.
.

Monday February 12, 2024 Day 25 Simple Gas Laws
Textbook Readings: 5.3: The Simple Gas Laws:        Boyle’s Law, Charles’s Law        and Avogadro’s Law	Course Lectures 5.2. pdf  Video*    Gas Laws
Gas Laws	Which gas equation do I use?
Objectives 1.  Correctly identify situations where the simple gas laws are used.  (Boyle's, Charles and Combined) 2.  Use the simple gas laws to determine gas parameters (V, T, P, and n) under new conditions.
Homework Problems 25.1  What simple gas law relates volume and temperature?           What must be assumed constant when using this law? 25.2  What simple gas law relates volume and pressure?            What must be assumed constant when using this law? 25.3   What simple gas law relates volume, pressure and temperature?            What must be assumed constant when using this law? 25.4  A container holds 500. mL of CO2 at 20.° C and 742 torr. What will be the volume of the CO2           if the pressure is increased to 795 torr? 25.5  A sample of neon occupies a volume of 461 mL at STP. What will be the volume of the neon          when the pressure is reduced to 93.3 kPa? 25.6  A balloon contains 7.2 L of He. The pressure is reduced to 2.00 atm and the balloon          expands to occupy a volume of 25.1 L.          What was the initial pressure exerted on the balloon? 25.7  A container holds 50.0 mL of nitrogen (N2) at 25° C and a pressure of 736 mm Hg.           What will be its volume if the temperature increases by 35° C? 25.8  Initially, 568 cm3 of chlorine gas is at 25° C.  What is the gas's new volume at -25° C          if the pressure and moles remain constant? 25.9  Initially a gas is at a pressure of 12 atm, a volume of 23 liters and a temperature of 200. K.          What is the new volume of the gas if the pressure is increased to 14 atm and the temperature           is increased to 300. K? 25.10  Initially, 2.9 L of a gas is at a pressure of 5.0 atm and a temperature of 50.0oC.             What will be the temperature of the gas if the volume is  decreased to 2.4 L            AND  the pressure decreased to 3.0 atm? Click and drag the region below for correct answers 25.1. Charles's Law. Volume and temperature are directly proportional to each other.  V1/T1 = V2/T2                                  Assumes pressure and moles of gas (n) are constant. 25.2. Boyle's Law.     Volume and pressure are inversely proportional to each other.  P1V1 = P2V2                                          Assumes temperature and moles of gas (n) are constant. 25.3. Combined gas law:  P1V1/T1 = P2V2/T2                                 Assumes moles of gas (n) are constant. 25.4    467 mL CO2(g) 25.5    501 mL Ne(g) 25.6     7.0 atm 25.7     55.9 mL 25.8    473 cm3 25.9     29.6 L 25.10    160. K

.
.

Tuesday February 13, 2024 Day 26 Ideal Gas Law and Gas Law Applications
Textbook Readings: 5.4: The Ideal Gas Law 5.5: Applications of the Ideal Gas Law:        Molar Volume, Density, Molar Mass of a Gas	Course Lectures 5.2. pdf  Video*    Gas Laws 5.4  pdf  Video*    Gas Laws, Density                               and Molecular Weights
Ideal Gas Law Introduction	Ideal Gas Law Practice Problems
Ideal Gas Law Practice Problems with Density	Ideal Gas Law Practice Problems with Molar Mass
Objectives
Homework Problems 26.1  How many moles of gas are contained in 890.0 mL at 21.0 °C and 750.0 mm Hg pressure? 26.2 What volume would 32.0 g of NO2 gas occupy at 3.12 atm and 18.0 °C? 26.3 An amount of an ideal gas at 290.9 K has a volume of 17.05 L at a pressure of 1.40 atm.         What is the pressure of this gas sample when the volume is halved and the absolute         temperature is multipled by four? 26.4 A gas consisting of only carbon and hydrogen has an empirical formula of CH2.         The gas has a density of 1.65 g/L at 27.0 °C and 734.0 torr.         Determine the molar mass and molecular formula of the gas. 26.5    13.9 grams of an unknown noble gas is placed in a 5.00 L container at 60.0 °C.            Under these conditions, gas's pressure is 58.6 kPa.           Determine thenoble gas's molar mass and identify it. Click and drag the region below for correct answers 26.1. 0.03639 moles 26.2. 5.33L 26.3. 11.2 atm 26.4. C3H6 26.5. Molecular weight = 131.4 g/mol.   Xenon.

