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KinBackup

This website contains an archive of files for the Acorn Electron, BBC Micro, Acorn Archimedes, Commodore 16 and Commodore 64 computers, which Dominic Ford has rescued from his private collection of floppy disks and cassettes.

Some of these files were originally commercial releases in the 1980s and 1990s, but they are now widely available online. I assume that copyright over them is no longer being asserted. If you own the copyright and would like files to be removed, please contact me.

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File contents
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�CEThe Kinetic Theory of Gases
�CEby M.T.Ford, January 1996

       All of our present understanding of the behaviour of materials is
based on one important fact: that all matter is composed of a finite
number of microscopic particles. This idea was first put forward by the
Greek philosopher Democritus in about 400 BC, but no experimental evidence
was found for his theory until the early nineteenth century. Until then,
it was assumed that if you took a small drop of sea water, and magnified
it, say, to the volume of the Mediterranean Sea, its properties would be
exactly the same as those of the water in a real sea. That is to say, they
thought the properties of a substance such as water were identical on any
scale. However, it is now known that the water in this experiment would
not appear like an ordinary sea, but would consist of hard rubbery
objects, a few centimetres across. Each of these objects would contain
three overlapping spheres: a large oxygen atom with two smaller hydrogen
atoms attached. These objects, called molecules, would be moving around
rapidly, sometimes squeezing past each other, so any object thrown into
the water would be jostled about randomly at high speed as the molecules
bombarded it.
	In 1808, Dalton carried out experiments on the volumes of gases
taking part in chemical reactions and he, together with Avogadro, came to
the conclusion that these molecules must exist. Later, in 1827, the
botanist Robert Brown observed that pollen grains, or fine grains of
quartz would undergo small random movements if placed on the surface of
water. Similarly, if smoke particles are observed through a microscope,
they are also seen to jiggle about in random directions. This is exactly
the motion that we would expect for an object in a "sea" of molecules: the
particle is continually being bombarded by air (or water) molecules on all
sides, but these collisions occur randomly. In a short time interval, the
particle may for example be struck by 100 molecules on one side, and only
97 molecules on the other. If the particle is sufficiently light, this
difference will result in an acceleration. In the next time interval, the
particle may accelerate in a different direction, so the effect we see is
a "jiggling" motion. This provided important evidence that there molecules
in the air or water, which are constantly moving in random directions,
even though the material may seem perfectly still in bulk.  Kinetic theory
is a way of explaining the bulk properties of gases  in terms of this
motion (the word "Kinetic" is derived from the Greek verb "to move").
	It can be difficult to visualise how this motion results in the
effects we see, due to the high speed and minute size of the particles.
For example, an oxygen molecule at room temperature travels at an average
speed of 440 metres per second, or 990 mph. Yet the length of the molecule
is only about one two-millionth of a millimetre. That means that when it
collides with a stationary molecule, it approaches at a rate of about one
million million times is own length every second (compare this with a
tennis ball, travelling at 100 mph, which strikes a wall at a speed of 600
times its own lengh per second - or a bullet, at 100000 times its own
length per second). So even if we could magnify a gas to the extent that
the molecules were no longer microscopically small, they would be
travelling thousands of millions of times too fast to be seen.
	The other problem with visualising a gas, is that the molecules
move and bounce off each other in three dimensions, while we are used to
using two-dimensional diagrams. For example, the path of a golf ball, or a
snooker ball is restricted to one plane (assuming there is no wind, in the
case of the golf ball), so we can represent their movement on a piece of
paper, and measure angles easily. In the case of a gas, this is not the
case, so it is best to visualise it using a two-dimensional analogy. We
represent the molecules by balls, moving about on a snooker table (without
pockets!). It is essential that the balls do not slow down when they roll
or hit each other if they are to represent a gas accurately, otherwise the
gas would just fall down to the ground and collapse into a solid.
Therefore we must imagine them as hard rubber balls, which would bounce
forever if they were dropped on to the table. This model is of limited use
for mathematical analysis but the general principles will be the same as
in a real gas.
