Kinetic Molecular Theory and Gas Laws Day 1

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Kinetic Molecular Theory and Gas LawsDay 1
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• Kinetic-Molecular Theory explains how particles
  in matter behave
• All matter is composed of small particles that
  are far apart. Gas is mostly empty space.
• Particles are in constant, random motion.
• Particles collide with each other and walls of
  their containers collisions create pressure
• 4. Collisions are elastic no KE lost
• 5. No attractive/repulsive
  forces between particles.
  Molecules move in straight
  lines.

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PHASE SOLID LIQUID GAS
Kinetic Energy Low Medium High
Shape Rigid Fluid Fluid fills container diffusion
Forces Strong Medium Weak
Volume Definite Definite Indefinite
Compressibility Incompressible Incompressible COMPRESSIBLE! squish!
Density High (particles close together) Medium Low (spread out)
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EXIT QUESTIONS

• 1) Each of these flasks contains the same number
  of molecules. Which container has the highest
  pressure? Explain your answer.

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EXIT QUESTIONS

• Which of the following changes to a system will
  NOT result in an increase in pressure? Explain
  why you chose your answer.
• Increasing the volume of container
• adding more gas molecules
• Decreasing the volume of the container
• Raising the temperature

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Kinetic Molecular Theory and Gas LawsDay 2
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Factors affecting gases

• Volume amount of space an object occupies
• Measured in milliliters (mL) or Liters (L)
• 1000 mL 1 L
• We already have heard
• that 1 mol 22.4 L _at_ STP
• The more moles we
• have the bigger the
• balloon will need to be!

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Example 1

• How many moles of nitrogen gas are in 89.6 L at
  STP?

1 mol N2
89.6 L N2
22.4 L N2

4.00 mol N2
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Example 2What volume does 76 grams of fluorine
(F2) occupy at STP (normal conditions)?
76 g F2
1 mole F2
22.4 L F2
37.996 g F2
1 mole F2
44.8 L F2 45 L

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• 2. Temperature
• Average kinetic energy of particles (how fast
  they go)
• Measured in Kelvin
• K oC 273
• Ex Convert 17oC to
• Kelvin
• 17oC 273 290 K

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• C 5/9 (F-32)
• F 9/5 (C) 32
• K C 273
• C K- 273

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3. Pressure

• Force exerted by a gas per unit area on a
  surface. Example Pounds/in2 or psi
• Results from the simultaneous collisions of
  billions of gas particles with the walls of the
  vessel containing the gas.

Standard pressure 760 mm Hg 1 atmosphere
101.3 kPa 29.92 in. Hg 14.7
psi 760 torr
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Measuring Atmospheric Pressure

• Measured with a barometer.
• A barometer uses a column of mercury that rises
  to an average height of 760 mmHg at sea level.
• 1 atmosphere (1 atm)

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Standard Temperature and Pressure (STP)

• The conditions of standard temperature and
  pressure are
• 1.0 atm pressure and
• 273 K (or 0?C).
• _at_STP 1 mole of gas 22.4 L of gas

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Example 1

• The atmospheric pressure in Denver, CO is 0.830
  atm on average. Express this pressure in mm Hg.

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• 0.83 atm ? mm Hg
• 1 atm 760 mm Hg

0.83atm
760 mm Hg
1 atm
630.8 mm Hg 631 mmHg
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Example 2

• Convert a pressure of 175 kPa to atmospheres.

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• 175 kPa ? atm
• 101.3 kPa 1 atm

175 kPa
1 atm
101.3 kPa
1.72 atm
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Gas Law Foldable

• Fold the left and right to the middle.
• Cut along solid lines (but only to the crack!)

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Daltons Law of Partial Pressure

• The pressure of a mixture of gases is equal to
  the sum of the partial pressures of the component
  gases.

PTotal P1 P2 P3.
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Example

• A balloon is filled with air (O2, CO2, N2) at a
  pressure of 1.3 atm.
• If PO2 0.4 atm and PCO2 0.3 atm, what is
  the partial pressure of the nitrogen gas?

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• PTotal P1 P2 P3.
• Ptotal PO2 PCO2 PN2
• 1.3 atm 0.4 atm 0.3 atm PN2
• PN2 0.6 atm