1 / 38

Chapter 10 Physical Characteristics of Gases

Kinetic-Molecular Theory

- Kinetic-molecular theory is based on the idea

that particles of matter are always in motion. - Ideal Gas an imaginary gas that perfectly fits

all the assumptions of the kinetic-molecular

theory.

Assumptions of Kinetic-Molecular Theory

- Gases consist of large numbers of tiny particles

that are far apart relative to their size. - Collisions between gas particles and between

particles and container walls are elastic

collisions. An elastic collision is one in which

there is no net loss of kinetic energy. - Gas particles are in continuous, rapid, random

motion. They therefore posses kinetic energy,

which is energy of motion.

Assumptions of Kinetic-Molecular Theory

- There are no forces of attraction or repulsion

between gas particles. - The average kinetic energy of gas particles

depends on the temperature of the gas. - KE ½ mv2 where KE is kinetic energy
- m is mass and
- v is velocity

Nature of Gases

- Expansion Gases have no definite shape or

definite volume so they expand to fill the

container. Gas particles move rapidly in all

directions (assumption 3) without significant

attraction or repulsion between them (assumption

4). - Fluidity Because the attractive forces between

gas particles are insignificant (assumption 4),

gas particles glide easily past one another.

Nature of Gases

- Low Density The density of a gas is about

1/1000 the density of the same substance in the

liquid or solid state. - Compressibility During compression, the gas

particles, which are initially very far apart

(assumption 1), are crowded closer together.

Under the influence of pressure, the volume of a

gas can be greatly decreased.

Nature of Gases

- Diffusion Gases spread out and mix with one

another, even without being stirred. This

spontaneous mixing of particles of two substances

caused by their random motion is called

diffusion. - Effusion is a process by which gas particles pass

through a tiny opening. Rates of effusion of

different gases are directly proportional to the

velocities of their particles.

Real Gases

- A real-gas is a gas that does not behave

completely according to the assumptions of the

kinetic-molecular theory. - Deviations from ideal-gas behavior usually occur

at high pressures and/or low temperatures where

the attractive forces between molecules begin to

play a role.

Chapter 10, Section 1 Review

- State the kinetic molecular theory of matter, and

describe how it explains certain properties of

matter. - List the five assumptions of the

kinetic-molecular theory of gases. - Define the terms ideal gas and real gas.

Chapter 10, Section 1 Review continued

- Describe each of the following characteristic

properties of gases expansion, density,

fluidity, compressibility, diffusion, and

effusion. - Describe the conditions under which a real gas

deviates from ideal behavior.

Pressure and Force

- Pressure (P) is defined as the force per unit

area on a surface. - pressure force/area
- The SI unit for force is the newton, abbreviated

N. It is the force that will increase the speed

of a one kilogram mass by 1 meter per second per

second that it is applied. - force mass x acceleration

Example Pressure on the Feet of a Ballet Dancer

- Acceleration of gravity is 9.8 m/s/s
- What is the force of a 51 Kg dancer on the floor?

- F m x a 51 kg x 9.8 m/s/s 500 N
- What is the pressure?
- Flat footed 500 N/325 cm2 1.5 N/cm2
- Two feet tip toes 500 N /13 cm2 38.5

N/cm2 - One foot tip toes 500 N /6.5 cm2 77 N/cm2

Units of Pressure

- A common unit of pressure is millimeters of

mercury, mm Hg. A pressure of 1 mm of Hg is now

called 1 torr in honor of Torricelli. - One atmosphere is defined as being exactly

equivalent to 760 mm Hg. - In SI units, pressure is expressed in derived

units called pascals. One pascal (Pa) is defined

as the pressure exerted by a force of one newton

(1 N) acting on an area of 1 square meter. One

atmosphere is 1.01325 x 105 Pa or 101.325 kPa.

Units of Pressure

Standard Temperature and Pressure

- For purposes of comparison, scientists have

agreed on standard conditions of exactly 1 atm

pressure and 0oC. These conditions are called

standard temperature and pressure and are

commonly abbreviated STP.

