The way that atoms bond together to form molecules has been a question asked since scientists came to the consensus that atoms do indeed exist. Work progressed rapidly after the turn of the 20th century from both theoretical and experimental breakthroughs. To keep the discussion easy to visualize, we shall consider only diatomic molecules, but the concepts are good for any number of atoms. One of the great advances was the development of the idea that chemical bonds can either be covalent, where each atom shares bonding electrons equally, or ionic, where one atom donates an electron to another atom entirely.
Actually, pure ionic bonds do not exist because all bonds have at least a little bit of covalent character. Pure covalent bonds are common, common examples being the nitrogen and oxygen of the atmosphere. There is a very cool way to predict where a particular bond falls in the covalent to ionic spectrum, and that is to use electronegativity values.
The brilliant physical chemist Linus Pauling came up with the electronegativity scale in 1932, just after the birth of quantum mechanics. Pauling was quite a fellow, having the distinction of being the only person ever to be awarded two Nobel prizes with no coawardees. His first one was in 1954 for chemistry and the second one in 1962 for peace, as he was very much the humanitarian. His story is fascinating, and the Wikipedia piece on him is pretty good.
Pauling came up with the idea about electronegativity because very often a single bond between two different atoms is stronger than the average of single bonds between two different kinds of atoms. He took measurements from the literature of bond dissociation energies for thousands of compounds and developed a scale for the behavior of individual atoms, except for the noble gases.
Here is a periodic table showing the electronegativities of most of the elements. I could not find a picture that would imbed legibly, so just open this link in a new window so you can flip back and forth.
In addition, this chart is useful:
EN Difference % Ionic Character
0.0 0
0.2 1
0.4 4
0.6 9
0.8 15
1.0 22
1.2 30
1.4 39
1.6 47
1.8 55
2.0 63
2.2 70
2.4 76
2.6 82
2.8 86
3.0 89
3.2 92
As an example, let us look at carbon monoxide, CO. Carbon has an electronegativity of 2.55 and oxygen of 3.44. The difference is 3.44 - 2.55, or 0.89, corresponding to ionic character of around 20%. That means that the oxygen carries a partial negative charge. The classic Lewis structure of :C:::O: does not completely represent the reality of the partial negative charge on the oxygen, but writing it as :C::O::- does. However, this structure only gives six electrons to the carbon, and as you might recall from high school chemistry, atoms are usually more stable with the completed octet. However, since physical measurements show that the molecule actually has a dipole moment support at least some contribution from the second structure.
This trend is general. For example, cesium fluoride has an electronegativity difference of 3.98 - 0.79 = 3.19, showing that it is over 90% ionic, indicating that the electron from cesium has been almost completely transferred to the fluorine. This is strictly true only in the gas phase, because due to stability gained from forming a crystal lattice in the solid phase causes the electron to be essentially completely transferred to the fluorine.
So, why do some elements “like” to lose electrons whilst others tend to gain them? It has to do with a combination of factors, including nuclear charge, largeness of the atom, and forming the most stable arrangement of electrons possible.
The nuclear charge is pretty clear cut, meaning that the more positive charge (protons) in the nucleus, all else being equal, the more tightly electrons are held and the more tendency to get more electrons from other atoms there is. So, looking at the second row of elements you see a steady increase in electronegativity from lithium to fluorine. This trend repeats in the third row and so on, but you can see that the numbers are smaller in general with each row.
This is because of the increase in size of the atoms as we move down the periodic table, because we add a new shell of electrons each time. Since the electrons that are further from the positive nucleus are less attracted by electrostatic forces, these larger atoms are not as attractive of outer electrons as are the smaller ones. There is a bit more, as well. Not only is the distance greater from the nucleus to the outer electrons, the ones between the outer ones and the nucleus effectively shield the influence of the positive nucleus in larger atoms.
These trends get wobbly the further down the periodic table you go, so there must be other things going on as well. For example, krypton has an electronegativity of 3.0, so it should be almost as reactive as chlorine. However, it is not nearly as reactive as chlorine. In fact, krypton has to be subjected to rather extreme conditions to get it to react with anything, and most of the handful of krypton compounds known must be kept quite cold or they decompose.
