Identify and label the Bronsted-Lowry acid, its conjugate base, the Bronsted- Lowry base, and its conjugate acid in each of the following equations: (a) \(\mathrm{NO}_{2}^{-}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{HNO}_{2}+\mathrm{OH}^{-}\) (b) \(\mathrm{HBr}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{Br}^{-}\) (c) \(\mathrm{HS}^{-}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{H}_{2} \mathrm{S}+\mathrm{OH}^{-}\) (d) \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}+\mathrm{OH}^{-} \longrightarrow \mathrm{HPO}_{4}^{2-}+\mathrm{H}_{2} \mathrm{O}\) (e) \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}+\mathrm{HCl} \longrightarrow \mathrm{H}_{3} \mathrm{PO}_{4}+\mathrm{Cl}^{-}\) (f) \(\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5}(\mathrm{OH})\right]^{2+}+\left[\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+} \longrightarrow\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}+\left[\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5}(\mathrm{OH})\right]^{2+}\) (g) \(\mathrm{CH}_{3} \mathrm{OH}+\mathrm{H}^{-} \longrightarrow \mathrm{CH}_{3} \mathrm{O}^{-}+\mathrm{H}_{2}\)

Step by step solution

01

Identifying Acid-Base Pairs in (a)

For reaction (a) \(\mathrm{NO}_{2}^{-}+\mathrm{H}_{2}O \longrightarrow \mathrm{HNO}_{2}+\mathrm{OH}^{-}\), \mathrm{NO}_{2}^{-} is the Bronsted-Lowry base as it accepts a proton \(\mathrm{H}^{+}\) to become \mathrm{HNO}_{2}, which is its conjugate acid. \mathrm{H}_{2}O donates a proton \(\mathrm{H}^{+}\) to become \mathrm{OH}^{-}, thus it is the Bronsted-Lowry acid and \mathrm{OH}^{-} is its conjugate base.

02

Identifying Acid-Base Pairs in (b)

For reaction (b) \(\mathrm{HBr}+\mathrm{H}_{2}O \longrightarrow \mathrm{H}_{3}O^{+}+\mathrm{Br}^{-}\), \mathrm{HBr} is the Bronsted-Lowry acid as it donates a proton to \mathrm{H}_{2}O, resulting in \mathrm{H}_{3}O^{+}\(\mathrm{H}_{3}O^{+}\) is the conjugate acid of water. \mathrm{Br}^{-} is the conjugate base of \mathrm{HBr}.

03

Identifying Acid-Base Pairs in (c)

For reaction (c) \(\mathrm{HS}^{-}+\mathrm{H}_{2}O \longrightarrow \mathrm{H}_{2}S+\mathrm{OH}^{-}\), \mathrm{HS}^{-} accepts a proton to become \mathrm{H}_{2}S, making it the Bronsted-Lowry base and \mathrm{H}_{2}S its conjugate acid. \mathrm{H}_{2}O donates a proton to become \mathrm{OH}^{-}, which makes water the Bronsted-Lowry acid and \mathrm{OH}^{-} the conjugate base.

04

Identifying Acid-Base Pairs in (d)

For reaction (d) \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}+\mathrm{OH}^{-} \longrightarrow \mathrm{HPO}_{4}^{2-}+\mathrm{H}_{2}O\), \mathrm{H}_{2}\mathrm{PO}_{4}^{-} donates a proton to become \mathrm{HPO}_{4}^{2-}, which is its conjugate base, making it the Bronsted-Lowry acid. \mathrm{OH}^{-} accepts a proton and becomes \mathrm{H}_{2}O, therefore it is the Bronsted-Lowry base and \mathrm{H}_{2}O is the conjugate acid.

05

Identifying Acid-Base Pairs in (e)

For reaction (e) \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}+\mathrm{HCl} \longrightarrow \mathrm{H}_{3}\mathrm{PO}_{4}+\mathrm{Cl}^{-}\), \mathrm{HCl} is the Bronsted-Lowry acid as it donates a proton to \mathrm{H}_{2}\mathrm{PO}_{4}^{-}, resulting in \mathrm{H}_{3}\mathrm{PO}_{4}, which is the conjugate acid, and \mathrm{Cl}^{-} is the conjugate base of \mathrm{HCl}.