	The easiest phenomenon to explain with this model is pressure. If
all the balls are set in random motion, they will frequently bounce
against each other and the sides of the table. As a ball strikes a cushion
(at any angle), it is reflected, like a beam of light from a mirror.
According to Newton's First Law of Motion, a body cannot change its speed
or direction of motion without a force acting on it, so the cushion must
have exerted some force on the ball. Now the speed (velocity, strictly) 
of the ball can be divided into two components, one parallel to the
cushion and one perpendicular to it - i.e. the rate that it is travelling
across the table, and the rate that it is travelling up or down the table.
But the component parallel to the cushion was unaffected by the collision,
so the cushion must have exerted a force on the ball in the direction
perpendicular to itself. According to Newton's Third Law, the cushion must
have experienced an equal force in the opposite direction (also
perpendicular to itself), away from the balls. This force is not in the
direction one might expect intuitively if the ball approaches at an angle.
But there are many of these collisions every second in a real gas - for
example a full stop on this page is struck about two thousand million
million million times by air molecules every second, so it is impossible
for us to see the effect of each collision individually. The result is a
spread out force, like that seen in an inflated balloon, where the air at
each point on the surface of the rubber is pushing at right angles to the
surface. Although the air exerts a pressure on all surfaces around us, of
about 100000 pascals (1 bar, or 15 pounds per square inch), we do not
normally feel it. This is because there is a similar pressure inside us
balancing out the external pressure. In a balloon, the pressure is greater
on the inside than on the outside, as the inside surface is struck by more
molecules every second than the outside surface. So it is really only the
pressure difference that we feel in a balloon. 
	Returning to the analogy, suppose we double the length of the
table, without altering the number of balls, or the width of the table.
Now there are only half as many balls in each square metre of the table as
there were before. It follows from this that the end cushion of the table
(which is the same length as before) is only struck half as often - i.e.
the average force on each metre of the cushion has halved. This is
equivalent to saying that if you double the volume of a gas, you halve its
pressure (pressure in a gas is defined by the average force acting on a
square metre of its container - this is the three-dimensional equivalent
of the average force per metre of a cushion). This is a statement of
Boyle's Law, which was discovered in 1662. 
	Temperature can also be represented in this model. The simplest way
to understand the temperature of a gas (other than how hot it is!) is the
speed of its molecules - the faster they move, the higher the
temperatuere. In reality, the molecules do not all move at the same speed
- their speeds form a distribution - so we must measure the temperure by
their average speed, or more strictly, the average of the squares of their
speeds. We have to be careful about the scale we use for measuring
temperature in physics. We are used to measuring temperatures on the
Celsius, or Fahrenheit scales in everyday life, where temperatures both
above and below zero are possible. Of course the speeds of the molecules
can never be below zero, so is sensible to choose a scale where the
molecules are stationary at zero degrees. This scale is called the
absolute or Kelvin scale, and is obtained by simply taking a temperature
on the Celsius scale and adding 273.15 degrees. On this scale, doubling a
temperature actually means doubling the squares of the speeds of the
molecules.
	Suppose we start the balls moving again, and then double their
speeds. The squares of their speeds, and hence their temperatures (on the
Kelvin scale) increase by a factor of four. The balls will strike the
cushions moving twice as fast as before, on average, and so the force
exerted on the cushion in each collision will be doubled. But the balls
also strike the cushions twice as often as before (imagine filming the
balls, and then playing the film back at double speed - you will see the
same number of collisions in half the time). So the pressure on the
cushions, which depends on the force and frequency of the collisions, will
increase by a factor of four. Since the temperature also increased by a
factor of four, we can deduce that the pressure of a gas varies in
proportion to its temperature (on the Kelvin scale), provided its volume
is not changed. This, combined mathematically with Boyle's law, leads to
the conclusion that the volume of a gas varies in proportion to its
temperature (i.e. if you double the temperature you will double the
volume), provided the pressure is not changed. This is called Charles's
Law, and was not discovered until over a century later, in 1787. It allows
us to understand the expansion of gases when they are heated at
atmospheric pressure, and the important phenomenon of convection, which
carries air up from a heating element, drawing up cooler air from below.
