Introduction – Electronegativity

by Adam Le Gresley, PhD

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    00:02 In the previous lecture, we talked about the affinity of nuclear positive charge for electrons.

    00:10 We talked about it in the context of ionic bonding and we also talked about it in the context of covalent bonding. And this relates to the Pauling scale that we alluded to earlier as well.

    00:23 Here is just a quick revision of that scale. If you look at the periodic table representation shown here on the board, you’ll see that the increase in electronegativity follows an increase in atomic number running from group I to Group VII and that increase in electronegativity is associated with the smallest number of shells running from the highest period to the lowest.

    00:49 Large atoms have a low affinity because the electrons are a long way from a positive nucleus.

    00:56 The core electrons shield the outer electrons from the nuclear charge. And, going across the periodic table, the affinity for bonding electrons increases because nuclear charge increases without a significant change in the distance of the electrons from the nucleus.

    01:16 And this is okay in explaining how we observe bond polarisation in covalent bonds and also where we see ionic bonds forming.

    01:26 Fluorine is the most electronegative atoms with a value of 4.0 on the Pauling scale.

    01:34 And remember that Group 18, or Group 8 depending on your periodic table elements, do not react: they are the noble gases and therefore their shells are already full and they are unreactive.

    01:47 All other elements have smaller values on the Pauling scale. And in bonds between atoms of different elements, the bonding electrons can be shared unequally.

    01:58 Let’s have a look for example at carbon bonded to oxygen. Note I’ve shown you a sigma bonding: a single carbon-to-oxygen bond.

    02:09 If we look at the individual values on the Pauling scale for carbon and for oxygen, we see they are 2.5 and 3.5 respectively. The difference between the two is 1. And what that means is we have a polar covalent bond. What that means is that the probability density or the chance of finding an electron within the sigma bond is greatest when we move towards the more electronegative element, in this case the oxygen. This is so called polarised.

    02:45 It also – the scale – allows us to determine whether or not we’re likely to see an ionic or a covalent interaction.

    02:52 If we look at one of the carbon-hydrogen sigma bonds of our friend methane from previous lectures, you’ll see that the difference in electronegativity between the two is only 0, based on 2.5 – 2.1. This gives us a relatively non-polar covalent bond where the electrons, or the chances of finding the electrons, are likely to be mostly in the centre between the two.

    03:22 If on the other hand we look at the interaction between sodium and chlorine and we form an ionic sodium and chlorine bond, we see the difference is quite different: chlorine has an electronegativity on the Pauling scale of 3, sodium of 0.9. The difference between the two is 2.1 and the electrons, because they are much more strongly attracted to the chlorine, result in the formation of an ionic compound with formal loss of an electron from the sodium to the atomic orbital of the chlorine. So, in sodium chloride, as when you see other interactions between Group-1 and Group-2 and Group-6 and 7 atoms, you will observe an ionic compound being formed. Bear in mind ionic compounds exist as formula units. You would never ever refer to a salt or ionic compound as a molecule. Please be aware of that distinction.

    About the Lecture

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

    Included Quiz Questions

    1. The electronegativity increases when we move down the period in the periodic table
    2. Electronegativity represents the magnitude of the affinity of an atom in a molecule for the electrons in a chemical bond
    3. Fluorine is the most electronegative element in the periodic table
    4. The lower Ionization energy is, the more readily the atom becomes a cation
    5. The electronegativity is directly related to an atomic number of the element
    1. The outer shell electrons feel fewer forces of attraction from the nucleus due to shielding effecting of core electrons
    2. The outer shell electrons revolve at faster speeds, which leads to reduction in electronegativity values
    3. In the larger atoms, the strong attraction between the outermost electrons and nucleus leads to the low electronegativities
    4. In larger atoms, an increase in electrons is not balanced with an equivalent increase in the positive charges in the nucleus
    5. The outer shell electrons revolve at slower speeds, which cause a reduction in electronegativity value
    1. The nature of the bond
    2. The number of electrons involved in the chemical bond
    3. The type of nuclear forces playing a role in the chemical bond
    4. The magnitude of van der Waal’s forces participating in the chemical bond
    5. The type of electrons participating in the chemical bond
    1. Ionic
    2. Non-polar covalent
    3. Polar covalent
    4. Hydrophobic
    5. Hydrophilic

    Author of lecture Introduction – Electronegativity

     Adam Le Gresley, PhD

    Adam Le Gresley, PhD

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