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CHAPTER 12 GASES AND THEIR PROPERTIES

- That this was not the case, I attributed to the

force of prejudice, which, unknown to ourselves,

biases not only our judgments, properly so

called, but even the perceptions of our senses

for we may take a maxim so strongly for granted,

that the plainest evidence of sense will not

entirely change, and often hardly modify, our

persuasions, and the more ingenious a man is the

more effectually he is entangled in his errors

his ingenuity only helping him to deceive

himself, by evading the force of truth. - -Joseph Priestley (1790) recounting the

discovery of oxygen/gases

12.0 OBJECTIVES

- Describe the properties of gases (volume, amount,

pressure, temperature), units of measurement, and

instruments used to measure these quantities. - Understand and use the ideal gas law to solve a

variety of problems. - Apply gas laws equations to stoichiometry

problems. - Understand the uses of Daltons Law and Grahams

Law equations. - Use the kinetic molecular theory to describe the

behavior of gases at the molecular level. - Compare the behavior of real gases to ideal gases.

HOMEWORK

- HW1 11, 15, 17, 19, 21, 25, 65
- Conversions, Boyles, Charles, Combined
- HW2 27, 29, 31, 33, 37, 63, 75, 99
- Ideal Gas Law, Gas Density
- HW3 41, 45, 71, 85, 105
- Gas Stoichiometry
- HW4 47, 49, 97, 103
- Partial Pressures
- HW5 51, 55, 59, 61
- Kinetic-Molecular, Diffusion, Nonideal Gases

12.1 PROPERTIES OF GASES

- 1. Study of gases as a separate unit in

chemistry - In the gas phase, all substances are remarkably

similar and easily described by the kinetic

molecular theory. - Factors that affect gases are easy to describe

and measure. - Universal simple mathematical relationships apply

to all gases.

12.1 PROPERTIES OF GASES

- 2. Four inter-related variables associated with

gases - a. Volume V
- Gases take up a large volume
- Usually measured in L
- 1 mol 22.4 L _at_ STP
- b. Amount of gasn
- Measured in moles

12.1 PROPERTIES OF GASES

- c. TemperatureT
- ENERGY! (avg. kinetic energy of particles)
- Impacts properties of gas
- Measured in K (Kelvin) KC273.15

12.1 PROPERTIES OF GASES

- d. PressureP
- 1. Cause of pressure
- Atmosphere composed of gas molecules
- Mass of all of these cause pressure (Force/area)

12.1 PROPERTIES OF GASES

- 2. Units of pressure
- Torr (mmHg), Atmosphere (atm), Pascal (Pa)

N/m2, Bar, inHg - 1 atm 760 torr 101.3 kPa 1.013 bar 29.92

inHg - STP
- Standard Temperature and Pressure
- 1 atm (760 torr) 273 K (0oC)

12.1 PROPERTIES OF GASES

- 3. Measuring Pressure
- a. Manometer and Barometer
- Device used to measure pressure

12.1 PROPERTIES OF GASES

12.1 PROPERTIES OF GASES

- Manometer Problems
- Just follow along ?

12.1 PROPERTIES OF GASES

12.1 PROPERTIES OF GASES

12.1 PROPERTIES OF GASES

12.1 PROPERTIES OF GASES

- b. Altitude and barometric pressure
- What do YOU think the trend is?

12.1 PROPERTIES OF GASES

- 4. Ex12.1 Express 753 mmHg in atm, kPa, and

bars.

12.2 GAS LAWS EXPERIMENTAL BASIS

- 1. Proportions
- a. Direct
- As one term increases, so does the other
- X 1
- Y
- b. Inverse
- As one term increases, the other decreases
- XY 1

12.2 GAS LAWS EXPERIMENTAL BASIS

- 2. Volume and Pressure, T and n constant -

Boyle's Law - For a given amount of gas at constant

temperature, the volume is inversely proportional

to the Pressure - P1V1 P2V2

12.2 GAS LAWS EXPERIMENTAL BASIS

- Ex12.2 When an auto airbag inflates as a result

of an accident, the gases inside are at a final

volume of 25.0L and pressure of just over

atmospheric pressure, 780mmHg. What is the

pressure in the uninflated bag with a volume of

1.00L?

