Question

Indicate whether each of the following statements is correct or incorrect. (a) Every Brønsted-Lowry acid is also a Lewis acid. (b) Every Lewis acid is also a Bronsted-Lowry acid. (c) Conjugate acids of weak bases produce more acidic solutions than conjugate acids of strong bases. (d) \(\mathrm{K}^{+}\)ion is acidic in water because it causes hydrating water molecules to become more acidic. (e) The percent ionization of a weak acid in water increases as the concentration of acid decreases.

Short answer

(a) Correct. Every Brønsted-Lowry acid can be a Lewis acid because when a Brønsted-Lowry acid donates a proton (H⁺), it is also accepting a pair of electrons. (b) Incorrect. Not every Lewis acid is a Brønsted-Lowry acid. Lewis acids accept electron pairs and can be any molecule or ion that has an empty orbital available for electron sharing. However, Brønsted-Lowry acids specifically donate a proton (H⁺). (c) Correct. Conjugate acids of weak bases produce more acidic solutions than conjugate acids of strong bases because weak bases have a stronger affinity for protons, making their conjugate acids more likely to donate protons back to the solution. (d) Incorrect. The \(\mathrm{K}^{+}\) ion is not acidic in water. In fact, it is the conjugate acid of a very strong base (\(\mathrm{KOH}\)), which makes it a neutral ion in water. It does not cause hydrating water molecules to become more acidic. (e) Correct. The percent ionization of a weak acid in water increases as the concentration of the acid decreases. As the concentration of the acid decreases, the equilibrium between protonated and deprotonated species shifts towards ionization, resulting in a higher percentage of ionized molecules in the solution.

Step-by-step solution

Statement (a) Analysis

To analyze statement (a), we need to know the definitions of a Brønsted-Lowry acid and a Lewis acid. A Brønsted-Lowry acid is a substance that donates a proton (H⁺) when it dissociates in a solution. A Lewis acid, on the other hand, is a substance that can accept an electron pair. So, let's check if a Brønsted-Lowry acid can also be a Lewis acid.

Statement (a) Solution

Yes, statement (a) is correct. Every Brønsted-Lowry acid can be a Lewis acid because when a Brønsted-Lowry acid donates a proton (H⁺), it is also accepting a pair of electrons.

Statement (b) Analysis

Now we need to analyze statement (b). Is every Lewis acid also a Brønsted-Lowry acid?

Statement (b) Solution

Statement (b) is incorrect. Not every Lewis acid is a Brønsted-Lowry acid. Lewis acids accept electron pairs and can be any molecule or ion that has an empty orbital available for electron sharing. However, Brønsted-Lowry acids specifically donate a proton (H⁺).

Statement (c) Analysis

For statement (c), we need to understand what conjugate acids are and the difference between weak bases and strong bases.

Statement (c) Solution

Statement (c) is correct. Conjugate acids of weak bases produce more acidic solutions than conjugate acids of strong bases because weak bases have a stronger affinity for protons, making their conjugate acids more likely to donate protons back to the solution.

Statement (d) Analysis

In statement (d), we analyze the behavior of K⁺ ions in water and their effect on water molecules.

Statement (d) Solution

Statement (d) is incorrect. The \(\mathrm{K}^{+}\) ion is not acidic in water. In fact, it is the conjugate acid of a very strong base (\(\mathrm{KOH}\)), which makes it a neutral ion in water. It does not cause hydrating water molecules to become more acidic.

Statement (e) Analysis

For statement (e), let's analyze the relationship between the percent ionization of a weak acid and the concentration of the acid in water.

Statement (e) Solution

Statement (e) is correct. The percent ionization of a weak acid in water increases as the concentration of the acid decreases. As the concentration of the acid decreases, the equilibrium between protonated and deprotonated species shifts towards ionization, resulting in a higher percentage of ionized molecules in the solution.

Key concepts

Acid-Base Theories

To understand the behavior of substances in various chemical contexts, it is crucial to grasp the fundamental theories behind acid-base chemistry. The two primary acid-base theories cited in many chemistry problems are the Brønsted-Lowry and Lewis theories.

The Brønsted-Lowry theory defines an acid as a proton (hydrogen ion, H⁺) donor and a base as a proton acceptor. This proton transfer concept leads to the understanding of conjugate acid-base pairs: when an acid donates a proton, it becomes a conjugate base; conversely, when a base accepts a proton, it becomes a conjugate acid.

The Lewis theory goes beyond just protons and focuses on the transfer of electron pairs. A Lewis acid is defined as an electron pair acceptor, while a Lewis base is an electron pair donor. The Lewis definition is broader as it is not limited to hydrogen-containing species and applies to a wider range of reactions, such as complexation and organic chemistry contexts.

Both theories are integral to predicting the outcomes of acid-base reactions and are often used in tandem to provide a comprehensive understanding of chemical processes.

Conjugate Acids and Bases

Conjugate acids and bases are a pair of substances that transform into each other by gaining or losing a proton according to the Brønsted-Lowry theory. They are directly related to each other; for example, when ammonia (NH₃) accepts a proton, it becomes ammonium (NH₄⁺) - its conjugate acid.

Strength of Conjugates

The strength of a conjugate acid or base is inversely related to its counterpart. A strong acid will have a weak conjugate base since the strong acid donates its proton readily, leaving behind a base that has little affinity for a proton. Conversely, a weak acid's conjugate base will be comparatively stronger, as the acid doesn't lose its proton as easily, and thus the base has a higher tendency to regain it.

Understanding how conjugate pairs work is essential for predicting the behavior of substances in solution and for various analytical purposes including pH calculation, buffering systems, and titration processes.

Percent Ionization

Percent ionization is a measure of the strength of an acid or base in solution. It can be thought of as the percentage of a substance that has ionized (formed ions) compared to the total amount of the substance initially present.

The higher the percent ionization, the stronger the acid or base. For weak acids, an important trait to remember is that as we dilute the solution (decrease the concentration), the percent ionization increases. This behavior illustrates the dynamic nature of acid-base equilibrium; even a weak acid can show significant ionization under the right conditions, influencing the pH of the solution and the reactivity of the acid in chemical reactions.

This concept is essential not only in understanding acid-base chemistry but also in applications such as drug formulation, where ionization can greatly influence the medication's absorption and distribution within the body.

Acid-Base Equilibrium

Acid-base equilibrium is the state in which the rates of the forward and reverse reactions are equal, resulting in no overall change in the concentrations of the reactants and products. This concept is central to acid-base chemistry since it applies to how acids and bases behave in water.

According to the Le Châtelier's principle, any change in concentration, pressure, or temperature will shift the equilibrium position to counteract the imposed change. This principle explains why decreasing the concentration of a weak acid in water increases its percent ionization; the equilibrium shifts to restore balance by producing more ions. Calculating the equilibrium constant (Ka for acids, Kb for bases) can provide insights into the strength of acids and bases in solution.

Grasping acid-base equilibrium helps in understanding buffer systems, solubility, and biochemical systems where maintaining a certain pH is vital for proper functioning, such as in blood pH regulation.