06

Identifying Acid-Base Pairs in (f)

For reaction (f) \([\mathrm{Fe}(\mathrm{H}_{2}\mathrm{O})_{5}(\mathrm{OH})]^{2+}+[\mathrm{Al}(\mathrm{H}_{2}\mathrm{O})_{6}]^{3+} \longrightarrow [\mathrm{Fe}(\mathrm{H}_{2}\mathrm{O})_{6}]^{3+}+[\mathrm{Al}(\mathrm{H}_{2}\mathrm{O})_{5}(\mathrm{OH})]^{2+}\), the complex ion containing \(\mathrm{OH}^{-}\) as \(\mathrm{Fe}(\mathrm{H}_{2}O)_{5}(\mathrm{OH})^{2+}\) is acting as a Bronsted-Lowry base, it donates \(\mathrm{OH}^{-}\) and becomes its conjugate acid \([\mathrm{Fe}(\mathrm{H}_{2}O)_{6}]^{3+}\) The complex ion \([\mathrm{Al}(\mathrm{H}_{2}O)_{6}]^{3+}\) is the Bronsted-Lowry acid as it accepts \(\mathrm{OH}^{-}\) and becomes its conjugate base \([\mathrm{Al}(\mathrm{H}_{2}O)_{5}(\mathrm{OH})]^{2+}\).

07

Identifying Acid-Base Pairs in (g)

For reaction (g) \(\mathrm{CH}_{3}\mathrm{OH}+\mathrm{H}^{-} \longrightarrow \mathrm{CH}_{3}\mathrm{O}^{-}+\mathrm{H}_{2}\), \mathrm{CH}_{3}\mathrm{OH} is the Bronsted-Lowry acid as it donates a proton to \mathrm{H}^{-}, resulting in \mathrm{H}_{2}, which is the conjugate acid of the hydride ion. \mathrm{CH}_{3}O^{-} is the conjugate base of \mathrm{CH}_{3}\mathrm{OH}.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Acid-Base Reactions

Acid-base reactions are fundamental processes in chemistry where an acid and a base react with each other. At the heart of these reactions lies the Bronsted-Lowry acid-base theory, which explains acid-base interactions in terms of proton (hydrogen ion, \(H^+\)) transfer.

According to this theory, a Bronsted-Lowry acid is a substance that can donate a proton, whereas a Bronsted-Lowry base is one that can accept a proton. This proton transfer results in the formation of new substances: the conjugate base of the acid and the conjugate acid of the base. For instance, in the reaction \(HBr + H_2O \rightarrow H_3O^+ + Br^-\), \(HBr\) is the acid that donates a proton to \(H_2O\), forming the hydronium ion \(H_3O^+\) and the bromide ion \(Br^-\), which is the conjugate base.

Understanding these reactions requires recognizing the acid-base pairs involved and the direction of proton transfer. This knowledge will enable you to identify the species remaining after the reaction and predict possible shifts in reaction equilibrium under different conditions.

Conjugate Acid-Base Pairs

The concept of conjugate acid-base pairs is integral to understanding how acids and bases behave in a reaction. In any acid-base reaction, an acid donates a proton to a base, forming a conjugate base and a conjugate acid.

For example, if we look at the reaction \(HS^- + H_2O \rightarrow H_2S + OH^-\), \(HS^-\) accepts a proton and forms \(H_2S\); therefore, \(HS^- \text{ and } H_2S\) form a conjugate acid-base pair. Similarly, \(H_2O\) loses a proton to become \(OH^-\), so \(H_2O \text{ and } OH^-\) also create a conjugate acid-base pair.

Conjugate pairs are always found on opposite sides of the reaction equation, and it's important to note that the strength of an acid or base is inversely related to its conjugate partner. Thus, a strong acid will have a weak conjugate base, and a strong base will have a weak conjugate acid. Recognizing these pairs enhances your ability to address acid-base equilibria and predict the outcome of chemical reactions.

Chemical Equilibrium

Chemical equilibrium is a dynamic state in which the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in the concentrations of reactants and products over time. It's essential in understanding how acid-base reactions reach a state of balance.

Consider the reaction where \(H_2PO_4^-\) acts as a Bronsted-Lowry acid: \(H_2PO_4^- + OH^- \rightarrow HPO_4^{2-} + H_2O\). Initially, reactants are converted into products. Over time, however, the products can also react to regenerate the reactants. Equilibrium is achieved when these two processes occur at identical rates.

At equilibrium, the concentrations of acid, base, and their conjugate counterparts remain constant, though they may not be equal. The equilibrium constant (\(K_{eq}\)) provides a quantitative measure of the ratio of product concentrations to reactant concentrations at equilibrium. In acid-base reactions, manipulating the equilibrium by changing concentration, pressure, or temperature can shift the equilibrium position and alter the pH of the system, a concept known as Le Chatelier's principle.