-1
$\begingroup$

enter image description here

In general, going down a group Zeff initially increases but then becomes approximately constant, while electrons are in higher n orbitals, hence valence electrons on average further from nucleus, thereby size of atom increases, and electronegativity decreases. An exception is present in groups 13-16, going from row 2 to 3, where atoms of elements in row 3 have greater electronegativity than atoms of elements in row 2. This is because the row 3 element follows the d-block, where 10 extra electrons have been added in the next n down, so don't fully shield the valence electrons from the nuclear charge. That's a lot of electrons not doing much shielding, yet the nuclear charge has increased by +10 with 10 extra protons. Accordingly Zeff increases from row 2 to 3 more than expected, so electronegativity increases.

Why then, by the time you get to group 17, does this trend no longer hold i.e. electronegativity decreases from Cl to Br (see graph above)?

$\endgroup$
  • 2
    $\begingroup$ I "smell a rat" here. What's the source of your data and what electronegativity scale? $\endgroup$ – Mithoron Apr 22 at 20:55
  • $\begingroup$ The source is my lecturer. I don't know which electronegativity scale he used to create the graph. $\endgroup$ – ETS Apr 23 at 13:32
  • 2
    $\begingroup$ Then you should check valid data. Different scales use different properties for calculations. For example as far as Pauling (most common) scale goes S is slightly more electronegative then Se. Another point is for F it should be precisely 4. $\endgroup$ – Mithoron Apr 23 at 13:52
0
$\begingroup$

This can be simply explained by the electronic configuration. The electronic configuration of bromine is $ 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^5 $ Thus we can see it has 5 electrons in the 4p subshell which provides shielding effect so it's electronegativity is less than that of chlorine . It has also decreased due to the increase in size from chlorine to bromine.

$\endgroup$
-1
$\begingroup$

The Effective Nuclear Charge of an atom (Zeff) is the net charge an electron experiences in an atom with multiple electrons.5 It increases down a period (as an extra proton is being added for each incremental move down a period) and generally stays constant down a group (as the previous shell is now shielding some charge as well). I think you meant to say down a group rather than down a period. However, this change in Effective Nuclear Charge down a group is relatively insignificant.(1)

As one moves down a group, higher energy levels are being occupied by the valence electrons, the number of protons increases and there are more valence electrons as well. Consequently, not only is the atom more dense, but the atomic radii increases too, due to the increase in occupied shells and the increase in electrostatic repulsion by the increased number of valence electrons.(1)

Electronegativity is a measure of an element's ability to attract electrons in a covalent bond. According to periodic trends, it increase up a group and across a period (from left to right).(1) So in theory, the electronegativity values should decrease from Si to Ge (although Ge is not shown in your graph) or from Al to Ga, but this is not the case. Why is this not the case? Well, as you said, it has something to do with the d-block. To be more concise, it is due d-block contraction, which, in our case, refers to the "unexpectedly small increase (or even decrease) in atomic/ionic radius when moving down a group". (2)(3) For example Al and Ga, who have atomic radii of 184 pm and 187 pm respectively.(4) One would expect this increase in radius to be much larger as Gallium occupies an additional energy level. However, Ga's electron configuration is [Ar] 3d10 4s2 4p1. As one can see, it has the 3d orbitals occupied with electrons– these electrons are ineffective at shielding . Hence, the valence electrons of Ga experience a stronger attractive force from their nucleus. Consequently, Ga has a smaller atomic radius than expected and as this is the case, its electronegativity increases as well as a result of the nucleus having a stronger attraction for electrons than if it would have a larger atomic radius.

All in all, these inconsistencies in periodic trends for electronegativity can be attributed to the d-orbitals' electrons being ineffective at shielding, resulting in a stronger attractive force from a nucleus than expected. Thus, I think your question could be reformulated to why is there this disruption in electronegativity trends instead of why group 17 does not follow these trends– although in the end, these questions will result in the same answer.

Note: I recommend going through the references I have listed below to gain a better idea of what is being discussed.

References:

  1. Pearson Higher Level Chemistry Textbook, 2nd Edition. By Catrin Brown and Mike Ford. Chapter 3: Periodicity

  2. https://www.quora.com/What-is-a-detailed-explanation-of-d-block-contraction

  3. https://en.wikipedia.org/wiki/D-block_contraction#/media/File:D-block_contraction--EN.png

  4. https://en.wikipedia.org/wiki/Atomic_radii_of_the_elements_(data_page)

  5. https://www.thoughtco.com/definition-of-effective-nuclear-charge-605056

Other helpful references:

https://www.quora.com/What-causes-Ga-and-Ge-to-have-higher-electronegativity-than-the-elements-over-them

$\endgroup$
  • 1
    $\begingroup$ Your definition of effective nuclear charge is incorrect. It is in fact the net positive charge experienced by any electron in an atom. $\endgroup$ – ETS Apr 21 at 15:16
  • $\begingroup$ Good point. Changed. $\endgroup$ – Liam Apr 21 at 15:26

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.