Electronegativity – Chemical Bonding

by Adam Le Gresley, PhD

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    00:01 So as I mentioned to you before, covalent bonds between atoms of the same type tend to facilitate electrons being in the center between the two hydrogens. Or rather, the probability of that happening—at the highest 95%—is within those two atoms. However, different electronegativity can polarize a covalent bond. That is to say electrons can be moved more to one atom than to the other, depending upon the electron charge density of the atoms involved. Large atoms tend to have low affinity for electrons, and this is for a number of reasons. The electrons are a long way from the positive nucleus.

    00:48 The core electrons actually shield the outer electrons from the nuclear charge, allowing them to be pushed further away. Going across the periodic table, as you'll see in a moment, the affinity for bonding electrons increases from group 1 through to group 7, because the nuclear charge increases. And as we'll see, as we go down the periods from period 1, 2, and 3, as we increase the number of shells, we are pushing the electrons further and further away and decreasing the effective nuclear charge that can be brought to bear in polarizing a sigma bond. Here we have a representation of the periodic table. Note the arrows. As we go across from group 1 to group 7, we're increasing electronegativity, and we reduce electronegativity as we go from the lowest period to the highest period.

    01:52 Let's have a look at another sigma bond. This one, however, involves 2p orbitals. Remember, as we said before when looking at HF: We need to have in-phase overlap. That is to say, we need to have lobes of the right phase interacting with each other in order for a bonding orbital to be formed. So, as we discussed before, we have two atomic orbitals: the 2s and 2p set, and particularly, I want to look at the 2p set for fluorine. Because what's happening here, as you can see, when the orbitals are correctly aligned with the correct geometry, it's possible for the in-phase overlap of the two lobes of 2p orbitals from two fluorine atoms to occur, and the result is the same: A sigma molecular orbital is formed. Again, I draw your attention to the fact that the outer-phase lobes have been reduced in size.

    About the Lecture

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

    Included Quiz Questions

    1. HF
    2. PH3
    3. NH3
    4. H2S
    5. HI
    1. The larger atoms have low electron affinities because the effective nuclear charges get reduced due to the shielding effect of core electrons
    2. The larger atoms have high electron affinities because of higher effective nuclear charges
    3. The nuclei of larger atoms repel the nuclei of smaller atoms due to repulsive forces between their positive charges
    4. The nuclei of larger atoms attract the nuclei of smaller atoms due to attractive forces between their positive charges
    5. The larger atoms can not overlap effectively in a sigma bond formation due to larger nuclear radii
    1. Increase due to increasing electronegativities of the atoms
    2. Decrease due to increasing electronegativities of the atoms
    3. Remain unaffected as the increase in electronegativity of an atom is compensated by an equal increase in the nuclear charge
    4. Show irregular trend which can not be explained using present theories
    5. Decrease due to decreasing electronegativities of the atoms

    Author of lecture Electronegativity – Chemical Bonding

     Adam Le Gresley, PhD

    Adam Le Gresley, PhD

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