	Now imagine we place a moveable cushion across the middle of the
table, representing a piston in a gas jar. The balls can make the cushion
move when they strike it, but they can never cross to the other side. We
now place an equal number of balls on each side of the cushion, and set
them all in motion. We now double the speeds of the balls in one half
(thus raising their temperature and pressure), and leave the other half
unchanged. The cushion will now experience a greater pressure on one side
than the other, and will move, so the the area of the "hot" side
increases, and that of the "cold" side decreases. Eventually it will reach
a position where the two pressures are equal, and it experiences no net
force either way - this is the equilibrium position. But it is still being
continually bombarded by balls on both sides, so it will undergo small
random motions, like the pollen grains observed by Robert Brown. When one
of the fast "hot" molecules strikes the cushion, it will tend to make the
cushion move fast towards the slow "cold" molecules on the other side. The
fast molecule will be slowed down in this interaction, as it has given
part of its energy to the cushion. The next time a slow molecule strikes
it, it will receive a hard "kick" from the cushion, and will probably move
away faster than it came in. As this process continues, the result on
average is that the fast molecules slow down, and the slow molecules speed
up. Eventually the average speeds on both sides of the cushion will
equalise, and they will be at the same temperature. They are now said to
be in thermal equilibrium, as there is no net transfer of heat either way.
This is a property of temperature which we experience in everyday life: if
we heat the air in a house to, say, 20XC, when the air outside is at 10XC,
then switch the heaters off, the heat will gradually escape through the
windows, making the inside colder, and the outside warmer (by a very, very
small amount, as there is a much larger amount of air outside than
inside). The windows are, of course, much more complicated structures than
the cushion used in the analogy, as they contain many molecules between
one side and the other. But if each molecule in the window is treated as
an individual "cushion", it will pass the heat on to the next molecule,
and eventually through to the other side. The heat flow stops when the
house temperature has reached  10XC, and the house is then in thermal
equilibrium with the air outside, and indeed with anything else that
happens to be at 10XC. This statement is so obvious that it can easily be
overlooked, and had to be named the Zeroth Law of Thermodynamics, after
the First and Second Laws had already been formulated.
       One of the arguments used in this explanation was that each time a
ball strikes the partitiom, it pushes the partition back, and in doing so
cools itself. This principle is used in the internal combustion engine:
fuel gases are compressed inside the cylinder, then a spark causes them to
explode so the temperature and pressure suddenly increase as the chemical
energy is released. The molecules then strike the piston at high speed,
accelerating the piston, and cooling the gases. Thus the energy from the
fuel is converted first into heat, then into motion. But it is important
that this process occurs at very high speed, since while the gases are
hot, there will be a tendency for heat to leak out through the side of the
cylinder into the bodywork and atmosphere, and be wasted. If the expansion
occurs at such high speed that no energy has time to escape, it is called
an adiabatic expansion.
	We also make use of expansion in refrigerators. A liquid
refrigerant is pumped into the cold box, and then allowed to expand and
vaporise. As it does so, the molecules are forced to push back the vapour
occupying the pipes, just as the gases in a cylinder push back a piston.
However, in this instance, we actually want heat to leak through the sides
of the pipe, in this case from the outside (i.e. the cold box) to the
inside. So we carry out the expansion quite slowly, to give time for this
to occur. If an expansion occurs so slowly that the temperature doesn't
change at all (because it is continually equalising itself with the
outside temperature), then it is called an isothermal expansion.
