10.1 CHARACTERISTICS OF GASES

Air behaves physically as one gaseous material

•N2 (78%), O2 (21%) and Ar (0.9%)

Only a few elements exist as gases under standard conditions

•H2, N2, O2, F2, and Cl2, the noble gases (He, Ne, Ar, Kr, Xe)

Gas molecules are relatively far apart

•Each molecule behaves largely as though the others are

not present

•Readily compressible and expansible

•Forms homogeneous mixtures with other gases

•Low density

10.1 CHARACTERISTICS OF GASES

10.2 PRESSURE

Pressure is defined as:

ATMOSPHERIC PRESSURE AND THE

BAROMETER

10.2 PRESSURE

In the 17th century, people believed

that the atmosphere had no weight

Torricelli’s experiment• Proved the atmosphere has weight

Pascal’s experiment• Measured the height of the mercury

column at two different places• Supported Torricelli’s explanation

Standard atmospheric pressure

ATMOSPHERIC PRESSURE AND THE

BAROMETER

Figure 10.2 A mercury barometer invented by Torricelli

Manometer•This device is used

to measure the differ-ence in pressure be-tween atmospheric pressure and that of a gas in a vessel.

ATMOSPHERIC PRESSURE AND THE

BAROMETER

10.2 PRESSURE

10.3 THE GAS LAW

Hypertension is abnormally high blood pressure. The usual

criterion is a blood pressure greater than 140/90.

mercury manometer or related device

closed, air-filled cuff

stethoscope

Blood pressure is re-

ported by two values• Systolic pressure: maximum

pressure (pumping)• Diastolic pressure: resting

pressure

10.3 THE GAS LAW

Pressure-volume relationship The volume of a fixed quantity of gas at constant temperature

is inversely proportional to the pressure

BOYLE’S LAW

BOYLE’S LAW10.3 THE GAS LAW

For a fixed quantity of gas at constant temperature, the vol-

ume of the gas is inversely proportional to its pressure

CHARLES’S LAW Temperature-volume relationship

The volume of a fixed amount of gas

at constant pressure is directly pro-

portional to its absolute temperature.

10.3 THE GAS LAW

AVOGADRO’S LAW Quantity-volume relationship Equal volumes of gases at the same temperature and pres-

sure contain equal numbers of molecules The volume of a gas at constant temperature and pressure is

directly proportional to the number of moles of the gas

10.3 THE GAS LAW

AVOGADRO’S LAW10.3 THE GAS LAW

At the same volume, pressure and temperature, samples of

different gases have the same number of molecules but dif-

ferent masses

10.4 THE IDEAL-GAS EQUATION

The term R is called the gas constant

R = 0.08206 L-atm/mol-K = 8.314 J/mol-K

Molar volume: the volume occupied by one mole of ideal gas

at STP (273.15K and 1 atm), 22.41 L

One mole of an ideal gas at STP occupies a volume of

22.41 L. One mole of various real gases at STP occupies

close to this ideal volume

10.4 THE IDEAL-GAS EQUATION

▲ Figure 10.11 Comparison of molar volumes at STP

Sample Exercise 10.4 Using the Ideal-Gas equation

Calcium carbonate, CaCO3(s), decomposes upon heating to give CaO(s) and

CO2(g). A sample of CaCO3 is decomposed, and the carbon dioxide is col-

lected in a 250-mL flask. After the decomposition is complete, the gas has a

pressure of 1.3 atm at a temperature of 31 °C. How many moles of CO2 gas

were generated?

Sample Exercise 10.5

The gas pressure in an aerosol can is 1.5 atm at 25 °C. Assuming that

the gas inside obeys the ideal-gas equation, what would the pressure

be if the can were heated to 450 °C?

GAS DENSITIES AND MOLAR MASS

10.5 FURTHER APPLICA-TIONS

GAS DENSITIES AND MOLAR MASS

10.5 FURTHER APPLICA-TIONS

▲ Figure 10.12 Carbon dioxide gas flows downhill be-cause it is denser than air.

GAS DENSITIES AND MOLAR MASS

10.5 FURTHER APPLICA-TIONS

VOLUMES OF GASES IN CHEMICAL RE-ACTIONS

10.5 FURTHER APPLICA-TIONS

10.6 GAS MIXTURES AND PARTIAL PRES-SURES The total pressure of a mixture of gases equals the sum of

the pressures that each would exert if it were present alone.

- Dalton’s law of partial pressure Partial pressure

•The pressure exerted by a particular component of a mix-ture of gases

Sample Exercise 10.10 Applying Dalton’s Law to Partial Pressures

A gaseous mixture made from 6.00 g O2 and 9.00 g CH4 is placed in a 15.0-L ves-

sel at 0 °C. What is the partial pressure of each gas, and what is the total pressure

in the vessel?

PARTIAL PRESSURE AND MOLE FRAC-TIONS

Each gas in a mixture behaves independently We can relate the amount of a given gas in a mixture to its

partial pressure

10.6 GAS MIXTURES AND PARTIAL PRESSURES

Sample Exercise 10.11 Relating Mole Fractions to Partial Pressures

A study of the effects of certain gases on plant growth

requires a synthetic atmosphere composed of 1.5 mol% CO2,

18.0 mol% O2, and 80.5 mol% Ar.

