Gravity Chapter
3 continued
Why did just two atoms join in some cases, and not more, as in other
cases, and also, as the bond between the two elemental atoms must be
strong enough to withstand innumerable, random, high velocity collisions
without separating, how could this strong bond then be broken in the
process of the change of state to a compound?
This would apply particularly in the case of Hydrogen Chloride,
where di-atomic molecules of chlorine, of relative atomic
mass of 35.5, on the approach of another di-atomic molecule of hydrogen
of mass 1, divide and each single atom joins with one of the hydrogen
atoms to form a HCl molecule. One must assume that this chemical
bonding occurs ultimately as a result of molecular collisions between
the different molecules, but as to how the hydrogen molecules, having
masses 1/35.5th of the chlorine molecules, can break the proportionately
greater attractive forces of the chlorine molecule and split it in
two, it is difficult to imagine.
This idea therefore not only contradicted the earlier assumptions
of kinetic-atomic theory but was also illogical and accordingly
was rejected by most of the eminent chemists of the time.
Other ideas were of course proposed, but none were capable of providing
practically workable laws or formulae.
Brownian Motion
In 1827 the British botanist Robert Brown discovered
the phenomenon now called Brownian motion, which is the
observed random movement of microscopic particles suspended in
a gas or liquid. This motion of, for example, grains of
pollen or smoke particles in air, appears to be completely random in
both direction and duration.
This was later put forward as a visual manifestation of
the effect of the kinetic motion of atoms/molecules colliding
with these particles. It was suggested that the inherent motion of
the molecules causes the molecules to strike the suspended particles
at random and that the impact makes these particles move.
Clerk Maxwell
Debate continued on the merits of one or the other theories
of matter during the first half of the 1800’s but the
next significant development for kinetic-atomic theory
came when Clerk Maxwell in 1859 put forward his ‘Law of Distribution
of Velocities’ as a statistical or mathematical explanation of the
distribution of kinetic molecular velocities in gases.
The importance of the Maxwell distribution function, and
of the later, more general Maxwell-Boltzmann distribution
function ‘is that they contain all the information
necessary to calculate any measurable variable of a gas’ such as the pressure, temperature,
or volume.
In other words this enabled kinetic atomic theory to be
put to some practical use, for example by chemists and
engineers, to enable them to predict (albeit approximately) the behaviour
of gases in various conditions of pressure and temperature.
Clerk Maxwell based his statistics on the following assumptions.
1) Molecules are perfectly elastic balls of atomic dimensions
that are in perpetual random motion.
2) The average kinetic energy of the molecules is proportional
to the absolute temperature of the gas.
3) The molecules do not exert any appreciable attraction
on each other.
4) The volume of the molecules is infinitesimal when compared
to the volume of the gas.
5) The time spent in collisions is small compared with
the time between collisions
(Inherent in assumption 4 is the concept of ‘empty space’,
however Clerk Maxwell did not define this, either as a
pure vacuum or as an ‘aether’, however he did not accept that a vacuum
was possible and always believed in the concept of a media (an aether)
through which electromagnetic radiation travelled)
Kinetic Theory
Some principles of kinetic-atomic theory are described
as follows: -
Atoms in a gas, within a container, are ‘rushing
around at different velocities and bouncing off each other and
the walls like a three-dimensional game of billiards’
and ‘are moving in random directions, and because as many move in one
direction as another, the average velocity of the molecules is zero’ - in other
words the gas as a whole is not moving or producing unequal pressure
on any inside surface of the container.
‘Pressure arises from the multiple collisions the atoms
of a gas have with the walls that contain the gas’ and
‘heat applied to a gas results in an increase in the velocity
of the atoms and a corresponding increase in collisions with the
walls’
Also ‘when the fast moving atoms of a hot gas
collide with slower moving atoms of a cooler gas, kinetic energy is
transferred from the ‘hot’ to the ‘cold’ atoms’.
‘The collisions between atoms/molecules are completely
elastic’, or in other words no energy of motion or ‘kinetic’
energy is lost as a result of any collision with other atoms/molecules
of the gas or of the container.
‘The duration of collisions of atoms is about one thousandth
of the time between collisions. Atoms spend the overwhelming
part of their time in free motion, and collisions are a
rare event in their life.’
In addition the theory suggests that the atoms of a gas
only take up a minute proportion of the actual space the
gas occupies. ‘An atom generally takes up only
1/1000th of the volume available to it and if we were to scale atoms
to the size of human beings with a radius of 0.5 m, they would be
spaced some 10m apart.’
In other words in any given volume of a gas only about 0.1% is matter
in the form of atoms.
To put this in some sort of perspective 1000 cubic centimetres
(one Litre) of gas contains a total volume of atomic matter
that could be fitted into 1 cc while the remaining 999 ccs are empty
‘space’.
With this spacing the atoms, on average, have to go some
distance before colliding with another and the theory states
that ‘the mean free path of an atom is some 3000 times greater than
the diameter of the atom itself’.
Note: Quotations above are extracts from University level
textbooks.
 Figure
3
The diagram above represents the distribution of nitrogen and oxygen
in the atmosphere at sea level.
According to the Laws of Distribution, the velocity of
individual gas atoms cannot be zero and the average velocity of the
molecules of air at normal climatic temperatures and pressures is in
the region of 500 metres per second. The theory suggests that in air
the large majority of molecules in are travelling in the region of
these velocities.
The graph below shows the Maxwell distribution for helium
at 300K giving an average velocity of 1300 m/s.
Figure 4
Kinetic atomic theory at this stage then provided the means by which
the characteristics of gases, compound or elemental, could be predicted
in a limited* range of conditions of pressure temperature and volume.
(*limited due to technical capabilities available at this time.)
Karlsruhe Conference
The next significant event, for our purposes,
was the Karlsruhe (Germany) conference of chemists in 1860.
At this time there was no common system of chemical notation, some
chemists, for example, using a bar over a letter denoting an element
to indicate it’s characteristics in a compound, others using different
symbols to denote the same characteristic. So there was no universal,
internationally recognised method of communication amongst chemists,
or in other words no common chemical language.
This of course was confusing and seriously inhibited progress,
not only in chemistry but also in science in general, and
one of the main purposes of the meeting at Karlsruhe was to discuss
the need for a standardisation of notation and weights.
Avogadro died in 1856, but a compatriot Cannizzaro presented
Avogadro’s hypotheses and, while they were not immediately
accepted at the conference, many chemists in the interests of standardisation
later ignored the illogicalities of the original assumptions
and in the subsequent period these were generally accepted and ‘The
great value of Avogadro’s hypothesis lay in its use to
determine atomic weights.’3
So if equal volumes of different gases contain the same
number of atoms or molecules, ‘the molecular weight is
proportional to the weight of unit volume of that substance in the
state of gas or vapour.’
Continued >
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