	Returning to the snooker analogy, suppose we introduce a set of
lighter balls on to the table, in addition to the set that was used
previously (leaving out the cushion across the middle). When the balls are
set in motion, all at the same speed, the heavy balls will collide with
the light ones. In these collisions, Newton's Third Law dictates that the
same force must act on both balls in opposite directions. But this force
will have a greater effect on the light balls than the heavy ones (imagine
a collision between a car and a lorry - the lorry may continue along its
original path almost unimpeded, while the car is thrown violently across
the road). The result is that the light balls move faster on average than
the heavy ones. But eventually the balls must reach thermal equilibrium,
and we know from the Zeroth Law that they must be at the same temperature,
despite the fact that the two sets of balls move at different speeds. This
suggests that in general, temperature is not just a measure of the speeds
of the molecules - or even the squares of the speeds - but involves their
mass as well. In fact it turns out that temperature is a measure of the
energy of the molecules. All motion has a quantity of energy associated
with it, called its kinetic energy, and when two different sets of
molecules are in thermal equilibrium, it is not necessarily their speeds,
but their energies which are equal. This explains why temperature is such
a useful concept in thermodynamics - the study of energy and heat changes.  
	To illustrate this principle using real gases, consider helium and
radon. In both of these two gases, the molecules all consist of a single
spherical atom, but in radon, the atoms are 56 times heavier than those in
helium. This means that at any particular temperature, the atoms in helium
travel about 7 times faster than those in radon. This can easily be tested
experimentally by measuring the speed of sound in each gas. Sound waves
are carried through a gas by the motion of molecules, so the speed of
sound acts as a good measure of the speeds of the molecules themselves.
	Something we have left out of the model is the force of gravity.
The molecules in a gas are all acting as projectiles - when they are
moving upwards they are getting slower, and when they move downwards they
get faster. So why doesn't the air in a room all concentrate at the
bottom, leaving a vacuum at the top? The answer lies in the high speed of
the molecules. If you were to take an evacuated room, with oxygen
molecules held against the floor, then suddenly released them all, they
would simultaneously start rising towards the ceiling. Now all molecules
travelling with an initial speed greater than about 7 metres per second
will reach the ceiling (assuming the room is about 2.5 metres high). The
rest will rise to a certain height then fall back, like a ball thrown
gently into the air. But the average speed of an oxygen molecule is about
440 metres per second, so the vast majority will have speeds much greater
than 7 metres per second, and will reach the ceiling. Hence, nearly every
molecule at the bottom of a room full of air has enough energy to reach
the top, and so there are nearly as many molecules at the top of a room as
there are at the bottom. When we look at the Earth's atmosphere as a
whole, there is a very striking difference in air density at different
heights, because the molecules would then need a much greater energy to
travel from the bottom to the top. But our snooker table is only supposed
to represent a small quantity of gas, so gravity can be neglected. (In
reality, of course, molecules do not travel straight from the bottom to
the top of a room, because they undergo many collisions on the way.
However, collisions are just as likely to reduce the energy of a molecule
as to increase it, so their effects average out to zero, and the above
line of argument is still valid). 
	Another important effect that we have left out is the attraction
that molecules have for one another. If we were to suddenly stop the
motion of every molecule in a gas, they would not just sit still, like the
balls on a snooker table, even without the influence of gravity. They
would all collapse together and form a solid. If we then gradually
increased the temperature of this solid, thus speeding the molecules up,
they would eventually resist their attractive forces sufficiently to form
a liquid, then a gas again. But in changing from a liquid to a gas, the
average distance between adjacent molecules increases by a factor of 10,
and the attractive forces become so weak that they are almost negligible
(this is like pulling two magnets apart: when they are close together,
they attract each other strongly, but the attraction quickly gets weaker
as they are pulled apart). It is convenient to neglect these forces when
studying gases, because the laws governing their behaviour become very
much simpler. These approximate "laws" are called the laws of "perfect
gases" or "ideal gases", which are defined to be gases with no attractive
forces, and infinitely small molecules. There is no such thing as an ideal
gas in reality, but most gases at atmospheric pressure fit the ideal gas
laws quite closely, so the absence of attractive forces in our model is
not a major problem.
