Ch. 12 The Behavior of Gases Ch The Properties of Gases Ch Factors Affecting Gas Pressure Ch The Gas Laws Ch Ideal Gases Ch Gas Molecules: Mixtures and Movements

Ch The Properties of Gases Kinetic theory revisited Gas particles are so small in relation to the distances between them that their mass is considered to be insignificant Explains the importance of gas compressibility No attractive or repulsive force exists between gas particles Explains why gases expand to fill their containers Gas particles move in constant, random motion; in straight paths and independently of each other Also, all collisions are perfectly elastic

Ch The Properties of Gases Variables that describe a gas Pressure (P) – measured in kPa Volume (V) – measured in liters Temperature (T) – measured in Kelvin's Number of moles (n) – measured in moles

Ch Factors Affecting Gas Pressure Amount of gas The greater the amount of gas, the greater the number of particles that can collide Increasing temperature or pressure will increase the number of collisions Volume Reducing volume increases pressure (a linear relationship)

Ch Factors Affecting Gas Pressure Temperature Raising the temperature of a gas increases the pressure The faster moving particles collide with the walls of the container more frequently Temperature and pressure is also a linear relationship In contrast, decreasing the temperature of a gas decreases the pressure

Ch The Gas Laws The pressure-volume relationship Boyle’s law Describes the effect of pressure on the volume of a contained gas while temperature remains constant As pressure goes up, volume goes down and as pressure goes down, volume goes up P 1 x V 1 = P 2 x V 2 The resulting graph shows an inverse relationship (inversely proportional)

Ch The Gas Laws The temperature-volume relationship Charles’ Law Describes the effect of temperature on the volume of a gas, while pressure remains constant As temperature goes up, volume goes up, and as temperature goes down, volume goes down The resulting graph shows direct relationship (directly proportional) Can only be measured over a limited range because at low temperatures gases condense into liquids (importance of this was recognized by Lord Kelvin) V 1 T 2 = V 2 T 1

Ch The Gas Laws The temperature-pressure relationship Gay-Lussac’s Law The pressure of a gas is directly proportional to the Kelvin temperature if the volume remains constant P 1 T 2 = P 2 T 1 The combined gas law Combines the three gas laws The other laws can be obtained from this law by holding one quantity constant P 1 V 1 T 2 = P 2 V 2 T 1

Ch Ideal Gases Ideal gas law Involves all 4 variables that affect a gas Temperature, pressure, volume, and number of moles T, P, and V all depend on the number of moles in the sample of gas (P x V)(T x n) = a gas constant This constancy holds for “ideal gases” This constant is known as R R = 8.31 (L x kPa) / (mol x K) R =.0821 (L x atm) / (mol x K)

Ch Ideal Gases Ideal gas law The formula for the ideal gas law is PV=nRT The ideal gas law and kinetic theory The kinetic theory and the gas laws assume that all gases are ideal gases True “ideal gases” do not exist Particles could have no volume There could be absolutely no force between molecules

Ch Ideal Gases Departures from the ideal gas law Real gases can be liquefied and sometimes solidified, ideal gases cannot (based on the assumption that gases have no volume) Real gases behave like ideal gases except at very low temperatures or very high pressures At low temperatures, the particles slow down, allowing intermolecular forces to play a role At high pressures, the particles are forced together to the point that they can no longer be compressed

Ch Gas Molecules: Mixtures and Movements Avagadro’s hypothesis Equal volumes of gases at the same temperature and pressure contain equal numbers of particles Size of the particles does not matter, since there is a large amount of empty space between particles Also takes into account the fact that these gases will have the same kinetic energy and are contained in equal volumes

Ch Gas Molecules: Mixtures and Movements Dalton’s law of partial pressures At constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases P total = P 1 + P 2 + P 3 + … The fractional contribution to pressure exerted by each gas in a mixture does not change as the temperature, volume or pressure changes

Ch Gas Molecules: Mixtures and Movements Graham’s Law Diffusion is the tendency of molecules to move from an area of greater concentration to an area of lower concentration, until equilibrium is reached Thomas Graham did work on diffusion, as well as effusion Effusion is the process by which a gas escapes through a tiny hole in it’s container

Ch Gas Molecules: Mixtures and Movements Graham’s law of effusion The rate of effusion of a gas is inversely proportional to the square root of the gas’s molar mass Related to the KE of an object If two bodies of different masses have same kinetic energy, the lighter body will move faster (at the same temperature) Gases with a lower molar mass will effuse faster Rate A / Rate B = square root of molar mass B / square root of molar mass A