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Chapter 16: Problem 28

(a) Write a chemical equation that illustrates the autoionization of water. (b) Write the expression for the ion-product constant for water, \(K_{w}\). Why is \(\left[\mathrm{H}_{2} \mathrm{O}\right]\) absent from this expression? (c) A solution is described as basic. What does this statement mean?

Short Answer

Expert verified
(a) The chemical equation for the autoionization of water is: \(\mathrm{H_2O} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^+} + \mathrm{OH^-}\). (b) The ion-product constant for water, \(K_w\), is given by: \(K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}]\). The concentration of water, \([\mathrm{H_2O}]\), is not included in the expression because it is a liquid. (c) A solution is basic if its pH is greater than 7, meaning the concentration of hydroxide ions (\(\mathrm{OH^-}\)) is greater than the concentration of hydronium ions (\(\mathrm{H_3O^+}\)).

Step by step solution

01

(a) Autoionization of Water Chemical Equation

In the autoionization of water, two water molecules react to form a hydronium ion (\(\mathrm{H_3O^+}\)) and a hydroxide ion (\(\mathrm{OH^-}\)). The chemical equation for this process is: \[ \mathrm{H_2O} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^+} + \mathrm{OH^-} \]
02

(b) Ion-Product Constant for Water

The ion-product constant for water (\(K_w\)) is the equilibrium constant for the autoionization of water. This constant helps us understand the relationship between the concentration of hydronium ions, \([\mathrm{H_3O^+}]\), and hydroxide ions, \([\mathrm{OH^-}]\), in water. The expression for \(K_w\) is: \[ K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}] \] Since water is a liquid, its concentration is not included in the equilibrium constant expression.
03

(c) Basic Solution Definition

A solution is described as basic if it has a pH greater than 7. This means that the concentration of hydroxide ions (\(\mathrm{OH^-}\)) in the solution is greater than the concentration of hydronium ions (\(\mathrm{H_3O^+}\)). In a basic solution, \(K_w\) remains unchanged; however, the ratio of hydroxide ions to hydronium ions shifts in favor of hydroxide ions.

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Most popular questions from this chapter

Designate the Bronsted-Lowry acid and the Bronsted-Lowry base on the left side of each of the following equations, and also designate the conjugate acid and conjugate base of each on the right side: $$ \text { (a) } \mathrm{NH}_{4}^{+}(a q)+\mathrm{CN}^{-}(a q) \rightleftharpoons \mathrm{HCN}(a q)+\mathrm{NH}_{3}(a q) $$ (b) \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{~N}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons\) $$ \left(\mathrm{CH}_{3}\right)_{3} \mathrm{NH}^{+}(a q)+\mathrm{OH}^{-}(a q) $$ (c) \(\mathrm{HCOOH}(a q)+\mathrm{PO}_{4}^{3-}(a q) \rightleftharpoons\) $$ \mathrm{HCOO}^{-}(a q)+\mathrm{HPO}_{4}^{2-}(a q) $$

The hypochlorite ion, \(\mathrm{ClO}^{-},\) acts as a weak base. (a) Is \(\mathrm{ClO}^{-}\) a stronger or weaker base than hydroxylamine? (b) When ClO \(^{-}\) acts as a base, which atom, Cl or \(\mathrm{O}\), acts as the proton acceptor? (c) Can you use formal charges to rationalize your answer to part (b)?

(a) Write an equation for the reaction in which \(\mathrm{H}_{2} \mathrm{C}_{6} \mathrm{H}_{7} \mathrm{O}_{5}^{-}(a q)\) acts as a base in \(\mathrm{H}_{2} \mathrm{O}(l)\) (b) Write an equation for the reaction in which \(\mathrm{H}_{2} \mathrm{C}_{6} \mathrm{H}_{7} \mathrm{O}_{5}^{-}(a q)\) acts as an acid in \(\mathrm{H}_{2} \mathrm{O}(l) .(\mathrm{c})\) What is the conjugate acid of \(\mathrm{H}_{2} \mathrm{C}_{6} \mathrm{H}_{7} \mathrm{O}_{5}^{-}(a q) ?\) What is its conjugate base?

Calculate \(\left[\mathrm{H}^{+}\right]\) for each of the following solutions, and indicate whether the solution is acidic, basic, or neutral: (a) \(\left[\mathrm{OH}^{-}\right]=0.00045 \mathrm{M} ;\) (b) \(\left[\mathrm{OH}^{-}\right]=8.8 \times 10^{-9} \mathrm{M} ;(\mathrm{c})\) a so- lution in which \(\left[\mathrm{OH}^{-}\right]\) is 100 times greater than \(\left[\mathrm{H}^{+}\right]\).

Codeine \(\left(\mathrm{C}_{18} \mathrm{H}_{21} \mathrm{NO}_{3}\right)\) is a weak organic base. A \(5.0 \times 10^{-3} \mathrm{M}\) solution of codeine has a \(\mathrm{pH}\) of \(9.95 .\) Calculate the value of \(K_{b}\) for this substance. What is the \(\mathrm{p} K_{b}\) for this base?
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