Explain why (a) the pH decreases when lactic acid is added...

Explain why (a) the $\mathrm{pH}$ decreases when lactic acid is added to a sodium lactate solution. (b) the $\mathrm{pH}$ of $0.1 \mathrm{M} \mathrm{NH}_{3}$ is less than 13.0 . (c) a buffer resists changes in $\mathrm{pH}$ caused by the addition of $\mathrm{H}^{+}$ or $\mathrm{OH}^{-}$. (d) a solution with a low $\mathrm{pH}$ is not necessarily a strong acid solution.

Explain why
(a) the $\mathrm{pH}$ decreases when lactic acid is added to a sodium lactate solution.
(b) the $\mathrm{pH}$ of $0.1 \mathrm{M} \mathrm{NH}_{3}$ is less than 13.0 .
(c) a buffer resists changes in $\mathrm{pH}$ caused by the addition of $\mathrm{H}^{+}$ or $\mathrm{OH}^{-}$.
(d) a solution with a low $\mathrm{pH}$ is not necessarily a strong acid solution.

Key Concepts

Buffer Systems

Buffer systems consist of a weak acid and its conjugate base, or a weak base and its conjugate acid. They are designed to maintain a relatively constant pH upon the addition of small amounts of acids or bases. This occurs because the weak acid can react with any added base, and the conjugate base can react with any added acid, thereby minimizing changes in the pH of the solution.

Common Ion Effect

The common ion effect is the phenomenon where the addition of an ion that is already present in a solution shifts the equilibrium position, reducing the dissociation of a weak acid or base. This principle explains why adding lactic acid to a sodium lactate solution causes the equilibrium to shift, increasing the concentration of the undissociated acid and thus lowering the pH.

Weak vs Strong Acids and Bases

The strength of an acid or base is determined by its degree of dissociation in water. A strong acid or base dissociates completely, while a weak one does not. Importantly, the pH of a solution depends on both the concentration and the degree of ionization of the acid or base present. Consequently, a solution with a low pH does not necessarily imply the presence of a strong acid; it may contain a high concentration of a weak acid that partially dissociates.

Equilibrium and Le Chatelier's Principle

Chemical equilibria, such as those in acid-base reactions, can shift when a concentration change is imposed by the addition of reactants or products. Le Chatelier's principle predicts that the system will adjust to counterbalance the imposed change. This principle is key to understanding how the addition of acids or bases to buffers, or the introduction of common ions, influences the pH of the solution.

Base Dissociation in Aqueous Solutions

When bases like ammonia are dissolved in water, they undergo partial reaction with water to form their corresponding conjugate acids and hydroxide ions, rather than complete dissociation. This limited dissociation, defined by the base dissociation constant (Kb), results in a pH that reflects the equilibrium between the unreacted base and its ion products, explaining why the pH of a weak base solution is lower than that of a solution containing a completely dissociated strong base.

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