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Gases. Exploring Gases. Make a table: Demo # Prediction Observation. Kinetic Theory (Observing Properties of Gases ). Gases are tiny particles, that have mass but a small volume (Volume assumed to = 0 ) Gases in constant random motion

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Exploring gases
Exploring Gases

Make a table:

Demo #PredictionObservation

Kinetic theory observing properties of gases
Kinetic Theory(Observing Properties of Gases)

  • Gases are tiny particles, that have mass but a small volume

    • (Volume assumed to = 0)

  • Gases in constant random motion

    • Collisions are elastic with walls and each other (No attraction between molecules)

  • Kinetic energy (movement) depends on temperature – (high temperature is more movement)

  • Measuring gases
    Measuring Gases

    • Amount (n)

      moles (number of)

    • Volume (V)


    • Temperature (T)

      Kelvin K= C + 273

    • Pressure (P)

      mm Hg or atm


    Particles colliding with objects

    • Pressure=force (Pascal)


      Gas pressure

      gas particles colliding with objects

      Atmospheric pressure

      air particles colliding with objects


    Measures atmospheric pressure

    Measuring pressure
    Measuring pressure


    760 mm Hg = 1 atmosphere

    = 101,300 Pascals

    =14.7 lb/in2


    Standard temperature 0oC

    Standard pressure 1 atm

    Pressure problems
    Pressure problems

    If 760 mm Hg = 1 atm

    • Convert 793 mm Hg to atm

    • Convert 3.5 atm to mm Hg

    Manometer problems
    Manometer problems

    Measures pressure in a closed container




    755 mmHg



    755 mmHg



    855 mmHg


    ? mmHg


    100 mm Hg

    Dalton s law of partial pressures
    Dalton’s Law of Partial Pressures

    Each of the components of a gas mixture contributes some of the collisions

    Each component contributes part of the total pressure.


    PT = P1 + P2 + P3…..

    What is temperature
    What is temperature?

    • Measure’s average kinetic energy of particles.

      • Higher temp means higher energy

        • More energy means faster particles

      • Lower temp means lower energy

        • Less energy means slower particles

    • When particles move faster, they collide more often and with more force.

    Temperature and energy revisited
    Temperature and Energy Revisited

    • If there are a fixed number of gas particles in a container

    • And it has a fixed pressure

    • What happens when it is heated up?

    • The particles go faster

    • They collide more

    • The volume goes up

    Combined gas law
    Combined Gas law


    • Boyles,

    • Charles,

    • Gay-Lussac law

      P1V1 = P2V2

      T1 T2

    Boyles law p v
    Boyles Law(P,V)

    At a constant T,N, the volume varies indirectly with the pressure

    P1V1 = P2V2

    Lab boyles law
    Lab: Boyles Law

    • Purpose:

      To observe changes in pressure with the volume changes

    Charles law t v
    Charles Law(T,V)

    • At a constant P, N the volume varies directly with the Kelvin temperature

      V1 = V2

      T1 T2

    Gay lussacs law p t
    Gay-Lussacs Law(P,T)

    At a constant V,N the pressure varies directly with the Kelvin temperature.

    P1 = P2

    T1 T2

    Lab pressure temperature
    Lab: Pressure/temperature

    Purpose: To determine the absolute zero using Gay-Lussac’s Law

    Absolute zero demonstration
    Absolute zero demonstration

    Temperature at which all molecules stop moving

    0 K

    -273 oC

    Avagadro s law
    Avagadro’s Law

    • If you hold pressure and temperature constant

      • Like at standard temperature and pressure

        • Which are?

    • Volume and moles are related



    Lab combined gas law
    Lab: combined gas law

    Purpose: To determine the volume of 1 mole of a gas using the combined gas law

    Reaction: Hydrochloric acid and Mg

    Ideal gas law
    Ideal Gas Law

    At STP 1 mole=22.4L

    If not at STP use Ideal Gas Law

    P V = n R T

    R=ideal gas constant

    (0.0821 L atm)

    moles K

    Ideal gas law1
    Ideal Gas Law

    • All the gas laws are related.

    • By the pressure, volume, temperature and number of particles (moles or n)

      PV = constant

      V/T = constant

      n/V = constant

      P/T = constant

    Gas proportions
    Gas Proportions

    • There are four variables in the Gas Laws

      • Pressure

      • Volume

      • Temperature

      • Moles

    • We can intuit each gas law using KMT

    • For example:

      • If moles and temperature are constant

      • How does the volume and pressure compare?

    Gas stoichiometry
    Gas Stoichiometry

    • C8H18, octane, combusts in your car’s engine. If the cylinder is 0.500 L and the oxygen intake is at 45oC and 1.05 atm, how many grams of octane are needed to completely react with the oxygen?

    Gas collected over water
    Gas Collected Over Water

    If the water level in the flask is equal to the surrounding water, than the inside pressure is equal to the outside pressure.

