Chapter 16

When ozone is bubbled through an aqueous solution of \(\mathrm{Fe}^{2+}\) ions, the ions are oxidized to \(\mathrm{Fe}^{3+}\) ions. How does the oxidation process affect the \(\mathrm{pH}\) of the solution?

As an aqueous solution of KOH is slowly added to a stirred solution of \(\mathrm{AlCl}_{3},\) the mixture becomes cloudy but then clears when more KOH is added. a. Explain the chemical changes responsible for the changes in the appearance of the mixture. b. Would you expect to observe the same changes if KOH were added to a solution of \(\mathrm{FeCl}_{3} ?\) Explain why or why not.

Chromium(III) hydroxide is amphiprotic. Write chemical equations showing how an aqueous suspension of this compound reacts to the addition of a strong acid and a strong base.

Zinc hydroxide is amphiprotic. Write chemical equations showing how an aqueous suspension of this compound reacts to the addition of a strong acid and a strong base.

To remove impurities such as calcium and magnesium carbonates and iron(III) oxides from aluminum ore (which is mostly \(\mathrm{Al}_{2} \mathrm{O}_{3}\) ), the ore is treated with a strongly basic solution. In this treatment, \(\mathrm{Al}^{3+}\) dissolves but the other metal ions do not. Why?

What is the pH of \(0.25 M \mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3} ?\)

What is the pH of \(0.50 M \mathrm{CrCl}_{3} ?\)

What is the pH of \(1.00 M \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2} ?\)

What is the difference between molar solubility and solubility product?

Give an example of how the common-ion effect limits the dissolution of a sparingly soluble ionic compound.

Which cation will precipitate first as a carbonate mineral from an equimolar solution of \(\mathrm{Mg}^{2+}, \mathrm{Ca}^{2+},\) and \(\mathrm{Sr}^{2+} ?\)

If the solubility of a compound increases with increasing temperature, does \(K_{\mathrm{sp}}\) increase or decrease?

The \(K_{\text {sp }}\) of strontium sulfate increases from \(2.8 \times 10^{-7}\) at \(37^{\circ} \mathrm{C}\) to \(3.8 \times 10^{-7}\) at \(77^{\circ} \mathrm{C} .\) Is the dissolution of strontium sulfate endothermic or exothermic?

Tooth enamel is composed of a mineral known as hydroxyapatite with the formula \(\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3}(\mathrm{OH}) .\) Explain why tooth enamel can be eroded by acidic substances released by bacteria growing in the mouth.

Fluoride ions in drinking water and toothpaste convert hydroxyapatite in tooth enamel into fluorapatite: $$\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3}(\mathrm{OH})(s)+\mathrm{F}^{-}(a q) \rightleftharpoons \mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{F}(s)+\mathrm{OH}^{-}(a q)$$ Why is fluorapatite less susceptible than hydroxyapatite to erosion by acids?

At a particular temperature the \(\left[\mathrm{Ba}^{2+}\right]\) in a saturated solution of barium sulfate is \(1.04 \times 10^{-5} M .\) Starting with this information, calculate the \(K_{\mathrm{sp}}\) value of barium sulfate at this temperature.

If only \(0.160 \mathrm{g}\) of \(\mathrm{Ca}(\mathrm{OH})_{2}\) dissolves in \(0.100 \mathrm{L}\) of water, what is the \(K_{\mathrm{sp}}\) value for calcium hydroxide at that temperature?

What are the equilibrium concentrations of \(\mathrm{Pb}^{2+}\) and \(\mathrm{F}^{-}\) in a saturated solution of lead fluoride at \(25^{\circ} \mathrm{C} ?\)

What is the solubility of calcite \(\left(\mathrm{CaCO}_{3}\right)\) in grams per milliliter at a temperature at which its \(K_{\mathrm{sp}}=9.9 \times 10^{-9} ?\)

Suppose you have \(100 \mathrm{mL}\) of each of the following solutions. In which will the most \(\mathrm{CaCO}_{3}\) dissolve? (a) \(0.1 M \mathrm{NaCl} ;(\mathrm{b}) 0.1 \mathrm{M} \mathrm{Na}_{2} \mathrm{CO}_{3} ;(\mathrm{c}) 0.1 \mathrm{M} \mathrm{NaOH}\) (d) \(0.1 M\) HCl

