Chemistry Name:_________________________________ Hood River Valley High School Unit 8 Note Packet and Goals Period:_________

Unit 8 – States of Matter and the Behavior of Gases Unit Goals: As you work through this unit, you should be able to:

1. Describe, at the molecular level, the difference between a gas, liquid, and solid phase. 2. Explain gas behavior using the kinetic molecular theory. 3. Relate attractive forces to boiling points and vapor pressure. 4. Interpret a phase diagram and describe the significance of the triple point. 5. Describe the four variables that define a gaseous system quantitatively 6. Use STP to determine the amount of gas used or produced in a chemical reaction. 7. Use ideal gas behavior and solve the following gas laws:

Charles's Law Combined Gas Law Boyle's Law Ideal Gas Law Gay-Lussac’s Law

8. Understand how gas mixtures and movements are described by the following gas laws: Dalton’s Law of Partial Pressures & Graham’s Law

Read: Chapter 13: Sections 13.1, 13.2, 13.4

Chapter 14: Sections 14.1 – 14.4

Homework: Classwork and Labs Description Goals 5 4 0

HW 1

Kinetic Molecular Theory WS

1 – 2

Unit 8 Demo Packet.

Charles Law Lab

HW 2

Relationship between gases and liquids WS

2 – 4

Ideal Gas Law Candle Lab

Molar Volume Lab

HW 3

Boyle’s, Charles, Gay-Lussac’s Law WS

5-7

Grahams Law Lab

Unit 8 test

HW 4

Combined and Ideal Gas Law WS

5-7

HW 5

Dalton’s and Graham’s Law WS

8

HW 6

HW 7

Kinetic Theory Notepack Name:___________________________

Period:__________

Chapter 13: The States of Matter A. What are the three states of matter chemists work with?___________________________________

B. We assume _______ moles of gas equal ________ at STP. What does this mean?

C. What four factors can affect the behavior of a gas, thus disrupting the assumption above? List in

each box below and diagram an example.

Chapter 13.1: The Nature of Gases A. We will explain the behavior of gases using the Kinetic Molecular Theory:

1. What is KINETIC ENERGY? The energy of an object because of ___________________

2. Define the kinetic theory as it pertains to gases: _________ or ___________ are in constant

______________, ___________ motion

3. There are 3 basic assumptions of the kinetic theory of gases. Summarize them here:

a. The particles in a gas are considered to be ________, ___________________ with an insignificant

volume: ______ attractive or ________________ forces exist between them.

b. The particles of a gas travel in ______________, constant, and _______________ motion. They

travel in _______________line paths that are __________________of each other.

c. All collisions between particles are perfectly _______________(no kinetic energy is ________).

The total ______________energy of the gas particles remains ________________.

4. If we opened a perfume bottle in Washington D.C., the gas molecules should reach Mexico City in about

90 minutes. Why don’t they ever make it?

B. Gas Pressure, pg. 386

1. Define gas pressure – _____________ of a gas per _________________. Think tire pressure,

________. Pressure is caused by _______________ on the __________ of anything.

2. Define a vacuum – ________________ space with _____ _______________ or ________________

3. Define Atmospheric Pressure – Results from collision from ________ or ___________ in normal

air. The top of a mountain has _________ atomospheric pressure than the bottom.

4. What is a barometer ?

5. What are the units used to measure pressure? (There are 3 commonly used)

6. Define standard conditions, both in temperature and pressure, when working with gases (STP)?

What pressure it this for each type of unt?

C. Kinetic Energy and Kelvin Temperature Page 388-389

1. Why do gas molecules contain kinetic energy?

2. What happens to the amount of kinetic energy an object contains as it is heated?

3. The average of the particles of a substance is proportional to the

______________________________ of the substance.

4. Particles of all substances at the same have the same

average_______________. Bigger molecules move _____________, Smaller move ____________

5. Theoretically, there is no ___ _ to which a substance’s temperature can be raised.

By contrast, however, there is a_______________________.

