Question

Fill in the blanks in the given table with the appropriate choice.$$ \begin{array}{|c|c|c|} \hline \text { Species } & \text { Conjugate acid } & \text { Conjugate base } \\\ \hline \mathrm{HCO}_{3}^{-} & \rho & \mathrm{CO}_{3}^{2-} \\ \hline \mathrm{HSO}_{4}^{-} & \mathrm{H}_{2} \mathrm{SO}_{4} & q \\ \hline \mathrm{NH}_{3} & r & -s \\ \hline \mathrm{H}_{2} \mathrm{O} & t & \mathrm{OH}^{-} \\ \hline \end{array} $$ (a) \(\mathrm{H}_{2} \mathrm{CO}_{3} \quad \mathrm{SO}_{4}^{2-}\) \(\begin{array}{lll}\mathrm{NH}_{4}^{*} & \mathrm{NH}_{2}^{-} & \mathrm{H}_{3} \mathrm{O}^{*}\end{array}\) \(\begin{array}{lllll}\text { (b) } & \mathrm{HCO}_{3}^{-} & \mathrm{H}_{2} \mathrm{SO}_{3} & \mathrm{NH}_{2}^{*} & \mathrm{NH}_{4}^{*}\end{array}\) \(\mathrm{H}_{3} \mathrm{O}^{+}\) (c) \(\begin{array}{lllll}\mathrm{H}_{2} \mathrm{CO}_{3} & \mathrm{HSO}_{4}^{-} & \mathrm{NH}_{4}^{+} & \mathrm{NH}_{2}^{-} & \mathrm{H}_{2} \mathrm{O}\end{array}\) \(\begin{array}{lllll}\text { (d) } \mathrm{HCO}_{3}^{-} & \mathrm{H}_{2} \mathrm{SO}_{4} & \mathrm{NH}_{2}^{+} & \mathrm{NH}_{2}^{-} & \mathrm{OH}^{-}\end{array}\)

Short answer

p = H₂CO₃, q = SO₄²⁻, r = NH₄⁺, t = H₃O⁺

Step-by-step solution

Identify Conjugate Acid-Base Pairs

Conjugate acid-base pairs differ by one H⁺ ion. The species that donates a proton (H⁺) becomes a conjugate base, and the species that gains a proton becomes a conjugate acid.

Fill in the First Blank

The conjugate acid of HCO₃⁻ is H₂CO₃ since it can gain one proton (H⁺). The given choice for the conjugate acid of HCO₃⁻ is H₂CO₃.

Fill in the Second Blank

The conjugate base of HSO₄⁻ is SO₄²⁻ as it can lose a proton. The choice given for conjugate base of HSO₄⁻ is SO₄²⁻.

Fill in the Third Blank

NH₃ can gain a proton to form its conjugate acid NH₄⁺. The given choice for the conjugate acid of NH₃ is NH₄⁺.

Fill in the Fourth Blank

The conjugate acid of H₂O is H₃O⁺ because H₂O can gain a proton. The given choice for the conjugate acid of H₂O is H₃O⁺.

Key concepts

Chemical Equilibrium

Chemical equilibrium is a concept that plays a central role in acid-base chemistry. It refers to the state where the concentrations of all reactants and products remain constant over time, indicating that the forward and reverse reactions are occurring at the same rate.

For acid-base reactions, equilibrium comes into play when an acid donates a proton to a base and the conjugate pairs form. These pairs, a conjugate acid and a conjugate base, are in equilibrium with each other. Understanding how they interact provides key insights into the strength of acids and bases and their behavior in different solutions.

As a learner, it's crucial to grasp that equilibrium doesn't mean the amounts of reactants and products are equal, but that their rates of formation are in balance. This understanding can be applied to predict the direction of a reaction or to calculate the equilibrium constant, a number that expresses the ratio of the concentration of products to reactants at equilibrium.

Bronsted-Lowry Acids and Bases

The Bronsted-Lowry theory defines an acid as a proton (H⁺) donor and a base as a proton (H⁺) acceptor. This theory goes beyond the concept of acids and bases simply being substances that increase hydrogen or hydroxide ion concentration in water.

The strength of an acid or base can be assessed by how readily it donates or accepts protons. Strong acids, for example, easily donate protons, while weak acids do so less readily. Similarly, strong bases readily accept protons, compared to weak bases. The concept of conjugate acid-base pairs is essential in this theory; when an acid donates a proton, it forms its conjugate base, and vice versa.

Understanding these pairs helps explain the nuances of acid-base reactions beyond the pH scale. It also fosters a deeper comprehension of why certain reactions occur the way they do, making it easier to predict outcomes in chemical processes.

Acid-Base Chemistry

Acid-base chemistry is a fundamental topic within chemical equilibrium that encompasses the study of the properties and reactions of acids and bases. It involves balancing chemical equations, understanding neutralization reactions, and using indicators to determine acidity or basicity.

One crucial component of this topic is the pH scale, which measures how acidic or basic a solution is. The scale ranges from 0 to 14, with 7 being neutral. Solutions with a pH less than 7 are acidic, while those with a pH greater than 7 are basic. The calculation of pH is based on the concentration of hydrogen ions (H⁺) in a solution.

In context with the exercise, filling in the blanks for conjugate acid-base pairs supports the fundamentals of acid-base chemistry. By examining how molecules like HCO₃⁻ and NH₃ react with acids and bases, students can better understand the dynamic nature of these reactions and hone their ability to solve more complex chemical equations.