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Chapter 14: Problem 9

The common ion effect for weak acids is to significantly decrease the dissociation of the acid in water. Explain the common ion effect.

Short Answer

Expert verified
The common ion effect is the suppression of weak acid dissociation in water when a solution already contains an ion involved in the equilibrium. According to Le Chatelier's Principle, when a common ion is added, the system shifts the equilibrium to counteract the change, leading to decreased dissociation of the weak acid and a smaller hydrogen ion concentration, ultimately affecting the pH of the solution.

Step by step solution

01

Understand Weak Acid Dissociation

A weak acid is an acid that does not fully dissociate in solution, meaning that not all of its acidic hydrogen atoms are donated to water. The dissociation of a weak acid (HA) can be represented as: \(HA(aq) \rightleftharpoons H^{+}(aq) + A^{-}(aq)\) The double arrow indicates that the reaction is reversible and reaches an equilibrium.
02

Know the Concept of Equilibrium

Equilibrium is reached when the rate of the forward reaction (products forming) is equal to the rate of the reverse reaction (reactants forming). At equilibrium, the concentrations of reactants and products remain constant. For weak acid dissociation, this equilibrium can be represented using the acid-dissociation constant (Ka): \(K_a = \dfrac{[H^{+}][A^{-}]}{[HA]}\) Where [H⁺], [A⁻], and [HA] represent the equilibrium concentrations of hydronium ions, conjugate base anions, and the weak acid, respectively.
03

Understand Common Ion Effect

The common ion effect occurs when a solution already contains an ion that is involved in the dissociation of a weak acid. This effect influences the equilibrium by suppressing the dissociation of the weak acid.
04

Apply Le Chatelier's Principle to the Common Ion Effect

Le Chatelier's Principle states that if a change is made to a system at equilibrium, the system will shift its equilibrium position to counteract that change. When a common ion is added to a weak acid solution, the system tries to reduce the stress of the added ion by shifting the equilibrium position. In the case of weak acid dissociation: \(HA(aq) + H_2O(l) \rightleftharpoons H^{+}(aq) + A^{-}(aq)\) When a common ion (A⁻) is added, the system shifts the equilibrium to the left in order to counteract the increased concentration of A⁻ ions. This leads to decreased dissociation of the weak acid and a smaller hydrogen ion concentration, which causes the pH to rise.
05

Conclusion

The common ion effect is the suppression of weak acid dissociation in water when a solution already contains an ion involved in the equilibrium. The equilibrium shifts in response to the added ion, according to Le Chatelier's Principle, leading to a change in the system's pH.

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Most popular questions from this chapter

A student dissolves 0.0100 mole of an unknown weak base in \(100.0 \mathrm{mL}\) water and titrates the solution with \(0.100 \mathrm{M} \mathrm{HNO}_{3}\) After \(40.0 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{HNO}_{3}\) was added, the \(\mathrm{pH}\) of the resulting solution was \(8.00 .\) Calculate the \(K_{\mathrm{b}}\) value for the weak base.

In the titration of \(50.0 \mathrm{mL}\) of \(1.0 \mathrm{M}\) methylamine, \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) \(\left(K_{\mathrm{b}}=4.4 \times 10^{-4}\right),\) with \(0.50 M\) HCl, calculate the pH under the following conditions. a. after \(50.0 \mathrm{mL}\) of \(0.50 \mathrm{M}\) HCl has been added b. at the stoichiometric point

A certain indicator HIn has a \(\mathrm{p} K_{\mathrm{a}}\) of 3.00 and a color change becomes visible when \(7.00 \%\) of the indicator has been converted to \(\operatorname{In}^{-} .\) At what \(\mathrm{pH}\) is this color change visible?

Calculate the \(\mathrm{pH}\) of each of the following buffered solutions. a. \(0.10 M\) acetic acid/0.25 \(M\) sodium acetate b. \(0.25 M\) acetic acid/0.10 \(M\) sodium acetate c. \(0.080 M\) acetic acid/0.20 \(M\) sodium acetate d. \(0.20 M\) acetic acid/0.080 \(M\) sodium acetate

Acid-base indicators mark the end point of titrations by "magically" turning a different color. Explain the "magic" behind acid-base indicators.
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