In which of the following solutions will \(\mathrm{CaF}_{2}\) be most soluble? (a) \(0.010 M \mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2} ;\) (b) \(0.01 M \mathrm{NaF}\) (c) \(0.001 M\) NaF; (d) \(0.10 M \mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}\)

Some scientists have proposed adding iron(III) compounds to large expanses of the open ocean to promote the growth of phytoplankton that would in turn remove \(\mathrm{CO}_{2}\) from the atmosphere through photosynthesis. Assuming the average pH of open ocean water is \(8.13,\) what is the maximum value of \(\left[\mathrm{Fe}^{3+}\right]\) in seawater if the \(K_{\mathrm{sp}}\) value of \(\mathrm{Fe}(\mathrm{OH})_{3}\) is \(1.1 \times 10^{-36} ?\)

Will calcium fluoride precipitate when \(125 \mathrm{mL}\) of \(0.375 \mathrm{M}\) \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}\) is added to \(245 \mathrm{mL}\) of \(0.255 M \mathrm{NaF}\) at \(25^{\circ} \mathrm{C} ?\)

Will lead(II) chloride precipitate if \(185 \mathrm{mL}\) of \(0.025 M\) sodium chloride is added to \(235 \mathrm{mL}\) of \(0.165 M\) lead (II) perchlorate at \(25^{\circ} \mathrm{C} ?\)

Solution \(A\) is \(0.0200 M\) in \(A g^{+}\) ions and \(P b^{2+}\) ions. You have access to two other solutions: (B) \(0.250 M\) NaCl and (C) \(0.250 M\) NaBr. a. Which solution, B or \(\mathrm{C}\), would be the better one to add to Solution \(A\) to separate \(A g^{+}\) ions from \(P b^{2+}\) by selective precipitation? b. Using the solution you selected in part (a), is the separation of the two ions complete?

Between 1993 and \(1995,\) sodium phosphate was added to Seathwaite Tarn in the English Lake District to increase its pH. Explain why addition of this compound increased pH.

A cook dissolves a teaspoon of baking soda (NaHCO \(_{3}\) ) in a cup of water and then discovers that the recipe calls for a tablespoon, not a teaspoon. So the cook adds two more teaspoons of baking soda to make up the difference. Does the additional baking soda change the pH of the solution? Explain why or why not.

Antacids contain a variety of bases such as \(\mathrm{NaHCO}_{3}, \mathrm{MgCO}_{3}, \mathrm{CaCO}_{3},\) and \(\mathrm{Mg}(\mathrm{OH})_{2} .\) Only NaHCO \(_{3}\) has appreciable solubility in water. a. Write a net ionic equation for the reaction of each base with aqueous HCl. b. Explain how substances insoluble in water can act as effective antacids.

When silver oxide dissolves in water, the following reaction occurs: $$ \mathrm{Ag}_{2} \mathrm{O}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow 2 \mathrm{Ag}^{+}(a q)+2 \mathrm{OH}^{-}(a q) $$ If a saturated aqueous solution of silver oxide is \(1.6 \times 10^{-4} M\) in hydroxide ion, what is the \(K_{\mathrm{sp}}\) of silver oxide?

Why does adding \(\mathrm{CaCl}_{2}\) to a \(\mathrm{HPO}_{4}^{2-} / \mathrm{PO}_{4}^{3-}\) buffer increase the ratio of \(\mathrm{HPO}_{4}^{2-}\) ions to \(\mathrm{PO}_{4}^{3-}\) ions?

Some climate models predict the pH of the oceans will decrease by as much as 0.77 pH units as a result of increases in atmospheric carbon dioxide. a. Explain, by using the appropriate chemical reactions and equilibria, how an increase in atmospheric \(\mathrm{CO}_{2}\) could produce a decrease in oceanic \(\mathrm{pH}\). b. How much more acidic (in terms of \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) ) would the oceans be if their \(\mathrm{pH}\) dropped this much? c. Oceanographers are concerned about how a drop in oceanic pH would affect the survival of oysters. Why?