6. A 0.4550 g sample of an unknown organic compound with the empirical formula CH2O was placed into a 150.0 ml vesseland was vaporized into a gas. At 175.0°C, the pressure of the vaporized compound was measured at 941.1 torr.(a) Determine the molar mass of the compound(b) Determine the molecular formula of the compound.

Given:Mass of the gas = m = 0.4550 g (4 sig fig)…

Answered: 6. A 0.4550 g sample of an unknown…

Chemistry for Engineering Students

4th Edition

ISBN:9781337398909

Author:Lawrence S. Brown, Tom Holme

Publisher:Lawrence S. Brown, Tom Holme

Chapter5: Gases

Section: Chapter Questions

Problem 5.101PAE

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You have two pressure-proof steel cylinders of equal volume, one containing 1.0 kg of CO and the other containing 1.0 kg of acetylene, C2H2. (a) In which cylinder is the pressure greater at 25 C? (b) Which cylinder contains the greater number of molecules?
You have a gas, one of the three known phosphorus-fluorine compounds (PF3, PF3, and P2F4). To find out which, you have decided to measure its molar mass. (a) First, yon determine that the density of the gas is 5.60 g/L at a pressure of 0.971 atm and a temperature of 18.2 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 rales of effusion of the unknown gas and CO2. You find that CO2 effuses at a rate of 0.050 mol/min, whereas the unknown phosphorus fluoride effuses at a rate of 0.028 mol/min. Calculate the molar mass of the unknown gas based on these results.
93 The complete combustion of octane can be used as a model for the burning of gasoline: 2C8H18+25O216CO2+18H2O Assuming that this equation provides a reasonable model of the actual combustion process, what volume of air at 1.0 atm and 25°C must be taken into an engine to burn 1 gallon of gasoline? (The partial pressure of oxygen in air is 0.21 atm and the density of liquid octane is 0.70 g/mL.)
Liquid oxygen was first prepared by heating potassium chlorate, KClO3, in a closed vessel to obtain oxygen at high pressure. The oxygen was cooled until it liquefied. 2KClO3(s)2KCl(s)+3O2(g) If 171 g of potassium chlorate reacts in a 2.70-L vessel, which was initially evacuated, what pressure of oxygen will be attained when the temperature is finally cooled to 25C? Use the preceding chemical equation and ignore the volume of solid product.
A typical barometric pressure in Redding. California, is about 750 mm Hg. Calculate this pressure in atm and kPa.
105 The decomposition of mercury(II) thiocyanate produces an odd brown snake-like mass that is so unusual the process was once used in fireworks displays. There are actually several reactions that take place when the solid Hg(SCN)2 is ignited: 2Hg(SCN)2(s)2HgS(s)+CS2(s)+C3N4(s)CS2(s)+3O2(g)CO2(g)+2SO2(g)2C3N4(s)3(CN)2(g)+N2(g)HgS(s)+O2(g)Hg(l)+SO2(g) A 42.4-g sample of Hg(SCN)2 is placed into a 2.4-L vessel at 21°C. The vessel also contains air at a pressure of 758 torr. The container is sealed and the mixture is ignited, causing the reaction sequence above to occur. Once the reaction is complete, the container is cooled back to the original temperature of 21°C. (a) Without doing numerical calculations, predict whether the final pressure in the vessel will be greater than, less than, or equal to the initial pressure. Explain your answer. (b) Calculate the final pressure and compare your result with your prediction. (Assume that the mole fraction of O2 in air is 0.21.)
A typical barometric pressure in Kansas City is 740 torr. What is this pressure in atmospheres, in millimeters of mercury, and in kilopascals?
5-114 Carbon dioxide gas, saturated with water vapor, can be produced by the addition of aqueous acid to calcium carbonate based on the following balanced net ionic equation: (a) How many moles of wet CO (g), collected at 60.°C and 774 torr total pressure, are produced by the complete reaction of 10.0 g of CaCO3 with excess acid? (b) What volume does this wet CO2 occupy? (c) What volume would the CO2 occupy at 774 torr if a desiccant (a chemical drying agent) were added to remove the water? The vapor pressure of water at 60.°C is 149.4 mm Hg.
5-107 If 60.0 g of NH3 occupies 35.1 L under a pressure of 77.2 in. Hg, what is the temperature of the gas, in °C?
Many nitrate salts can be decomposed by heating. For example, blue, anhydrous copper(II) nitrate produces the gases nitrogen dioxide and oxygen when heated. In the laboratory, you find that a sample of this salt produced a 0.195-g mixture of gaseous NO2 and O2 with a total pressure of 725 mm Hg at 35 C in a 125-mL flask (and black, solid CuO was left as a residue). What is the average molar mass of the gas mixture? What are the mole fractions of NO2 and O2 in the mixture? What amount of each gas b in the mixture? Do these amounts reflect the relative amounts of NO2 and O2 expected based on the balanced equation? Is it possible that the fact that some NO2 molecules combine to give N2O4 plays a role? Heating copper(II) nitrate produces nitrogen dioxide and oxygen gas and leaves a residue of copper(ll) oxide.
If equal masses of O2 and N2 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 N2 is greater than that in the flask containing O2. (b) There are more molecules in the flask containing O2 than in the flask containing N2.
You have an equimolar mixture of the gases SO2 and O2, along with some He, in a container fitted with a piston. The density of this mixture at STP is 1.924 g/L. Assume ideal behavior and constant temperature and pressure. a. What is the mole fraction of He in the original mixture? b. The SO2 and O2 react to completion to form SO3. What is the density of the gas mixture after the reaction is complete?

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