Suppose \(2.0\) moles of sodium acetate, \(\mathrm{NaO}_{2} \mathrm{CCH}_{3}\), are dissolved in some water and then \(1.0 \mathrm{~L}\) of a \(1.0 \mathrm{M} \mathrm{HCl}\) solution is added. (a) Write the chemical reaction that occurs. (b) After the reaction, what are the predominant species in solution? How many moles of each species are there? (c) Is the resulting solution a buffer? If yes, explain why.

Stoichiometry is foundational in understanding chemical reactions and it concerns with the quantification of reactants and products in a chemical process. It allows us to predict the amounts of products and reactants that are involved in a reaction based on balanced chemical equations.

When applying stoichiometry to an exercise such as the reaction between sodium acetate and hydrochloric acid, it’s crucial to use a balanced chemical equation. In the provided exercise, the stoichiometry is straightforward because the reaction occurs in a 1:1 ratio: one mole of sodium acetate reacts with one mole of hydrochloric acid to produce one mole of acetic acid and one mole of sodium chloride. Understanding this concept helps students gauge the exact amount of each substance that will be present at the end of the reaction.

Acid-base reactions are a type of chemical reaction that involve the transfer of protons (H⁺) from an acid to a base. One of the most common types of acid-base reactions is a neutralization reaction where an acid and a base combine to form water and a salt.

In our exercise, hydrochloric acid (HCl), a strong acid, donates a proton to the base sodium acetate, which forms acetic acid and sodium chloride (a salt). This specific reaction is important in buffer solution chemistry because it showcases how weak acids and bases behave in solution. Their ability to donate or accept protons underpins the functioning of buffer solutions, which is to maintain a stable pH in a system.

A conjugate acid-base pair consists of two species that differ by the presence or absence of a proton. When an acid donates a proton, it becomes its conjugate base, and when a base accepts a proton, it becomes its conjugate acid. The ability to interconvert like this allows the substance to either buffer against pH changes or drive specific chemical reactions forward.

In the context of our exercise, after sodium acetate reacts with HCl, the acetic acid produced is the conjugate acid of the acetate ion. It's essential to understand that the presence of both members of this pair in the solution is a prerequisite for buffer action. This concept is vital as it plays a direct role in governing the solution’s ability to resist pH changes upon the addition of more acid or base.

Chemical equilibrium occurs when a chemical reaction and its reverse reaction proceed at the same rate, leading to a stable ratio of products and reactants. It's a dynamic state, meaning that the reaction hasn’t stopped, but there is no net change in the concentrations of the reactants and products over time.

In buffers, the equilibrium between the weak acid or base and its conjugate counterpart is what allows for the buffering action. The conjugate pairs can react with added acids or bases to shift the equilibrium minimally, thus stabilizing pH. This concept is intimately linked to understanding buffer solutions like the one formed in the exercise, where the weak acid, acetic acid, and its conjugate base, the acetate ion, establish an equilibrium that resists drastic changes in pH.