	Another important effect we often observe in a gas is viscosity. If
we move rapidly through a gas, by standing in the wind, or by driving
through it in a car, we feel a resisting force. This can partially be
explained by the fact that air is "bunching up" in front of us, and not
catching up behind us as we move. This generates a high pressure region in
front and a low pressure region behind, so the net result is a backward
force.
	This explanation seems to solve the problem completely, but
consider the following situation. A solid cylinder is rotated at high
speed by an electric motor, then a hollow cylinder is placed around the
rotating cylinder, with an air space of about one millimetre between them.
The outer cylinder experiences a force, trying to rotate it in the same
direction as the motion of the inner cylinder. One would imagine
intuitively that this would only happen in gases where the molecules are
attracted to one another, but in fact the effect occurs even in ideal
gases. The force cannot be explained in the same way as the air resistance
on a car, since there is no air being "pushed out of the way", and no
regions of high and low pressure. In fact the reason for the gas's
resistance to this motion is as follows:
	While the outer cylinder is being held stationary, the air
immediately next to it acts as a stationary air mass. That is to say, the
molecules are moving, but there are as many moving one way as another, so
there is no bulk flow. The air immediately next to the inner cylinder
though, is rotating around at the speed of the cylinder, since each time a
molecule strikes the cylinder, it receives a "kick" to keep it moving
round. Thus the air between the cylinders can be divided into thin
cylindrical layers, each layer moving at a slightly different speed from
the layers next to it. Now imagine two adjacent layers of air as two
trains moving on parallel rails at slightly different speeds. The engines
are both switched off, and the trains coast along on perfectly smooth
rails without losing any speed. Now the trains are both full of people,
representing air molecules within the layers. The people are moving around
blindly inside the trains in completely random directions, and sometimes
one of them will step from one train on to the other. If a person steps
from the slow train to the fast one, the fast train will have to exert a
force on him to bring him up to the right speed. But the train will
experience a force in the opposite direction, slowing itself down
slightly. Similarly, if a person steps from the fast train to the slow
one, the slow train will speed up slightly. As people continue to step
from one train to the other, the slow train speeds up and the fast train
slows down, until they are both travelling at the same speed. So each
layer of air in the cylinder experiment is exerting a force on the next
layer, trying to equalise their speeds. This force is conducted through to
the outer cylinder, which tries to equalise its speed with the inner one.
To stop it moving, we have to push it the other way, so we are effectively
acting as the "engines" of the trains, which keep the fast trains moving
faster than the slow trains.
	We have seen that most of the important phenomena occurring in
gases can be explained using simple analogies, which can help by reducing
the number of dimensions involved, and by changing the scales and orders
of magnitude. It is important to realise though, that these analogies do
not provide rigorous proofs of the reasons behind an effect - they just
help us to visualise the possible explanations. It is possible to imagine
the behaviour of solids and liquids by similar methods, but in these cases
it is much more difficult to find a suitable analogy. For example in a
liquid, we must consider gravity, intermolecular forces, and the
liquid-vapour interface at the surface, to explain its behaviour fully. It
is the very simple approximations of the ideal (perfect) gas model that
allow us to reach conclusions so easily. 
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00005f10  74 61 62 6c 65 1a 20 61  6e 61 6c 6f 67 79 2e 1a  |table. analogy..|
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00005fe0  61 70 70 72 6f 78 69 6d  61 74 69 6f 6e 73 20 6f  |approximations o|
00005ff0  66 20 74 68 65 20 69 64  65 61 6c 20 28 70 65 72  |f the ideal (per|
00006000  66 65 63 74 29 20 67 61  73 20 6d 6f 64 65 6c 20  |fect) gas model |
00006010  74 68 61 74 0d 61 6c 6c  6f 77 20 75 73 20 74 6f  |that.allow us to|
00006020  20 72 65 61 63 68 20 63  6f 6e 63 6c 75 73 69 6f  | reach conclusio|
00006030  6e 73 20 73 6f 20 65 61  73 69 6c 79 2e 20 0d     |ns so easily. .|
0000603f
KinBackup.m0
KinBackup.m1
KinBackup.m2
KinBackup.m4
KinBackup.m5