chemistry i honors

Chemistry I Honors

Unit 9: Gases

Objectives #1-3: Introduction to the Kinetic Theory of Gases

I. The Kinetic TheoryA. Assumptions of the Kinetic Theory

Gases are composed of tiny particles that are arranged far apart from each other

Gases are composed of individual atoms such as in the element neon or molecules such as in the element oxygen


Collisions that may occur between gas particles are elastic with no net loss of energy

Gas particles are in constant, random motion

No attractive forces exist between gas particles

The temperature of a gas depends on the average kinetic energy of the particles

(video clip Ch. 3 “Common Gas Properties”)


II. The Kinetic Theory and its Implications on the Properties of GasesThe temperature of a gas is directly related to its kinetic energyThe temperature at which no kinetic energy is present is called absolute zero; this temperature is 0 on the Kelvin scale and -273o on the Celsius scaleGases are able to expand freely due to their random motion and the lack of attractive forces between the particles of a gas Gas particles are tiny and can move past each other with ease due to the lack of attractive forces between gas particles


The density of gases is generally less than other substances due to the ability of gases to move freely

Gases can be compressed due to the great distances between gas particles

The ability of gases to spread out spontaneously or diffuse is due to the rapid motion of gas particles; particles with less mass and faster velocities spread out at a faster rate than heavier, slower gas particles (This is why you smell cookies baking! )


Due to their small size, gas particles can be made to pass through small openings; this characteristic of effusion depends on the velocity and molar mass of the gas particles present; the rate of effusion is directly proportional to the velocity of the gas particles and inversely proportional to the molar mass of the gas particles

When gas particles collide against a surface they exhibit pressure; which is defined as the force per unit area

Objectives #4-9 : Properties of Gases

Common Units of Pressure:

Unit Symbol Standard Value

Pascal/Kilopascal Pa/kPa 101325/101.3

Millimeters of Hg mm Hg 760

Atmospheres atm 1

Torr torr 760

Objectives #4-9: Properties of Gases

Pressure Conversion Examples:• Convert 2.5 atm to mmHg

2.5 atm x 760 mmHg = 1900 mmHg

1 atm• Convert 300. Pa to atm

300. Pa x 1 atm_____ = .00296 atm 101325 Pa


Dalton’s Law of Partial PressureThe total pressure of a mixture of gases is equal

to the sum of the partial pressures of each of the individual gases in the mixture

PT = P1 + P2 + P3 + ….This concept must be kept in mind when gases are

collected over water in the laboratory; the vapor pressure of water must be subtracted from the measured pressure of the gas in order to obtain the true pressure of the gas being collected

Pgas = Patm - Pwater


The vapor pressure due to water increases with increasing temperature (see attached chart)

Molar Volumeat standard temperature and pressure

(STP), which are defined as O o C and 1 atm, 1 mole of any gas occupies 22.4 L; this is referred to as molar volume

*Interpreting Vapor Pressure Charts

Examples:1. Which substance has the highest vapor pressure?2. Which substance has the lowest vapor pressure?3. Boiling occurs when the vapor pressure of a

substance equals the external atmospheric pressure. At what external pressure will ethanol boil at 40oC?

4. A substance is volatile when it readily evaporates. Which substance is the most volatile?

5. Which substance has the strongest intermolecular forces?

6. Which substance has the weakest intermolecular forces?


Relationships Among Gas CharacteristicsA. Amount of Gas vs. Pressure

(assumes temperature and volume are held constant)

What is happening at the particle level:


As the amount of gas in a container increases, the pressure increases

This illustrates a direct relationship between the amt. of the gas and the pressure of the gas


B. Pressure vs. Volume of Gases (assumes constant temperature)



As the pressure of a fixed amount of gas increases, its volume decreases

This illustrates an inverse relationship


C. Temperature vs. Volume (assumes constant pressure)



As the Kelvin temperature of a fixed amount of gas increases, its volume increases

This illustrates a direct relationship:


D.Temperature vs. Pressure (assuming constant volume)



As the Kelvin temperature of a fixed amount of gas increases, its pressure increases

This illustrates a direct relationship


Ideal vs. Real GasesIdeal gases always follow the kinetic

theory under any conditionsReal gas particles do have

attractive forces among each otherReal gases no longer act as ideal

gases under conditions of high pressure and extremely low temperature

Objective #10 Solving problems Involving the Gas Laws

*the Gas Laws*the Gas Laws are mathematical formulas

based on the relationships discussed in the previous section of notes

P1V1 / T1 = P2V2 / T2 (Combined Gas Law)

PV = nRT (Ideal Gas Law)(video clip “Ch. 4 Ideal Gas Law)I. Combined Gas Law*examples:

1. To what pressure must a gas be compressed to get it into a 9.00 L tank if it occupies 90.0 L at 1.00 atm?

Can a variable be eliminated?P1V1 = P2V2

P1V1 / V2 = P2

(1.00 atm) (90.0 L) / 9.00 L = P2

10.0 atm = P2

2. A container with a movable piston contains .89 L of methane gas at 100.50C. If the temperature of the gas drops to 11.3oC, what is the new volume of the gas?

