Chapter 5

Following are some laboratory methods occasionally used for the preparation of small quantities of chemicals. Write a balanced equation for each. (a) preparation of ${\mathrm{H}}_2\mathrm{S}(\mathrm{g}):\mathrm{HCl}(\mathrm{aq})$ is heated with $\mathrm{FeS}(\mathrm{s})$ (b) preparation of ${\mathrm{Cl}}_2(\mathrm{g}):\mathrm{HCl}(\mathrm{aq})$ is heated with ${\mathrm{MnO}}_2(\mathrm{s});{\mathrm{MnCl}}_2(\mathrm{aq})$ and ${\mathrm{H}}_2\mathrm{O}(1)$ are other products (c) preparation of ${\mathrm{N}}_2:{\mathrm{Br}}_2$ and ${\mathrm{NH}}_3$ react in aqueous solution; ${\mathrm{NH}}_4\mathrm{Br}$ is another product (d) preparation of chlorous acid: an aqueous suspension of solid barium chlorite is treated with dilute ${\mathrm{H}}_2{\mathrm{SO}}_4(\mathrm{aq})$

When concentrated ${\mathrm{CaCl}}_2(\mathrm{aq})$ is added to ${\mathrm{Na}}_2{\mathrm{HPO}}_4(\mathrm{aq}),$ a white precipitate forms that is $38.7\mathrm{\%}$ Ca by mass. Write a net ionic equation representing the probable reaction that occurs.

Sodium hydroxide used to make standard $\mathrm{NaOH}(\mathrm{aq})$ solutions for acid-base titrations is invariably contaminated with some sodium carbonate. (a) Explain why, except in the most precise work, the presence of this sodium carbonate generally does not seriously affect the results obtained, for example, when $\mathrm{NaOH}(\mathrm{aq})$ is used to titrate HCl(aq). (b) Conversely, show that if ${\mathrm{Na}}_2{\mathrm{CO}}_3$ comprises more than $1\mathrm{\%}$ to $2\mathrm{\%}$ of the solute in NaOH(aq), the titration results are affected.

A $110.520\mathrm{g}$ sample of mineral water is analyzed for its magnesium content. The ${\mathrm{Mg}}^{2+}$ in the sample is first precipitated as ${\mathrm{MgNH}}_4{\mathrm{PO}}_4,$ and this precipitate is then converted to ${\mathrm{Mg}}_2{\mathrm{P}}_2{\mathrm{O}}_7,$ which is found to weigh 0.0549 g. Express the quantity of magnesium in the sample in parts per million (that is, in grams of $\mathrm{Mg}$ per million grams of ${\mathrm{H}}_2\mathrm{O}$ ).

What volume of $0.248\mathrm{M}{\mathrm{CaCl}}_2$ must be added to $335\mathrm{mL}$ of $0.186\mathrm{M}\mathrm{KCl}$ to produce a solution with a concentration of $0.250\mathrm{M}{\mathrm{Cl}}^{-2}$ Assume that the solution volumes are additive.

An unknown white solid consists of two compounds, each containing a different cation. As suggested in the illustration, the unknown is partially soluble in water. The solution is treated with $\mathrm{NaOH}(\mathrm{aq})$ and yields a white precipitate. The part of the original solid that is insoluble in water dissolves in $\mathrm{HCl}(\mathrm{aq})$ with the evolution of a gas. The resulting solution is then treated with ${\left({\mathrm{NH}}_4\right)}_2{\mathrm{SO}}_4(\mathrm{aq})$ and yields a white precipitate. (a) Is it possible that any of the cations $M{g}^{2+},C{u}^{2+}$ ${\mathrm{Ba}}^{2+},{\mathrm{Na}}^{+},$ or ${\mathrm{NH}}_4^{+}$ were present in the original unknown? Explain your reasoning. (b) What compounds could be in the unknown mixture (that is, what anions might be present)?