6. What is the term used to describe the temperature at which the motion of particles ceases and

therefore it has no kinetic energy?_____________________

7. What are its values in Celsius___________? Kelvin___________?

8. The ___________ scale is used to directly measure the kinetic energy of an object, NOT the

___________scale. Why?

13.2 The nature of Liquids

A. Particles in a liquid have enough ______________ to __________past one another. However, the particles

in a liquid have enough ________________ __________________ to keep them close to each other.

1. Conversion of liquid to a gas is called_______________________

2. Conversion of gas to a liquid is called _______________________

3. When this occurs at the surface only, it is called ___________________________

4. The force exerted by the vaporized gas above the __________ is called ____________ ___________

5. An increase in ________________ will increase the ________________ ____________________

6. The Boiling Point is the temperature at which the _________________ _____________ becomes

equal to the ________________ _______________ pushing down on the liquid. Therefore, the

liquid becomes a _______ even below the surface.

7. More ________________ _______________ equates to _____________ vapor pressure.

Vapor Pressure

Boiling Point

8. In a closed container, an increase in external ________________ will cause an _________________

in ______________ _____________ because they must be equal to ________ ______________

13.3 The nature of solids Solids have such _________ energy the particles can only ______________. The ______________

______________ forces are too strong to break.

Conversion of a solid to a liquid is called ______________________.

Converstion of liquid to a solid is called ___________________.

Conversion of a solid to a gas is called __________________

13.4 Phase changes (Changes of State) Phase diagram

14.1 Factors Affecting Gas Pressure

1._____________________

Copy Fig. 14.4, pg 415

How does adding molecules into a container, keeping volume and temperature constant, affect the pressure

inside the container?

How does removing molecules from a container, keeping volume and temperature constant, affect the pressure

inside the container?

2._________________________________

Copy Fig 14.6, pg 416 What happens to the pressure inside of a container

when we reduce the volume, keeping temperature

constant?

What happens to the pressure inside of a container

when we enlarge the volume, keeping the temperature

constant?

3._____________________________

Copy Fig 4.7, pg. 417 How does raising the temperature of a gas affect the

pressure if the volume is kept constant?

How does lowering the temperature of a gas affect the

pressure if the volume is kept constant?

14.2 The Gas Laws Pages 418-425

A. Boyle’s Law for Pressure Volume changes

State Boyle’s Law in words –

1. Show Boyle’s Law in an equation –

2. Show the graph illustrating Boyle’s Law-

Boyle’s Example 1 : A high-altitude balloon contains 30.0 L of helium gas at 103 kPa. What

is the volume when the balloon rises to an altitude where the pressure is only 25.0 kPa?

(Assume that the temperature remains constant.)

Boyle’s Example 2: A gas with a volume of 4.0 L at a pressure of 205 kPa is allowed to

expand to a volume of 12.0 L. What is the pressure in the container if the temperature remains

constant?

B. Charles’ Law for Temperature-Volume Changes

State Charles’ Law in words –

1. Show Charles’ Law in an equation –

2. Show the graph illustrating Charles’ Law –

Charles’ Law Example 1: A balloon inflated in a room at 24 C has a volume of 4.0 L. The balloon is

then heated to a temperature of 58 C. What is the new volume if the pressure remains constant?

Charles’ Law Example 2: If a sample of gas in a balloon occupies 6.8 L at 325 C, what is the

temperature if the volume later is found to be 4.2L if the pressure does not change?

C. Gay-Lussac’s Law for Temperature-Pressure Changes

State Gay-Lussac’s Law in words –

1. Show Gay-Lussac’s Law in an equation –

2. Show the graph illustrating Gay-Lussac’s Law –

Gay-Lussac’s Example 1: A gas left in a used aerosol can is at a pressure of 103 kPa at 25C. If the

can is thrown in fire, what is the pressure of the gas when its temperature reaches 928C ?

Gay-Lussac’s Example 2: A gas has a pressure of 6.58 kPa at 539 K. What will be the temperature if

pressure raises to 101.3 kPa if the volume does not change?

D. The Combined Gas Law

Show this law in an equation:

Combined Gas Law Example 1: The volume of a gas-filled balloon is 30.0 L at 40.0 C and 153 kPa

pressure. What volume will the balloon have at standard temperature and pressure?

