A mixture of cyclopropane gas, C3H6, and oxygen, O2, in a 1.00:4.00 mole ratio is used as an anaesthetic. The anaesthetic is in a 2.00 L bulb at 22.5°C with total pressure 1.060 × 105 Pa. Determine the partial pressure of O2 in Pa.

A 9.00 L tank at 4.22 °C is filled with 3.83 g of sulfur tetrafluoride gas and 7.58 g of sulfur hexafluoride gas. You can assume both gases behave as ideal gases under these conditions. Calculate the mole fraction and partial pressure of each gas, and the total pressure in the tank. Be sure your answers have the correct number of significant digits. mole fraction: sulfur tetrafluoride partial pressure: || atm mole fraction: sulfur hexafluoride partial pressure: atm Total pressure in tank: atm

The temperature of your water was 22.4 degrees Celsius. The volume of hydrogen collected was 35.3 mL. The atmospheric pressure in the lab room was 29.60 inches Hg. The difference in the water level between the beaker and the burette is 20.0 cm. What was the mass of the magnesium ribbon used? Hint: 2 HCl(aq) + Mg(aq) = H2(g) + MgCl2(aq)

1) a) Sulfuric acid, the industrial chemical produced in the greatest quantity each year, is prepared by the reaction of sulfur in the presence of oxygen and water as described in the balanced chemical equation below. Calculate the volume, in ml, of 18.0 M sulfuric acid solution that can be prepared from the reaction of 1850 g of sulfur with 6.50E4ML oxygen gas at 25°C and 0.965atm and plenty of water. 25(solid) + 302 (gas) + 2H20(liquid) > 2H:SO4(aqueous) b. Which substance is the limiting reactant? Explain how you know.

Calculate the mole fraction and partial pressure of each gas as well as the total pressure in the tank using the information below.

When hydrogen peroxide decomposes, oxygen is produced: 2H,02(aq) → 2H,0 + 0,(g)What volume of oxygen gas at 25°C and 1.00 atm is produced from the decomposition of 25.00 mL of a 30.0% (by mass) solution of hydrogen peroxide (d = 1.05g/mL)?

Chlorine pentafluoride gas is collected at -17.0 °C in an evacuated flask with a measured volume of 50.0 L. When all the gas has been collected, the pressure in the flask is measured to be 0.330 atm . Calculate the mass and number of moles of chlorine pentafluoride gas that were collected. Round your answer to 3 significant digits. mass: g x10 mole: || mol

A 8.00 L tank at 28.6 °C is filled with 11.1 g of sulfur hexafluoride gas and 6.50 g of carbon monoxide gas. You can assume both gases behave as ideal gases under these conditions. Calculate the mole fraction and partial pressure of each gas, and the total pressure in the tank. Be sure your answers have the correct number of significant digits. sulfur hexafluoride carbon monoxide mole fraction: partial pressure: mole fraction: partial pressure: Total pressure in tank: 0 atm atm atm x10 X Ś

A 6.00 L tank at 16.8 °C is filled with 2.62 g of carbon dioxide gas and 5.02 g of sulfur tetrafluoride gas. You can assume both gases behave as ideal gases under these conditions. Calculate the mole fraction of each gas. Round each of your answers to 3 significant digits. gas carbon dioxide sulfur tetrafluoride mole fraction 0 X

A 9.00 L tank at 12.6 °C is filled with 7.50 g of carbon dioxide gas and 13.2 g of dinitrogen difluoride gas. You can assume both gases behave as ideal gases under these conditions. Calculate the mole fraction and partial pressure of each gas, and the total pressure in the tank. Round each of your answers to 3 significant digits. mole fraction: x10 carbon dioxide partial pressure: || atm ? mole fraction: dinitrogen difluoride partial pressure: atm Total pressure in tank: || atm