Download The Ideal Gas Laws

The Ideal Gas Laws
Chapter 14
Expectations
After this chapter, students will:
Know what a “mole” is
Understand and apply atomic mass, the atomic
mass unit, and Avogadro’s number
Understand how an ideal gas differs from real
ones
Use the ideal gas equation, Boyle’s Law, and
Charles’ Law, to solve problems
Expectations
After this chapter, students will:
understand the connection between the
macroscopic properties of gases and the
microscopic mechanics of gas molecules
Preliminaries: the Mole
A mole is a very large number of discrete objects,
such as atoms, molecules, or sand grains.
Specifically, it is Avogadro’s Number (NA) of such
things: 6.022×1023 of them.
The mole (“mol”) is not a dimensional unit; it is a
label.
Amadeo Avogadro
1776 – 1856
Native of Turin, Italy
Hypothesized that equal volumes
of gases at the same temperature
and pressure contained equal
numbers of molecules.
(He was correct, too.)
The Mole and Atomic Mass
Mathematical definition: 12 g of C12 contains one
mole of carbon-12 atoms.
Mass of one
C12
atom:
12 g
.012 kg
- 26


1
.
993

10
kg
23
N A 6.022 10
The mass of one C12 atom is also 12 atomic mass
units (amu), so:
1.993 10-26 kg
1 amu 
 1.66110-27 kg
12
The Mole and Atomic Mass
Atomic masses for the elements may be found in the
periodic table of the elements, located inside the
back cover of your textbook.
These are often erroneously called “atomic
weights.”
Atomic masses may be added to calculate molecular
masses for chemical compounds (or diatomic
elements).
The Mole: Calculations
If we have N particles, how many moles is that?
number of moles
N
n
NA
If we have a given mass of something, how many
moles do we have?
mass of sample
mass of sample
n

mass per mole
atomic or molecular mass
The Ideal Gas
The notion of an “ideal” gas developed from the
efforts of scientists in the 18th and 19th centuries
to link the macroscopic behavior of gases
(volume, temperature, and pressure) to the
Newtonian mechanics of the tiny particles that
were increasingly seen as the microscopic
constituents of gases.
The Ideal Gas
An ideal gas was one whose particles are wellbehaved, in terms of the Newtonian theory of
collisions: elastic collisions and the impulsemomentum theorem.
An ideal gas is one in which the particles have no
interaction, except for perfectly-elastic collisions
with each other, and with the walls of their
container.
The Ideal Gas
An ideal gas has no chemistry. That is, the particles
(atoms or molecules) have no tendency to “stick”
to other particles through chemical bonds.
Inert gases (He, Ne, Ar, Kr, Xe, Rn) at low densities
are very good approximations to the ideal gas.
Our analytic model of the ideal gas gives us insights
into the properties of many real gases, inert or not.
The Ideal Gas Equation
Observations from experience
The pressure of a gas is directly proportional to the
number of moles of particles in a given space.
Example: blow up a balloon, and you’re adding to
n, the number of moles of molecules.
Conclusion:
Pn
The Ideal Gas Equation
Observations from experience
The pressure of a gas is directly proportional to its
temperature. Example: toss a spray can into a fire
(no, wait, really, don’t do it, just think about it).
Increasing pressure will cause the can to fail
catastrophically.
Conclusion:
P T
The Ideal Gas Equation
Observations from experience
The pressure of a gas is inversely proportional to its
volume. Example: squeeze the air in a half-filled
balloon down to one end and squeeze it tighter.
Increased pressure makes the balloon’s skin tight.
Conclusion:
1
P
V
The Ideal Gas Equation
Combine the observations
nT
P
V
A constant of proportionality, R, makes this an
equation:
nT
PR
or PV  nRT
V
The Ideal Gas Equation
The constant of proportionality, R, is called the
universal gas constant. Its value and units
depend on the units used for P, V, and T.
pressure
volume
PV  nRT
absolute
temperature
universal gas constant
number of moles
Value and SI units of R: 8.31 J / (mol K)
The Ideal Gas Equation
We can also write the ideal gas equation in terms of
the number of particles, N, instead of the number
of moles, n.
Since N = n·NA, we can both multiply and divide the
right-hand side by NA:
 R 
T
PV  nN A 
 NA 
Boltzmann’s constant
R
- 23
PV  NkT where k 
 1.38 10 J/K
NA
Ludwig Boltzmann
Austrian physicist
1844 – 1906
Boyle’s Law
Suppose we hold both n and T constant: how are P
and V related?
PV  nRT

P1V1  P2V2
This is called Boyle’s Law.
PV  constant
Robert Boyle
Irish mathematician
1627 – 1691
Charles’ Law
Suppose we hold both n and P constant: how are T
and V related?
PV  nRT

V nR

 constant
T
P
V1 V2

T1 T2
This is called Charles’ Law.
Jacques Alexandre Cesar Charles
French scientist
1746 – 1823
Built and flew the first
large hydrogen-filled
balloon.
Kinetic Theory of the Ideal Gas
Macroscopic properties of a gas: temperature,
pressure, volume, density
Microscopic properties of the particles making up
the gas: mass, velocity, momentum, kinetic
energy
How are they related?
Kinetic Theory of the Ideal Gas
Consider a gas molecule contained in a cube having
edge length L.
The molecule’s mass is m, and
its velocity (in the X direction
only) is v.
Time between collisions with the
right-hand wall:
2L
t
v
Kinetic Theory of the Ideal Gas
The time between collisions with the right-hand wall is
just the round-trip time:
2L
t
v
From the impulse-momentum
theorem, we can calculate the
average force exerted on the
particle by the wall:
J  Ft  p  p f  p0  m v   mv
Kinetic Theory of the Ideal Gas
Substitute for the time and simplify:
J  F t  p  p f  p0  m v   mv
2L
F
 2mv
v
2
 mv
F
L
By Newton’s third law, the average
force exerted on the wall is
mv
F
L
2
Kinetic Theory of the Ideal Gas
The average force on the wall from one particle is
mv
F
L
2
If there are N particles, and
their directions are random, we
could expect 1/3 of them to be
moving in the X direction.
Total force on the wall:
N mv
F
3 L
2
Kinetic Theory of the Ideal Gas
Average pressure on the wall:
N mv 2
2
F
N
m
v
P   3 2L 
3
A
L
3 L
But
V L
N mv
P
3 V
3
2
So:
N
 PV  mv 2  NkT
3
Kinetic Theory of the Ideal Gas
Substituting kinetic energy:
1 2 1 1
2
2
kT  mv   2 mv   KE
3
3 2
 3
3
KE  kT
2
So, we see that for an ideal gas,
the average molecular kinetic energy
is directly proportional to the
absolute temperature.
Kinetic Theory of the Ideal Gas
3
KE  kT This result is true for any ideal gas.
2
By a similar argument, if an ideal gas is monatomic
(the gas particles are single atoms), the internal
energy of n moles of the gas at an absolute
temperature T is
3
U  nRT
2