Chapter 11

A You have a sample of helium gas at \(-33^{\circ} \mathrm{C},\) and you want to increase the rms speed of helium atoms by \(10.0 \% .\) To what temperature should the gas be heated to accomplish this?

A bicycle tire has an internal volume of \(1.52 \mathrm{L}\) and contains 0.406 mol of air. The tire will burst if its internal pressure reaches 7.25 atm. To what temperature, in degrees Celsius, does the air in the tire need to be heated to cause a blowout?

The temperature of the atmosphere on Mars can be as high as \(27^{\circ} \mathrm{C}\) at the equator at noon, and the atmospheric pressure is about \(8 \mathrm{mm}\) Hg. If a spacecraft could collect \(10 . \mathrm{m}^{3}\) of this atmosphere, compress it to a small volume, and send it back to Earth, how many moles would the sample contain?

If you place \(2.25 \mathrm{g}\) of solid silicon in a \(6.56-\mathrm{L}\). flask that contains \(\mathrm{CH}_{3} \mathrm{Cl}\) with a pressure of \(585 \mathrm{mm} \mathrm{Hg}\) at \(25^{\circ} \mathrm{C}\) what mass of dimethyldichlorosilane, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{SiCl}_{2}(\mathrm{g})\) can be formed? $$\mathrm{Si}(\mathrm{s})+2 \mathrm{CH}_{3} \mathrm{Cl}(\mathrm{g}) \rightarrow\left(\mathrm{CH}_{3}\right)_{2} \mathrm{SiCl}_{2}(\mathrm{g})$$ What pressure of \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{SiCl}_{2}(\mathrm{g})\) would you expect in this same flask at \(95^{\circ} \mathrm{C}\) on completion of the reaction? (Dimethyldichlorosilane is one starting material used to make silicones, polymeric substances used as lubricants, antistick agents, and water-proofing caulk.)

\(\mathrm{Ni}(\mathrm{CO})_{4}\) can be made by reacting finely divided nickel with gaseous CO. If you have CO in a \(1.50-\mathrm{L}\). flask at a pressure of \(418 \mathrm{mm}\) Hg at \(25.0^{\circ} \mathrm{C},\) along with \(0.450 \mathrm{g}\) of Ni powder, what is the theoretical yield of \(\mathrm{Ni}(\mathrm{CO})_{4} ?\)

Ethane burns in air to give \(\mathrm{H}_{2} \mathrm{O}\) and \(\mathrm{CO}_{2}\) $$2 \mathrm{C}_{2} \mathrm{H}_{6}(\mathrm{g})+7 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{CO}_{2}(\mathrm{g})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})$$ (a) Four gases are involved in this reaction. Place them in order of increasing rms speed. (Assume all are at the same temperature.) (b) A \(3.26-\) - flask contains \(C_{2} H_{6}\) at a pressure of \(256 \mathrm{mm} \mathrm{Hg}\) and a temperature of \(25^{\circ} \mathrm{C} .\) Suppose \(\mathrm{O}_{2}\) gas is added to the flask until \(\mathrm{C}_{2} \mathrm{H}_{6}\) and \(\mathrm{O}_{2}\) are in the correct stoichiometric ratio for the combustion reaction. At this point, what is the partial pressure of \(\mathrm{O}_{2}\) and what is the total pressure in the flask?

You have four gas samples: 1\. \(1.0 \mathrm{L}\) of \(\mathrm{H}_{2}\) at \(\mathrm{STP}\) 2\. \(1.0 \mathrm{L}\) of \(\mathrm{Ar}\) at \(\mathrm{STP}\) 3\. \(1.0 \mathrm{L}\) of \(\mathrm{H}_{2}\) at \(27^{\circ} \mathrm{C}\) and \(760 \mathrm{mm}\) Hg 4\. \(1.0 \mathrm{L}\) of He at \(0^{\circ} \mathrm{C}\) and \(900 \mathrm{mm} \mathrm{Hg}\) (a) Which sample has the largest number of gas particles (atoms or molecules)? (b) Which sample contains the smallest number of particles? (c) Which sample represents the largest mass?

