As time has gone by within the last 100 to
150 years, the scientific knowledge of mankind
has increased and matter has been found to
be made up of smaller fundamental particles.
In particular, if we look at subatomic particles
which you may be familiar with such as protons,
neutrons and electrons, they themselves are
actually made up of even smaller particles
falling into the quark and lepton class.
But as we will see when we are discussing
chemistry, we are principally looking at the
movements and the interactions of electrons.
Anything which indeed goes beneath that in
terms of size, typically speaking we leave
in the realms of physics. Chemistry is about
the movement of electrons. Nuclear physics
is about how nuclear particles interact with
each other. And therefore a knowledge of the
latest developments in this discovery of these
fundamental particles isn’t necessarily
essential for an appreciation of how electrons
move around and how ionic covalent molecules
and formerly units can be formed respectively.
As indicated in the previous slide, there
are certain small fundamental particles such
as leptons and quarks which are really unnecessary
at this level for understanding about chemistry.
Chemistry in its very heart, as I mentioned
before, is about electrons and not necessarily
about the nucleus.
If you look at the board, you’ll be able
to see three fundamental particles that you
are expected to be familiar with. They are
the proton, the neutron and the electron.
The proton and the neutron are both nucleons,
that is to say they are subatomic particles
which reside within the nucleus of an atom.
Protons have a charge of +1, neutrons have
a charge of 0. And they both have a mass of
1 atomic unit.
1 atomic unit, as you can see at the bottom
of the board, is given as 1,67 × 10^-24 grams.
And what is worthy of note, if you look at
the table, is the third entry: the electron.
Electrons have a charge of -1 but they have
a substantially smaller mass. The mass of
an electron is given relative to an AMU of
5,48 × 10^-4. However, the reality of this
in kilograms is that an electron has a mass
of 9,11 × 10^-31 kilograms. So very much smaller
than those nucleons.
Most of the atom is actually empty space with
the protons and the neutrons clustered together
in the centre. And this is something that
was actually detected experimentally by Thomson
The electrons seemingly form a cloud around
the central nucleus. And it is these electrons
which engage with each other to form matter
as we currently understand it, whether it’s
ionic or whether it is indeed covalent.
And what we’re going to be going through
is how these individual arrangements of electrons,
protons and neutrons come to form the elements
that we see within the periodic table, a very
important index for the elements that we find
on this planet.
The analogy which is often used for the structure
of an atom is that the nucleus is the ball
on the centre spot of a football field with
the electrons actually being the tiny specks
of dust blowing around the stands. Since within
an atom there are equal numbers of protons
and electrons, the overall charge of an atom
as its element is 0. It has to be 0 because
the number of protons with a charge of +1
equals the number of electrons with a charge of -1.
So only certain combinations of fundamental
particles can form stable atoms. And this
goes back to what I was saying about the nucleus.
Whilst it is not necessarily essential for
us to understand it in the context of compound
formation, it is important to be aware of
the existence of things called isotopes.
You have probably heard of the term ‘radioactive’
isotope and this, for example, would be where
you have an unstable configuration of protons
and neutrons within the nucleus that are liable
to undergo a disintegration involving the
loss of more of one of these fundamental particles.
And it is the nucleus that decays in this
case. It is also the origin, as we will see,
of isotopes. These are elements which have
the same chemical characteristics but a different
number of neutrons within the nucleus. And
radioactive decay, to give you an example,
would be not too dissimilar to that which
you observe in the decay of uranium to thorium.
Or used in the fission process – in nuclear
fission – the breakdown of uranium-235 to
barium and krypton. And this process would
be known as a nuclear reaction and the previous
one I just mentioned is radioactive decay.
So let’s get back to where we were originally:
talking about the atoms and talking about
the elements that we see in the periodic table.
There is some nomenclature that you should
also be familiar with. And that is shown here
on the board: Z, N and A.
Z correlates to the atomic number or element
number and this relates to the number of protons
in the nucleus and defines which element within
the periodic table the atom actually is.
N is the number of neutrons which is, obviously
as you would expect, the neutron number.
And finally A, which is the combination of
neutron number, N, and atomic number, Z. So
this gives you the mass number.
Since, as we’ve indicated earlier, electrons
have a very, very, very small mass they are
largely ignored from the overall mass of an
atom. Instead, we tend to look at the combination
of protons and neutrons when considering the
So here we have an example of an atom. To
properly identify it, it is written thus.
Note we have the chemical symbol for this
particular element – Cl. This correlates
to chlorine. As you will see if you interrogate
the periodic table, you will often see elements
which, ostensibly, don’t make any sense
in English or indeed in any other European
language because they are actually derived
from the Latin or the Greek. So, for example,
Cl – chlorine, chlóros – comes from the
Greek, meaning green. And, as we will see
a little later on, there are a number of other
elements which also don’t make sense in
the context of their English name or their
standard IUPAC names.
So anyway, as I was saying, if you look here
we have an example of the element chlorine.
Note the larger number at the top is A. This
is the atomic mass number. The lower number
is Z, which is the atomic number. And whenever
you’re looking at this if you’re getting
confused as to what is an atomic number and
what is an atomic mass number, the atomic
mass number is always larger than the atomic
number. So, if you can’t remember whether
it’s top or bottom, don’t worry. Just
find the largest number: that is the mass,
which correlates to the number of protons
and also the number of neutrons. Z is the
atomic number, which correlates to the number
of protons, as we indicated, but also just
as important in elemental form must therefore
correlate to the number of electrons in the
shells of that atom in order for it to have
a charge of 0.
Now I’ve shown chlorine here for a good
reason because it is one of those elements
which exists as two stable isotopes within
the periodic table. This is isotopes where
you have the same chemical activity, because
you have the same number of protons and therefore
electrons in an atom, but a different number
of neutrons. And, as you can see here, we
have – or I’m showing you here – three
different isotopes, of which one is actually
unstable. They are 35Cl and 37CL and 36Cl,
which is the unstable radioisotope with a
half-life of 308,000 years and is negligible
in concentration within the environment. The
ones which are stable are 35 and 37.
As you’ll see sometimes with periodic tables,
the Z is often omitted because the chemical
symbol of a particular element automatically
defined in the periodic table further defines
the number of protons and electrons it must
If we look at 35Cl as a stable isotope, it
is found in 75% of all chlorine in the environment.
37Cl, on the other hand, is found in 25% of
chlorine in the environment. And so therefore,
when we are calculating the relative overall
atomic mass, we need to take into consideration
the natural occurrence of both of those isotopes.
And we’ll come onto an equation that deals
with this a little later on.