(a) Calculate the partial pressure of O2 in the mixture if

the total pressure of the atmosphere is to be 745 torr. (b) If this atmosphere is to be held in a 121-L space at

295 K, how many moles of O2 are needed?

COLLECTING GASES OVER WATER How to measure the amount of gases generated from a

chemical reaction

10.6 GAS MIXTURES AND PARTIAL PRESSURES

▲ Figure 10.15 Collecting water-insoluble gas over water.

COLLECTING GASES OVER WATER

10.6 GAS MIXTURES AND PARTIAL PRESSURES

10.6 GAS MIXTURES AND PARTIAL PRESSURES

122.6 g/mol

1. Gases consist of large numbers of molecules that are in continuous, random motion

2. The combined volume of all the molecules of the gas is negligible relative to the to-tal volume in which the gas is contained

3. Attractive and repulsive forces between gas mole-cules are negligible

10.7 KINETIC-MOLECULAR THEORY This is a model that aids in our understanding of what hap-

pens to gas particles as environmental conditions change. Summaries of the theory

4. Energy can be transferred between molecules during collisions, but the average kinetic energy of the mole-cules does not change with time, as long as the temperature of the gas remains constant

5. The average kinetic energy of the molecules is proportional to the ab-solute temperature

The pressure of a gas is caused by col-

lisions of the molecules with the walls of

the container

10.7 KINETIC-MOLECULAR THEORY This is a model that aids in our understanding of what hap-

pens to gas particles as environmental conditions change. Summaries of the theory

Although the molecules in a sample of gas have an average ki-

netic energy and hence an average speed, the Individual mole-

cules move at varying speeds

DISTRIBUTIONS OF MOLECULAR SPEED

10.7 KINETIC-MOLECULAR THEORY

mp: most probable speedav: average speedrms: root-mean-square speed

Effect of a volume increase at constant temperature

•If the volume is increased, the molecules must move a longer

distance between collisions

→ pressure decreases

Effect of a temperature increase at constant volume

•An increase in T means an increase the average kinetic E of

the molecule and thus increase in u

•If there is no change in volume, there will be more collisions

with the walls per unit time

→ pressure increases

APPLICATIONS TO THE GAS LAWS

10.7 KINETIC-MOLECULAR THEORY

10.7 KINETIC-MOLECULAR THEORY

(a) Constant (b) Constant (c) Increase (d) Increase

10.8 MOLECULAR EFFUSION AND DIFFU-SION At the same T, two gases have the same KEave, m(μrms)2

Therefore, the particles of the lighter gas must have a higher rms

speed than the particles of the heavier one.

▲ Figure 10.19 The effect of molecular mass on molecular speeds.

GRAHAM’S LAW OF EFFUSION

Effusion ( 유출 ) is the escape of gas molecules through a tiny hole into an evacuated space.

10.8 MOLECULAR EFFUSION AND DIFFUSION

▲ Figure 10.19 Effusion. Gas molecules in top half effuse through pinhole only when they happen to hit pinhole

GRAHAM’S LAW OF EFFUSION10.8 MOLECULAR EFFUSION AND DIFFUSION

DIFFUSION AND MEAN FREE PATH

Diffusion ( 확산 ) is the spread of one substance throughout a space or throughout a second substance

10.8 MOLECULAR EFFUSION AND DIFFUSION

The diffusion of gases is much slower than molecular speeds because of molecular collisions

The mean free path of a mole-cule is the average distance traveled by the molecule be-tween collisions

The mean free path for air molecule•60 nm at sea level•10 cm at 100 km in altitude

10.9 REAL GASES

Although the ideal-gas equation is a very useful description of

gases, all real gases fail to obey the relationship to some degree

▲ Figure 10.24 Gases behave more ideally at low pressure than at high pressure. The volume of gas mole-cules is not negligible at high pressure.

▲ Figure 10.25 In any real gas, at-tractive intermolecular forces reduce pressure to values lower than in an ideal gas.

10.9 REAL GASESAt high P, gas volumes are not negligible

Attractive forces between mole-cules reduce the pressure

▲ Figure 10.22 The effect of pressure on the behavior of several real gases at constant T. The deviations increases with increasing P.

10.9 REAL GASES

Cooling a gas increase

the chance for mole-

cules to interact with

each other

▲ Figure 10.23 The effect of temperature and pressure on the be-havior of nitrogen gas. The deviations increase with decreasing T.

THE VAN DER WAALS EQUATION

The ideal-gas equation can be adjusted to take the devia-

tions from ideal behavior into account The van der Waals Equation

10.9 REAL GASES

= Pideal

= Videal

THE VAN DER WAALS EQUATION a and b values increase with mass of the molecule and

the complexity of its structure

10.9 REAL GASES

10.9 REAL GASES

10.9 REAL GASES

10.9 REAL GASES

Homework

Practice Exercisesp397, 399, 402, and 412

Due on 06-13 (Thur)