Iron carbonyl can be made by the direct reaction of iron metal and carbon monoxide. $$ \mathrm{Fe}(\mathrm{s})+5 \mathrm{CO}(\mathrm{g}) \rightarrow \mathrm{Fe}(\mathrm{CO})_{5}(\ell) $$ What is the theoretical yield of \(\mathrm{Fe}(\mathrm{CO})_{5}\), if \(3.52 \mathrm{g}\) of iron is treated with CO gas having a pressure of \(732 \mathrm{mm} \mathrm{Hg}\) in a \(5.50-\mathrm{L}\). flask at \(23^{\circ} \mathrm{C} ?\)

Analysis of a gaseous chlorofluorocarbon, \(\mathrm{CCl}_{x} \mathrm{F}_{y}\), shows that it contains \(11.79 \%\) C and \(69.57 \%\) Cl. In another experiment, you find that \(0.107 \mathrm{g}\) of the compound fills a 458 -mL. flask at \(25^{\circ} \mathrm{C}\) with a pressure of \(21.3 \mathrm{mm}\) Hg. What is the molecular formula of the compound?

There are five compounds in the family of sulfur-fluorine compounds with the general formula \(\mathrm{S}_{x} \mathrm{F}_{y}\). One of these compounds is \(25.23 \%\) S. If you place \(0.0955 \mathrm{g}\) of the compound in a \(89-\mathrm{mL}\). flask at \(45^{\circ} \mathrm{C},\) the pressure of the gas is 83.8 mm Hg. What is the molecular formula of \(\mathrm{S}_{x} \mathrm{F}_{y} ?\)

The density of air \(20 \mathrm{km}\) above Earth's surface is \(92 \mathrm{g} / \mathrm{m}^{3} .\) The pressure of the atmosphere is \(42 \mathrm{mm} \mathrm{Hg}\) and the temperature is \(-63^{\circ} \mathrm{C}\) (a) What is the average molar mass of the atmosphere at this altitude? (b) If the atmosphere at this altitude consists of only \(\mathrm{O}_{2}\) and \(\mathrm{N}_{2},\) what is the mole fraction of each gas?

Chlorine dioxide, \(\mathrm{ClO}_{2},\) reacts with fluorine to give a new gas that contains \(\mathrm{Cl}\), \(\mathrm{O}\), and \(\mathrm{F}\). In an experiment, you find that \(0.150 \mathrm{g}\) of this new gas has a pressure of \(17.2 \mathrm{mm}\) Hg in a \(1850-\mathrm{mL}\). flask at \(21^{\circ} \mathrm{C} .\) What is the identity of the unknown gas?

A xenon fluoride can be prepared by heating a mixture of \(\mathrm{Xe}\) and \(\mathrm{F}_{2}\) gases to a high temperature in a pressureproof container. Assume that xenon gas was added to a \(0.25-\mathrm{L}\) container until its pressure reached \(0.12 \mathrm{atm}\) at \(0.0^{\circ} \mathrm{C} .\) Fluorine gas was then added until the total pressure reached 0.72 atm at \(0.0^{\circ} \mathrm{C}\). After the reaction was complete, the xenon was consumed completely, and the pressure of the \(\mathrm{F}_{2}\) remaining in the container was 0.36 atm at \(0.0^{\circ} \mathrm{C} .\) What is the empirical formula of the xenon fluoride?

Which of the following is not correct? (a) Diffusion of gases occurs more rapidly at higher temperatures. (b) Effusion of \(\mathrm{H}_{2}\) is faster than effusion of He (assume similar conditions and a rate expressed in units of \(\mathrm{mol} / \mathrm{h})\) (c) Diffusion will occur faster at low pressure than at high pressure. (d) The rate of effusion of a gas \((\mathrm{mol} / \mathrm{h})\) is directly proportional to molar mass.

