# Nuclear Properties

by Jared Rovny

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00:02 Now though we've finished optics and how to discuss optical instruments and the basics of how they work.

00:07 We're going to zoom to the very small world and start by talking about the atomic nucleus.

00:12 We'll first start with the atomic nucleus and nuclear properties of that nucleus.

00:18 In the atomic nucleus we find a few objects and these are called the "Nucleons." So we have protons and neutrons.

00:24 And those are built in the nucleus of an atom.

00:26 Where the atom includes not just a nucleus but also the electrons that are orbiting the nucleus.

00:33 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.

00:42 So the number of protons in nucleus, in fact defines for us the identity of that nucleus.

00:48 So carbon for example is defined as the element which has 6 protons in it's nucleus.

00:54 On the other hand, we also have the number of neutrons in that nucleus.

00:59 And the number of neutrons is not going to change the identity of the element.

01:03 Instead we call them isotopes when we have a different number of neutrons in the nucleus.

01:08 So for carbon for example, there might be 6 neutrons in the nucleus.

01:12 But there might be other numbers of neutrons in the nucleus.

01:14 In those case we would say that this is isotopes of carbon.

01:17 While the identity of carbon remains the same because the number of protons are still 6 protons.

01:23 We can always find the total number of nucleons or the things in the nucleus which includes protons as well as neutrons.

01:31 By simply adding the number of protons and neutrons.

01:34 So for example, if we have 6 protons and 6 neutrons, we would refer to this element as carbon 12.

01:41 Which means that there are 12 total nucleons in the carbon.

01:44 While the number of protons which defines the carbon is still 6.

01:49 On the other hand we have an isotope of carbon which has a different number of neutrons.

01:54 In that case the total number of nucleons would be 13.

01:58 And we refer to it as carbon 13 instead.

02:00 It's also important to know that with isotopes, many isotopes with these extra numbers of neutrons are unstable isotopes.

02:09 So we can't just add as many neutrons we would like into a nucleus and have it still be a stable nucleus.

02:15 Eventually, and pretty quickly usually, they will become unstable and decay as we call it.

02:20 Losing neutrons and becoming may be some other atom altogether.

02:26 We can define few more properties and sort of relate them on to each other.

02:30 The mass number is the total number of nucleons.

02:34 This includes both protons and neutrons.

02:36 The atomic number defines the atom itself.

02:40 And this is the number of protons.

02:42 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.

03:00 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.

03:11 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.

03:18 And then we write the number of protons here, the atomic number on the bottom right.

03:23 So carbon 13 has of course a total atomic number of 13.

03:28 Meaning it has 6 protons and 7 neutrons.

03:32 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.

03:42 And so by saying carbon 13, I already know in my head that it's actually got to have 6 protons because we called it carbon.

03:49 And I can figure it out how many neutrons are there by myself just using the mass number.

03:55 Each stable isotope has a particular natural abundance.

03:59 And it's important that we talk about stable isotopes.

04:02 Because the unstable isotopes are the ones that we said break down and they don't stick around for very long.

04:06 But isotopes are stable meaning that they do stick around.

04:11 Long enough to be measured by us.

04:13 So that if I scooped up may be some carbon from the ground, I can measure how many different types of isotopes are there.

04:18 And that would be referred to as a natural abundance.

04:21 So for example, if I scooped up some carbon right now, over 98% of that carbon would be carbon 12.

04:27 That would be natural abundance of carbon 12.

04:29 And only a very small percentage of it would be carbon 13.

04:34 And that's the natural abundance of carbon 13.

04:36 The atomic weight is something that we've referred to already.

04:41 In a periodic table setting, we listed below the element.

04:44 And this number tells us how many grams of that element would be in one mole of that element.

04:50 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.

04:59 We mentioned before that because carbon isn't all carbon 12.

05:02 Some small percentage of it is actually carbon 13, slightly heavier.

05:06 The atomic weight of carbon is slightly higher because of the presence of this natural abundance of carbon 13.

05:14 The atomic weight, don't forget will include this carbon 13 which would just be only in the ground.

05:20 So if I was talking about carbon 12, I'm not talking about a sample I took from the ground.

05:25 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.

The lecture Nuclear Properties by Jared Rovny is from the course Atomic Nucleus.

### Included Quiz Questions

1. Protons, Neutrons
2. Protons, Neutrons, Electrons
3. Neutrons, Electrons
4. Neutrons, Photons
5. Photons, Protons, Neutrons, Electrons
1. 7 Protons, 8 Neutrons
2. 8 Protons, 7 Neutrons
3. 15 Protons, 7 Neutrons
4. 7 Protons, 15 Neutrons
5. 7 Protons, 1 Neutron
1. Mass number 14, Atomic number 6, Neutron number 8
2. Mass number 6, Atomic number 14, Neutron number 8
3. Mass number 10, Atomic number 6, Neutron number 8
4. Mass number 8, Atomic number 10, Neutron number 6
5. Mass number 10, Atomic number 4, Neutron number 6
1. The weight depends on the number of nucleons, which vary around 12 in natural abundance.
2. The weight is always twice the atomic number, plus a contribution from the binding energy of the atoms.
3. The weight is the weight of the most common isotope found in nature, which has a weight of more than 12.
4. In reality, physical quantities will never be exactly the theoretical values.
5. Carbon always appears in nature as a molecule of two Carbon atoms.

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