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Chapter 10 & 11: Gases Chapter 10: Page 300-330 Chapter 11: Page 332-359 Chlorine gas was used as a weapon in WWI – PowerPoint PPT presentation

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Title: Chapter%2010%20


1
Chapter 10 11 Gases
Chapter 10 Page 300-330 Chapter 11 Page 332-359
Chlorine gas was used as a weapon in WWI
2
Kinetic Molecular Theory
  • all matter is made up of particles (atoms) in
    random and constant motion. (colliding)
  • Gases have very low density
  • particles are spaced far apart.
  • Gases are compressible.
  • Extreme pressures-gases will compress until they
    become liquids (or solids, CO2).
  • Adding heat to a system
  • increases the temperature
  • Temperature measure of the average kinetic
    energy of the particles.
  • Increasing the pressure of a gas,
  • increases the density of the gas - the number of
    particles in a given space.

3
Ideal and Real Gases
?Ideal gasses
  • Ideal Gas
  • Imaginary, perfect gas makes calculations
    easier
  • Real Gas
  • Gas that actually behaves in reality
  • When compressed, real gases will form liquids,
    and even exhibit liquid-like behaviors when still
    in gas form.
  • Real gas molecules interact with each other
    -causing them to travel in non-linear paths and
    collide inelastically.
  • With real gases, the size of the gas molecules
    effects their behavior.

Real Gasses ?
4
Gases in our atmosphere
  • Nitrogen-78
  • Oxygen-21
  • Argon-lt1
  • Trace amounts of CO2, Ne, He, CH4, Kr, H2, O3,
    and others.
  • Some gases function as greenhouse gases, and work
    to hold heat on the earths surface.
  • Some gases function to block harmful UV radiation
    energy from the sun.

5
The Greenhouse Effect
  • The suns energy travels through space and warms
    the surface of the earth.
  • Some of the energy is reflected back into space.
  • Greenhouse Gases
  • trap heat that would leave the atmosphere.
  • H2O, CH4, and CO2 are common greenhouse gases.
  • Global Warming
  • Theory that increasing levels of Greenhouse
    gasses is causing the global average temps to
    increase.

6
The Ozone Layer (O3)
Page 778 for more info
  • Ozone is
  • a corrosive poison in the Troposphere (where we
    live)
  • frequently created and given off from free
    electrical ionization.
  • Ozone
  • Absorbs harmful Ultraviolet (UV) energy in the
    stratosphere, 11km (6 miles) above us.
  • Note the Ozone layer is less than 1mm thick!
  • It is always moving, like a cloud, due to weather
    patterns and climate variations.

7
Pascals Principle and Pressure
F Force a area
  • French physician, Blaise Pascal, showed that
  • fluids (including gasses) exert a uniform
    pressure on all the surfaces that they contact.
  • Exerting a force on the top surface of a gas,
    causes that force (pressure) to be exerted on all
    the walls of its container.
  • Pressure is due to the particles of a gas
    striking a surface. We can detect pressure from
    billions upon billions of gas molecules striking
    a surface at any point in time.

? Which exerts a greater pressure? ?
8
Atmospheric pressure
  • The weight of the atmosphere above us exerts a
    pressure of 10 Newtons per cm2,
  • equivalent to the weight of a small brick on a
    surface no larger than your fingertip!
  • To feel the effect of doubling the pressure due
    to atmospheric pressure, you must swim down to a
    depth of roughly 10 metersouch!
  • During aircraft flight, the cabins of jetliners
    are pressurized to just below 1 atmosphere were
    they not, we would never be able to fly at 30,000
    feet. ( WHY? )

9
Pressure units
  • The SI unit of pressure is the Pascal, Pa,
    equaling one newton per square meter.
  • Earths air pressure at sea level 100,000 Pa.
    100kPa
  • PSI (US)
  • Pound per square inch. Atmospheric pressure at
    sea level is about 14.5 PSI.
  • mmHg (EU, Asia) (AKA Torr)
  • Millimeters of mercury.
  • Atmospheric pressure is 760 mmHg at sea level.
  • This has to due with the height of a column of
    liquid mercury raised in a barometer.
  • inHg
  • Inches of mercury. Used only in meteorology.
  • Atmospheric pressure is apx 30inHg.