The ideal gas law is least accurate under conditions of high pressure and low temperature. In those situations, using the van der Waals equation is advisable. (a) Calculate the pressure exerted by \(12.0 \mathrm{g}\) of \(\mathrm{CO}_{2}\) in a \(500-\mathrm{mL}\) vessel at \(298 \mathrm{K},\) using the ideal gas equation. Then, recalculate the pressure using the van der Waals equation. Assuming the pressure calculated from van der Waal's equation is correct, what is the percent error in the answer when using the ideal gas equation? (b) Next, cool this sample to \(-70^{\circ} \mathrm{C}\). Then perform the same calculation for the pressure exerted by \(\mathrm{CO}_{2}\) at this new temperature, using both the ideal gas law and the van der Waals equation. Again, what is the percent error when using the ideal gas equation?

Carbon dioxide, \(\mathrm{CO}_{2},\) was shown to effuse through a porous plate at the rate of \(0.033 \mathrm{mol} / \mathrm{min.}\) The same quantity of an unknown gas, 0.033 moles, is found to effuse through the same porous barrier in 104 seconds. Calculate the molar mass of the unknown gas.

If you have a sample of water in a closed container, some of the water will evaporate until the pressure of the water vapor, at \(25^{\circ} \mathrm{C},\) is \(23.8 \mathrm{mm}\) Hg. How many molecules of water per cubic centimeter exist in the vapor phase?

You are given \(1.56 \mathrm{g}\) of a mixture of \(\mathrm{KClO}_{3}\) and \(\mathrm{KCl}\). When heated, the KClOg decomposes to KCl and \(\mathrm{O}_{2}\), $$ 2 \mathrm{KClO}_{3}(\mathrm{s}) \rightarrow 2 \mathrm{KCl}(\mathrm{s})+3 \mathrm{O}_{2}(\mathrm{g}) $$ and \(327 \mathrm{mL}\) of \(\mathrm{O}_{2}\) with a pressure of \(735 \mathrm{mm}\) Hg is collected at \(19^{\circ} \mathrm{C}\). What is the weight percentage of \(\mathrm{KClO}_{3}\) in the sample?

A A study of climbers who reached the summit of Mount Everest without supplemental oxygen showed that the partial pressures of \(\mathrm{O}_{2}\) and \(\mathrm{CO}_{2}\) in their lungs were 35 mm Hg and 7.5 mm Hg, respectively. The barometric pressure at the summit was 253 mm Hg. Assume the lung gases are saturated with moisture at a body temperature of \(37^{\circ} \mathrm{C}\) [which means the partial pressure of water vapor in the lungs is \(\left.P\left(\mathrm{H}_{2} \mathrm{O}\right)=47.1 \mathrm{mm} \mathrm{Hg}\right]\). If you assume the lung gases consist of only \(\mathbf{O}_{2}, \mathbf{N}_{2}, \mathbf{C O}_{2},\) and \(\mathrm{H}_{2} \mathrm{O},\) what is the partial pressure of \(\mathrm{N}_{2} ?\)

Nitrogen monoxide reacts with oxygen to give nitrogen dioxide: $$ 2 \mathrm{NO}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{NO}_{2}(\mathrm{g}) $$ (a) Place the three gases in order of increasing rms speed at \(298 \mathrm{K}\) (b) If you mix \(\mathrm{NO}\) and \(\mathrm{O}_{2}\) in the correct stoichiometric ratio and NO has a partial pressure of \(150 \mathrm{mm} \mathrm{Hg}\) what is the partial pressure of \(\mathrm{O}_{2} ?\) (c) After reaction between \(\mathrm{NO}\) and \(\mathrm{O}_{2}\) is complete, what is the pressure of \(\mathrm{NO}_{2}\) if the NO originally had a pressure of \(150 \mathrm{mm} \mathrm{Hg}\) and \(\mathrm{O}_{2}\) was added in the correct stoichiometric amount?

