CHAPTER 5 ZUMDAHL

GASES

5.1 PRESSURE

• A gas uniformly fills a container and mixes completely with any other gas• Pressure is one of a gases most obvious

properties• See collapsing can experiment• A device to measure atmospheric pressure is a

barometer invented in 1642 by Evangelista Torricelli made by inverting a column of Mercury. The units of mmHg is sometimes called Torr in honor of Torricelli

TORRICELLIAN BAROMETER

UNITS OF PRESSURE

• A manometer is commonly used to measure pressure based upon the height of mercury it can support.• 1atmosphere =

760mmHg = 760 torr = 101.3kPa • A pascal (Pa) is a

N/m2

5.2 THE GAS LAWS

• Boyles Law – Robert Boyle showed that PV=k. In other words there is an inverse relationship between pressure and volume

• Charles Law – Jacques Charles was the first person to fill a balloon with Hydrogen gas and make a solo flight. He found that volume is directly proportional to temperature. V = bT. Also if V and T are ploted on the Kelvin scale they all go back to the same temperature for volume. That temperature is absolute zero 0K = -273oC

• Avogadro’s Law. He showed that equal volumes of gas at the same temperature and pressure contain equal numbers of particle. V = an

BOYLES LAW

CHARLES LAW

AVOGADRO’S LAW

• Each of these balloons contains 1L of gas at 25oC and 1atm pressure.

5.3 THE IDEAL GAS LAW

• The three gas laws can be combined when all the proportionality constants are combined into one constant R (the universal gas constant)• The resulting equation is PV = nRT and is known

as the ideal gas law• Where pressure is in atm and volume in L then

the gas constant R = 0.08206 L atm/K mol

5.4 GAS STOICHIOMETRY

• The volume 1 mole of an ideal gas will occupy at STP is 22.4L

GAS DENSITY AND MOLAR MASS

• One very important use of the ideal gas law in in the calculation of molar mass of a gas from its measured density.

• Molar mass = dRT/P

5.5 DALTONS LAWS OF PARTIAL PRESSURES

• Pressure is directly affected by the number of moles of gas.• The total pressure

is the sum of each of the partial pressures

• PT = P1+P2 ……

MOLE FRACTION (Χ)

• The ratio of the number of moles in a given component in a mixture to the total number of moles in the mixture.

• So χ1= n1/ntotal = P1/Ptotal

• This can be rearranged to give

P1 = χ1 x Ptotal

COLLECTING A GAS OVER WATER

• Collecting over water involves a mixture of the product gas and water vapor. The amount of water is called the vapor pressure of water

5.6 KINETIC MOLECULAR THEORY

1. The particles are so small compared with the distances between them that the volume of individual particles is assumed to be negligible (zero)

2. The particles are in constant motion. The collisions with the walls of the container are the cause of the pressure

3. The particles are assumed to exert no forces on each other. They are assumed neither to attract nor to repel each other.

4. The average kinetic energy of a collection of gas particles is assumed to be directly proportional to the Kelvin temperture of the gas

PATH OF A GAS PARTICLE

KINETIC ENERGY AND TEMPERATURE

THE DISTRIBUTION OF VELOCITIES

5.7 EFFUSION AND DIFFUSION

• Diffusion describes the mixing of gases. A smell spreads throughout the room as a result of the movement of the gas particles spreading a far apart as they can from areas of high concentration to areas of low concentration.• Effusion describes the passage of a gas through a

tiny orifice into an evacuated chamber. The rate of effusion measures the speed at which the gas is transferred into the chamber

GRAHAMS LAW OF EFFUSION

DIFFUSION

5.8 REAL GASES

At different Pressures At different Temperatures

INTERACTIONS

• At high temperature collisions more elastic.• At low

concentration fewer interactions

VAN DER WAALS EQUATION

VAN DER WAALS CONSTANTS

5.10 CHEMISTRY OF THE ATMOSPHERE

• Make up of gases in the air• Air pollution• Photochemical smog• The effect of acid rain• Use of scrubbers to remove sulfur dioxide.