.
.
.

Wednesday February 14, 2024    Day 27 Gas Mixtures and Dalton's Law of Partial Pressures
Textbook Readings: 5.6: Mixtures of Gases and Partial Pressures	Course Lectures 5.3  pdf  Video*    Gas collection over water
Dalton's Law and Partial Pressures	Gas Collection Over Water
Objectives 1.  Use gram and mole amounts to determine      mole fractions and partial pressures of all      components of a gaseous mixture. 2.   Determine the vapor pressure of water at       arbitrary temperatures using the data table       at right and interpolation. 3.   Determine unknown gas pressure from       atmospheric pressure and water vapor       pressure date using Dahlton's Law. 4.   Calculate the molar mass of a gas collected       "over water."	Water Vapor Pressure (torr) At Various Temperatures (oC)
Homework Problems 27.1. A cylinder of compressed natural gas has a volume of 20.0 L and contains            1813 g of methane (CH4) and 336 g of ethane (C2H6).            Calculate the partial pressure of each gas at 22.0°C and the total pressure in the cylinder. 27.2 Venus is an inhospitable place, with a surface temperature of 560°C and a total surface            pressure of 90 atm.  The atmosphere consists of about 96 mol % CO2 and 3% mol N2,            with trace amounts of other gases, including water, sulfur dioxide, and sulfuric acid.            Calculate the partial pressures of CO2 and N2. 27.3  What is the total pressure exerted by a mixture of 2.00 g of H2  8.00 g of  N2  and 12.0 g of Ar           at 273 K in a 10.0 L vessel? 27.4  Butane gas from a flame lighter was collected over water. The loss of mass of the flame lighter          during the process was 128 mg and the volume of the collected gas was 0.0600 L.          The temperature of the room was 22.75 oC and the atmospheric pressure was 0.988 atm.              a. Use the Water Vapor Pressure table above to determine the water vapor pressure at             22.75 oC.  You will have to use pressures &  temps on either side of this value to             interpolate a pressure value.  Use slope and y = mx +b.     Guessing is not allowed.         b. Calculate the pressure of the butane gas using Dalton's law, atmospheric pressure,             and the water vapor  pressure you determined in part "a."         c.  Use the Ideal Gas Law to determine the moles of butane gas.         d.  Determine the molar mass of the gas using your mole and gram information available above.              How does your value compare to the known butane molar mass  (C4H10):  58.1222 g/mol 27.5    193 mL of O2 was collected over water on a day when the atmospheric  pressure was 762 mmHg.             The temperature for the experiment was 23.0 oC.             How many grams of oxygen gas were collected? Click and Drag below for answers 27.1    PCH4 = 137 atm,   PC2H6 = 13.4 atm,   Ptot = 151 atm           27.2    PCO2 = 86 atm    PN2 = 2.7 atm   27.3     3.53 atm         27.4     a. 20.9672 torr                    b. 729.9127 torr = 0.96041 atm                 c.  2.37327 x 10-3 mol  using T = 295.9 K             d. 53.9339 g/mol        27.5     2.48 x 10-1 g

.
.
.