Ammonia gas is synthesized by combining hydrogen and nitrogen: $$ 3 \mathrm{H}_{2}(\mathrm{g})+\mathrm{N}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{NH}_{3}(\mathrm{g}) $$ (a) If you want to produce 562 g of \(\mathrm{NH}_{3}\), what volume of \(\mathrm{H}_{2}\) gas, at \(56^{\circ} \mathrm{C}\) and \(745 \mathrm{mm}\) Hg, is required? (b) Nitrogen for this reaction will be obtained from air. What volume of air, measured at \(29^{\circ} \mathrm{C}\) and \(745 \mathrm{mm}\) Hg pressure, will be required to provide the nitrogen needed to produce \(562 \mathrm{g}\) of \(\mathrm{NH}_{3} ?\) Assume the sample of air contains 78.1 mole \(\%\) N \(_{2}\)

Nitrogen trifluoride is prepared by the reaction of ammonia and fluorine. $$ 4 \mathrm{NH}_{3}(\mathrm{g})+3 \mathrm{F}_{2}(\mathrm{g}) \rightarrow 3 \mathrm{NH}_{4} \mathrm{F}(\mathrm{s})+\mathrm{NF}_{3}(\mathrm{g}) $$ If you mix \(\mathrm{NH}_{3}\) with \(\mathrm{F}_{2}\) in the correct stoichiometric ratio, and if the total pressure of the mixture is \(120 \mathrm{mm} \mathrm{Hg}\), what are the partial pressures of \(\mathrm{NH}_{3}\) and \(\mathrm{F}_{2} ?\) When the reactants have been completely consumed, what is the total pressure in the flask? (Assume \(T\) is constant.)

Chlorine trifluoride, \(\mathrm{CIF}_{3}\), is a valuable reagent because it can be used to convert metal oxides to metal fluorides: \(6 \mathrm{NiO}(\mathrm{s})+4 \mathrm{ClF}_{3}(\mathrm{g}) \rightarrow 6 \mathrm{NiF}_{2}(\mathrm{s})+2 \mathrm{Cl}_{2}(\mathrm{g})+3 \mathrm{O}_{2}(\mathrm{g})\) (a) What mass of NiO will react with CIF \(_{3}\) gas if the gas has a pressure of \(250 \mathrm{mm} \mathrm{Hg}\) at \(20^{\circ} \mathrm{C}\) in a \(2.5-\mathrm{L}\) flask? (b) If the CIF a described in part (a) is completely consumed, what are the partial pressures of \(\mathrm{Cl}_{2}\) and of \(\mathrm{O}_{2}\) in the 2.5 -I. flask at \(20^{\circ} \mathrm{C}\) (in \(\mathrm{mm}\) Hg)? What is the total pressure in the flask?

A You have a \(550 .\) -mL. tank of gas with a pressure of 1.56 atm at \(24^{\circ} \mathrm{C} .\) You thought the gas was pure carbon monoxide gas, CO, but you later found it was contaminated by small quantities of gaseous \(\mathrm{CO}_{2}\) and \(\mathrm{O}_{2}\). Analysis shows that the tank pressure is 1.34 atm (at \(24^{\circ} \mathrm{C}\) ) if the \(\mathrm{CO}_{2}\) is removed. Another experiment shows that \(0.0870 \mathrm{g}\) of \(\mathrm{O}_{2}\) can be removed chemically. What are the masses of \(\mathrm{CO}\) and \(\mathrm{CO}_{2}\) in the tank, and what is the partial pressure of each of the three gases at \(25^{\circ} \mathrm{C} ?\)

