Unit 4 Acid-Base and Solubility Equilibria
4.1 Exercises
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Section 4.1 Exercises
- Write equations that show NH3 as both a conjugate acid and a conjugate base.
- Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid:
(a) H3O+
(b) HCl
(c) CH3NH4+
(d) CH3COOH
(e) NH4+
(f) HSO4− - Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry base:
(a) H2O
(b) OH−
(c) NH3
(d) CN−
(e) S2−
(f) H2PO4− - What is the conjugate acid of each of the following? What is the conjugate base of each?
(a) OH−
(b) H2O
(c) HCO3−
(d) NH3
(e) HSO4−
(f) H2O2
(g) HS−
(h) H5N2+ - Identify and label the Brønsted-Lowry acid, its conjugate base, the Brønsted-Lowry base, and its conjugate acid in each of the following equations:
(a) [latex]\text{HNO}_3\;+\;\text{H}_2\text{O}\;{\longrightarrow}\;\text{H}_3\text{O}^{+}\;+\;\text{NO}_3^{\;\;-}[/latex]
(b) [latex]\text{CN}^{-}\;+\;\text{H}_2\text{O}\;{\longrightarrow}\;\text{HCN}\;+\;\text{OH}^{-}[/latex]
(c) [latex]\text{H}_2\text{SO}_4\;+\;\text{Cl}^{-}\;{\longrightarrow}\;\text{HCl}\;+\;\text{HSO}_4^{\;\;-}[/latex]
(d) [latex]\text{HSO}_4^{\;\;-}\;+\;\text{OH}^{-}\;{\longrightarrow}\;\text{SO}_4^{\;\;2-}\;+\;\text{H}_2\text{O}[/latex]
(e) [latex]\text{O}^{2-}\;+\;\text{H}_2\text{O}\;{\longrightarrow}\;2\text{OH}^{-}[/latex]
(f) [latex]\text{H}_2\text{S}\;+\;\text{NH}_2^{\;\;-}\;{\longrightarrow}\;\text{HS}^{-}\;+\;\text{NH}_3[/latex] - What are amphiprotic species? Illustrate with suitable equations.
- State which of the following species are amphiprotic and write chemical equations illustrating the amphiprotic character of these species.
(a) NH3
(b) HPO42−
(c) Br−
(d) NH4+
(e) AsO43-
Solutions
- One example for NH3 as a conjugate acid: [latex]\text{NH}_2^{\;\;-}\;+\;\text{H}^{+}\;{\longrightarrow}\;\text{NH}_3[/latex]; as a conjugate base: [latex]\text{NH}_4^{\;\;+}(aq)\;+\;\text{OH}^{-}(aq)\;{\longrightarrow}\;\text{NH}_3(aq)\;+\;\text{H}_2\text{O}(l)[/latex]
- (a) [latex]\text{H}_3\text{O}^{+}(aq)\;+\text{ H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_2\text{O}(l)\;+\;\text{H}_3\text{O}^+(aq)[/latex];
(b) [latex]\text{HCl}(g)\;+\text{ H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_3\text{O}^{+}(aq)\;+\;\text{Cl}^{-}(aq)[/latex];
(c) [latex]\text{CH}_3\text{NH}_4^+(aq)\;+\text{ H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_3\text{O}^{+}(aq)\;+\;\text{CH}_3\text{NH}_3(aq)[/latex];
(d) [latex]\text{CH}_3\text{COO}\text{H}(aq)\;+\text{ H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_3\text{O}^{+}(aq)\;+\;\text{CH}_3\text{COO}^{\;\;-}(aq)[/latex];
(e) [latex]\text{NH}_4^{\;\;+}(aq)\;+\text{ H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_3\text{O}^{+}(aq)\;+\;\text{NH}_3(aq)[/latex];
(f) [latex]\text{HSO}_4^{\;\;-}(aq)\;+\text{ H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_3\text{O}^{+}(aq)\;+\;\text{SO}_4^{\;\;2-}(aq)[/latex]
