Chapter 21

One material needed to make silicones is dichlorodimethylsilane, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{SiCl}_{2} .\) It is made by treating silicon powder at about \(300^{\circ} \mathrm{C}\) with \(\mathrm{CH}_{3} \mathrm{Cl}\) in the presence of a copper-containing catalyst. (a) Write a balanced equation for the reaction. (b) Assume you carry out the reaction on a small scale with \(2.65 \mathrm{g}\) of silicon. To measure the \(\mathrm{CH}_{3} \mathrm{Cl}\) gas, you fill a \(5.60-\mathrm{L}\) flask at \(24.5^{\circ} \mathrm{C} .\) What pressure of \(\mathrm{CH}_{3} \mathrm{Cl}\) gas must you have in the flask to have the stoichiometrically correct amount of the compound? (c) What mass of \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{SiCl}_{2}\) can be produced from \(2.65 \mathrm{g}\) of \(\mathrm{Si}\) and excess \(\mathrm{CH}_{3} \mathrm{Cl} ?\)

Sodium borohydride, \(\mathrm{NaBH}_{4},\) reduces many metal ions to the metal. (a) Write a balanced equation for the reaction of \(\mathrm{NaBH}_{4}\) with \(\mathrm{AgNO}_{3}\) in water to give silver metal, \(\mathrm{H}_{2}\) gas, boric acid, and sodium nitrate. (The chemistry of \(\mathrm{NaBH}_{4}\) is described on pages \(982-983 .\) ) (b) What mass of silver can be produced from \(575 \mathrm{mL}\) of \(0.011 \mathrm{M} \mathrm{AgNO}_{3}\) and \(13.0 \mathrm{g}\) of \(\mathrm{NaBH}_{4} ?\)

A common analytical method for hydrazine involves its oxidation with iodate ion, \(\mathrm{IO}_{3}^{-},\) in acid solution. In the process, hydrazine acts as a four-electron reducing agent. \(\mathrm{N}_{2}(\mathrm{g})+5 \mathrm{H}_{3} \mathrm{O}^{+}(\mathrm{aq})+4 \mathrm{e}^{-} \rightarrow\) $$\mathrm{N}_{2} \mathrm{H}_{5}^{+}(\mathrm{aq})+5 \mathrm{H}_{2} \mathrm{O}(\ell) \quad E^{\circ}=-0.23 \mathrm{V}$$ Write the balanced equation for the reaction of hydrazine in acid solution \(\left(\mathrm{N}_{2} \mathrm{H}_{5}^{+}\right)\) with \(\mathrm{IO}_{3}^{-}(\mathrm{aq})\) to give \(\mathrm{N}_{2}\) and I. Calculate \(E^{\circ}\) for this reaction.

When \(1.00 \mathrm{g}\) of a white solid \(\mathrm{A}\) is strongly heated, you obtain another white solid, \(\mathrm{B},\) and a gas. An experiment is carried out on the gas, showing that it exerts a pressure of \(209 \mathrm{mm}\) Hg in a 450 -mL flask at \(25^{\circ} \mathrm{C}\) Bubbling the gas into a solution of \(\mathrm{Ca}(\mathrm{OH})_{2}\) gives another white solid, C. If the white solid B is added to water, the resulting solution turns red litmus paper blue. Addition of aqueous HCl to the solution of \(B\) and evaporation of the resulting solution to dryness yield 1.055 g of a white solid D. When \(\mathrm{D}\) is placed in a Bunsen burner flame, it colors the flame green. Finally, if the aqueous solution of \(B\) is treated with sulfuric acid, a white precipitate, \(\mathbf{E},\) forms. Identify the lettered compounds in the reaction scheme.

In \(1937,\) R. Schwartz and M. Schmiesser prepared a yellow-orange bromine oxide (BrO,) by treating Br with ozone in a fluorocarbon solvent. Many years later, J. Pascal found that, on heating, this oxide decomposed to two other oxides, a less volatile golden yellow oxide (A) and a more volatile deep brown oxide (B). Oxide B was later identified as \(\mathrm{Br}_{2} \mathrm{O}\). To determine the formula for oxide \(\mathrm{A},\) a sample was treated with sodium iodide. The reaction liberated iodine, which was titrated to an equivalence point with \(17.7 \mathrm{mL}\) of 0.065 M sodium thiosulfate. $$\mathrm{I}_{2}(\mathrm{aq})+2 \mathrm{S}_{2} \mathrm{O}_{3}^{2-}(\mathrm{aq}) \rightarrow 2 \mathrm{I}^{-}(\mathrm{aq})+\mathrm{S}_{4} \mathrm{O}_{6}^{2-}(\mathrm{aq})$$ Compound A was also treated with AgNO \(_{3},\) and 14.4 mL of 0.020 M AgNO \(_{3}\) was required to completely precipitate the bromine from the sample. (a) What is the formula of the unknown bromine oxide A? (b) Draw Lewis structures for \(A\) and \(B r_{2} O .\) Speculate on their molecular geometry.

