Gases

1
Gases

• Kinetic Molecular Theory of Gases

2
A gas consists of small particles
(atoms/molecules) that move randomly with rapid
velocities

• Further Information
• They move faster when heated.

3
The attractive forces between particles of a gas
can be neglected

• Do you think this is accurate?
• Why would this be important for calculations?

4
The actual volume occupied by a gas molecule is
extremely small compared to the volume that gas
occupies.

• Is this true in the real world?
• Why would this be helpful with calculations?

5
The average kinetic energy of a gas molecule is
proportional to Kelvin temperature

• What is kinetic energy?
• Why Kelvin temperature and not Celsius or
  Fahrenheit?
• What does proportional mean?

6
Gas particles are in constant motion, moving
rapidly in straight paths.

• Is this true?
• What do we know about their motion?
• Why would the real situation make the
  calculations more difficult?

7
Ideal Gases

• An imaginary gas that perfectly fits all the
  assumptions of the kinetic molecular theory
  (KMT).

8
Expansion

• Gases do not have definite shape or volume.
• The expand to any container they are enclosed in.
• A gas in a 1 L container is then put into a 2 L
  container. How much volume does it have now?

9
Fluidity

• In an ideal gas, the gas particles glide past
  each other.
• This feature allows gases to be referred to as
  fluids just like liquids.

10
Low Density

• Density of a gas substance is only about 1/1000
  of the same substance in liquid or solid state.
• Why is this true?

11
Compressibility

• This is a crowding effect of gases when the
  volume is decreased

12
Diffusion

• Spontaneous (does not require energy) mixing of
  particles of two substances caused by their
  random motion

13
Properties of a Gas

• Units of Measure

14
Pressure

• Pressure is not the same as force.
• Pressure is a force over an area.
• Example psi Pounds per in2

15
Measuring Pressure

• A barometer measures atmospheric pressure.

16
Units of Pressure

• kPa, atm, mm of Hg, torr
• Helpful Conversions
• 1 atm 760 mm Hg
• 1 atm 760 torr
• 1 mm Hg 1 torr
• 1 atm 101.325 kPa

17
Volume

• L, mL or cm3
• Helpful conversions
• 1000 mL 1 L
• 1 mL 1 cm3

18
Temperature

• 0C or K
• Helpful conversions
• 0C K 273
• K 0C 273

19
moles

• Number of moles n
• If you are given grams, how would you convert to
  moles?

20
Standard Temperature and Pressure (STP)

• Standard Temperature is 00C or 273 K
• Standard Pressure is 101.3 kPa or 1 atm

21
Boyles Law

• Pressure and volume are inversely proportional
• P1V1 P2V2

22
Boyles Law
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23
Charles Law

• Temperature and Volume are directly proportional
• V1 / T1 V2 / T2

24
Gay-Lussacs Law

• Pressure and Temperature are directly
  proportional
• P1/T1 P2/T2

25
Combined Gas Law

• P1V1 P2V2
• T1 T2
• If you remember this law, hold constant the other
  variables not used and you have all the gas laws
  weve used so far.

26
Molar Volume

• 1 mole 22.4 L of a gas at STP

27
Daltons Law of Partial Pressure

• Be sure all units of pressure are the same.
• If not, convert all units to the same unit of
  measure
• The total pressure is equal to the sum of the
  partial pressures

28
Avagadros Law

• V1 / n1 V2/n2
• Where n number of moles
• How do you convert grams to moles?

29
Ideal Gas Law

• PVnRT
• P Pressure (atm)
• V volume (L)
• n number of moles
• R 0.0821 atm x L / moles x K
• T temperature (K)
• You must use these units for the R constant to be
  correct.

30
Name the Law!

• You will be given a series of laws and asked to
  name the law or you will be given the name and be
  asked to come up with the formula!

31
PV nRT

• Ideal Gas Law

32
V1 / T1 V2 / T2

• Charles Law

33
Boyles Law

• P1V1 P2V2

34
Combined Gas Law

• P1V1 P2V2
• T1 T2

35
Gay-Lussacs Law

• P1/T1 P2/T2

36
Avagadros Law

• V1 / n1 V2/n2