10
And finally
  • And, finallythe atmosphere, atm
  • the pressure exerted by the atmosphere at sea
    level, at 00C. (This creates STP)
  • Standard Temperature and Pressure
  • STP
  • usually used when referring to reactions with
    gases. STP is defined as
  • 1 atm and 273.15 K
  • 101 kPa and 273.15 K
  • 760 mmHg and 273.15 K

When doing work with gases, select the STP that
matches the pressure you are using. (atm in this
class)
11
Avogadros.Law(?)
  • It was Amedeo Avogadro that investigated the
    relationship between gas volume and the number of
    particles. He found that
  • all the gases that he used, when measured out to
    their molecular weight, had a volume of about
    22.4 liters!
  • That is equal volumes of gases have equal
    number of particles
  • 6.02x1023 molecules of almost any gas occupies
    22.4 liters of space!
  • This led to the formation of a gas constant

More about this later
12
Charles Law
Simulation. constant volume
  • French chemist, Jacque Charles, showed that at
    constant pressure,
  • temperature and volume varied proportionally.
    That is
  • V / Tk (k some constant )
  • We tend to write Charles Law as the volumes and
    temperatures under two conditions

c 1780s
13
Boyles Law
Were leaving one law out can you guess what it
is?
  • A young, adventurous, British aristocrat named
    Robert Boyle found that
  • when temperature is kept constant, volume varies
    inversely proportional with pressure. That is
  • P V k (constant)
  • We tend to write Boyles Law as the volumes and
    pressures under two conditions

c 1660s
14
Charles Law Boyles Law Avogadros Law
THE IDEAL GAS LAW
  • R is the gas constant and numerically depends
    upon the pressure units used.

Pressure
Volume (in Liters)
Moles
Constant
Temperature (in Kelvin)
15
Gas Law Summary
16
The Gas Constant
  • The Gas Constant is the numerical bridge between
    number of moles of a gas, its temperature, and
    volume or pressure.
  • R 8.314 L?kPa / mol?K
  • R 0.0821 L?atm / mol?K
  • Note that the first constant is in KILO Pascals.
    When given Pascals, you must first convert to
    kilopascals.
  • Our calculations will be done in Atm

17
Daltons Law of Partial Pressures
  • The total pressure in a system is the sum of the
    individual pressures exerted by each gas.
  • So, if gas A exerts a pressure of 2 units, and
    gas B exerts a pressure of 3 units, the total
    pressure of a system of equal parts of A and B,
    would be ?
  • Total A B .. 2 3 5 units.
  • In our atmosphere, Oxygen is about 21. If we
    have a sample of air at 1 atm, what is the
    pressure due to oxygen?

18
Grahams Law of Gas Effusion
  • Effusion
  • motion of a gas through an opening in a
    container.
  • not Diffusion - dispersing from higher
    concentration to lower concentration.
  • Rates (speeds) of effusion are related to the
    molar mass of a gas.
  • The higher the molar mass, the slower the gas
    will effuse.
  • This is a property of real gases

19
Grahams Law of Gas Effusion
  • At the same temperature
  • The higher the molar mass, the slower the gas
    will effuse.
  • Grahams Law of Effusion

Molar mass
Molar mass
velocity
velocity
Gas A vs Gas B
20
Vapor Pressure
Page 324
  • All liquids exert a vapor pressure.
  • Vapor pressure liquids molecules ? gas phase.
  • Higher temperatures ? greater molecular speed ?
    greater vapor pressure.
  • More volatile liquids exert a greater vapor
    pressure than do less volatile liquids.
  • Can you explain why this is?
  • In lab we collect gasses over water.
    There is a small amount of water vapor
    in our gas samples, due
    to waters vapor pressure.

21
Phase Diagrams
Example on page 381
  • Phase diagrams
  • predict if a substance will be a solid, liquid or
    gas
  • depends upon the pressure and temperature of the
    substance.
  • Triple Point
  • point where solid, liquid, and gas all exist
    for water, 00C.

Notice, that as you increase pressure, the
boiling point of water increases-this is why a
pressure cooker works. What about Denver, the
mile-high city?
End of Gases lecture, Chapters 10,11, problems
following
22
In-chapter problems
  • Page 327, 5,7,8 What is Pressure?
  • Page 327, 11-14 Pressure Units
  • Page 327, 17-19 Pressure Conversions
  • Page 330, 20-24e Boyles Law
  • Page 330, 25-27 Charles Law
  • Page 330, 31-35o Combined Law
  • Page 331, 39,40 Daltons Law of Partial
    Pressures
  • Page 357, 9-13o Avogadros Molar Gasses
  • Page 358, 17-20 Ideal Gas Law
  • Page 358, 23-29o Ideal Gas Law and Stoichiometry
  • Page 359, 39-42 Grahams Law of Gas Effusion

End of Gases Unit, Chapters 10,11
23
CCSD Syllabus Objectives
  • 11.1 Kinetic Molecular Theory
  • 11.2 Physical Properties of Gasses
  • 11.3 STP
  • 11.4 Volume-Temp relationships
  • 11.5 Volume-Pressure relationships
  • 11.6 Density-Volume-Pressure-Temperature
  • 11.10 Ideal Gas Law
  • 11.11 Grahams Law
  • 11.12 Ideal Gas vs Real Gas
  • 12.3 Evaporation, Condensation, Sublimation
  • 21.1 Environmental Chemistry
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