Unit 6 Gases, Phase Changes and Introduction to Thermochemistry

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Unit 6Gases, Phase Changes and Introduction to
Thermochemistry
Part I Gases

• Characteristics of Gases
• Pressure
• Kinetic-Molecular Theory
• The Gas Laws
• Partial Pressures
• Effusion and Diffusion
• Real Gases

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Properties of Gases

• Three phases of matter
• solid
• liquid
• gas

Definite shape and volume
Definite volume, shape of container
Shape and volume of container
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Properties of Gases

• A gas is a collection of molecules that are very
  far apart on average.
• In air, gas molecules occupy only 0.1 of the
  total volume.
• In liquids, molecules occupy 70 of the total
  space.

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Properties of Gases

• Gases are highly compressible.
• Volume decreases when pressure is applied.
• Gases form homogeneous mixtures with each other
  regardless of the identities or relative
  proportions of the different gases.
• Water and gasoline heterogeneous mixture.
• Water vapor and gasoline vapor homogeneous
  mixture.

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Properties of Gases

• Chemical properties of gases vary depending on
  their composition.
• Air 78 N2 and 21 O2
• CO2 colorless, odorless
• CO colorless, odorless, highly toxic
• NO2 toxic, red-brown, irritant
• N2O colorless, sweet odor (laughing gas)

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Pressure

• Four quantities are commonly needed to describe a
  gas
• amount of gas
• temperature
• volume
• pressure

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Pressure

• Gases exert pressure on the objects in their
  surroundings.
• Pressure is caused by collisions between the gas
  molecules and objects with which they are in
  contact.
• Pressure the force exerted on a unit area
• P F
• A

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Pressure

• Atmospheric pressure the pressure exerted by
  gas molecules in the air on all objects exposed
  to the atmosphere
• Atmospheric pressure varies with altitude.

Altitude (ft above sea level) Atmospheric Pressure Atmospheric Pressure Atmospheric Pressure
Altitude (ft above sea level) in. Hg Torr psi
0 29.92 760 14.7
5000 24.9 632.5 12.23
10,000 20.58 522.7 10.1
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Pressure
Why does atmospheric pressure decrease with
increasing altitude?

• Gravity decreases
• Density of gas decreases
• Fewer gas molecules
• Fewer collisions
• Lower pressure

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Pressure

• Many different units used to report pressure.
• millimeters of Hg (mm Hg)
• inches of Hg (in. Hg)
• pounds per square inch (psi)
• atmosphere (atm)
• torr (torr)
• pascal (Pa) SI base unit
• kilopascal (kPa)

Must know units and abbreviations!!
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Pressure

• Relationships between different pressure units

1 atm 760 mm Hg 760 torr 29.92 in. Hg
14.7 psi 1.01325 x 105 Pa 101.325 kPa
Must be able to interconvert between units.
Memorize the ones in redIll give you the others
except kPa
You must know that 1 kPa 1000 Pa
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Pressure

• Example The measured pressure inside the eye of
  a hurricane was 669 torr. What was the pressure
  in atm?

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Pressure

• Example On a nice sunny day in Chicago the
  barometric pressure was 30.45 in. Hg. What was
  the pressure in Pa?

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Pressure

• Example On Titan, the largest moon of Saturn,
  the atmospheric pressure is 1.631 Pa. What is
  the pressure in atm?

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Kinetic Molecular Theory

• The behavior of gases can be described and
  explained using kinetic molecular theory.
• the theory of moving molecules
• You must know the basic ideas that are part of
  kinetic molecular theory.

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Kinetic Molecular Theory

• Gases consist of large numbers of molecules that
  are in continuous, random motion.
• The combined volume of all the molecules of the
  gas is negligible compared to the total volume in
  which the gas is contained.
• i.e. the molecules are very far apart on average

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Kinetic Molecular Theory

• Attractive and repulsive forces between gas
  molecules are negligible.
• Energy can be transferred between molecules
  during collisions, but the average kinetic energy
  of the molecules does not change as long as the
  temperature remains constant.
• Collisions are perfectly elastic.

