PHYSICAL CHARACTERISTICS OF GASES. Chapter 10

Transcription

1 PHYSICAL CHARACTERISTICS OF GASES Chapter 10

2 Section 10-1 Objectives: State the kinetic-molecular theory of matter, and describe how it explains certain properties of matter. List the 5 assumptions of the kinetic-molecular theory of gases. Define the terms ideal gas and real gas. Describe these characteristic properties of gases: n Expansion n Density n Fluidity n Compressibility n Diffusion n Effusion Describe the conditions under which a real gas deviates from ideal behavior.

3 Terms in this Chapter Expansion Pressure Barometer Manometer Density Fluidity Diffusion Effusion Condensation temperature

4 Kinetic-Molecular Theory Particles of matter are always in motion! For gases, the theory provides a model of what is called an ideal gas n An imaginary gas that perfectly fits all the assumptions of the kinetic-molecular theory.

5 Background Kinetic-Molecular Theory Johann van Helmont 1662 Invented the term, gas, to describe the most energetic phase of matter. The term, gas, was derived from the Greek chaos, meaning original matter of the earth. Vapor a gas produced from a substance that is a solid or liquid under normal conditions. It is formed if the vapor pressure is equal to that of the atmosphere.

6 Kinetic-Molecular Theory of Gases Based on 5 assumptions 1. Gases consist of large numbers of tiny particles that are far apart relative to their size. Most of the volume occupied by a gas is empty space 2. Collisions between gas particles and between particles and container walls are elastic collisions. An elastic collision is one in which there is NO net loss of kinetic energy. Kinetic energy is transferred between 2 particles during collisions and is conserved as long as temperature is constant.

7 K-M Theory of Gases (assumptions) 3. Gas particles are in continuous, rapid, random motion. Therefore, they possess kinetic energy, which is energy of motion. 4. There are no forces of attraction or repulsion between gas particles. (behave as billiard balls) 5. The average kinetic energy of gas particles is constant and depends on the temperature of the gas. m=mass KE = ½ mv 2 v=speed

8 Kinetic-Molecular Theory and the Nature of Gases The kinetic-molecular theory applies only to ideal gases. Although ideal gases do not actually exist, many gases behave nearly ideally if pressure is not very high or temperature is not very low.

9 Expansion Gases have no definite shape or volume. They completely fill any container and take its shape. A gas transferred from a 1-liter vessel to a 2-liter vessel will quickly expand to fill the entire 2-liter volume. Gas particles move rapidly in all directions, without significant attraction or repulsion between them.

10 Fluidity Gas particles glide easily past one another. This ability to flow causes gases to behave similarly to liquids. Because liquids and gases flow, they are both referred to as fluids.

11 Low Density The density of a substance in the gas state is about 1/1000 the density of the same substance in the liquid or solid state. For gases, density is a measure of mass (number of molecules) per volume, g/l.

12 Compressibility During compression, the gas particles, which are initially very far apart, are crowded close together. The volume of a gas can be greatly decreased. Steel cylinders containing gases under pressure are widely used in industry. When they are full, they may contain 100 times as many particles of gas as would be contained under non-pressurized conditions.

13 Diffusion and Effusion Gases spread out and mix with one another, even without being stirred. Such spontaneous mixing of the particles of two substances caused by their random motion is called diffusion. Rate of diffusion depends on 3 things Speed, diameter, and the attractive forces between the particles n Hydrogen gas would diffuse faster because its particles are smaller and move faster than molecules of other gases. Effusion is a process by which gas particles pass through a tiny opening. Rates of effusion are directly proportional to the velocities of their particles. Molecules of low mass effuse faster than molecules of high mass.

14 Condensation Temperature The temperature (point) at which gas molecules form a liquid.

15 Deviations of Real Gases from Ideal Behavior Real gas a gas that does not behave completely according to the assumptions of kinetic-molecular theory 1. in real gases, a weak intermolecular attraction occurs between some gas molecules 2. London dispersion forces n Caused by motion of e - Johannes Van Der Waals -1873, and increases as number of e - increases 3. dipole-dipole interactions electrostatic attractions based on polarity of molecules.

16 Van der Waal forces Particles of real gases occupy space and exert attractive forces on each other. At very high pressures and low temperatures, the deviation may be considerable. Under these conditions, particles are closer together and their kinetic energy will be insufficient to completely overcome the attractive forces. Noble gases show ideal gas behavior. They are monatomic and nonpolar. N 2 and H 2 also exhibit ideal gas behavior. The more polar a gas s molecules are, the greater the attractive forces between them and the more the gas will deviate from ideal gas behavior. (NH 3 and H 2 O vapor)

17 TO DESCRIBE A GAS FULLY, YOU MUST STATE 4 MEASURABLE QUANTITIES: VOLUME TEMPERATURE NUMBER OF PARTICLES PRESSURE

18 Units to Know! 1. Newton (force) = kg-m/sec 2 2. Pressure as pascal = N/m 2 3. Kilopascal = 1000 pascal 4. Pressure units include: inches Hg, mm Hg, torr, atm, kpa 5. Standard = in Hg = 760 mm Hg =760 torr = 1 atm = kpa 6. STP standard temperature (0 0 C, 273 K) and pressure ( kpa) 7. K = 0 C + 273

