In an experimental setup, chlorine gas and nitrogen gas are in separate containers separated by a closed stopcock. One hour after the stopcock is opened, the gases have completely mixed. Five minutes after the stopcock is opened, which gas will have traveled farther, the nitrogen or the chlorine?

Nitric acid is produced commercially by the Ostwald process. In the first step ammonia is oxidized to nitric oxide:

$$4 \mathrm{NH}_3(g)+5 \mathrm{O}_2(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_2 \mathrm{O}(g)$$

The stopcock between the two reaction containers is opened, and the reaction proceeds using proper catalysts. Calculate the partial pressure of $\mathrm{NO}$ after the reaction is complete. Assume $100 \%$ yield for the reaction, assume the final container volume is $3.00 \mathrm{~L}$, and assume the temperature is constant.

Nitric acid, a major industrial and laboratory acid, is produced commercially by the multistep Ostwald process, which begins with the oxidation of ammonia:
Step 1. $4 \mathrm{NH}_3(g)+5 \mathrm{O}_2(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_2 \mathrm{O}(l)$
Step 2. $2 \mathrm{NO}(g)+\mathrm{O}_2(g) \longrightarrow 2 \mathrm{NO}_2(g)$ Step 3. $3 \mathrm{NO}_2(g)+\mathrm{H}_2 \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_3(l)+\mathrm{NO}(g)$
(a) What are the oxidizing and reducing agents in each step?

Nitric acid is produced from nitric oxide, $NO$, which in turn is prepared from ammonia by the Ostwald process:

$$4NH_3(g) + 5O_2(g) \longrightarrow 4NO(g) + 6H_2O(g)$$

What volume of oxygen at $35^\circ C$ and 2.15 atm is needed to produce 50.0 g of nitric oxide?

Nitric acid is produced commercially by the Ostwald process, represented by the following equations:

$$\begin{array} { l } { 4 \mathrm { NH } _ { 3 } ( g ) + 5 \mathrm { O } _ { 2 } ( g ) \longrightarrow 4 \mathrm { NO } ( g ) + 6 \mathrm { H } _ { 2 } \mathrm { O } ( g ) } \\ { 2 \mathrm { NO } ( g ) + \mathrm { O } _ { 2 } ( g ) \longrightarrow 2 \mathrm { NO } _ { 2 } ( g ) } \\ { 3 \mathrm { NO } _ { 2 } ( g ) + \mathrm { H } _ { 2 } \mathrm { O } ( l ) \longrightarrow 2 \mathrm { HNO } _ { 3 } ( a q ) + \mathrm { NO } ( g ) } \end{array}$$

What mass of $\mathrm { NH } _ { 3 }$ must be used to produce $1.0 \times 10 ^ { 6 } \mathrm { kg }$ $\mathrm {HNO} _ { 3 }$ by the Ostrwald process? Assume 100% yield in each reaction, and assume that the NO produced in the third step is not recycled.

Nitric acid is produced commercially by the Ostwald process, represented by the following equations:

$$4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g) \\ 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g) \\ 3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(a q)+\mathrm{NO}(g)$$

What mass of $\mathrm{NH}_{3}$ must be used to produce $1.0 \times 10^{6} \mathrm{kg}\ \mathrm{HNO}_{3}$ by the Ostwald process? Assume 100% yield in each reaction, and assume that the NO produced in the third step is not recycled.

The Ostwald process is a series of three reactions used for the industrial production of nitric acid from ammonia.

If NO is not recycled, how many moles of NH3 are consumed per mole of HNO3 produced?

The Ostwald process is used commercially to produce nitric acid, which is, in turn, used in many modern chemical processes. In the first step of the Ostwald process, ammonia is reacted with oxygen gas to produce nitric oxide and water.

What is the maximum mass of $\ce{H_2O}$ that can be produced by combining $\ce{66.1 g}$ of each reactant?

$$\ce{4NH3(g) +5O2(g) → 4NO(g) +6H2O(g)}$$

A mixture of 1.00 g $\mathrm { H } _ { 2 }$ and 1.00 g He is placed in a 1.00-L container at $27 ^ { \circ } \mathrm { C }.$ Calculate the partial pressure of each gas and the total pressure.

Two flasks of equal volume and at the same temperature contain different gases. One flask contains 10.0 g of Ne and the other flask contains 10.0 g of He. Which of the following statements are correct? Explain.

a. Both flasks contain the same number of atoms.

b. The pressures in the flasks are the same.

c. The flask that contains helium has a higher pressure than the flask that contains neon.

d. The densities of the gases are the same.

Consider the flasks in the following diagram. What are the final partial pressures of $\mathrm{H}_2$ and $\mathrm{N}_2$ after the stopcock between the two flasks is opened? (Assume the final volume is $3.00 \mathrm{~L}$.) What is the total pressure (in torr)?

Examine the sensitivity of the final pressure to the containment room volume in earlier problem. Solve for the volume for a range of final pressures, $100-250 \mathrm{kPa}$, and sketch the pressure versus volume curve.

Imagine some helium in a cylinder with an initial volume of 1 liter and an initial pressure of 1 atm. Somehow the helium is made to expand to a final volume of 3 liters, in such a way that its pressure rises in direct proportion to its volume. (a) Sketch a graph of pressure vs. volume for this process. (b) Calculate the work done on the gas during this process, assuming that there are no "other" types of work being done. (c) Calculate the change in the helium's energy content during this process. (d) Calculate the amount of heat added to or removed from the helium during this process. (e) Describe what you might do to cause the pressure to rise as the helium expands.

When a gas is compressed adiabatically, there is a pressure rise and a temperature rise. The ratio of final temperature to initial temperature is less than the ratio of final pressure to initial pressure. Will the final density be (a) less or (b) greater than the initial density?