Basics of Gas Laws and Kinetic Theory
The fundamental concepts of gas laws, the nature of gases, kinetic theory, ideal gases, gas pressure, and more. Learn about the behavior of gases, ideal gas properties, gas pressure measurements, and standard conditions.
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Presentation Transcript
The Nature of Gases Gases expand to fill their containers Gases are fluid they flow Gases have low density 1/1000 the density of the equivalent liquid or solid Gases are compressible Gases effuse and diffuse
Kinetic Theory Gases are composed of small, separate particles called molecules Move in constant motion All collisions between particles are perfectly elastic The molecules of a gas display no attraction or repulsion The average kinetic energy of the molecules is directly proportional to Kelvin temperature of the gas
Ideal Gases Ideal gases are imaginary gases that perfectly fit all of the assumptions of the kinetic molecular theory. Gases consist of tiny particles that are far apart relative to their size. Collisions between gas particles and between particles and the walls of the container are elastic collisions No kinetic energy is lost in elastic collisions
Ideal Gases (continued) Gas particles are in constant, rapid motion. They therefore possess kinetic energy, the energy of motion There are no forces of attraction between gas particles The average kinetic energy of gas particles depends on temperature, not on the identity of the particle.
Gas Pressure Pressure= Force/Area Atmospheric pressure: The pressure the earth s atmosphere exerts due to its weight Barometer: Instrument used to measure atmospheric pressure Invented by Toricelli Baro= weight Meter= measure Normal Atmospheric Pressure: Also called standard pressure 1 atm = 760 mm Hg = 760 torr 1 atm= 101.3 kPa
Gas Pressure Barometer: Instrument used to measure atmospheric pressure Invented by Toricelli 760 mm Hg at 1 atm Normal Atmospheric Pressure: Also called standard pressure 760 mmHg 760 torr 1 atm 101.3 kPa 273 K
Gas Pressure STP Standard temp and pressure P= 1 atm, 760 torr T= 0 C, 273 K Molar volume of ideal gas is 22.4 L at STP Manometer Instrument used to measure gas pressure U-shaped tube partially filled with mercury One end open to confined gas One end open to atmosphere
Converting Celsius to Kelvin Gas law problems involving temperature require that the temperature be in KELVINS! Kelvins = C + 273 C = Kelvins - 273
Boyles Law When temperature is held constant, the pressure and volume of a gas are inversely proportional If P goes up, V goes down
Charless Law When pressure is held constant, the volume and temperature of a gas are directly proportional If V goes up, T goes up
Gay Lussacs Law When volume is held constant, the pressure and temperature of a gas are directly proportional If P goes up, T goes up
The Combined Gas Law The combined gas law expresses the relationship between pressure, volume and temperature of a fixed amount of gas. P V P V = 1 T 1 2 T 2 1 2 Boyle s law, Gay-Lussac s law, and Charles law are all derived from this by holding a variable constant.
Health Note When a scuba diver is several hundred feet under water, the high pressures cause N2 from the tank air to dissolve in the blood. If the diver rises too fast, the dissolved N2 will form bubbles in the blood, a dangerous and painful condition called "the bends". Helium, which is inert, less dense, and does not dissolve in the blood, is mixed with O2 in scuba tanks used for deep descents.
Daltons Law of Partial Pressures Each gas exerts the same pressure it would if it alone was present at the same temperature Gas collected over water- pressure in the container is the sum of the vapor pressure of the gas and the water s vapor pressure Subtract the water vapor pressure from the total pressure to obtain the pressure of the gas alone
Solve This! A student collects some hydrogen gas over water at 20 degrees C and 768 torr. What is the pressure of the H2 gas? 768 torr 17.5 torr = 750.5 torr
Ideal Gas Law PV = nRT P = pressure in atm V = volume in liters n = moles R = proportionality constant = 0.08206 L atm/ mol T = temperature in Kelvins Holds closely at P < 1 atm
Avogadros Law Equal volumes of different gases, at the same temp and pressure contain the same number of molecules How would the number of molecules in 2 liters of hydrogen gas compare with the number of molecules in 2 liters of oxygen gas at the same temperature and pressure? Why is 22.4 liters called the molar volume of gas?
Avogadros Law Equal volumes of gases at the same T and P have the same number of molecules. V = n (RT/P) V and n are directly related. twice as many molecules
Gas Density molar mass molar volume mass volume = = Density so at STP molar mass 22.4 L Density =
Density and the Ideal Gas Law Combining the formula for density with the Ideal Gas law, substituting and rearranging algebraically: MP D RT T = Temperature in Kelvins M = Molar Mass P = Pressure R = Gas Constant =
