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Chapter 16: Problem 65

How does the common ion effect influence solubility equilibria? Use Le Châtelier's principle to explain the decrease in solubility of \(\mathrm{CaCO}_{3}\) in a \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) solution.

Short Answer

Expert verified
The common ion effect reduces the solubility of \(\mathrm{CaCO}_{3}\) in a \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) solution. This is because the \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) introduces more Carbonate ions to the solution, causing the equilibrium to shift to the left in order to reduce this increased concentration, thus resulting in decreased solubility of \(\mathrm{CaCO}_{3}\).

Step by step solution

01

Understand the Solubility Equilibrium of Calcium Carbonate

The solubility equilibrium of Calcium Carbonate (\(\mathrm{CaCO}_{3}\)) in water can be written as: \[\mathrm{CaCO}_{3}(s) \rightleftharpoons \mathrm{Ca}^{2+}(aq) + \mathrm{CO}_{3}^{2-}(aq)\] This equation shows that when Calcium Carbonate dissolves, it breaks down into Calcium ions and Carbonate ions.
02

Understand the Common Ion Effect

The common ion effect states that the solubility of a salt is reduced when it is dissolved in a solution that contains a common ion. In this case, the Sodium Carbonate (\(\mathrm{Na}_{2} \mathrm{CO}_{3}\)) solution provides a common ion, namely the Carbonate ion (\(\mathrm{CO}_{3}^{2-}\)).
03

Apply Le Châtelier's Principle

According to Le Châtelier's Principle, a system in equilibrium will adjust to counteract a change in the conditions. Here, the introduction of more \(\mathrm{CO}_{3}^{2-}\) ions from the \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) increases the concentration of \(\mathrm{CO}_{3}^{2-}\) ions. To counteract this, the equilibrium shifts to the left, reducing the solubility of the \(\mathrm{CaCO}_{3}\).

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Most popular questions from this chapter

The molar solubility of \(\mathrm{Pb}\left(\mathrm{IO}_{3}\right)_{2}\) in a \(0.10 \mathrm{M} \mathrm{NaIO}_{3}\) solution is \(2.4 \times 10^{-11} \mathrm{~mol} / \mathrm{L} .\) What is \(K_{\mathrm{sp}}\) for \(\mathrm{Pb}\left(\mathrm{IO}_{3}\right)_{2} ?\)

A \(25.0-\mathrm{mL}\) of \(0.20 \mathrm{M}\) HF solution is titrated with a \(0.20 M\) NaOH solution. Calculate the volume of \(\mathrm{NaOH}\) solution added when the \(\mathrm{pH}\) of the solution is (a) \(2.85,\) (b) \(3.15,\) (c) \(11.89 .\) Ignore salt hydrolysis.

Determine the \(\mathrm{pH}\) of (a) a \(0.40 \mathrm{MCH}_{3} \mathrm{COOH}\) solution, (b) a solution that is \(0.40 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}\) and \(0.20 \mathrm{M} \mathrm{CH}_{3} \mathrm{COONa} .\)

How many grams of \(\mathrm{CaCO}_{3}\) will dissolve in \(3.0 \times\) \(10^{2} \mathrm{~mL}\) of \(0.050 \mathrm{M} \mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2} ?\)

A student titrates an unknown monoprotic acid with a \(\mathrm{NaOH}\) solution from a buret. After the addition of \(12.35 \mathrm{~mL}\) of \(\mathrm{NaOH},\) the \(\mathrm{pH}\) of the solution read 5.22. The equivalence point is reached at \(24.70 \mathrm{~mL}\) of \(\mathrm{NaOH}\). What is the \(K_{\mathrm{a}}\) of the acid?
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