A Methane is burned in a laboratory Bunsen burner to give \(\mathrm{CO}_{2}\) and water vapor. Methane gas is supplied to the burner at the rate of \(5.0 \mathrm{L} / \mathrm{min}\) (at a temperature of \(28^{\circ} \mathrm{C}\) and a pressure of \(773 \mathrm{mm} \mathrm{Hg}\) ). At what rate must oxygen be supplied to the burner (at a pressure of \(\left.742 \mathrm{mm} \mathrm{Hg} \text { and a temperature of } 26^{\circ} \mathrm{C}\right) ?\)

A Iron forms a series of compounds of the type \(\mathrm{Fe}_{x}(\mathrm{CO})_{x}\). In air, these compounds are oxidized to \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) and \(\mathrm{CO}_{2}\) gas. After heating a 0.142 -g sample of \(\mathrm{Fe}_{x}(\mathrm{CO})_{y}\) in air, you isolate the \(\mathrm{CO}_{2}\) in a \(1.50-\mathrm{L}\). flask at \(25^{\circ} \mathrm{C} .\) The pressure of the gas is \(44.9 \mathrm{mm}\) Hg. What is the empirical formula of \(\mathrm{Fe}_{x}(\mathrm{CO})_{y} ?\)

A Group 2 A metal carbonates are decomposed to the metal oxide and \(\mathrm{CO}_{2}\) on heating: $$ \mathrm{MCO}_{3}(\mathrm{s}) \rightarrow \mathrm{MO}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{g}) $$ You heat \(0.158 \mathrm{g}\) of a white, solid carbonate of a Group 2A metal (M) and find that the evolved CO \(_{2}\) has a pressure of \(69.8 \mathrm{mm}\) Hg in a \(285-\mathrm{mL}\). flask at \(25^{\circ} \mathrm{C}\) Identify M.

One way to synthesize diborane, \(\mathrm{B}_{2} \mathrm{H}_{6},\) is the reaction \(2 \mathrm{NaBH}_{4}(\mathrm{s})+2 \mathrm{H}_{3} \mathrm{PO}_{4}(\ell) \rightarrow\) $$ \mathrm{B}_{2} \mathrm{H}_{6}(\mathrm{g})+2 \mathrm{NaH}_{2} \mathrm{PO}_{4}(\mathrm{s})+2 \mathrm{H}_{2}(\mathrm{g}) $$ (a) If you have \(0.136 \mathrm{g}\) of \(\mathrm{NaBH}_{4}\) and excess \(\mathrm{H}_{3} \mathrm{PO}_{4}\) and you collect the \(\mathrm{B}_{2} \mathrm{H}_{6}\) in a \(2.75-\mathrm{L}\). flask at \(25^{\circ} \mathrm{C}\) what is the pressure of the \(\mathrm{B}_{2} \mathrm{H}_{6}\) in the flask? (b) A by-product of the reaction is \(\mathrm{H}_{2}\) gas. If both \(\mathrm{B}_{2} \mathrm{H}_{6}\) and \(\mathrm{H}_{2}\) gas come from this reaction, what is the total pressure in the \(2.75-\mathrm{L}\). flask (after reaction of \(0.136 \mathrm{g}\) of \(\mathrm{NaBH}_{4}\) with excess \(\mathrm{H}_{3} \mathrm{PO}_{4}\) ) at \(25^{\circ} \mathrm{C} ?\)

You are given a solid mixture of \(\mathrm{NaNO}_{2}\) and \(\mathrm{NaCl}\) and are asked to analyze it for the amount of \(\mathrm{NaNO}_{2}\) present. To do so, you allow the mixture to react with sulfamic acid, HSO \(_{3} \mathrm{NH}_{2}\), in water according to the equation $$\begin{aligned} \mathrm{NaNO}_{2}(\mathrm{aq})+\mathrm{HSO}_{3} \mathrm{NH}_{2}(\mathrm{aq}) & \rightarrow \\ & \mathrm{NaHSO}_{4}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{N}_{2}(\mathrm{g}) \end{aligned}$$ What is the weight percentage of \(\mathrm{NaNO}_{2}\) in \(1.232 \mathrm{g}\) of the solid mixture if reaction with sulfamic acid produces \(295 \mathrm{mL}\) of dry \(\mathrm{N}_{2}\) gas with a pressure of \(713 \mathrm{mm}\) Hg at \(21.0^{10} \mathrm{C} ?\)

A You have \(1.249 \mathrm{g}\) of a mixture of \(\mathrm{NaHCO}_{3}\) and \(\mathrm{Na}_{2} \mathrm{CO}_{3} .\) You find that \(12.0 \mathrm{mL}\) of \(1.50 \mathrm{M} \mathrm{HCl}\) is required to convert the sample completely to \(\mathrm{NaCl}\), \(\mathrm{H}_{2} \mathrm{O},\) and \(\mathrm{CO}_{2}\) \(\mathrm{NaHCO}_{3}(\mathrm{aq})+\mathrm{HCl}(\mathrm{aq}) \rightarrow\) $$ \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{CO}_{2}(\mathrm{g}) $$ \(\mathrm{Na}_{2} \mathrm{CO}_{3}(\mathrm{aq})+2 \mathrm{HCl}(\mathrm{aq}) \rightarrow\) $$ 2 \mathrm{NaCl}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{CO}_{2}(\mathrm{g}) $$ What volume of \(\mathrm{CO}_{2}\) is evolved at \(745 \mathrm{mm}\) Hg and \(25^{\circ} \mathrm{C} ?\)

A compound containing \(\mathrm{C}, \mathrm{H}, \mathrm{N},\) and \(\mathrm{O}\) is burned in excess oxygen. The gases produced by burning \(0.1152 \mathrm{g}\) are first treated to convert the nitrogen-containing product gases into \(\mathrm{N}_{2},\) and then the resulting mixture of \(\mathrm{CO}_{2}, \mathrm{H}_{2} \mathrm{O}, \mathrm{N}_{2},\) and excess \(\mathrm{O}_{2}\) is passed through a bed of \(\mathrm{CaCl}_{2}\) to absorb the water. The \(\mathrm{CaCl}_{2}\) increases in mass by \(0.09912 \mathrm{g} .\) The remaining gases are bubbled into water to form \(\mathrm{H}_{2} \mathrm{CO}_{3},\) and this solution is titrated with \(0.3283 \mathrm{M} \mathrm{NaOH} ; 28.81 \mathrm{mL}\) is required to achieve the second equivalence point. The excess \(\mathbf{O}_{2}\) gas is removed by reaction with copper metal (to give CuO). Finally, the \(\mathrm{N}_{2}\) gas is collected in a 225.0 -mL. flask, where it has a pressure of \(65.12 \mathrm{mm} \mathrm{Hg}\) at \(25^{\circ} \mathrm{C} .\) In a separate experiment, the unknown compound is found to have a molar mass of \(150 \mathrm{g} / \mathrm{mol}\). What are the empirical and molecular formulas of the unknown compound?

You have a gas, one of the three known phosphorusfluorine compounds \(\left(\mathrm{PF}_{3}, \mathrm{PF}_{5}, \text { and } \mathrm{P}_{2} \mathrm{F}_{4}\right) .\) To find out which, you have decided to measure its molar mass. (a) First, you determine that the density of the gas is \(5.60 \mathrm{g} / \mathrm{L}\) at a pressure of \(0.971 \mathrm{atm}\) and a temperature of \(18.2^{\circ} \mathrm{C} .\) Calculate the molar mass and identify the compound. (b) To check the results from part (a), you decide to measure the molar mass based on the relative rates of effusion of the unknown gas and \(\mathrm{CO}_{2} .\) You find that \(\mathrm{CO}_{2}\) effuses at a rate of \(0.050 \mathrm{mol} / \mathrm{min}\) whereas the unknown phosphorus fluoride effuses at a rate of \(0.028 \mathrm{mol} / \mathrm{min.}\) Calculate the molar mass of the unknown gas based on these results.

