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Atom – Introduction to Chemistry

by Adam Le Gresley, PhD
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    00:00 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.

    00:13 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.

    00:32 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 [Inaudible 0:02:57] units can be formed respectively.

    01:17 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.

    01:30 Chemistry in its very heart, as I mentioned before, is about electrons and not necessarily about the nucleus.

    01:37 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.

    01:49 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.

    02:11 1 atomic unit, as you can see at the bottom of the board, is given as 1,67 × 10^-24 grams.

    02:19 And what is worthy of note, if you look at the table, is the third entry: the electron.

    02:27 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.

    02:56 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 and Rutherford.

    03:13 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.

    03:31 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.

    03:52 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.

    04:32 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.

    04:53 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.

    05:12 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.

    05:38 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.

    05:57 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.

    06:03 There is some nomenclature that you should also be familiar with. And that is shown here on the board: Z, N and A.

    06:13 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.

    06:29 N is the number of neutrons which is, obviously as you would expect, the neutron number.

    06:35 And finally A, which is the combination of neutron number, N, and atomic number, Z. So this gives you the mass number.

    06:47 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 atomic mass.

    07:07 So here we have an example of an atom. To properly identify it, it is written thus.

    07:13 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.

    08:00 So anyway, as I was saying, if you look here we have an example of the element chlorine.

    08:04 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.

    08:54 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.

    09:46 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 possess.

    10:04 If we look at 35Cl as a stable isotope, it is found in 75% of all chlorine in the environment.

    10:12 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.


    About the Lecture

    The lecture Atom – Introduction to Chemistry by Adam Le Gresley, PhD is from the course Chemistry: Introduction.


    Included Quiz Questions

    1. The study of movements and interactions of electrons
    2. The role of metals in the synthesis of organic compounds
    3. The role of hydrocarbons in fossil fuel formation
    4. The production of fossil fuels via degradation of plant waste
    5. The design, chemical synthesis, and development of fossil fuels
    1. A nucleon can be either a proton or a neutron located in the orbitals of an atom along with electrons
    2. Protons have +1 charge, whereas neutrons are neutral
    3. Both protons and neutrons are composite particles, each having a mass of 1 atomic unit
    4. A proton is made up of two up and one down quark
    5. A neutron is comprised of one up and two down quarks
    1. …1.67 × 10^-18 micrograms
    2. …1.67 × 10^-18 grams
    3. …16.7 × 10^-18 micrograms
    4. …16.7 × 10^-18 grams
    5. …16.7 × 10^-24 grams
    1. An electron can be divided into quarks, bosons and photon particles
    2. The mass of an electron is 9.11 × 10^-31 kilograms
    3. Electrons are the actual particles in an atom which participate in ionic or covalent bond formations
    4. An electron is an elementary sub-atomic particle with -1.6 × 10^-19 coulombs of elementary charge on it
    5. Electrons were discovered by J.J. Thomson in 1897 during the cathode ray experiment
    1. …their atomic masses due to a different number of neutrons in their nuclei.
    2. …their atomic masses due to a different number of protons in their nuclei
    3. …their atomic charges due to a different number of protons in their nuclei
    4. …their atomic charges due to a different number of electrons in their shells
    5. …their atomic mass due to a different number of protons in their shells
    1. …loss of one or more fundamental particles to achieve a stable state
    2. …loss of one electron to achieve a stable state
    3. …loss of atomic nucleus to achieve a stable state
    4. …loss of outermost electron shell to achieve a stable state
    5. …loss of an atomic orbital to achieve a stable state
    1. Uranium
    2. Francium
    3. Technetium
    4. Protactinium
    5. Radon

    Author of lecture Atom – Introduction to Chemistry

     Adam Le Gresley, PhD

    Adam Le Gresley, PhD


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    All in all it is a brilliant lecture.
    By peter n. on 23. September 2017 for Atom – Introduction to Chemistry

    A fantastic lecture and it is very helpful for me thank you very much.