19 Unit of Pressure Symbol Definition/ Units of Pressure relationship Pascal Pa SI pressure unit 1 Pa = 1 N/ m 2 Millimeters of mercury- mm Hg A pressure of 1 mm Hg is now called 1 torr Millimeter of mercury mm Hg Pressure that supports a 1 mm Atm atmosphere of pressure; defined as equal to 760 mercury column in a mm Hg barometer Pascal (Pa) SI unit of pressure; named for Blaise Pascal, French mathematician and philosopher who studied pressure during the 17 th century. Torr torr 1 torr = 1 mm Hg Atmosphere atm Average atmospheric pressure 1 Pascal (Pa) = the pressure exered by a force of 1 N acting on an area of 1 m 2. at sea level and 0 0 C 1 atm = 760 mm Hg Kilopascals (kpa) The standard atmosphere (1 atm) = x 10 5 Pa, or kpa. = 760 torr = x 10 5 Pa = kpa

20 Conversion of Units-Practice Calculate kpa for an 800 N force over a 0.1 meter by 0.2 meter area. N/m 2 Convert 850 mm Hg to atm. 760 mm Hg = 1 atm Convert kpa to torr. 760 torr = kpa Convert C to K K = 0 C + 273

21 Pressure The force per unit area on a surface. (P) PRESSURE = FORCE/AREA SI Unit for FORCE is the Newton. (N) It is the force that will increase the speed of a 1-kg mass by 1 meter per second each second it is applied. 1 Newton = 1kg x m/s 2 At Earth s surface, each kg of mass exerts 9.8 N of force, due to gravity. Ex. A mass of 51 kg exerts a force of 500 N (51 x 9.8) on Earth s surface.

22 Atmospheric Pressure The atmosphere, because of its mass, exerts pressure on Earth. At sea level, atmospheric pressure is about equal to 1.03 kg mass per cm 2 of surface, or 10.1 N/cm 2 or 760 mm Hg at 0 degrees Celsius. Since the atmosphere is made up of about 78% N 2, 21% O 2, and 1% other gases such as CO 2 and Argon n Atmospheric pressure is the sum of all the individual pressures of the various gases.

23 Measuring Pressure A barometer is a device used to measure atmospheric pressure. 1 st type of barometer was introduced by Evangelista Torricelli in the early 1600s.

24 Standard Temperature & Pressure Necessary to know the temperature and pressure at which volumes of gases are measured in order to compare volumes of gases. STP - Standard conditions (determined by scientists) are exactly 1 atm pressure and 0 0 C.

25 GAS LAWS Chapter 10

26 Gas Laws Simple mathematical relationships between the volume, temperature, pressure, and amount of a gas. Boyles s Law Charles s Law Gay-Lussac s Law Combined Gas Law Dalton s Law of Partial Pressures

27 Boyle s Law Pressure-Volume Relationship 1662-Robert Boyle discovered that gas pressure and volume are related mathematically. Doubling the pressure on a sample of gas at constant temperature reduces its volume by ½. Likewise, reducing the pressure on a gas by ½ allows the volume of the gas to double. Boyle s Law-a volume of a fixed mass of gas varies inversely with the pressure at constant temperature. n V = k x 1/P OR PV = k n P 1 V 1 = P 2 V 2 ( used to compare changing conditions for a gas where P 1 and V 1 stand for initial conditions, and P 2 and V 2 stand for new conditions.)

28 Boyle s Law example A sample of oxygen gas has a volume of 150. ml when its pressure is atm. What will the volume of the gas be at a pressure of atm if the temperature remains constant?

29 Charles s Law: Volume-Temperature Relationship 1787-French scientist Jacques Charles showed that all gases expand to the same extent when heated through the same temperature interval. He found that the volume changes by 1/273 of the original volume for each Celsius degree. Ex. Raising the temp. to 1 0 C causes the gas volume to increase by 1/273 of the volume it had a 0 0 C.

30 Charles s Law The volume of a fixed mass of gas at constant pressure varies directly with the Kelvin temperature. V 1 and T 1 represent initial conditions V 2 and T 2 new conditions Absolute zero = C or 0 Kelvin. K = C

31 Sample Problem Problem A sample of neon gas occupies a volume of 752 ml at 25 0 C. What volume will the gas occupy at 50 0 C if the pressure remains constant?

32 Gay-Lussac s Law: Pressure-Temperature Relationship Joseph Gay-Lussac recognized in 1802 that for every kelvin of temperature change, the pressure of a confined gas changes by 1/273 of the pressure at 0 0 C. Gay-Lussac s Law the pressure of a fixed mass of gas at constant volume varies directly with the Kelvin temperature.

33 Sample Problem The gas in an aerosol can is at a pressure of 3.00 atm at 25 0 C. Directions on the can warn the user not to keep the can in a place where the temperature exceeds 52 0 C. What would the gas pressure in the can be at 52 0 C? Remember to convert from 0 C to K.

34 Combined Gas Law All of the formulas we have learned assume that either temperature, pressure, or volume remains constant depending on the conditions and the type of problem we are trying to solve. But a gas sample often undergoes changes in temperature, pressure, and volume all at the same time, causing us to deal with 3 variables at the same time. CGL expresses the relationship between pressure, volume, and temperature of a fixed amount of gas.

35 Combined Gas Law

36 Sample Problem A helium filled balloon has a volume of 50.0 L at 25 0 C and 1.08 atm. What volume will it have at atm and C?

37 Dalton s Law of Partial Pressures John Dalton found that in the absence of a chemical reaction, the pressure of a gas mixture is the sum of the individual pressures of each gas alone. Dalton s Law of Partial Pressures the total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases. P T = P 1 + P 2 + P 3 +.

38 Gases Collected by Water Displacement P atm = P gas + P H2O Sample Problem Oxygen gas from the decomposition of potassium chlorate, KClO 3, was collected by water displacement. The barometric pressure and the temperature during the experiment were torr and C, respectively. What was the partial pressure fo the oxygen collected?