Electronegativity is one of the most useful concepts in chemistry. In a nutshell, it is a measure of how “hungry” an element is for electrons, which is a function of the number of electrons in the valence shell and the effective nuclear charge felt by them. What makes it great is its simplicity: like GPA, it is a single number that gives an immediate mental impression of an element’s personality. Fluorine, the “Tiger of Chemistry“, is head of the class at 4.0 . At the bottom left corner of the periodic table lies Cesium, which lacks a catchy nickname [but has attracted a strangely devoted following] with an electronegativity of 0.79.
Like in many human relationships, when elements get together to form bonds, a sort of hierarchy is established: in this case, the more electronegative element gets an unequal share of the electrons. The resulting dipoles play a tremendous role in organic chemistry. If you are just starting out in organic chemistry, let me say it now: electronegativity rears its head in organic chemistry in more ways than you can currently imagine. The direct and indirect effects of electronegativity dictate important properties such as solubility, acidity, hydrogen bonding,melting and boiling points, chemical reactivity and many more.
In organic chemistry, we tend to focus on elements that form covalent bonds to carbon and hydrogen. Here is a table of the electronegativities of the 12 elements you will likely encounter the most [Pauling scale]. From the concepts that spring from the facts in this table come many, many potential exam questions.
Source: Wikipedia [includes link to a more comprehensive list]. Just knowing these 12 numbers will take you a long way. This is a good example of something that is worth the time to brute-force memorize. I’ve shared a quiz I set up for this at memorize.com [no account required].
A few notes:
- Carbon is more electronegative than you think. While the electronegativity difference for the C-F bond is large (1.4), it goes down very rapidly as one goes down to C-Cl (0.6), C-Br (0.4) and C-I (0.1).
- Halides have similar properties, but be careful when drawing analogies between oxygen/sulfur [chalcogens] and nitrogen/phosphorus [pnictides]: the C and S electronegativities are essentially identical [2.6] while C-P is actually polarized toward carbon [0.4].[Note 1]
- Similarly, while the O-H bond is highly polarized (1.2), there is very little polarization in the S-H bond (0.4). The upshot of this is that there is no hydrogen bonding: if you want liquid H2S, you have to condense it at -60 °C .[Note 2]
- Hydrogen halides (HF, HCl, HBr, HCl) are all polarized toward the halide, but by the time you get to HI the difference in electronegativities isn’t that great (0.5).
- Check out boron: it’s less electronegative than hydrogen [2.0 vs. 2.2]. This is the core reason why the hydroboration reaction is so-called “Anti-Markovnikoff” – the hydrogen is partially negative and boron partially positive.
[Note 1] One disclaimer: electronegativities are for the native elements themselves. The electronegativity of phosphorus as phosphine (PH3) is a lot lower than the electronegativity of phosphorus as phosphate [O=P(OH)3]. The “effective electronegativities” of functional groups is worth a separate post.
[Note 2] Having done this personally, and nervously warming the resulting liquid H2S to room temperature in a pressure vessel, I can report that this can be done without becoming a social pariah for the following few hours provided it is done in a well-functioning fume hood. Old time chemists like Scheele who worked with it without modern stench-fighting apparatus deserve our undying admiration: H2S is not to be trifled with.