A mixture of \(\mathrm{PCl}_{5}(12.41 \mathrm{g})\) and excess \(\mathrm{NH}_{4} \mathrm{Cl}\) was heated at \(145^{\circ} \mathrm{C}\) for 6 hours. The two reacted in equimolar amounts and evolved \(5.14 \mathrm{L}\) of \(\mathrm{HCl}\) (at STP). Three substances \((A, B, \text { and } C)\) were isolated from the reaction mixture. The three substances had the same elemental composition but differed in their molar mass. Substance A had a molar mass of \(347.7 \mathrm{g} / \mathrm{mol}\) and \(\mathrm{B}\) had a molar mass of \(463.5 \mathrm{g} / \mathrm{mol}\). Give the empirical and molecular formulas for \(\mathrm{A}\) and \(\mathrm{B}\) and draw a reasonable Lewis structure for A.

The density of lead is \(11.350 \mathrm{g} / \mathrm{cm}^{3},\) and the metal crystallizes in a face-centered cubic unit cell. Estimate the radius of a lead atom.

You have a 1.0 - \(L\) flask that contains a mixture of argon and hydrogen. The pressure inside the flask is \(745 \mathrm{mm}\) Hg, and the temperature is \(22^{\circ} \mathrm{C} .\) Describe an experiment that you could use to determine the percentage of hydrogen in this mixture.

The boron atom in boric acid, \(\mathrm{B}(\mathrm{OH})_{3},\) is bonded to three - OH groups. In the solid state, the \(-\mathrm{OH}\) groups are in turn hydrogen-bonded to - OH groups in neighboring molecules. (a) Draw the Lewis structure for boric acid. (b) What is the hybridization of the boron atom in the acid? (c) Sketch a picture showing how hydrogen bonding can occur between neighboring molecules.

How would you extinguish a sodium fire in the laboratory? What is the worst thing you could do?

You are given a stoppered flask that contains hydrogen, nitrogen, or oxygen. Suggest an experiment you could do to identify the gas.

Sodium metal is produced by electrolysis of molten sodium chloride. The cell operates at \(7.0 \mathrm{V}\) with a current of \(25 \times 10^{3} \mathrm{A}\) (a) What mass of sodium can be produced in 1 hour? (b) How many kilowatt-hours of electricity are used to produce \(1.00 \mathrm{kg}\) of sodium metal \((1 \mathrm{kWh}=\) \(\left.3.6 \times 10^{6} \mathrm{J}\right) ?\)

(a) Magnesium is obtained from sea water. If the concentration of \(\mathrm{Mg}^{2+}\) in sea water is \(0.050 \mathrm{M},\) what volume of sea water (in liters) must be treated to obtain \(1.00 \mathrm{kg}\) of magnesium metal? What mass of lime (CaO; in kilograms) must be used to precipitate the magnesium in this volume of sea water? (b) When \(1.2 \times 10^{3} \mathrm{kg}\) of molten \(\mathrm{MgCl}_{2}\) is electrolyzed to produce magnesium, what mass (in kilograms) of metal is produced at the cathode? What is produced at the anode? What is the mass of this product? What is the total number of Faradays of electricity used in the process? (c) One industrial process has an energy consumption of \(18.5 \mathrm{kWh} / \mathrm{kg}\) of \(\mathrm{Mg} .\) How many joules are required per mole ( \(1 \mathrm{kWh}=1\) kilowatt-hour \(=\) \(\left.3.6 \times 10^{6} \mathrm{J}\right) ?\) How does this energy compare with the energy of the following process? $$\mathrm{MgCl}_{2}(\mathrm{s}) \rightarrow \mathrm{Mg}(\mathrm{s})+\mathrm{Cl}_{2}(\mathrm{g})$$