Combined Gas Law Example 2 : A gas at 155 kPa and 25 C occupies a container with an initial

volume of 0.010 L. By changing the temperature, the pressure of a gas increases to 605 kPa and the

volume increases to 2.4 L. What is the new temperature?

Combined Gas Law Example 3: A 5.00 L air sample at a temperature of – 50 C has a pressure of 107

kPa. What is the new pressure if the temperature is raised to 102 C and the volume expands to 7.00 L?

Graphing all three laws

Direct Relationships:

Indirect Relationships:

14.3 The Ideal Gases Pages 426-431

A. This equation considers a fourth variable that affects the behavior of gases… NUMBER OF

_______________!

1. This equation incorporates Avogadro’s hypotheses (1 mole of any gas = 22.4 L at STP) into

the combined gas law from above: Write the Ideal Gas Law Equation here:

2. Substitute the values of standard condition and solve for “R”

a. Using Pressure in a kPa

b. Using Pressure in an ATM

c. Using Pressure in a mmHg

**The advantage of the ideal gas law over the combined gas law is that it permits us to solve for the

number _________________ of a contained gas when P, V, and T are known.

PV=nRT Example 1: You fill a rigid steel cylinder that has a volume of 20.0 L with nitrogen gas to a

final pressure of 20,000.0 kPa at 28 C. How many moles of nitrogen gas does the cylinder contain?

PV=nRT Example 2: What volume will 12 grams of oxygen gas occupy at 25 C and a pressure of 52.7

kPa?

PV=nRT Example 3: When the temperature of a rigid hollow sphere containing 68.5 moles of helium

gas is held at 621 L, the pressure of the gas is 1.89 x 103 kPa. What temperature is this container of

helium?

PV=nRT Example 4: What pressure will be exerted by 0.45 mol of a gas at 25 C if it is contained in a

0.65 L vessel?

PV=nRT Example 5: A child has a lung capacity of 2.2 L. How many grams of air do her lungs hold

at a pressure of 102 kPa and a normal body temperature of 37 C? Air is a mixture, but you may assume

it has a molar mass of 28 g/mole.

14.4 Gas Mixtures and Movements, pgs 432-437

Dalton’s Law of Partial Pressures

A. Many gases, including air, are ___________________. Remember, the particles in a gas at the

same ________________ have the same average __________________.

B. Gas pressure depends only on the__________________ of gas particles in a given

_____________ and their average _____________________. The kind of particle is

_____________________.

1. Define partial pressure –

2. Define Dalton’s law of partial pressures, include equation. –

Partial Pressures Example 1: Air contains oxygen, nitrogen, carbon dioxide, and trace amounts of

other gases. What is the partial pressure of oxygen (PO2) at 101.3 kPa if the partial pressure of

nitrogen, carbon dioxide, and other gases are 79.1 kPa, 0.04 kPa, and 0.94 kPa, respectively.

Partial Pressures Example 2: Determine the total pressure of a gas mixture that contains oxygen,

nitrogen, and helium if the partial pressure of the gases are as follows: Oxygen = 20.0 kPa, Nitrogen =

46.7 kPa, Helium = 26.7 kPa

Mole Fractions and %’s when dealing with Partial Pressures

i.e. A balloon has 5 moles CO2, 3 moles N2, and 2 moles O2. What is the pressure of each gas in the

balloon if the total pressure is 1.10 atm?

Diffusion and Graham’s Law Page 352

A. Define diffusion –

B. Define effusion –

C. State Graham’s Law of Diffusion –

1. Graham’s law may be understood by an examination of the relationship of the mass and

speed of a moving body to the force the body exerts when it strikes a stationary object:

Show the mathematical proof that derives Graham’s Law:

2. Restate Graham’s law of diffusion showing the relationship to the lighter gas and it’s velocity

as compared to the heavier gas.

Graham’s law Example 1: Compare the rate of diffusion of nitrogen gas to helium gas:

Graham’s law Example 2: Compare the rate of diffusion of argon to neon gas: