# Chemistry 1000 A Fall 2004 Fourth assignment

advertisement ```Chemistry 1000 A
Fall 2004
Fourth assignment
Gases and the Kinetic Molecular Theory
Responsible sections of Oxtoby, 4th edition: Chapter 5, in its entirety.
Self-Study Assignment
Review Questions
Learn the Ideal Gas Law and how to manipulate it for various kinds of problems. Know how to choose and use values of the
constant R.
Learn how the Gas Law can be combined with the formulas for molar mass and density. Practice problems that relate gas
densities to their identity via their molar mass.
Practice problems that combine stoichiometry with the ideal gas law to combine solids, solutions and gases into
comprehensive stoichiometric relationships.
Learn Dalton’s law of partial pressures and practice problems involving gas mixtures. Learn the definition and use of the
mole fraction, a new unit of concentration.
Learn the postulates of the kinetic molecular theory of gases, and the derivation of the ideal gas law from the kinetic theory.
Learn the meaning of temperature.
Know the concept of mean free path for gas molecules and its use in calculating gas effusion.
Learn the van der Waals equation of state and be able to solve problems involving real gases.
Study Questions from the end of Chapter 5
a) From the main section of Problems, all the problems are relevant. As usual, you are strongly advised to try at least
several different-looking problems from each section.
b) From the Additional Problems, all are relevant but are fairly advanced (good for review!)
c) Do the Cumulative Problem as it ties together much of what we have learned in the course to date.
Problem Set #4
1. The density of air 20 km above the earth’s surface is 92 g m–3. The pressure of the atmosphere is 42 mm Hg and the
temperature is –63 &deg;C. (a) What is the average molar mass of the atmosphere at this altitude? (b) If the atmosphere at
this altitude is only O2 and N2, what is the mole fraction of each gas?
2.
A xenon fluoride can be prepared by heating a mixture of Xe and F2 gases to a high temperature in a pressure-proof
container. Assume that xenon gas is added to a 0.25-L container until its pressure is 0.12 atm at 0.0 &deg;C. Fluorine gas is
then added until the total pressure is 0.72 atm at 0.0 &deg;C. After the reaction is complete, the xenon is consumed
completely and the pressure of the F2 remaining in the container is 0.36 atm at 0 &deg;C. What is the empirical formula of
the compound prepared from Xe and F2?
3.
Acetylene can be made by allowing calcium carbide to react with water:
CaC2(s) + 2 H2O(l) &AElig; C2H2(g) + Ca(OH)2(s)
You react 2.65 g of CaC2 with excess water. If you collect the acetylene over water, as shown in the diagram below,
and find that the gas (acetylene and water vapor) has a volume of 795 mL at 25.2 &deg;C at a barometric pressure of 735.2
mm Hg, what is the percent yield of acetylene?
1
4.
A sample of gas was contained in a 245-mL flask at a temperature of 23.5 &deg;C; the gas pressure is found to be 48.5 mm
Hg. The gas is moved to a new flask, which is immersed in ice and water, and which has a volume of 68 mL. What is
the pressure of the gas in the smaller flask at the new temperature?
5.
Suppose you have two pressure-proof steel cylinders of equal volume, one containing CO and the other acetylene,
C2H2.
(a) If you have 1 kg of each compound, in which cylinder is the pressure greater at 25&deg;C?
(b) Now suppose the CO cylinder has twice the pressure of the acetylene cylinder at 25&deg;C. Which cylinder contains the
greater number of molecules?
6.
Chlorine trifluoride, one of the most reactive compounds known, is made by the reaction of chlorine and fluorine
according to the reaction:
Cl2 (g) + 3 F2(g) &AElig; 2 ClF3 (g)
Assume that you mix 0.71 g of Cl2 with 1.00 g of F2 in a 258-mL flask at 23 &deg;C. What are the partial pressures of each
of the two reactants before reaction? What is the partial pressure of the product and any leftover reactant after reaction
in the 258-mL flask at 23&deg;C? What is the total pressure in the flask at the end of the reaction? Is the final total pressure
more or less than the initial total pressure?
7.
You are given a solid mixture of (reactive) sodium nitrite and (inert) sodium chloride and are asked to analyze for the
amount of sodium nitrite present. To do so you dissolve the mixture of solids, allow the solution to react with sulfamic
acid, HSO3NH2, in water according to the equation:
NaNO2 (aq) + HSO3NH2 (aq) → NaHSO4 (aq) + H2O (l) + N2 (g)
What is the weight percent of sodium nitrite in 1.232 g of the solid mixture of sodium nitrite and sodium chloride if
reaction with sulfamic acid produces 295 mL of N2 gas? The gas is collected over water at a temperature of 21&deg;C and
with the barometric pressure equal to 736.0 mmHg.
8.
You have two flasks of equal volume. Flask A contains H2 at 0&deg;C and 1 atm pressure. Flask B contains CO2 gas at 25
&deg;C and 2 atm pressure. Compare the two gases with respect to each of the following:
(a) Average kinetic energy per molecule
(b) Average molecular velocity
(c) Number of molecules
(d) Mass of gas.
9.
Calculate the rms speed for CO molecules at 25 &deg;C. What is the ratio of this speed to that of Ar atoms at the same
temperature?
10. A sample of uranium fluoride is found to effuse at the rate of 17.7 mg/h. Under comparable conditions, gaseous I2
effuses at the rate of 15.0 mg/h. What is the molar mass of the uranium fluoride? Identify the molecular formula of the
fluoride.
11. You want to store 165 g of CO2 gas in a 12.5 L tank at room temperature (25&deg;C). Calculate the pressure the gas would
have using (a) the ideal gas law and (b) the van der Waals equation.
2
```