The density of a mixture of fluorine and chlorine gases is 1.77 g/L at 14 degree C and 0.893 atm. Calculate the mass percent of the gases.

A quantity of 0.225 g of a metal M (molar mass = 27.0 g/mol) liberated 0.303 L of molecular hydrogen (measured at 17 °C and 741 mmHg) from an excess of hydrochloric acid. Deduce from these data the corresponding equation, and write formulas for the oxide and sulfate of M.

Methane, the principal component of natural gas, is used for heating and cooking. The combustion process is:

$$\begin{align*} \mathrm{CH_4}(g)+2\;\mathrm{O_2}(g)&\rightarrow\mathrm{CO_2}(g)+2\;\mathrm{H_2O}(l) \end{align*}$$

If 15.0 moles of $\mathrm{CH_4}$ react with oxygen, what is the volume of $\mathrm{CO_2}$ (in liters) produced at 23.0 °C and 0.985 atm?

Find the total pressure for a mixture that contains four gases with partial pressures of 5.00 kPa, 4.56 kPa, 3.02 kPa, and 1.20 kPa.

The ${ }^{235} \mathrm{U}$ isotope undergoes fission when bombarded with neutrons. However, its natural abundance is only $0.72$ percent. To separate it from the more abundant ${ }^{238} \mathrm{U}$ isotope, uranium is first converted to $\mathrm{UF}_6$, which is easily vaporized above room temperature. The mixture of the ${ }^{235} \mathrm{UF}_6$ and ${ }^{238} \mathrm{UF}_6$ gases is then subjected to many stages of effusion. Calculate how much faster ${ }^{235} \mathrm{UF}_6$ effuses than ${ }^{238} \mathrm{UF}_6$.

Chlorine gas reacts with fluorine gas to form chlorine trifluoride. $\mathrm{Cl}_{2}(g)+3 \mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{ClF}_{3}(g)$ A 2.00-L reaction vessel, initially at 298 K, contains chlorine gas at a partial pressure of 337 mmHg and fluorine gas at a partial pressure of 729 mm Hg. Identify the limiting reactant and determine the theoretical yield of $ClF_3$ in grams.

A compound has the empirical formula SF4. At 20 degree C, 0.100 g of the gaseous compound occupies a volume of 22.1 mL and exerts a pressure of 1.02 atm. What is the molecular formula of the gas?

What pressure will be required for neon at 30 degree C to have the same density as nitrogen at 20 degree C and 1.0 atm?

At 27 degree C, 10.0 moles of a gas in a 1.50-L container exert a pressure of 130 atm. Is this an ideal gas?

A 2.10-L vessel contains 4.65 g of a gas at 1.00 atm and 27.0 degree C. (a) Calculate the density of the gas in grams per liter. (b) What is the molar mass of the gas?

Which halogens are gases at room temperature and pressure? (1) chlorine and fluorine (2) chlorine and bromine (3) iodine and fluorine (4) iodine and bromine

A compound has the empirical formula $\mathrm{SF_4}$. At 20 °C, 0.100 g of the gaseous compound occupies a volume of 22.1 mL and exerts a pressure of 1.02 atm. What is the molecular formula of the gas?

Dry ice is solid carbon dioxide. A 0.050-g sample of dry ice is placed in an evacuated 4.6-L vessel at 30 °C. Calculate the pressure inside the vessel after all the dry ice has been converted to $\mathrm{CO_2}$ gas.

The temperature of 2.5 L of a gas initially at STP is raised to 250 degree C at constant volume. Calculate the final pressure of the gas in atm.