Pressure Units Conversions

- Convert 0.830 atm to mm of Hg and kPa
- 0.830 atm x 760 mm of Hg/atm 631 mm

of Hg - 0.830 atm x 101.325 kPa/atm
- 84.1 kPa

Chapter 10, Section 2 Review

- Define pressure and relate it to force.
- Describe how pressure is measured.
- Convert units of pressure.
- State the standard conditions of temperature and

pressure.

Gas Laws

- Gas laws are simple mathematical relationships

between the volume, temperature, pressure, and

amount of a gas. - Boyles Law states that the volume of a fixed

mass of gas varies inversely with the pressure at

constant temperature.

Illustration of Boyles Law

Mathematical Expression of Boyles Law

- V k/P or PV k
- P1V1 k P2V2 k
- P1V1 P2V2
- P1V1 / V2 P2

Volume-Pressure Data

Boyles Law Example Problem

- V 150. mL of O2 at 0.947 atm.
- What is the volume at 0.987 atm (at constant

temperature)? - Formula P1V1 / P2 V2
- 0.947 atm x 150. mL / 0.987 atm 144 mL of O2

Absolute Zero

- Absolute zero, -273.15 oC., is the lowest

temperature possible. This is assigned a value

of zero on the Kelvin scale. - To convert from Celsius to Kelvin
- K 273.15 oC.
- To convert from Kelvin to Celsius
- oC. K 273.15

Charless Law

- Charless Law states that the volume of a fixed

mass of gas at constant pressure varies directly

with the Kelvin temperature. - V kT or V/T k
- V1 / T1 V2 / T2

Plot of Volume vs. Temperature

Charless Law Example Problem

- A sample of Ne gas has a volume of 752 mL at 25

oC. What volume will the gas occupy at 50 oC. if

the pressure remains constant? - Convert temperatures to Kelvin
- 25 oC. 273 298 K.
- 50 oC. 273 323 K.

Charless Law Sample Problem

- V2 V1 x T2 / T1
- V2 752 mL Ne x 323 K. / 298 K
- 815 mL of Ne

Gay-Lussacs Law

- Gay-Lussacs Law The pressure of a fixed mass of

gas at constant volume varies directly with the

Kelvin temperature. - P kT or P/T k
- P1/T1 P2/T2

Gay-Lussacs Law Example Problem

- Gas in an aerosol can is 3.00 atm at 25oC.
- What is the pressure at 52oC.?
- P2 P1 x T2/T1
- Convert temperatures to Kelvin
- 25oC. 273 298 K.
- 52oC. 273 325 K.

Gay-Lussacs Law Example Problem

P2 3.00 atm x 325 K / 298 K 3.27 atm

Combined Gas Law

- The combined gas law expresses the relationship

between pressure, volume, and temperature of a

fixed amount of gas. - PV/T k
- Or
- P1 V1/T1 P2 V2 / T2

Combined Gas Law Example Problem

- A helium-filled balloon has a volume of 50.0 L at

25 oC. and 1.08 atm. What volume will it have at

0.855 atm and 10 oC.? - Convert the temperatures to Kelvin
- 25 oC. 273 298 K
- 10 oC. 273 283 K.
- V2 P1 V1 T2 /( P2 T1)

Combined Gas Law Example Problem

V2 P1 V1 T2 /( P2 T1) V2 1.08 atm x50 L

Hex283 K/(0.855atm x298 K) 60.0 L

Daltons Law of Partial Pressures

- Daltons law of partial pressures states that the

total pressure of a mixture of gases is equal to

the sum of the partial pressures of the component

gases. - PT P1 P2 P3

Partial Pressure Example

- Oxygen gas is collected over water. The

barometric pressure was 731 torr and the

temperature was 20.0oC. What was the partial

pressure of the oxygen collected? - PT 731 torr
- Pwater 17.5 torr at 20.0oC.

Partial Pressure Example

Poxygen PT Pwater Poxygen 731

torr 17.5 torr 713.5 torr

Chapter 10, Section 3, Review

- Use the kinetic-molecular theory to explain the

relationship between gas volume, temperature, and

pressure. - Use Boyles law to calculate volume-pressure

changes at constant temperature. - Use Gay-Lussacs law to calculate

pressure-temperature changes at constant volume.

Chapter 10, Section 3, Review continued

- Use the combined gas law to calculate

volume-temperature-pressure changes. - Use Daltons law of partial pressures to

calculate the partial pressures and total

pressures.