A 1.50 L constant volume calorimeter (Figure 5.12 ) contains \(\mathrm{C}_{3} \mathrm{H}_{8}(\mathrm{g})\) and \(\mathrm{O}_{2}(\mathrm{g}) .\) The partial pressure of \(\mathrm{C}_{3} \mathrm{H}_{8}\) is 0.10 atm and the partial pressure of \(\mathrm{O}_{2}\) is 5.0 atm. The temperature is \(20.0^{\circ} \mathrm{C}\). A reaction occurs between the two compounds, forming \(\mathrm{CO}_{2}(\mathrm{g})\) and \(\mathrm{H}_{2} \mathrm{O}(\ell) .\) The heat from the reaction causes the temperature to rise to \(23.2^{\circ} \mathrm{C}\) (a) Write a balanced chemical equation for the reaction. (b) How many moles of \(\mathrm{C}_{3} \mathrm{H}_{8}(\mathrm{g})\) are present in the flask initially? (c) What is the mole fraction of \(\mathrm{C}_{3} \mathrm{H}_{8}(\mathrm{g})\) in the flask before reaction? (d) After the reaction, the flask contains excess oxygen and the products of the reaction, \(\mathrm{CO}_{2}(\mathrm{g})\) and \(\mathrm{H}_{2} \mathrm{O}(\ell) .\) What amount of unreacted \(\mathrm{O}_{2}(\mathrm{g})\) remains? (e) After the reaction, what is the partial pressure exerted by the \(\mathrm{CO}_{2}(\mathrm{g})\) in this system? (f) What is the partial pressure exerted by the excess oxygen remaining after the reaction?

A \(1.0-\) - flask contains 10.0 g each of \(\mathrm{O}_{2}\) and \(\mathrm{CO}_{2}\) at \(25^{\circ} \mathrm{C}\) (a) Which gas has the greater partial pressure, \(\mathbf{O}_{2}\) or \(\mathrm{CO}_{2},\) or are they the same? (b) Which molecules have the greater rms speed, or are they the same? (c) Which molecules have the greater average kinetic energy, or are they the same?

If equal masses of \(\mathrm{O}_{2}\) and \(\mathrm{N}_{2}\) are placed in separate containers of equal volume at the same temperature, which of the following statements is true? If false, explain why it is false. (a) The pressure in the flask containing \(\mathrm{N}_{2}\) is greater than that in the flask containing \(\mathbf{O}_{2}\) (b) There are more molecules in the flask containing \(\mathrm{O}_{2}\) than in the flask containing \(\mathrm{N}_{2}\)

You have two pressure-proof steel cylinders of equal volume, one containing \(1.0 \mathrm{kg}\) of \(\mathrm{CO}\) and the other containing \(1.0 \mathrm{kg}\) of acetylene, \(\mathrm{C}_{2} \mathrm{H}_{2}\) (a) In which cylinder is the pressure greater at \(25^{\circ} \mathrm{C} ?\) (b) Which cylinder contains the greater number of molecules?

Two flasks, each with a volume of \(1.00 \mathrm{L},\) contain \(\mathrm{O}_{2}\) gas with a pressure of 380 mm Hg. Flask \(A\) is at \(25^{\circ} \mathrm{C}\) and flask \(\mathrm{B}\) is at \(0^{\circ} \mathrm{C}\). Which flask contains the greater number of \(\mathrm{O}_{2}\) molecules?

Each of four flasks is filled with a different gas. Each flask has the same volume, and each is filled to the same pressure, \(3.0 \mathrm{atm},\) at \(25^{\circ} \mathrm{C} .\) Flask \(\mathrm{A}\) contains \(116 \mathrm{g}\) of air, flask \(\mathrm{B}\) has \(80.7 \mathrm{g}\) of neon, flask \(\mathrm{C}\) has \(16.0 \mathrm{g}\) of helium, and flask \(\mathrm{C}\) has \(160 . \mathrm{g}\) of an unknown gas. (a) Do all four flasks contain the same number of gas molecules? If not, which one has the greatest number of molecules? (b) How many times heavier is a molecule of the unknown gas than an atom of helium? (c) In which flask do the molecules have the largest kinetic energy? The highest rms speed?