    Pin = PO2 + PH2O = P atmospheric

    PH2O = 21 torr

    Pressure of collected gas
    Pressure of Collected Gas

    • The vapor pressure of water @ 20.0 C is 17.54 mmHg

    • How many mmH2O is this?

      • What data do you need?

      • Mercury d = 13.7 g./ml

      • 240. mmH2O

    • If 100.0 ml of oxygen is collected over 20.0 C

      • If the atmospheric pressure is 739 mmHg, what is the pressure of oxygen?

      • How many moles of Oxygen gas?

      • How many atoms of oxygen

    Pressure equalization
    Pressure Equalization

    Which pressure is


    Pressure equalization1
    Pressure Equalization

    Which pressure is


    Motion of gases
    Motion of Gases

    • At the same temperature, two samples of gas have the same average kinetic energy

    • What has more kinetic energy, a bus moving at 5 mph or a baby on a tricycle moving at 10 mph?

    • If mass is important let’s consider molecular motion.


    • Gases at the same temp have the same average KE

    • More massive gases must be moving slower than less massive gases at the same temp

    • If I let out a smelly gas of a flask, how does it get to your nose? What path does it take?

    • When a gas spreads out or dissolves into the air, we call this diffusion

    Graham s law
    Graham’s Law

    • The rate of diffusion is directly proportional the speed of the molecule

    • The bigger the molecule, the ______ the speed of the molecule (at the same temp)

    • The “bigger” really means molar mass.

    • We can compare rates or velocities

    Graham s law1
    Graham’s Law

    va = Mb

    vb Ma

    The ratio of the velocities of gas molecules is proportional to the Inverse square root of their molar masses


    • The rate of diffusion of an unknown gas is four times faster than the rate of oxygen gas. Calculate the molar mass of the unknown gas and identify it.

      va = 1 = Mb vb 4 32 g/mol

      Mb = 2 What gas has a molar mass of 2?

    Dalton s law
    Dalton’s law

    • The total pressure equals the sum of the partial pressures of the gases in the container

      PT = P1 + P2 + P3 + ……..

    Try this
    Try this

    Air contains oxygen, nitrogen, carbon dioxide and other gases. What is the pressure due to oxygen in mm Hg if

    PT= 1 atm

    PN=593.4 mm Hg

    PCO2 = 46.78 mm Hg


    • The pressure on 2.50 L of anesthetic gas is changed from 765 mm Hg to 304 mm Hg. What is the new volume if the temperature is constant?


    • A balloon inflated in air conditioning at 27oC has a volume of 4.0L. It is heated to 57oC. What is the new volume?


    A gas has a volume of 17.3 mL at 3.5 atm. What is the volume if the pressure is increased to 6.7 atm?

    A can contains a gas at 50oC and has a volume of .5L. When released what is its new volume at 20oC?

    Try this1
    Try this

    If 87.6 mL of hydrogen gas is collected at a room temperature of 23oc and room pressure of 742 mmHg, what will the volume be at STP?


    • A gas has a volume of 6.8L at 327oC. What is its volume at 36oC?

    Molar volume
    Molar Volume

    1 Mole = 22.4 L

    At STP

    Standard temp 0oC

    Standard pressure 1 atm

    Pressure temperature
    Pressure /Temperature

    • In a sealed container, with a fixed volume and fixed number of particles

    • What happens to the pressure, if the temperature of the system is increased? Why?

    • The pressure and the temperature vary directly. Just like in Charles Law.

    • 1. Why should the thistle tube be under the water level?

    • 2. Why was the first bottle “let go”?

    • 3. Why were the bottles placed upside down on the lab bench?

    • 4. What was this method called for collecting gas

    • using a pneumatic trough and pushing water out?

    • 5. Why did the splint go out inside the bottle?

    • 6. What was the clear, colorless liquid produced?

    • 7. Write a chemical reaction for its production.

    Lab combined gas law1
    Lab: Combined gas law

    1.Take room temperature and pressure

    2. Get about 5 cm (or less) Mg and mass. Tie onto copper wire

    3. Pour 15 ml of HCl into eudiometer and fill to top with water

    4. Put Mg into top of eudiometer. Stopper.

    5. Put finger over the hole, turn upside down and place into big beaker of water.

    6. When reaction is complete, put finger over hole and transfer to large graduated cylinder to measure volume of gas collected.

    Kinetic theory real vs ideal
    Kinetic Theory Real vs Ideal

    Valid if not at

    Low temperatures

    – attractive forces apply

    high pressures

    – volume of particles

    - attractive forces apply


    • What happens to the energy of the particles of gas when you put the flask into cold water?

    • Why do we use Kelvin when calculating gas law problems? (Hint – is the Celsius temperature directly proportional to pressure below zero degrees?)

    • Predict the volume of the gas at 0 K from your data. You do this by getting an equation an plugging in the numbers.

    • Compare your answer with the real answer. Why are they different? What could affect your results?

    • A flask has a pressure of 1.0 atm at 25 C. What is the pressure at –40 C? (remember to convert to Kelvins)