Alternatively, rather than donating a proton to water, the given acids could donate a proton to a generic base B to form HB+. Outside the context of acid-base chemistry, such as in Unit 6, we tend to use H+ for convenience, in place of H3O+. - (a) [latex]\text{H}_2\text{O}(l)\;+\;\text{H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_3\text{O}^{+}(aq)\;+\;\text{OH}^-(aq)[/latex];
(b) [latex]\text{OH}^{-}(aq)\;+\;\text{H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_2\text{O}(l)\;+\;\text{OH}^-(aq)[/latex];
(c) [latex]\text{NH}_3(aq)\;+\;\text{H}_2\text{O}(l)\;{\longrightarrow}\;\text{NH}_4^{\;\;+}(aq)\;+\;\text{OH}^-(aq)[/latex];
(d) [latex]\text{CN}^{-}(aq)\;+\;\text{H}_2\text{O}(l)\;{\longrightarrow}\;\text{HCN}(aq)\;+\;\text{OH}^-(aq)[/latex];
(e) [latex]\text{S}^{2-}(aq)\;+\;\text{H}_2\text{O}(l)\;{\longrightarrow}\;\text{HS}^{-}(aq)\;+\;\text{OH}^-(aq)[/latex];
(f) [latex]\text{H}_2\text{PO}_4^{\;\;-}(aq)\;+\;\text{H}_2\text{O}(l)\;{\longrightarrow}\;\text{H}_3\text{PO}_4(aq)\;+\;\text{OH}^-(aq)[/latex]
Alternatively, rather than accepting a proton from water, the given bases could accept a proton from a generic acid HA to form A−. Outside the context of acid-base chemistry, such as in Unit 6, we tend to use H+ for convenience, in place of H3O+. - (a) H2O, O2−
(b) H3O+, OH−
(c) H2CO3, [latex]\text{CO}_3^{\;\;2-}[/latex]
(d) [latex]\text{NH}_4^{\;\;+}[/latex], [latex]\text{NH}_2^{\;\;-}[/latex]
(e) H2SO4, [latex]\text{SO}_4^{\;\;2-}[/latex]
(f) [latex]\text{H}_3\text{O}_2^{\;\;+}[/latex], [latex]\text{HO}_2^{\;\;-}[/latex]
(g) H2S; S2−
(h) [latex]\text{H}_6\text{N}_2^{\;\;2+}[/latex], H4N2 - The labels are Brønsted-Lowry acid = BA; its conjugate base = CB; Brønsted-Lowry base = BB; its conjugate acid = CA.
(a) HNO3(BA), H2O(BB), H3O+(CA), [latex]\text{NO}_3^{\;\;-}(\text{CB})[/latex]
(b) CN−(BB), H2O(BA), HCN(CA), OH−(CB)
(c) H2SO4(BA), Cl−(BB), HCl(CA), [latex]\text{HSO}_4^{\;\;-}(\text{CB})[/latex]
(d) [latex]\text{HSO}_4^{\;\;-}(\text{BA})[/latex], OH−(BB), [latex]\text{SO}_4^{\;\;2-}(\text{CB})[/latex], H2O(CA)
(e) O2−(BB), H2O(BA) OH−(CB and CA)
(f) H2S(BA), [latex]\text{NH}_2^{\;\;-}(\text{BB})[/latex], HS−(CB), NH3(CA) - Amphiprotic species may either gain or lose a proton in a chemical reaction, thus acting as a base or an acid. An example is H2O. As an acid:
[latex]\text{H}_2\text{O}(aq)\;+\;\text{NH}_3(aq)\;{\rightleftharpoons}\;\text{NH}_4^{\;\;+}(aq)\;+\;\text{OH}^{-}(aq)[/latex]. As a base: [latex]\text{H}_2\text{O}(aq)\;+\;\text{HCl}(aq)\;{\rightleftharpoons}\;\text{H}_3\text{O}^{+}(aq)\;+\;\text{Cl}^{-}(aq)[/latex] - amphiprotic:
(a) [latex]\text{NH}_3\;+\;\text{H}_3\text{O}^{+}\;{\longrightarrow}\;\text{NH}_4^+\;+\;\text{H}_2\text{O}\text{, }\text{NH}_3\;+\;\text{OCH}_3^{\;\;-}\;{\longrightarrow}\;\text{NH}_2^{\;\;-}\;+\;\text{CH}_3\text{OH}[/latex]
(b) [latex]\text{HPO}_4^{\;\;2-}\;+\;\text{OH}^{-}\;{\longrightarrow}\;\text{PO}_4^{\;\;3-}\;+\;\text{H}_2\text{O}\text{, }\text{HPO}_4^{\;\;2-}\;+\;\text{HClO}_4\;{\longrightarrow}\;\text{H}_2\text{PO}_4^{\;\;-}\;+\;\text{ClO}_4^{\;\;-}[/latex]
not amphiprotic:
(c) Br−; (d) NH4+; (e) AsO43−
