A 14.22-g aluminum soda can reacts with hydrochloric acid according to the reaction:

$$2 \mathrm{Al}(\mathrm{s})+6 \mathrm{HCl}(a q) \longrightarrow 3 \mathrm{H}_2(g)+2 \mathrm{AlCl}_3(a q)$$

The hydrogen gas is collected by the displacement of water. What volume of hydrogen is collected if the external pressure is $749 \mathrm{~mm} \mathrm{Hg}$ and the temperature is $31{ }^{\circ} \mathrm{C}$, assuming that there is no water vapor present? What is the vapor pressure of water at this temperature? What volume do the hydrogen and water vapor occupy under these conditions?

Aerosol cans carry clear warnings against incineration because of the high pressures that can develop upon heating. Suppose a can contains a residual amount of gas at a pressure of $755 \mathrm{~mm} \mathrm{Hg}$ and a temperature of $25^{\circ} \mathrm{C}$. What would the pressure be if the can were heated to $1155^{\circ} \mathrm{C}$ ?

In a common classroom demonstration, a balloon is filled with air and submerged into liquid nitrogen. The balloon contracts as the gases within the balloon cool. Suppose the balloon initially contains $2.95 \mathrm{~L}$ of air at a temperature of $25.0^{\circ} \mathrm{C}$ and a pressure of $0.998 \mathrm{~atm}$. Calculate the expected volume of the balloon upon cooling to $-196^{\circ} \mathrm{C}$ (the boiling point of liquid nitrogen). When the demonstration is carried out, the actual volume of the balloon decreases to $0.61$ L. How well does the observed volume of the balloon compare to your calculated value? Can you explain the difference?

Olympic cyclists fill their tires with helium to make them lighter. Calculate the mass of air in an air-filled tire and the mass of helium in a helium-filled tire. What is the mass difference between the two? Assume that the volume of the tire is $855 \mathrm{~mL}$, that it is filled with a total pressure of $125 \mathrm{psi}$, and that the temperature is $25^{\circ} \mathrm{C}$. Also, assume an average molar mass for air of $28.8 \mathrm{~g} / \mathrm{mol}$.

Automobile air bags inflate following a serious impact. The impact triggers the chemical reaction:

$$2 \mathrm{NaN}_3(s) \rightarrow 2 \mathrm{Na}(s)+3 \mathrm{~N}_2(g)$$

If an automobile air bag has a volume of $11.8 \mathrm{~L}$, how much $\mathrm{NaN}_3$ in grams is required to fully inflate the air bag upon impact? Assume STP conditions.

Consider the reaction:

$$\mathrm{CO}(g)+2 \mathrm{H}_2(g) \longrightarrow \mathrm{CH}_3 \mathrm{OH}(g)$$

A reaction flask initially contains $112$ torr of $\mathrm{CO}$ and $282$ torr of $\mathrm{H}_2$. The reaction is allowed to proceed until the pressure stops changing, at which point the total pressure is $196$ torr. Determine the percent yield for the reaction. Assume a constant temperature and that no other reactions occur other than the one indicated.

Consider the reaction:

$$2 \mathrm{SO}_2(g)+\mathrm{O}_2(g) \longrightarrow 2 \mathrm{SO}_3(g)$$

A reaction flask initially contains $0.10 \mathrm{~atm}$ of $\mathrm{SO}_2$ and $0.10 \mathrm{~atm}$ of $\mathrm{O}_2$. What is the total pressure in the flask once the limiting reactant is completely consumed? Assume a constant temperature and volume and a $100 \%$ reaction yield.

A mixture containing $4.33 \mathrm{~g}$ of $\mathrm{CO}_2$ and $3.11 \mathrm{~g}$ of $\mathrm{CH}_4$ has a total pressure of $1.09 \mathrm{~atm}$. What is the partial pressure of $\mathrm{CO}_2$ in the mixture?

What is the combined gas law? When is it useful?

A mixture containing $235 \mathrm{mg}$ of helium and $325 \mathrm{mg}$ of neon has a total pressure of $453$ torr. What is the partial pressure of helium in the mixture?

Ammonium nitrate decomposes explosively upon heating according to the balanced equation:

$$2 \mathrm{NH}_4 \mathrm{NO}_3(s) \longrightarrow 2 \mathrm{~N}_2(g)+\mathrm{O}_2(g)+4 \mathrm{H}_2 \mathrm{O}(g)$$

Calculate the total volume of gas (at $25^{\circ} \mathrm{C}$ and $748 \mathrm{~mm} \mathrm{Hg}$ ) produced by the complete decomposition of $1.55 \mathrm{~kg}$ of ammonium nitrate.

Ammonium carbonate decomposes upon heating according to the balanced equation:

$$\left(\mathrm{NH}_4\right)_2 \mathrm{CO}_3(s) \longrightarrow 2 \mathrm{NH}_3(g)+\mathrm{CO}_2(g)+\mathrm{H}_2 \mathrm{O}(g)$$

Calculate the total volume of gas produced at $22{ }^{\circ} \mathrm{C}$ and $1.02 \mathrm{~atm}$ by the complete decomposition of $11.83 \mathrm{~g}$ of ammonium carbonate.

Consider the reaction for the production of $\mathrm{NO}_2$ from NO:

$$2 \mathrm{NO}(g)+\mathrm{O}_2(g) \longrightarrow 2 \mathrm{NO}_2(g)$$

(a) If $84.8 \mathrm{~L}^2$ of $\mathrm{O}_2(\mathrm{~g})$, measured at $35^{\circ} \mathrm{C}$ and $632 \mathrm{~mm}$ $\mathrm{Hg}$, is allowed to react with $158.2 \mathrm{~g}$ of $\mathrm{NO}$, find the limiting reagent.

(b) If $97.3 \mathrm{~L}$ of $\mathrm{NO}_2$ forms, measured at $35^{\circ} \mathrm{C}$ and $632 \mathrm{~mm} \mathrm{Hg}$, what is the percent yield? n n

Consider the reaction for the synthesis of nitric acid: $3 \mathrm{NO}_2(g)+\mathrm{H}_2 \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_3(a q)+\mathrm{NO}(g)$

(a) If $12.8 \mathrm{~L}$ of $\mathrm{NO}_2(\mathrm{~g})$, measured at STP, is allowed to react with $14.9 \mathrm{~g}$ of water, find the limiting reagent and the theoretical yield of $\mathrm{HNO}_3$ in grams.

(b) If $14.8 \mathrm{~g}$ of $\mathrm{HNO}_3$ forms, what is the percent yield?