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Kinetic Molecular Theory

• The average kinetic energy of the molecules is
  proportional to the absolute temperature.
• At any given temperature all molecules of a gas
  have the same average kinetic energy.
• As T (in K) increases,
• KE increases.

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Gas Laws

• Four variables are needed to define the physical
  condition or state of any gas
• Temperature (T)
• Pressure (P)
• Volume (V)
• Amount of gas (moles n)
• Equations relating these variables are known as
  the gas laws.

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Gas Laws

• Consider a fixed amount of gas that is confined
  to a container with a certain volume.

At a specific temperature, the gas sample will
exert a certain pressure on the container.
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Gas Laws
What will happen to the pressure if the volume is
decreased?
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Gas Laws

• As the volume of a fixed quantity of gas
  decreases, the pressure increases because
• gas molecules are more tightly packed together
• i.e. denser
• more collisions between gas molecules and the
  container
• greater pressure

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Gas Laws

• Boyles Law
• The volume of a fixed quantity of gas maintained
  at constant temperature is inversely proportional
  to the pressure.

• Mathematically,
• V k x 1 or PV k or P1V1 P2V2
• P
• at constant temperature and quantity of gas

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Gas Laws

• As liquid nitrogen (-196oC) is poured over a
  balloon, the volume of the balloon decreases.

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Gas Laws

• Charles Law
• The volume of a fixed amount of gas maintained at
  constant pressure is directly proportional to its
  absolute temperature.

V k x T or V k or V1
V2 T T1 T2 At constant pressure and quantity
of gas Remember T must be in Kelvin
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Gas Laws

• On a molecular level, as the temperature of a gas
  maintained at constant pressure decreases,
• KE decreases
• fewer collisions between gas molecules and the
  environment (i.e. container)
• volume decreases in order to maintain constant
  pressure

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Gas Laws
What happens when you blow up a balloon?
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Gas Laws
What happens when you blow up a balloon?
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Gas Laws
What happens when you blow up a balloon?
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Gas Laws
What happens when you blow up a balloon?
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Gas Laws
What happens when you blow up a balloon?

• the number of moles of gas (n) increases
• and
• the volume of the gas (balloon) increases

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Gas Laws

• Avogadros Law
• The volume of a gas maintained at constant
  temperature and constant pressure is directly
  proportional to the number of moles of the gas.
• Mathematically,
• V constant x n
• At constant temperature and pressure

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Gas Laws

• At any given temperature and pressure, as the
  amount of gas increases,
• the number of gas molecules increases
• the number of collisions between gas molecules
  and the environment (container) increases
• the volume must increase in order to maintain
  constant pressure

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Gas Laws

• In a chemical reaction, we use the coefficients
  to tell us how many moles or molecules are used
  or produced in a chemical reaction.
• N2 (g) 3 H2 (g) ? 2 NH3 (g)
• 1 mole of nitrogen reacts with 3 moles of
  hydrogen to produce 2 moles of ammonia

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Gas Laws

• Since the volume of a gas is directly
  proportional to the number of moles of gas at
  constant temperature and pressure, we can also
  use the coefficients to represent the volume of
  a gas involved in a reaction. (Avogadros
  Hypothesis)
• N2 (g) 3 H2 (g) ? 2 NH3 (g)
• 1 liter of nitrogen reacts with 3 liters of
  hydrogen to produce 2 liters of ammonia

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Gas Laws

• Boyles Law, Charles Law, and Avogadros Law can
  be combined to make a more general gas law
• Ideal Gas Law
• PV nRT
• where P pressure
• V volume
• n moles
• T temperature (K)
• R gas constant

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Gas Laws

• The value of the gas constant (R) depends on the
  units of P, V, n, and T.
• T must always be in Kelvin
• n is usually in moles
• If P (atm) and V (L),
• then R 0.08206 atm.L
• mol.K
• If P (torr) and V (L),
• then R 62.36 L.torr
• mol.K

I will give you these on the test.