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## Chapter 13 - GASES

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**Properties of gases**• 1. are compressible • 2. occupy all available volume • 3. one mole of gas at 0oC and 1 atm pressure occupies 22.4 liters • 4. gases have mass • 5. can move (diffuse) through each other • 6. gases exert pressure • 7. A gases pressure depends on temperature**The properties of gases are explained by the**Kinetic-Molecular Theory of Gas • Gases are composed of small particles that have mass. • The particles are far apart (compared to the volume of the particles). • They are in rapid motion and have perfectly elastic collisions**The collisions of gas molecules against the sides of**a container create pressure. • The kinetic energy of a gas molecule is given by the equation KE = 1/2mv2 • Temperature is a measure of the average KE of gas molecules. • Lighter gas molecules, such as H2 move faster than heavier gas molecules (O2) if they are at the same temperature.**Measuring Gas - page 424**• The four variables needed to completely describe a sample of gas are the amount of gas (moles), volume, temperature and pressure. • All types of gas behave the same so we do not need to know what type of molecule the gas is.**Atmospheric pressure - the weight of the atmosphere pushing**on us. • 1 atm = 14.7 lb/in2 = 101,325 N/m2 (Pa) = 760 mm Hg • 760 mm Hg corresponds to the weight per square inch that atmospheric pressure can support**If mercury is 13.6 times more dense than water, what is the**maximum water height that the atmosphere can support? • A well any deeper than this must use a submersible pump. • Start assignment (practice problems 1 & 2 page 427)**Enclosed gases**• Absolute pressure - the true pressure of a gas (barometric pressure is an example) • Gage pressure - the difference in pressure between the trapped gas and the atmosphere (tire pressure is an example).**Manometer**• A manometer can give the true pressure of a gas by adding or subtracting the height of the Hg column from the atmospheric pressure. • Draw a diagram and solve practice problems 3 and 4 on page 429.**Gas Laws**• One mole of gas = 22.4 liters at STP (The standard temperature and pressure for gases is 0oC and 1 atm.)**Boyle’s Law (pressure - volume**relationship) • If we have a given quantity of gas (moles) and the temperature is kept constant while the pressure and volume are changed, PV = constant (k). • P1V1 = P2V2 • The pressure times volume before the change is equal to the pressure times volume after the change.**Do sample problem 3 page 433**• Assignment continued (practice problems 5,6 page 434)**Charles’s Law (temperature - volume relationship)**• If the pressure is kept constant, the temperature and volume are directly proportional. As temperature increases, the volume also increases. • Experimental volume -temperature data can be graphed and the line extended (extrapolated) to zero volume. • V1T2 = V2T1**Absolute temperature scale**• There are no negative temperatures • Absolute zero is the coldest possible temp • (Kelvin scale, K) - nature’s temperature scale • K = oC + 273.15**As the temperature of a gas is changed, the volume of gas**will be changed by a ratio of the initial and final temperature in oK. • Example problem: What will be the new volume if 2 L of gas is heated from 100oC to 300oC? • V1T2 = V2T1**Show how to solve sample problem 4 on page 438.**• Assignment continued - page 438 practice problems 7,8**Avogadro’s Law**• Equal volumes of gases at the same pressure and temperature will have equal number of gas particles. • Three moles of gas will occupy three times the volume as one mole of gas.**Dalton’s Law of Partial Pressures**• The sum of the partial pressures of all the components in a gas mixture is equal to the total pressure of the gas mixture. • PT = pa + pb + pc + pd + - - - - - • Problem assignment continued (page 440 practice p. 9 & 10**THE IDEAL GAS LAW**• PV = nRT • R = 8.314 J/(mol K) • R = .0821 (atm L)/(mole K) • R = 8.314 (Pa m3)/(mol K) • Use the R that has the correct units for the problem • Problem assignment continued - page 443 practice 11, 12**Real gases**• Real gases can deviate from the ideal gas equation at very high pressure and at very low temperatures. This is because of the slight attractive forces between real gases.**“Free-style Gas Calculations**• Treat every problem as a conversion problem with the ratio of the change being a correction factor. • You must decide if each change will result in an increase or decrease.**Example Problem: What will be the new volume of 6 ft3 of**gas if the following changes are made? • 2.4 atm changed to 6.3 atm • Heated from 20 oC to 300 oC • 5.2 moles changed to 12.6 moles**Chapter 13 assignment**• Chapter questions on pages 452-453 (1– 19, 23, 24) • Problem Bank problems on pages 454 – 455 (27, 29, 31, 37, 39, 41, 42, 44, 52, 54) note that the answers are on page 944 • You must show your work for credit.