Key Concepts and Summary
A compound that can donate a proton (a hydrogen ion) to another compound is called a Brønsted-Lowry acid. The compound that accepts the proton is called a Brønsted-Lowry base. The species remaining after a Brønsted-Lowry acid has lost a proton is the conjugate base of the acid. The species formed when a Brønsted-Lowry base gains a proton is the conjugate acid of the base. Thus, an acid-base reaction occurs when a proton is transferred from an acid to a base, with formation of the conjugate base of the reactant acid and formation of the conjugate acid of the reactant base. Amphiprotic species can act as both proton donors and proton acceptors. Water is the most important amphiprotic species. It can form both the hydronium ion, H3O+, and the hydroxide ion, OH− when it undergoes autoionization:
$$2H_2O\;(l)⇌H_3O^+\;(aq)+OH^-\;(aq)$$The ion product of water, $K_w$ is the equilibrium constant for the autoionization reaction:
$$K_w=[H_3O^+][OH^-]=1.0×10^{-14}\quad at\quad 25°C$$Practice Questions
Write equations that show NH3 as both a conjugate acid and a conjugate base.
Solution
One example for NH3 as a conjugate acid:
$$NH_2^-+H^+⟶NH_3$$as a conjugate base:
$$NH_4^+\;(aq)+OH^-\;(aq)⟶NH_3\;(aq)+H_2O\;(l)$$
Write equations that show $H_2PO_4^-$ acting both as an acid and as a base.
Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid:
(a)$H_3O^+$
(b) HCl
(c) NH3
(d) CH3CO2H
(e)$NH_4^+$
(f)$HSO_4^-$
Solution
(a)$H_3O^+\;(aq)⟶H^+\;(aq)+H_2O\;(l)$
(b) $HCL\;(aq)⟶H^+\;(aq)+Cl^-\;(aq)$
(c) $NH_3\;(aq)⟶H^+\;(aq)+NH_2^-\;(aq)$
(d) $CH_3CO_2H\;(aq)⟶H^+\;(aq)+CH_3CO_2^-\;(aq)$
(e) $NH_4^+\;(aq)⟶H^+\;(aq)+NH_3\;(aq)$
(f) $HSO_4^-\;(aq)⟶H^+\;(aq)+SO_4^{2-}\;(aq)$
Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid:
(a) HNO3
(b) $PH_4^+$
(c) H2S
(d) CH3CH2COOH
(e) $H_2PO_4^-$
(f) HS−
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) $H_2PO_4^-$
Solution
(a) $H_2O\;(l)+H^+\;(aq)⟶H_3O^+\;(aq)$
(b) $OH^-\;(aq)+H^+\;(aq)⟶H_2O\;(l)$
(c) $NH_3\;(aq)+H^+\;(aq)⟶NH_4^+\;(aq)$
(d) $CN^-\;(aq)+H^+\;(aq)⟶HCN\;(aq)$
(e) $S^{2-}\;(aq)+H^+\;(aq)⟶HS^-\;(aq)$
(f) $H_2PO_4^-\;(aq)+H^+\;(aq)⟶H_3PO_4\;(aq)$
Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry base:
(a) HS−
(b) $PO_4^{3-}$
(c) $NH_2^-$
(d) C2H5OH
(e) O2−
(f) $H_2PO_4^-$
What is the conjugate acid of each of the following? What is the conjugate base of each?
(a) OH−
(b) H2O
(c) $HCO_3^-$
(d) NH3
(e) $HSO_4^-$
(f) H2O2
(g) HS−
(h) $H_5N_2^+$
Solution
(a) H2O, O2−; (b) H3O+, OH−; (c) H2CO3, $CO_3^{2-}$; (d) $NH_4^+$, $NH_2^-$; (e) H2SO4, $SO_4^{2-}$; (f) $H_3O_2^+$, $HO_2^-$;(g) H2S; S2−; (h) $H_6N_2^{2+}$,H4N2
What is the conjugate acid of each of the following? What is the conjugate base of each?