12.2 GAS LAWS EXPERIMENTAL BASIS

- Volume and Temperature, n and P constant -

Charles Law - The volume of a gas is directly proportional to

the temperature (in Kelvin) at constant pressure

and amount of gas - V1 V2
- T1 T2

12.2 GAS LAWS EXPERIMENTAL BASIS

- Volume and moles, T and P constant Gay-Lussacs

Law - Volume is directly proportional to of moles
- of molecules ? ? Volume ?
- So, coefficients in a balanced equation involving

gases are also the volume ratio

12.2 GAS LAWS EXPERIMENTAL BASIS

- 6. Temperature and Pressure, n and V constant

Amontons Law - The temperature is directly proportional to the

pressure - P1 P2
- T1 T2

12.2 GAS LAWS EXPERIMENTAL BASIS

- 7. Combined Gas Law
- _P1V1_ _P2V2_
- n1T1 n2T2

12.2 GAS LAWS EXPERIMENTAL BASIS

- Ex12.3 A gas occupies a volume of 7.50L at

300.mmHg and 200.0oC. What is its volume if the

same sample of gas is at a pressure of 1.50atm

and at a temperature of 22.0oC?

12.2 GAS LAWS EXPERIMENTAL BASIS

- 9. Avogadro's Hypothesis
- Equal volumes of gases at the same Temp. and

Pressure have the same number of molecules

12.2 GAS LAWS EXPERIMENTAL BASIS

- 10. Derivation of Volume-moles relationship
- Compared masses of equal volumes of different

gases and determined weight ratios? atomic

weights. - Avogadro molded together Daltons atomic theory

with Gay-Lussacs law of combining volumes - Avogadro threatened Daltons ideas about

atom/molar masses? work was ignored for 50 years!

12.2 GAS LAWS EXPERIMENTAL BASIS

- Ex12.4 Given the Haber reaction below a. What

volume of hydrogen is required to form 12.0L of

ammonia? b. What volume of nitrogen gas is

necessary to react completely with 1.41 Liters of

hydrogen gas? Assume constant T and P. - 3H2(g) N2(g) ? 2NH3(g)

12.3 IDEAL GAS LAW

- 1. Derivation of Ideal Gas Law
- Combination of Boyles, Charles, and Avogadros

Laws - PV nRT

12.3 IDEAL GAS LAW

- 2. Value of the gas law constant, R
- R .0821 Latm R 8.314 J
- molK molK

12.3 IDEAL GAS LAW

- 3. Ex12.5 Will it be safe to store 2500g of

oxygen gas in a 10.0L container at 20.0oC if the

container is built to a tolerance of 200atm?

12.3 IDEAL GAS LAW

- 4. Ex12.6 Calculate the number of moles of

ammonia present in a sample with a volume of

12.0L, at 22.0oC and 715mmHg.

12.3 IDEAL GAS LAW

- 5. Density of gas
- PV mass RT mass P(M.M.) D
- M.M. V RT

12.3 IDEAL GAS LAW

- 6. Ex12.7 What is the density of oxygen gas at

1.00 atm and 27.0oC?

12.3 IDEAL GAS LAW

- 7. Calculating molar mass
- PV mass RT
- M.M.

12.3 IDEAL GAS LAW

- 8. Ex12.8 What is the molar mass of a gas whose

density is 5.00g/L at 25.0oC and 1.00atm?

12.4 GAS LAWS AND CHEMICAL REACTIONS

- 1. Ex12.9 Hydrogen peroxide decomposes in the

presence of sunlight to produce oxygen gas and

water. Calculate the amount, in grams, of

hydrogen peroxide needed to produce 2.50L of

oxygen, measured at STP.

12.4 GAS LAWS AND CHEMICAL REACTIONS

- 2. Ex12.10 How many liters of oxygen gas at

1.00atm and 27.0oC are needed to burn 1.00g of

octane (C8H18)?

12.4 GAS LAWS AND CHEMICAL REACTIONS

- 3. Ex12.11 What mass in grams of potassium

chlorate must be used to produce 1.75L of oxygen

gas, measured at 18.0oC and 0.950atm according to

the following equation? - 2KClO3(s) ? 3O2(g) 2KCl(s)

12.5 GAS MIXTURES AND PARTIAL PRESSURES

- 1. Statement of Dalton's Law of Partial

Pressures - Total pressure of a mixture of gases is equal to

sum of the partial pressures of each component - Ptotal Pgas 1 Pgas 2 Pgas 3

12.5 GAS MIXTURES AND PARTIAL PRESSURES

- 2. Ex12.12 A gas mixture has a total pressure

of 1.50atm. If the mixture consists of 0.150mol

of methane and an unknown amount of ethane in an

8.50L vessel at 298K, what is the partial

pressure due to ethane?

