Identify the conjugate acid-base pairs in each of the following equations: (a) \(\mathrm{H}_{2} \mathrm{~S}+\mathrm{NH}_{3} \rightleftharpoons \mathrm{NH}_{4}^{+}+\mathrm{HS}^{-}\) (b) \(\mathrm{HSO}_{4}^{-}+\mathrm{NH}_{3} \rightleftharpoons \mathrm{SO}_{4}^{2-}+\mathrm{NH}_{4}^{+}\) (c) \(\mathrm{HBr}+\mathrm{CH}_{3} \mathrm{O}^{-} \rightleftharpoons \mathrm{Br}^{-}+\mathrm{CH}_{3} \mathrm{OH}\) (d) \(\mathrm{HNO}_{3}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{NO}_{3}^{-}+\mathrm{H}_{3} \mathrm{O}^{+}\)

Acid-base chemistry is a fundamental topic in chemistry, focusing on the reactions between acids and bases. An acid is a substance that can donate a proton (H+), while a base is one that can accept a proton. These substances interact in solutions to form conjugate pairs, which change based on the gain or loss of protons.
Acid-base reactions are crucial in numerous chemical processes such as metabolism in biology, industrial synthesis, and environmental chemistry. For example, acids like hydrochloric acid (HCl) and bases like sodium hydroxide (NaOH) play significant roles in laboratories and industries.
Understanding acid-base chemistry involves grasping the concepts of pH, pKa, and the different strength levels of acids and bases. Strong acids dissociate completely in water, providing a greater concentration of hydrogen ions, while weak acids only partially dissociate.

Conjugate acid-base pairs are pairs of compounds that differ by a single proton. When an acid donates a proton, it becomes its conjugate base. Conversely, when a base accepts a proton, it becomes its conjugate acid. These pairs help maintain the balance of reactions in acid-base chemistry.
For instance, in the reaction \(\text{H}_{2}S + \text{NH}_{3} \rightleftharpoons \text{NH}_{4}^{+} + \text{HS}^{-}\), \(\text{H}_{2}S\) is an acid that donates a proton to become \(\text{HS}^{-}\), its conjugate base. Similarly, \(\text{NH}_{3}\) accepts a proton to become \(\text{NH}_{4}^{+}\), its conjugate acid.
This concept is vital for predicting the outcome of acid-base reactions and understanding how buffer solutions work to resist pH changes. Conjugate pairs allow chemists to visualize and balance chemical equations effectively.

Chemical equations represent the chemical reactions where the reactants transform into products. They must be balanced to reflect the conservation of mass and charge. In the context of acid-base reactions, equations typically show how acids and bases form conjugate pairs.
For example, the equation \(\text{HNO}_{3} + \text{H}_{2}O \rightarrow \text{NO}_{3}^{-} + \text{H}_{3}O^{+}\) shows nitric acid (HNO3) donating a proton to water (H2O). Here, HNO3 is the acid, H2O is the base, \(\text{NO}_{3}^{-}\) is the conjugate base, and \(\text{H}_{3}O^{+}\) is the conjugate acid.
Writing and balancing these equations require recognizing the acids and bases and their respective conjugate forms. This helps in understanding the direction of the reaction and predicting the extent to which the reaction goes, assisting in understanding real-life processes like neutralization, titration, and buffer systems.

The Bronsted-Lowry theory is a concept in acid-base chemistry that defines acids and bases. According to this theory, an acid is a proton donor, while a base is a proton acceptor. This explanation broadens the understanding beyond just hydrogen ions in aqueous solutions, encompassing a variety of solvents and reactions.
For instance, in the equation \(\text{HBr} + \text{CH}_{3}\text{O}^{-} \rightleftharpoons \text{Br}^{-} + \text{CH}_{3}\text{OH}\), hydrobromic acid (\text{HBr}) donates a proton to methoxide (\text{CH}_{3}\text{O}^{-}). This reaction shows the formation of bromide (\text{Br}^{-}), the conjugate base, and methanol (\text{CH}_{3}\text{OH}), the conjugate acid.
The Bronsted-Lowry theory helps chemists understand a wider range of reactions and makes it easier to study acid-base behavior in non-aqueous environments. This theory allows for a more comprehensive approach to understanding the mechanisms of acid-base interactions in diverse chemical contexts.