(a) H2S
(b) $H_2PO_4^-$
(c) PH3
(d) HS−
(e) $HSO_3^-$
(f) $H_3O_2^+$
(g) H4N2
(h) CH3OH
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) $HNO_3+H_2O⟶H_3O^++NO_3^-$
(b) $CN^-+H_2O⟶HCN+OH^-$
(c) $H_2SO_4+Cl^-⟶HCl+HSO_4^-$
(d) $HSO_4^-+OH^-⟶SO_4^{2-}+H_2O$
(e) $O^{2-}+H_2O⟶2OH^-$
(f) $[Cu(H_2O)_3(OH)]^++[Al(H_2O)_6]^{3+}⟶[Cu(H_2O)_4]^{2+}+[Al(H_2O)_5(OH)]^{2+}$
(g) $H_2S+NH_2^-⟶HS^-+NH_3$
Solution
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), $NO_3^-$ (CB); (b) CN−(BB), H2O(BA), HCN(CA), OH−(CB); (c) H2SO4(BA), Cl−(BB), HCl(CA), $HSO_4^-$ (CB); (d) $HSO_4^-$ (BA), OH−(BB), $SO_4^{2-}$ (CB), H2O(CA); (e) O2−(BB), H2O(BA) OH−(CB and CA); (f) [Cu(H2O)3(OH)]+(BB), [Al(H2O)6]3+(BA), [Cu(H2O)4]2+(CA), [Al(H2O)5(OH)]2+(CB); (g) H2S(BA), $NH_2^-$ (BB), HS−(CB), NH3(CA)
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) $NO_2^-+H_2O⟶HNO_2+OH^-$
(b) $HBr+H_2O⟶H_3O^++Br^-$
(c) $HS^-+H_2O⟶H_2S+OH^-$
(d) $H_2PO_4^-+OH^-⟶HPO_4^-+H_2O$
(e) $H_2PO_4^-+HCl⟶H_3PO_4+Cl^-$
(f) $[Fe(H_2O)_5(OH)]^{2+}[Al(H_2O)_6]^{3+}⟶[Fe(H_2O)_6]^{3+}+[Al(H_2O)_5(OH)]^{2+}$
(g) $CH_3OH+H^-⟶CH_3O^-+H_2$
What are amphiprotic species? Illustrate with suitable equations.
Solution
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:
$$H_2O\;(aq)+NH_3\;(aq)⇌NH_4^+\;(aq)+OH^-\;(aq)$$As a base:
$$H_2O\;(aq)+HCl\;(aq)⇌H_3O^+\;(aq)+Cl^-\;(aq)$$
State which of the following species are amphiprotic and write chemical equations illustrating the amphiprotic character of these species:
(a) H2O
(b) $H_2PO_4^-$
(c) S2−
(d) $CO_3^{2-}$
(e) $HSO_4^-$
State which of the following species are amphiprotic and write chemical equations illustrating the amphiprotic character of these species.
(a) NH3
(b) $HPO_4^{2-}$
(c) Br−
(d) $NH_4^+$
(e) $AsO_4^{3-}$
Solution
amphiprotic: (a) $NH_3+H_3O^+⟶NH_4OH+H_2O$, $NH_3+CH_3O^-⟶NH_2^-+CH_3OH$;
(b) $HPO_4^{2-}+OH^-⟶PO_4^{3-}+H_2O$, $HPO_4^{2-}+HClO_4⟶H_2PO_4^-+ClO_4$;
not amphiprotic: (c) Br−; (d) $NH_4^+$; (e) $AsO_4^{3-}$
Is the self-ionization of water endothermic or exothermic? The ionization constant for water ($K_w$) is $2.9×10^{-14}$ at 40 °C and $9.3×10^{-14}$ at 60 °C.