In equation 1, water behaves as a Brønsted-Lowry acid. equation 1 $\text{H}_2\text{O} + \text{NO}_2^- \rightleftharpoons \text{HNO}_2 + \text{OH}^-$ Identify the two conjugate acid-base pairs in equation 1.
Water also acts as a Brønsted-Lowry acid when it dissolves $\text{CH}_3\text{NH}_2$. Explain the ability of $\text{CH}_3\text{NH}_2$ to act as a base.
Write an equation showing water acting as a base with $\text{CH}_3\text{COOH$.
The ionic product of water, $K_w$, indicates how far water dissociates. $\text{H}_2\text{O(l)} \rightleftharpoons \text{H}^+(\text{aq}) + \text{OH}^-(\text{aq})$ Fig. 2.1 illustrates how $K_w$ changes with temperature. Write an expression for $K_w$.
Use information from Fig. 2.1 to deduce whether the dissociation of water is an exothermic or an endothermic process. Explain your answer.
An aqueous solution has $\text{pH} = 7.00$ at $30\,^{\circ}\text{C}$. Use information from Fig. 2.1 to explain why this solution can be considered to be alkaline at $30\,^{\circ}\text{C}$.
The three physical states of $\text{H}_2\text{O}$ have different standard entropies, $S^{\circ}$, linked to them. Table 2.1 gives these $S^{\circ}$ values. Explain the difference between the $S^{\circ}$ values of $\text{H}_2\text{O(s)}$ and $\text{H}_2\text{O(l)}$.
Explain why the increase in $S^{\circ}$ is much greater when $\text{H}_2\text{O}$ boils than when it melts.
The energy changes for $\text{H}_2\text{O(s)} \rightarrow \text{H}_2\text{O(l)}$ are given below. $\Delta G = 0.00\,\text{kJ mol}^{-1}$ $\Delta H = +6.03\,\text{kJ mol}^{-1}$ Use these data to show that the melting point of $\text{H}_2\text{O(s)}$ is $0\,^{\circ}\text{C}$.
Calculate the standard cell potential, $E^\circ_{\text{cell}}$, of the zinc-air battery.
The zinc-air battery usually runs at $\text{pH}\,11$ and $298\,\text{K}$. The overall cell potential depends on $[\text{OH}^-]$. The Nernst equation shows how the electrode potential at the cathode varies with $[\text{OH}^-]$. $E = 0.40 - \left(\frac{0.059}{z}\right) \log([\text{OH}^-]^2)$ Calculate the electrode potential, $E$, at $\text{pH}\,11$.