Thursday February 15, 2024   Day 28 Kinetic Molecular Theory of Gases
Textbook Readings: 5.8: Kinetic Molecular Theory: A Model for Gases 5.9: Mean Free Path, Diffusion,        and Effusion of Gases	Course Lectures 5.5 pdf   Video*    Kinetic molecular theory Part 1 5.6  pdf  Video*    Kinetic molecular theory Part 2
Kinetic Molecular Theory and its Postulates	Passing Gases: Effusion, Diffusion, and the Velocity of a Gas
Objectives 1. Describe how gas phase molecular speeds     depend on temperature. 2. Describe how gas phase molecular speeds     depend on the molecular weight. 3. Calculate the root mean square velocity     of molecules at specific temperatures. 4. Determine diffusion/effusion rates for     molecular pairs 5.  Determine diffusion times for molecules      given molecular weights.	Graham's Law of Effusion Practice Problems and Examples
Homework Problems 28.1      List the five main postulates of the kinetic molecular theory.              28.2      a. What is the root mean square velocity (urms) of the molecules in a sample of oxygen gas at 0 °C                    and 100 °C in m/s and miles/hour units?               b.  How does increasing the temperature of a gas affect molecular velocity?                c.  At the same temperature, how does the urms of massive molecules compare to the                    urms of smaller, lighter molecules? 28.3      Under identical conditions of temperature and pressure, how many times faster will H2               effuse compared to CO2?                If the carbon dioxide takes 32 seconds to effuse, how long will the hydrogen take? 28.4       What is the relative rate of diffusion of NH3 compared to He?                Does NH3 diffuse faster or slower than He?               If the He takes 20.0 seconds to diffuse, how long will the NH3 take? 28.5       In your own words, what does this graph tell us about molecular velocities?               In your own words, what does this graph tell us about molecular velocities?               Click and drag the region below for correct answers 28.1   1.  Gases are made up of particles (atoms or molecules) that travel in straight lines unless                 they collide with something.  Collisions are elastic and no energy is lost.           2.   Gas is mostly empty space and the gas particles can be assumed to take up no space.           3.   Gas "pressure" is the result of particle collisions with the sides of the container.           4.   Gas phase particles don't interact (attract or stick) with eachother.           5.   The Kinetic Energy of gas particles is proportional to the Kelvin temperature of the gas. 28.2   At 0oC, urms = 461.3 m/s  (1031.9 miles/hour)    At 100 oC  urms = 539.2 m/s   (1206.7 miles/hour) 28.3   H2 effuses 4.7 times faster than CO2.     It will take the H2 gas 6.9 seconds. 28.4   Helium diffuses 2.06 X faster than NH3.     It will take NH3 41.2 seconds. 28.5    a. This graph shows how different gases have different speed distributions at the same                temperature (298 K).  All gases have the same average kinetic energy at the same temperature.                However, since KE depends on both speed and mass,  heavy gas phase particles will have                slower speeds than low mass particles.  This can be seen in the distributions as the peaks                for heavy atoms/molecules is shifted to lower speeds.                 The second graph demonstrates how the speed profile changes for a single gas species (N2)                 As the temperature goes up, the peak shifts right in the direction of higher speeds.                  Also, the peak broadens for higher temperatures.                 In all cases, the area under the curve is constant.

.
.

Friday February 16, 2024   Day 29 Real Gases and How They Differ from "Ideal" Gases
Textbook Readings: 5.10: Real Gases:       The Effects of Size and Intermolecular Forces
Non-Ideal Gases and the Van der Waals Equation
Homework Problems 29.1 Why does CCl4 have the largest "b" value in the table above? 29.2 Why does Xe have a smaller "b" value than CCl4? 29.3 Why does CCl4 have the largest "a" value in the table above? 29.4 Why does Xe have a smaller "a" value CCl4? 29.5  Which molecule is most "ideal" of those listed in the table above? 29.6  Consider  0.3000 mol of helium in a 0.2000 L container at -25 °C.            a. Should this gas behave as an ideal gas?            b. Using the. ideal gas law, calculate the gas pressure.            c.  Using the van der Waal's equation, calculate the gas pressure using the                  a and b constants in the table above.            d.  Are the pressures calculated in parts c and b above the same?                  What do the calculated pressures tell you about this gas as an "ideal gas?"   Click and drag the region below for correct answers 29.1 The CCl4 molecule is larger than any other on the table and takes     up significant space in a container resulting in non-ideal behavior. 29.2 Xe is smaller than CCl4 and takes up very little space in comparison. 29.3 Individual CCl4 molecules interact with each other more than any other     molecule on the table giving rise to non-ideal behavior. 29.4 Individual Xe molecules don't  interact with each other very much     resulting in more ideal gas-like behavior. 29.5.  He! 29. 6    a.  At low temperatures and small volumes, the gas particles are                  likely to interact with each other producing non-ideal gas behavior.                  However, helium is a very small and near ideal gas particle.              b. 30.55 atm              c. 31.59 atm

.
.
.
.
.
.
.