Question

Write the conjugate acid for each of the following bases. a. \(\mathrm{ClO}^{-}\) b. \(\mathrm{Cl}^{-}\) c. \(\mathrm{ClO}_{3}^{-}\) d. \(\mathrm{ClO}_{4}^{-}\)

Short answer

The conjugate acids for the given bases are: a. \(\mathrm{HClO}\) b. \(\mathrm{HCl}\) c. \(\mathrm{HClO}_{3}\) d. \(\mathrm{HClO}_{4}\)

Step-by-step solution

Identify conjugate acids for each base

To find the conjugate acid for each base, add \(\mathrm{H}^{+}\) to the given base and increase the charge by one. #a. For the base \(\mathrm{ClO}^{-}\): Add a hydrogen ion, \(\mathrm{H}^{+}\), to \(\mathrm{ClO}^{-}\): \(\mathrm{ClO}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{HClO}\) The conjugate acid is \(\mathrm{HClO}\). #b. For the base \(\mathrm{Cl}^{-}\): Add a hydrogen ion, \(\mathrm{H}^{+}\), to \(\mathrm{Cl}^{-}\): \(\mathrm{Cl}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{HCl}\) The conjugate acid is \(\mathrm{HCl}\). #c. For the base \(\mathrm{ClO}_{3}^{-}\): Add a hydrogen ion, \(\mathrm{H}^{+}\), to \(\mathrm{ClO}_{3}^{-}\): \(\mathrm{ClO}_{3}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{HClO}_{3}\) The conjugate acid is \(\mathrm{HClO}_{3}\). #d. For the base \(\mathrm{ClO}_{4}^{-}\): Add a hydrogen ion, \(\mathrm{H}^{+}\), to \(\mathrm{ClO}_{4}^{-}\): \(\mathrm{ClO}_{4}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{HClO}_{4}\) The conjugate acid is \(\mathrm{HClO}_{4}\).

Key concepts

Acid-Base Chemistry

Understanding the fundamentals of acid-base chemistry is crucial for studying many aspects of chemistry, biology, and medicine. At its core, acid-base chemistry involves the transfer of protons (hydrogen ions, \( \mathrm{H}^{+} \) ) between substances. Acids are substances that can donate a proton to other substances, whereas bases are those that can accept a proton.

When studying solutions in water (aqueous solutions), the presence of acids and bases alters the pH level, which measures the acidity or basicity of the solution. A low pH means high acidity, while a high pH indicates a basic or alkaline environment. The pH scale typically ranges from 0 to 14, with 7 being neutral. Acids have a pH less than 7, and bases have a pH greater than 7.

In the exercise provided, by determining the conjugate acid for each base, students practice the core principle of acid-base reactions—the exchange of protons. This exercise offers a straightforward application of acid-base interactions, enhancing students' comprehension of the chemical behaviors of different substances.

Conjugate Pairs

A fundamental notion in acid-base reactions is the concept of conjugate pairs. A conjugate acid-base pair consists of two species that differ by one \( \mathrm{H}^{+} \) ion. When a base accepts a proton, it transforms into its conjugate acid, just as an acid releases a proton to become its conjugate base.

This transformation is reversible, thereby establishing the pair's 'conjugate' nature. For example, when water (\( \mathrm{H_2O} \) ) acts as a base and acquires a proton, it becomes hydronium (\( \mathrm{H_3O}^{+} \) ), its conjugate acid. Conversely, when a hydronium ion donates a proton, it returns to being water, its conjugate base.

Understanding Conjugate Pairs in Solutions

In the exercise, the students are required to identify the conjugate acid of given bases. By adding a proton to the base, the new species is now able to donate that proton in a subsequent reaction, fulfilling the criteria of a conjugate acid. Recognizing these pairs is imperative for predicting the outcomes of acid-base reactions and for understanding the balance of pH in chemical solutions.

Brønsted-Lowry Theory

The Brønsted-Lowry theory, introduced in 1923 by Johannes Nicolaus Brønsted and Thomas Martin Lowry, expanded the definition of acids and bases beyond the substances that produce \( \mathrm{H}^{+} \) and \( \mathrm{OH}^{-} \) ions in water. According to this theory, an acid is a proton donor, and a base is a proton acceptor, which allows for the explanation of acid-base reactions in non-aqueous solutions as well.

The Brønsted-Lowry theory introduced the concept of conjugate acid-base pairs, as previously discussed. It emphasizes the role of proton transfer in acid-base reactions and provides a broader framework that can apply to a wide variety of chemical systems, not just those in aqueous solutions. For instance, ammonia (\( \mathrm{NH_3} \) ) is a base because it can accept a proton to become \( \mathrm{NH_4}^{+} \) , its conjugate acid.

With the Brønsted-Lowry theory in mind, the exercise in identifying conjugate acids of given bases becomes an application of this foundational principle, demonstrating how bases are transformed upon gaining a proton, a key point often tested in academic settings.