12.5 GAS MIXTURES AND PARTIAL PRESSURES

- 3. Gases collected by bubbling through water and

water vapor pressure - Collection over water? Ptotal Pgas PH2O
- When level of gas level of water, pressures are

equal - See. Pg. 13 in Reference Booklet

12.5 GAS MIXTURES AND PARTIAL PRESSURES

- 4. Ex12.13 30.0mL of hydrogen gas is collected

over water at a total pressure of 744mmHg and at

20.0oC. Calculate the pressure due to hydrogen

gas and the number of moles of hydrogen gas.

12.5 GAS MIXTURES AND PARTIAL PRESSURES

- 5. Mole fractions
- The ratio of the moles of a gas over the total

moles of a gas in a mixture of ideal gases - Xa na
- ntotal

12.5 GAS MIXTURES AND PARTIAL PRESSURES

- 6. Relationship between partial pressure, mole

fraction, and total pressure - Pa Xa
- Ptotal

12.5 GAS MIXTURES AND PARTIAL PRESSURES

- 7. Ex12.14 Calculate the mole fractions of

hydrogen and water vapor in the previous problem.

12.6 KINETIC MOLECULAR THEORY OF GASES

- 1. Basic statements of Kinetic Molecular Theory
- a. Gases consist of particles whose separation

is much greater than the size of the particles,

themselves. - b. The particles of a gas are in constant,

random, and rapid motion. - c. Gas particles constantly collide with one

another and with the walls of their container,

but they do so without loss of energy. - d. The average kinetic energy of a sample of gas

particles is proportional to the absolute

temperature of the gas. Therefore, all molecules

of gas, regardless of their mass, have the same

average kinetic energy at the same temperature.

12.6 KINETIC MOLECULAR THEORY OF GASES

- 2. The kinetic energy of a single molecule
- KE ½ mu2
- u speed
- Different molecules can have different speeds, so

only applies to a single molecule - 3. The average kinetic energy of a sample of gas

molecules depends on - Kelvin Temperature only

12.6 KINETIC MOLECULAR THEORY OF GASES

- 4. The average kinetic energy of the molecules

in a gas sample is related to average u2 - KE ½ mu2
- 5. The relationship between mass, average speed,

and temperature is - ?u2 3RT (Maxwells Eqn)
- M.M.

12.6 KINETIC MOLECULAR THEORY OF GASES

- 6. Maxwell-Boltzmann Distribution
- Distribution of speeds (KE) of molecules
- Areas under curves are the same

12.6 KINETIC MOLECULAR THEORY OF GASES

- 7. Ex12.15 Calculate the average velocity (rms

speed) of an oxygen molecule at 25.0oC.

12.6 KINETIC MOLECULAR THEORY OF GASES

- 8. Ex12.16 A professional tennis player can

serve a tennis ball at 45m/sec. At what

temperature will an oxygen molecule have the same

average speed?

12.7 DIFFUSION AND EFFUSION

- 1. Diffusion and Effusion
- Diffusion mixing of gases due to molecular

motion - Ex. Spread of aroma of a baking pie
- Effusion movement of gas through a tiny opening

in a container to another container of lower P - Ex. Punching a hole in a He balloon

12.7 DIFFUSION AND EFFUSION

- 2. Graham's Law relating molar mass, rate of

speed, and time - Rate of effusion of gas 1 M.M.-gas 2
- Rate of effusion of gas 2 ?M.M.-gas 1
- Rate of effusion of gas 1 rms for gas 1

? 3RT/(MM-gas 1) - Rate of effusion of gas 2 rms for gas 2

? 3RT/(MM-gas 2)

12.7 DIFFUSION AND EFFUSION

- 3. Ex12.17 It takes 40sec for a sample of

oxygen to effuse through a small opening into a

vacuum. Another gas takes only 10sec to effuse

under the same conditions. What is the molar

mass of the second gas?

12.7 DIFFUSION AND EFFUSION

- 4. Ex12.18 The ratio of the average rate of

effusion of SO2(g) to CH4(g) at 300K is

12.8 NON-IDEAL BEHAVIOR REAL GASES

- 1. Equations used to describe ideal gases are

based on assumptions of kinetic molecular theory. - Gases actually have a volume
- 1L container does not mean gas molecules have 1L

to move about - Elastic Collisions are not always observed
- When we approach the condensation point?

molecules MUST have some attraction

12.8 NON-IDEAL BEHAVIOR REAL GASES

- 2. Real gases deviate from ideal behavior under

two main conditions - Low Temperature (approaching condensation)
- High Pressure (molecular volume becomes

significant)

12.8 NON-IDEAL BEHAVIOR REAL GASES

- 3. van der Waal's Equation - a better predictor

of gas behavior under extreme conditions - Corrects for intermolecular forces and molecular

volume