Question

Indicate whether each of the following statements is correct or incorrect. (a) Every Bronsted-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 (b) Incorrect (c) Correct (d) Incorrect (e) Correct

Step-by-step solution

(a) Bronsted-Lowry and Lewis Acids

A Bronsted-Lowry acid is a substance that donates a proton (i.e., a hydrogen atom with a positive charge) to a Bronsted-Lowry base. On the other hand, a Lewis acid is a substance that accepts an electron pair. Now, when a Bronsted-Lowry acid donates a proton, it has a vacant electron orbit that can receive an electron pair. Therefore, every Bronsted-Lowry acid is also a Lewis acid. This statement is correct.

(b) Lewis Acids as Bronsted-Lowry Acids

We know that Lewis acids can accept electron pairs, but they don't necessarily have to donate protons. Thus not every Lewis acid is a Bronsted-Lowry acid. This statement is incorrect.

(c) Conjugate Acids of Weak and Strong Bases

Conjugate acids of weak bases are more acidic than the conjugate acids of strong bases. This is because strong bases have weaker conjugate acids, whereas weak bases have stronger conjugate acids. Therefore, conjugate acids of weak bases produce more acidic solutions than conjugate acids of strong bases. This statement is correct.

(d) \(\mathrm{K}^{+}\) Ion in Water

The \(\mathrm{K}^{+}\) ion itself does not cause hydrating water molecules to become more acidic. Rather, the presence of \(\mathrm{K}^{+}\) ions in water causes the water to shift equilibrium from the hydronium ion side to the hydroxide ion side, which results in a more neutral or slightly basic solution. Therefore, this statement is incorrect.

(e) Percent Ionization of Weak Acids

Percent ionization of a weak acid is the percentage of acid molecules that donate protons in an aqueous solution. When the concentration of a weak acid decreases, the ratio of dissociated ions to undissociated molecules increases, so the percent ionization also increases. This statement is correct.

Key concepts

Acid-Base Reactions

Understanding the basics of acid-base chemistry is crucial for grasping how substances interact in solutions. An acid-base reaction involves the transfer of a proton (hydrogen ion, H+) from an acid to a base. According to the Bronsted-Lowry theory, an acid is defined as a proton donor, while a base is a proton acceptor. This definition enlarges the number of potential acids and bases compared to the earlier Arrhenius definition, which was limited to aqueous solutions.

In these reactions, the formation of a water molecule from H+ (the proton) and OH- (hydroxide ion) is a classic example of an acid reacting with a base. This process is essential for understanding the chemical behavior of substances in biological, environmental, and industrial contexts.

Conjugate Acid-Base Pairs

A conjugate acid-base pair consists of two species that transform into each other by the gain or loss of a proton. When an acid donates a proton, it becomes its conjugate base; conversely, when a base accepts a proton, it becomes its conjugate acid. This concept is pivotal because it illustrates the reversible nature of acid-base reactions.

For instance, when ammonia (NH₃), a weak base, accepts a proton (H+), it forms its conjugate acid, the ammonium ion (NH₄+). The strength of a conjugate acid is inversely related to the strength of its base. This relationship explains why strong acids have weak conjugate bases and vice versa. Understanding conjugate pairs is essential to predicting the direction and extent of acid-base reactions in chemistry.

Percent Ionization

Percent ionization is a measure of the strength of an acid, reflecting the proportion of acid that dissociates into ions in an aqueous solution. For strong acids, this value is near 100%, meaning almost all the acid molecules dissociate. For weak acids, the percent ionization is much lower, indicating that fewer molecules dissociate.

The percent ionization increases as the acid dilutes, which can be explained by the principle of chemical equilibrium and Le Chatelier's Principle—when the concentration of an acid decreases, the equilibrium of the dissociation reaction shifts to produce more ions, increasing the ionization percentage. This concept is crucial in understanding the acidity and pH of solutions in chemical systems.

Proton Donors and Acceptors

In the Bronsted-Lowry theory, acids and bases are identified by their ability to donate or accept protons. Proton donors are Bronsted-Lowry acids because they release H+ ions into the solution. In contrast, proton acceptors are Bronsted-Lowry bases because they trap these H+ ions.

For a molecule to function as a Bronsted-Lowry acid, it must have a hydrogen atom that can easily be released as an H+ ion. Molecules with an atom that has a lone pair of electrons can be proton acceptors because those electrons can be used to bond with an H+ ion. Understanding the roles of proton donors and acceptors is essential for predicting the course of acid-base reactions and their consequences in various scientific fields.

Chemistry Equilibrium

Chemistry equilibrium is a state where the forward and reverse reactions occur at the same rate, resulting in no net change in the concentration of reactants and products over time. In the context of acid-base reactions, this concept is related to the dynamic balance between unreacted species and the ions produced by their dissociation.

The equilibrium can be represented by an equilibrium constant, K, which quantifies the ratio of the concentrations of products to reactants at equilibrium. Disturbances to this balance, such as changes in concentration, pressure, temperature, or the presence of a catalyst, will shift the equilibrium position according to Le Chatelier's Principle, adjusting the system to minimize the impact of the change. Understanding chemical equilibrium is fundamental in predicting how systems respond to external changes and for calculating the concentrations of substances in chemical reactions.