You have two gas-filled balloons, one containing He and the other containing \(\mathrm{H}_{2} .\) The \(\mathrm{H}_{2}\) balloon is twice the size of the He balloon. The pressure of gas in the \(\mathrm{H}_{2}\) balloon is \(1 \mathrm{atm},\) and that in the He balloon is 2 atm. The \(H_{2}\) balloon is outside in the snow \(\left(-5^{\circ} \mathrm{C}\right)\) and the He balloon is inside a warm building \(\left(23^{\circ} \mathrm{C}\right)\) (a) Which balloon contains the greater number of molecules? (b) Which balloon contains the greater mass of gas?

The sodium azide required for automobile air bags is made by the reaction of sodium metal with dinitrogen monoxide in liquid ammonia: $$\begin{aligned} 3 \mathrm{N}_{2} \mathrm{O}(\mathrm{g})+4 \mathrm{Na}(\mathrm{s})+\mathrm{NH}_{3}(\ell) & \rightarrow \\ & \mathrm{NaN}_{3}(\mathrm{s})+3 \mathrm{NaOH}(\mathrm{s})+2 \mathrm{N}_{2}(\mathrm{g}) \end{aligned}$$ (a) You have \(65.0 \mathrm{g}\) of sodium, a \(35.0-\mathrm{L}\). flask containing \(\mathrm{N}_{2} \mathrm{O}\) gas with a pressure of 2.12 atm at \(23^{\circ} \mathrm{C}\) and excess ammonia. What is the theoretical yield (in grams) of NaNg? (b) Draw a Lewis structure for the azide ion. Include all possible resonance structures. Which resonance structure is most likely? (c) What is the shape of the azide ion?

If the absolute temperature of a gas doubles, by how much does the rms speed of the gaseous molecules increase?

A Chlorine gas \(\left(\mathrm{Cl}_{2}\right)\) is used as a disinfectant in municipal water supplies, although chlorine dioxide \(\left(\mathrm{ClO}_{2}\right)\) and ozone are becoming more widely used. \(\mathrm{ClO}_{2}\) is a better choice than \(\mathrm{Cl}_{2}\) in this application because it leads to fewer chlorinated by-products, which are themselves pollutants. (a) How many valence electrons are in \(\mathrm{ClO}_{2} ?\) (b) The chlorite ion, \(\mathrm{ClO}_{2}^{-}\), is obtained by reducing ClO \(_{2} .\) Draw a possible electron dot structure for \(\mathrm{ClO}_{2}^{-} .\) (Cl is the central atom.) (c) What is the hybridization of the central Cl atom in \(\mathrm{ClO}_{2}^{-} ?\) What is the shape of the ion? (d) Which species has the larger bond angle, \(\mathbf{O}_{3}\) or \(\mathrm{ClO}_{2}^{-} ?\) Explain briefly. (e) Chlorine dioxide, \(\mathrm{ClO}_{2},\) a yellow-green gas, can be made by the reaction of chlorine with sodium chlorite: \(2 \mathrm{NaClO}_{2}(\mathrm{s})+\mathrm{Cl}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{NaCl}(\mathrm{s})+2 \mathrm{ClO}_{2}(\mathrm{g})\) Assume you react \(15.6 \mathrm{g}\) of \(\mathrm{NaClO}_{2}\) with chlorine gas, which has a pressure of \(1050 \mathrm{mm}\) Hg in a \(1.45-\mathrm{L}\). flask at \(22^{\circ} \mathrm{C}\). What mass of \(\mathrm{ClO}_{2}\) can be produced?