K = C + 273 = 100.5oC + 273 = 374 K

K = C + 273 = 11.3oC + 273 = 284 KCan a variable be eliminated?V1 / T1 = V2 / T2

V1T2 = T1V2

V1T2 / T1 = V2

(.89 L) (284 K) / 374 K = V2

.676 L = V2

3. A sample of gas occupies a volume of 5.0 L at a pressure of 650. torr and a temperature of 24oC . We want to put the gas in a 100. ml container that can only withstand a pressure of 3.0 atm. What temperature must be maintained so that the container does not explode?

5.0 L = 5000 ml650. torr = .855 atm24oC = 297 KP1V1 / T1 = P2V2 / T2

Can a variable be eliminated?P1V1T2 = T1P2V2

T2 = T1P2V2 / P1V1

T2 = T1P2V2 / P1V1

T2 = (297 K) (3.0 atm) (100. ml) /

= (.855 atm) (5000 ml) = 21 K4. A sample of gas occupies 2.00 L at

STP. What volume will it occupy at 27oC and 200. mmHg?

P1V1 / T1 = P2V2 / T2

Can a variable be eliminated?P1V1T2 = T1P2V2

P1V1T2 / T1P2 = V2

(760 mm Hg) (2.00 L) (300 K) /(273 K) (200. mm Hg) = V2

8.35 L = V2

II. Additional Problems*Recall that at STP conditions, 1 mole

of any gas = 22.4 L*examples:1. Calculate the volume of .55000

moles of gas at STP..55000 moles X 22.4 L / 1 mole =

12.320 L

2. Calculate the moles of gas contained in 350 L of gas at STP.

350 L X 1 mole / 22.4 L = 16 moles3. Calculate the mass in grams of 3.50

L of chlorine gas.3.50 L X 71.0 g / 22.4 L = 11.1 g

Objective #11 The Ideal Gas Law and Gas Stoichiometry

PV=nRT*ideal gas constant:R = PV / ntR = ( 1 atm) (22.4 L) / (1 mole) (273K) = .0821 L . atm. / mole.K*examples:1. What is the temperature of a .65 L sample

of fluorine gas at 620. torr which contains 1.3 mol fluorine?

T = PV / nR

*examples:1. What is the temperature of a .65 L

sample of fluorine gas at 620. torr which contains 1.3 mol fluorine?

T = PV / nR = (620 torr / 760 torr) (.65 L) / (1.3 mol) (.0821 L.atm / mol.K) = 5.0 K

2. A 25.0 gram sample of argon gas is placed inside a container with a volume of 10.0 L at a temperature of 65oC. What is the pressure of argon at this temperature?

PV=nRTP = nRT / V = (25.0 g / 39.9 g) (.0821) (338 K) / 10.0 L = 1.74 atm.

*there are two types of gas stochiometry problems:

at STPnon STP*examples:

1. Calculate the volume of hydrogen gas that can be produced from the reaction of 5.00 g of zinc reacted in an excess of hydrochloric acid. Assume STP conditions.

Zn + 2HCl --› H2 + ZnCl25.00 g Zn X 1 mole Zn / 65.4 g Zn X1 mole H2 / 1 mole Zn X 22.4 L / 1 mole H2

= 1.71 L

2. Calculate the volume in liters of oxygen gas that can be produced from the decomposition of 3.50 X 1024 formula units of potassium chlorate. Assume STP conditions.

2KClO3 --› 2KCl + 3O2

3.50 X 1024 formula units KClO3 X

1 mole KClO3 / 6.02 X 1023 f. units X

3 mole O2 / 2 mole KClO3 X 22.4 L / 1 mole =

195 L

3. Calculate the volume of hydrogen produced at 1.50 atm and 19oC by the reaction of 26.5 g of calcium metal with excess water. The vapor pressure of water is 16.5 mmHg.

Ca + 2H2O --› H2 + Ca(OH)2

*use amount of given reactant and stoichiometry to determine moles of gas desired in problem:

26.5 g Ca X 1 mole Ca / 40.1 g Ca X1 mole H2 / 1 mole Ca = .661 mole H2

*determine net pressure of gas:16.5 mmHg X 1atm/760 mmHg = .0217 atm1.50 atm - .0217 atm = 1.48 atm

*use moles of gas found in ideal gas law to calculate volume of gas:

PV=nRTV = nRT / P = (.661 moles) (.0821) (292 K) / 1.48 atm = 10.7 L

4. Calculate the volume of chlorine gas produced at 1.25 atm at 25oC from the reaction of 5.00 g of sodium chloride and an excess of fluorine.

2NaCl + F2 --› 2NaF + Cl25.00 g NaCl X 1 mole NaCl / 58.5 g

NaCl X1 mole Cl2 / 2 mole NaCl = .0427 mole

Cl2

PV = nRTV = nRT / P = (.0427 mole) (.0821) (298 K) /1.25

atm = .836 L

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