Transcript Physical Characteristics of Gases

Chapter Ten
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based on the idea that particles of matter are
always in motion
explains the properties of solids, liquids, and
gases in terms of the energy of particles and the
forces that act between them
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helps understand the behavior of gas molecules
and the physical properties of gases
provides a model of an ideal gas
Ideal Gas: an imaginary gas that perfectly fits all
the assumptions of the kinetic-molecular
theory
1. Gases consist of large #s of tiny particles
that are far apart relative to their size.
-volume 1000 times greater than that of
liquid or solid; much of volume is empty
space; lower densities; easily compressed
2. Collisions between gas particles and
between particles and container walls are
elastic.
Elastic collision – one in which there is no
net loss of kinetic energy. Kinetic energy
is transferred between particles during
collisions.
3. Gas particles are in continuous, rapid,
random motion.
-they therefore possess kinetic energy
4. There are no forces of attraction or
repulsion between gas particles
5. The average KE of gas particles depends on
the temp of the gas. Speeds and energies
increase with temp increase, and decrease
with temp decrease. ALL gases at the
same temp have the same avg KE.
*Many gases behave nearly ideally if
pressure is low and temp is high.
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Expansion
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No definite shape or
volume
Completely fill any
container and take
its shape
No significant
attraction or
repulsion
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Fluidity
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Because attractive forces
are insignificant, gas
particles glide easily past
one another. This makes
them like liquids.
Because gases and
liquids flow, they are
both referred to as fluids.
Fluids – gases and
liquids that flow
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Low Density
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Particles are very far
apart in gases, thus
they have very low
densities (~1/1000
that of same
substance as liquid or
solid)
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Compressibility
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During compression,
gas particles that are
very far apart are
crowded closer
together. The volume
of a given sample of a
gas can be greatly
decreased.
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Diffusion
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Spontaneous mixing of
the particles of 2
substances caused by
their random motion
Rate of diffusion depends
on 3 properties of the gas
particles: their speeds,
their diameters, and the
attractive forces between
them
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Effusion
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A process by which gas
particles pass through a
tiny opening
Rates of effusion directly
proportional to the
velocities of the particles
(molecules of low mass
effuse faster than
molecules of high mass)
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Real gas – a gas that does not behave
completely according to the assumptions of
the kinetic-molecular theory
Johannes van der Waals – pointed out that
particles of real gases occupy space and exert
attractive forces. At high pressures and low
temps the deviation from ideal behavior may
be considerable.
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KM theory more likely to hold true for gases
w/little attraction. He and Ne show ideal behavior
over wide range of temps and pressures. The more
polar the gas molecule, the greater the forces and
the less ideal its behavior.
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To describe a gas fully, you need to
state four measurable quantities:
volume, temperature, number of
molecules, and pressure.
Star this – you have to know it
-the force per unit area on a surface
-the smaller the area of contact, the greater the pressure
Newton – (SI) N – the force that will increase the speed of a one-kg
mass by one meter per second each second it is applied (at Earth’s
surface, each kg exerts 9.8 N of force due to gravity)
Atmospheric pressure – sum of individual pressures of various gases
in atmosphere (mostly N and O); at sea level is about 1.0 N/cm2
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Barometer
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A device used to
measure
atmospheric
pressure
First introduced by
Evangelista
Torricelli in the
1600s
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mercury in the tube falls
only until the pressure
exerted by its weight
equals the pressure
exerted by the
atmosphere
exact height of mercury
in tube depends on
atmospheric pressure
**at sea level and zero
degrees, average
atmospheric pressure can
support a 760 mm
column of mercury
1 atm = 760 mm Hg = 760 torr = 101.325 kPa = 1.01324 x 105 Pa
Simple mathematical relationships
between the volume, temperature,
pressure, and amount of a gas.
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Pressure is caused by moving particles hitting
the container walls. Decreasing volume
w/same # of particles increases collisions
therefore increasing pressure. Doubling
pressure – volume cut by one-half
Reducing pressure by one-half, volume
doubles
As one variable increases, the other decreases
Inverse relationship – graphs a curve
(hyperbola)
-states that the volume of a fixed mass of gas varies
inversely w/the pressure at constant temperature
PV = k
(k is constant for a given sample of gas)
If P changes, V will change, but k will remain constant.
P1V1 = k = P2V2
-most useful for changing conditions of
a given sample of a gas
Hint: We Boil Peas and Vegetables
Boyle’s Law Relationship
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If pressure is constant, gases expand when
heated
Relationship discovered by Jacques Charles in
1787
All gases expand to the same extent when
heated thru the same temp interval
Real gases can’t be cooled to -273C; before
reaching that temp, intermolecular forces
exceed KE of molecules and the gases
condense to form liquids and solids
Avg KE of gas molecules is most closely
related to Kelvin temp
-states that the volume of a fixed mass of gas at constant
pressure varies directly with the Kelvin temperature
Gas volume and Kelvin temperature are directly
proportional to each other.
V/T = k or V = kT
For changing conditions:
V1/V2 = T1/T2
Hint: Charlie’s angels are on TV
Charles’s Law Relationship
Directly proportional – graphs a straight line
Joseph Gay-Lussac recognized this relationship in
1802:
-for a fixed quantity of gas at constant volume,
pressure should be directly proportional to the
Kelvin temp, which depends directly on avg
KE
Gay-Lussac’s law: the pressure of a gas at constant
volume varies directly with the Kelvin temp
P = kT or P/T = k
Gay-Lussac’s Law Relationship
A gas sample often undergoes changes in T, P, and V
all at the same time. Boyle’s, Charles’s, and GayLussac’s laws can be combined into a single expression
that is useful in such situations.
-expresses the relationship between P, V, and T of a
fixed amount of gas
PV = K
V/T = k
P/T = k
Boyle’s
Charles’s
Gay-Lussac’s
therefore
PV/T = k
Combined
Hint: Peas and
Vegetables on the Table
Partial pressure – the pressure of each gas in a mixture
Dalton’s law of partial
pressures – states that
the total pressure of a
mixture of gases is
equal to the sum of the
partial pressures of the
component gases. True
for any # of gases
present.