Question

Write the formula of the conjugate acid of each base. (a) \(\mathrm{OH}^{-}\) (b) \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\) (c) \(\mathrm{HCO}_{3}^{-}\)

Short answer

The conjugate acid of \(\mathrm{OH}^{-}\) is \(\mathrm{H}_{2}\mathrm{O}\). The conjugate acid of \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\) is \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+}\). The conjugate acid of \(\mathrm{HCO}_{3}^{-}\) is \(\mathrm{H}_{2}\mathrm{CO}_{3}\)

Step-by-step solution

Find the Conjugate Acid of \(\mathrm{OH}^{-}\)

A conjugate acid is formed when a base gains a proton. Hence, the conjugate acid of \(\mathrm{OH}^{-}\) is formed when \(\mathrm{OH}^{-}\) gains a proton, resulting in \(\mathrm{H}_{2}\mathrm{O}\).

Find the Conjugate Acid of \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\)

Applying the same principle as above to the base \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\). This base will gain a proton forming its conjugate acid, resulting in \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+}\).

Find the Conjugate Acid of \(\mathrm{HCO}_{3}^{-}\)

The base \(\mathrm{HCO}_{3}^{-}\) will accept a proton resulting in its conjugate acid. The conjugate acid of \(\mathrm{HCO}_{3}^{-}\) is \(\mathrm{H}_{2}\mathrm{CO}_{3}\).

Key concepts

Brønsted-Lowry theory

The Brønsted-Lowry theory is a fundamental framework in the study of acid-base chemistry. It was proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. This theory redefined acids and bases in terms of proton (H+) transfer.

An acid, according to the Brønsted-Lowry theory, is a substance that can donate a proton to another substance, while a base is a substance that can accept a proton. This exchange of protons is what defines acid-base reactions under this theory. It’s a more generalized concept than the earlier theory proposed by Svante Arrhenius, which was limited to aqueous solutions and required acids and bases to form ions in water.

For example, in the exercises provided, elements like \(\mathrm{OH}^{-}\), \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\), and \(\mathrm{HCO}_{3}^{-}\) are considered bases as they have the capacity to accept protons. When they gain a proton, they form their conjugate acids: water (\(\mathrm{H}_{2}\mathrm{O}\)), anilinium ion (\(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+}\)) and carbonic acid (\(\mathrm{H}_{2}\mathrm{CO}_{3}\)), respectively.

Proton transfer reactions

Proton transfer reactions are the foundation of acid-base interactions. They encompass the Brønsted-Lowry acid-base definition by focusing on the movement of protons from one molecule to another.

During these reactions, a proton donor, which is the acid, interacts with a proton acceptor, the base, resulting in the formation of a conjugate base (from the acid) and a conjugate acid (from the base). These reactions are reversible, which means that the conjugate pairs can also function in reverse as acids and bases, showcasing the concept of chemical equilibrium.

Steps to Identify Proton Transfer

• Look for a compound that gains a proton - it acts as a base.
• Look for a compound that loses a proton - it acts as an acid.
• After the proton transfer, identify the resulting conjugate acid and conjugate base.

Following these steps in the homework exercise helps students identify the conjugate acids for \(\mathrm{OH}^{-}\), \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\), and \(\mathrm{HCO}_{3}^{-}\), providing a practical application of the theory.

Acid-base chemistry

Acid-base chemistry is a branch of chemical science that focuses on the properties and reactions of acids and bases. It's critical for understanding various biological systems, chemical reactions, and industrial processes.

In this context, the strength of an acid or base can affect reaction rates, equilibrium, and the overall outcome of chemical processes. Strong acids and bases disassociate completely in solution, yielding more protons or hydroxide ions, respectively. Weak acids and bases do not fully disassociate, and their equilibrium position lies far to the left, meaning more of the original compounds remain in solution.

Students often encounter these principles while dealing with pH calculations, buffer solutions, and titrations. Understanding the conjugate pairs, as seen in the provided exercises, is key as it illustrates the relationship between acids and bases and how they transform during reactions. By recognizing the conjugate acid of \(\mathrm{OH}^{-}\) as water, \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\) as \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+}\), and \(\mathrm{HCO}_{3}^{-}\) as \(\mathrm{H}_{2}\mathrm{CO}_{3}\), learners gain insight into how substances can vary in acidity or basicity in different environments.