Now though we've finished optics and how to discuss optical
instruments and the basics of how they work.
We're going to zoom to the very small world and start by
talking about the atomic nucleus.
We'll first start with the atomic nucleus and nuclear
properties of that nucleus.
In the atomic nucleus we find a few objects and these are called
the "Nucleons."
So we have protons and neutrons.
And those are built in the nucleus of an atom.
Where the atom includes not just a nucleus
but also the electrons that are orbiting the nucleus.
The way that we describe the atoms, specially when we are
talking about the periodic table is by talking about
the number of protons in that nucleus.
So the number of protons in nucleus, in fact defines for us
the identity of that nucleus.
So carbon for example is defined as the element
which has 6 protons in it's nucleus.
On the other hand, we also have the number of neutrons
in that nucleus.
And the number of neutrons is not going to change
the identity of the element.
Instead we call them isotopes when we have a different number
of neutrons in the nucleus.
So for carbon for example, there might be 6 neutrons
in the nucleus.
But there might be other numbers of neutrons in the nucleus.
In those case we would say that this is isotopes of carbon.
While the identity of carbon remains the same
because the number of protons are still 6 protons.
We can always find the total number of nucleons or the things
in the nucleus which includes protons as well as neutrons.
By simply adding the number of protons and neutrons.
So for example, if we have 6 protons and 6 neutrons,
we would refer to this element as carbon 12.
Which means that there are 12 total nucleons in the carbon.
While the number of protons which defines the carbon
is still 6.
On the other hand we have an isotope of carbon
which has a different number of neutrons.
In that case the total number of nucleons would be 13.
And we refer to it as carbon 13 instead.
It's also important to know that with isotopes,
many isotopes with these extra numbers of neutrons
are unstable isotopes.
So we can't just add as many neutrons we would like
into a nucleus and have it still be a stable nucleus.
Eventually, and pretty quickly usually, they will become
unstable and decay as we call it.
Losing neutrons and becoming may be some other atom
altogether.
We can define few more properties and sort of
relate them on to each other.
The mass number is the total number of nucleons.
This includes both protons and neutrons.
The atomic number defines the atom itself.
And this is the number of protons.
We refer to this as "Z."
The number of neutrons are called the "Neutron number."
We refer to it as "N."
And then as we just discussed, the total mass number,
which is the total number of nucleons "A,"
will simply be the sum of all the protons
and all the neutrons, Z + N.
In atomic notation we have a way of writing a particular atom
or particular element in a way that we can express
what all of these numbers are together.
So for example with carbon here, what we do,
is we write the mass number,
the total number of nucleons up on the top left.
And then we write the number of protons here,
the atomic number on the bottom right.
So carbon 13 has of course a total atomic number of 13.
Meaning it has 6 protons and 7 neutrons.
The reason we often refer it as just carbon 13 without the 6
written on the bottom left, is because since we know
it's carbon, we already know how many protons it must have.
And so by saying carbon 13 already know in my head that
it's actually got to have 6 protons
because we called it carbon.
And I can figure it out how many neutrons are there
by myself just using the mass number.
Each stable isotope has a particular natural abundance.
And it's important that we talk about stable isotopes.
Because the unstable isotopes are the ones that we said
break down and they don't stick around for very long.
But isotopes are stable meaning that they do stick around.
Long enough to be measured by us.
So that if I scooped up may be some carbon from the ground,
I can measure how many different types of isotopes are there.
And that would be referred to as a natural abundance.
So for example, if I scooped up some carbon right now,
over 98% of that carbon would be carbon 12.
That would be natural abundance of carbon 12.
And only a very small percentage of it would be carbon 13.
And that's the natural abundance of carbon 13.
The atomic weight is something that we've referred to already.
In a periodic table setting, we listed below the element.
And this number tells us how many grams of that element
would be in one mole of that element.
And remember that the mole is just a very, very large number
to find us the number of carbon 12 atoms
we would need to make up 12 grams of carbon.
We mentioned before that because carbon isn't all
carbon 12.
Some small percentage of it is actually carbon 13,
slightly heavier.
The atomic weight of carbon is slightly higher
because of the presence of this natural abundance of carbon 13.
The atomic weight, don't forget will include this carbon 13
which would just be only in the ground.
So if I was talking about carbon 12, I'm not talking about
a sample I took from the ground.
So this natural abundance is again what we would see in
nature if we just went out and found some sample of carbon
which would have a mix of both the carbon 12 and the carbon 13.