Balance these equations for reactions in acidic solution. (a) $\mathrm{IBr}+{\mathrm{BrO}}_3^{-}+{\mathrm{H}}^{+}⟶{\mathrm{IO}}_3^{-}+{\mathrm{Br}}^{-}+{\mathrm{H}}_2\mathrm{O}$ (b) ${\mathrm{C}}_2{\mathrm{H}}_5{\mathrm{NO}}_3+\mathrm{Sn}⟶$ ${\mathrm{NH}}_2\mathrm{OH}+{\mathrm{CH}}_3{\mathrm{CH}}_2\mathrm{OH}+{\mathrm{Sn}}^{2+}$ (c) ${\mathrm{As}}_2{\mathrm{S}}_3+{\mathrm{NO}}_3^{-}⟶{\mathrm{H}}_3{\mathrm{AsO}}_4+\mathrm{S}+\mathrm{NO}$ (d) ${\mathrm{H}}_5{\mathrm{IO}}_6+{\mathrm{I}}_2⟶{\mathrm{IO}}_3^{-}+{\mathrm{H}}^{+}+{\mathrm{H}}_2\mathrm{O}$ (e) ${\mathrm{S}}_2{\mathrm{F}}_2+{\mathrm{H}}_2\mathrm{O}⟶{\mathrm{S}}_8+{\mathrm{H}}_2{\mathrm{S}}_4{\mathrm{O}}_6+\mathrm{HF}$

A method of producing phosphine, ${\mathrm{PH}}_3$, from elemental phosphorus, $P_4$, involves heating the $P_4$ with ${\mathrm{H}}_2\mathrm{O}.$ An additional product is phosphoric acid, ${\mathrm{H}}_3{\mathrm{PO}}_4$ Write a balanced equation for this reaction.

Iron (Fe) is obtained from rock that is extracted from open pit mines and then crushed. The process used to obtain the pure metal from the crushed rock produces solid waste, called tailings, which are stored in disposal areas near the mines. The tailings pose a serious environmental risk because they contain sulfides, such as pyrite ( ${\mathrm{FeS}}_2$ ), which oxidize in air to produce metal ions and ${\mathrm{H}}^{+}$ ions that can enter into surface water or ground water. The oxidation of ${\mathrm{FeS}}_2$ to ${\mathrm{Fe}}^{3+}$ is described by the unbalanced chemical equation below. ${\mathrm{FeS}}_2(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})+{\mathrm{H}}_2\mathrm{O}(\mathrm{l})⟶$ ${\mathrm{Fe}}^{3+}(\mathrm{aq})+{\mathrm{SO}}_4^{2-}(\mathrm{aq})+{\mathrm{H}}^{+}(\mathrm{aq})(\text{ not balanced })$ Thus, the oxidation of pyrite produces ${\mathrm{Fe}}^{3+}$ and ${\mathrm{H}}^{+}$ ions that can leach into surface or ground water. The leaching of ${\mathrm{H}}^{+}$ ions causes the water to become very acidic. To prevent acidification of nearby ground or surface water, limestone $\left({\mathrm{CaCO}}_3\right)$ is added to the tailings to neutralize the ${\mathrm{H}}^{+}$ ions: ${\mathrm{CaCO}}_3(\mathrm{s})+2{\mathrm{H}}^{+}(\mathrm{aq})\underset{{\mathrm{Ca}}^{2+}}{⟶}(\mathrm{aq})+{\mathrm{H}}_2\mathrm{O}(\mathrm{l})+{\mathrm{CO}}_2(\mathrm{g})$ (a) Balance the equation above for the reaction of ${\mathrm{FeS}}_2$ and ${\mathrm{O}}_2$. [ Hint: Start with the half-equations ${\mathrm{FeS}}_2(\mathrm{s})\to$ ${\mathrm{Fe}}^{3+}(\mathrm{aq})+{\mathrm{SO}}_4^{2-}(\mathrm{aq})\text{ and }{\mathrm{O}}_2(\mathrm{g})\to {\mathrm{H}}_2\mathrm{O}(1).]$ (b) What is the minimum amount of ${\mathrm{CaCO}}_3(\mathrm{s})$ required, per kilogram of tailings, to prevent contamination if the tailings contain $3\mathrm{\%}$ S by mass? Assume that all the sulfur in the tailings is in the form ${\mathrm{FeS}}_2$.

A sample of battery acid is to be analyzed for its sulfuric acid content. A $1.00\mathrm{mL}$ sample weighs $1.239\mathrm{g}$. This $1.00\mathrm{mL}$ sample is diluted to $250.0\mathrm{mL}$, and $10.00\mathrm{mL}$ of this diluted acid requires $32.44\mathrm{mL}$ of $0.00498\mathrm{M}\mathrm{Ba}(\mathrm{OH}{)}_2$ for its titration. What is the mass percent of ${\mathrm{H}}_2{\mathrm{SO}}_4$ in the battery acid? (Assume that complete ionization and neutralization of the ${\mathrm{H}}_2{\mathrm{SO}}_4$ occurs.)