The reason for that is basically quantum mechanical, and has to do with the stability given to an atom with a complete subshell of electrons. A “full” subshell contains eight electrons for all elements except hydrogen and helium, that can only hold two. Krypton, like the other noble gases already has a full subshell, so there is no energy gain to be had by adding electrons or by losing them. Only extremely electronegative elements are known to react with krypton, notably oxygen, nitrogen, and fluorine, and the nitrogen and oxygen compounds also contain fluorine.
Since krypton and the other noble gases already contain full subshells, other elements tend to gain or lose electrons to attain an electron configuration like a noble gas, but at the cost of gaining electrical charge. Since there is no benefit to noble gases to gain or lose electrons for the most part, gaining an electrical charge is a huge cost in energy since it puts it in a state were it has not a filled subshell AND and electrical charge.
Let us take a very common material as an example. Table salt, sodium chloride, has an electronegativity difference of 3.16 - 0.93, or 2.23, for an ionic character of the bond in the gas phase of around 70%. In the solid phase, the ionic character is 100% (almost, but not completely because of the Heisenberg Uncertainty Principle) because of extra energy benefit for forming the crystal lattice of salt. Looking at the periodic table link, we see that if sodium loses an electron, its electronic structure attains the extremely stable configuration that the noble gas neon has. Neon forms no known compounds.
In addition, when chlorine gains an electron, its electronic configuration becomes that same as the noble gas argon. Even though each of those ions carries a charge, the attraction of the opposite charges (positive for sodium ion and negative for chloride ion allows them to pack together into a stable, but brittle, crystal lattice. The energy lost in forming the ions and the crystal lattice is huge (lower energy means greater stability), and sodium metal burns spontaneously and violently in an atmosphere of chlorine to form salt.
We mentioned a while ago that polar covalent bonds are often more stable than the average of the fully covalent bonds from each atom. Let us look at a common item, polytetrafluoroethene, aka Teflon. The F-F bond dissociation energy is 159 kJ/mol, the C-C is is around 350, depending on the exact structure, and the H-H one is 436 (quite high for a single bond, due to quantum effects). Now, the C-H bond dissociation energy is 410 kJ/mol and the C-F one is a whopping 490 kJ/mol. Carbon/hydrogen bonds are an exception to the rule, but as I said, that is because the H/H bond is anomalously strong. Note that the ionic character of the c/H bond is 2.55 - 2.20 = 0.35, or only about 3%.
However, with carbon/fluorine bonds, the ionic character is 3.98 - 2.55 = 1.43, or close to 40% ionic. This high ionic character increases the the bond dissociation energy from 410 kJ/mol for a C/H one to 490 kJ/mol for the C/F one. As a matter of fact, the C/F bond is the strongest single bond to carbon known.
There are other electronegativity scales, derived from other kinds of measurements , but none is in as widespread use as the Pauling one. However, the importance of the electronegativity scale is beginning to wane, and is now used only for semiquantative and qualitative purposes because of a better understanding of fundamental atomic interactions. It turns out that the concept of electronegativity is based on the valence bond theory of bonding in molecules, which has largely been supplanted by the molecular orbital theory of bonding, which has a stronger connexion with quantum mechanics than the VB theory.
It is still useful as a predictive tool, but is not treated except sort of in passing these days. As a matter of fact, I had to go back to my old copy of Pauling’s College Chemistry to find the table of ionic character as a function of electronegativity differences, because in a rather extensive search of the ‘net I could not find it.
Well, you have done it again! You have wasted many more einsteins of perfectly good photons reading this negative piece. Please keep those comments, questions, corrections, and other feedback coming because I always learn much more than I could possibly hope to teach in writing this series.
I shall stick around tonight for Comment Time as long as traffic warrants. Tomorrow night I shall return for Review Time to respond to late comments.
Warmest regards,
Doc, aka Dr. David W. Smith
Crossposted at Daily Kos,
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chemical bonding?
Warmest regards,
Doc