in the reaction between methylamine and acetic acid. During such reactions, the transfer of protons between reactants leads to the formation of products with varying acid and base strengths.

These reactions can be represented by equilibrium expressions, which show the reversible nature of the chemical process. The ability to identify the acid, base, and their respective conjugate pairs in a given reaction is a foundational skill in understanding acid-base chemistry. For instance, in the provided exercise, methylamine acts as a base by accepting a proton, whereas acetic acid donates a proton, acting as an acid.

Predicting the position of equilibrium in an acid-base reaction involves assessing the relative strengths of the acids and bases on both sides of the reaction. At equilibrium, a reaction has a balance of forward and reverse processes, with no net change in the concentration of reactants and products over time.

The position of equilibrium is influenced by the strengths of the reactants and products; a reaction will tend to favor the formation of the weaker acid and base. For example, because methylammonium ion is a stronger acid than acetic acid, and acetate ion is a weaker base than methylamine, the equilibrium in reaction (a) shifts to the right. Understanding these dynamics helps chemists to predict the outcomes of reactions and manipulate conditions to drive reactions in a desired direction.

When comparing acid and base strengths, we look at their tendencies to donate or accept protons. A strong acid readily donates its proton, while a strong base readily accepts a proton. The strength of an acid or base is often determined by the stability of its conjugate base or acid.

Factors Influencing Acid and Base Strength

• Electronegativity: Atoms with higher electronegativity can stabilize negative charge better, making for a stronger acid.
• Size: Larger atoms distribute negative charge over a greater volume, resulting in a more stable and thus stronger acid.
• Resonance: The ability of a charge to be delocalized by resonance can also stabilize the conjugate base, increasing the acid strength.

In essence, the stronger an acid, the weaker its conjugate base, and vice versa. Acid-base strengths play a pivotal role in driving the equilibrium position of a reaction.

The equilibrium constant, denoted as K_eq, quantifies the position of equilibrium in a reaction. It is defined as the ratio of the concentrations of products to reactants at equilibrium, each raised to the power of their stoichiometric coefficients. In acid-base reactions, we are often dealing with aqueous solutions and hence use the acid dissociation constant, K_a, and the base dissociation constant, K_b, to express the strength of an acid and base, respectively.

To calculate the equilibrium constant for a given reaction, we need to know the dissociation constants of the acids and bases involved. For instance, using the dissociation constants provided for acetic acid and methylamine, we calculated K_eq for reaction (a). In a similar way, we can calculate K_eq for any acid-base reaction as long as we have the dissociation constants of the reactants.