By Ron B. Davis Jr., Georgetown University
In 1916, Gilbert N. Lewis published his famous paper explaining what he called “the rule of eight”, later known as “the octet rule”. In the paper, Lewis noted that chemical bonding, especially among the p-block elements of the periodic table, seemed to be the result of electron-sharing between atoms seeking a noble gas configuration.
Irving Langmuir’s ‘Octet Rule’
Gilbert Lewis’s theory was adopted and championed by yet another famed chemist from that time named Irving Langmuir. Among Langmuir’s many talents was coming up with catchy-sounding names for chemical principles, some of which caught on and replaced terms suggested by Lewis. In fact, it was Langmuir who first coined the term “octet rule”, which quickly replaced “the rule of eight” in chemistry.
But, whichever name you choose to call it, strong support for the theory comes from the entire second row of the p-block, including boron, carbon, nitrogen, oxygen, fluorine, and neon. These six p-block elements are a stronghold of the octet rule because they seem to just about invariably follow it. But these elements differ from the larger atoms below them in a very important way.
How the Second Row Elements Reach the Octet
As the second row elements seek out octets through covalent bonding, they max out their principle shell (the second energy level) capacity at eight valence electrons, just as the octet rule predicts. For example, carbon can form four bonds to hydrogen atoms to make methane, nitrogen can make three bonds to hydrogen to form ammonia, oxygen makes two bonds to hydrogen in water, and fluorine can form hydrofluoric acid by bonding to just one hydrogen. Neon, of course, sits content with eight valence electrons already in its valence shell.
The changing strength of that covalent bonding, depending on where we are on the table, combined with the strict adherence to octet rule, explains how row two elements carbon and nitrogen can exhibit such starkly different phase behavior. This pattern on the table also explains why atmospheric nitrogen’s triple bond makes it so unreactive (and so difficult for most organisms to exploit), while oxygen’s double bond provides balanced reactivity (which many organisms exploit), while molecular fluorine’s single bond is so dangerously over-reactive.
This article comes directly from content in the video series Understanding the Periodic Table. Watch it now, on Wondrium.
Not All Compounds Follow the Octet Rule
So, when the central atom of a molecule is a second row nonmetal, eight valence electrons is the absolute maximum. There is simply nowhere to put anymore. So, the second row follows the octet rule invariably. There is no compound known today in which any of these elements contains more than four covalent bonds for a total of eight valence electrons.
But, the remaining p-block elements are not so cooperative, and this sparked a spirited debate between these two fathers of the octet rule—a debate that has continued to this day. Take the example of sulfur. In combination with hydrogen, it makes dihydrogen sulfide, perfectly analogous to water, which is dihydrogen oxide.
Like water, dihydrogen sulfide obeys the octet rule, with two bonds to hydrogen and two lone pairs around the central sulfur atom. A total of eight, so we are off to a good start. But in combination with fluorine, sulfur can form a very different sort of molecule called sulfur HEXA-fluoride.
Sulfur hexafluoride is a very stable gas and a good electrical insulator sometimes used in electronics. But, take a closer look—six covalently bonded fluorine atoms means twelve electrons are involved in the bonds to sulfur! Last time I checked twelve was more than eight. That would make the sulfur in SF6 ‘hypervalent’.
The Debate over the Rule of Eight
There is no denying that compounds like sulfur hexafluoride exist. Lewis and Langmuir argued aggressively and publicly about how compounds that appeared to violate the octet rule could possibly exist. Langmuir tried to preserve the octet rule by arguing that only four true covalent bonds exist in SF6, with the other two fluorine atoms held by bonds that are more ionic in nature. It was simply the larger size of sulfur that allowed more fluorine atoms to pack around it.
Lewis argued that additional covalent bonds might be possible if energetically nearby d-orbitals were recruited to form additional bonds. Is sulfur in SF6 genuinely violating the octet rule? Is it d-orbital participation or simply the larger size of the sulfur atom that accounts for this behavior? Believe it or not, a century later the answer to this question is not entirely resolved.
But what is certain is this—from the third row on, the remaining p-block elements have this amazing ability to take on what at least appears to be more than an octet of valence electrons.
Common Questions about the Octet Rule and the Debate over Its Accuracy
Fluorine, oxygen, nitrogen, and carbon belong to the second row of the p-block region. These elements strongly follow the octet rule in their compounds. For example, these elements, in combination with hydrogen, form from two to four bonds to satisfy their octet.
No, some of the p-block elements simply don’t follow the octet rule. For instance, sulfur, in combination with hydrogen, forms a molecule called dihydrogen sulfide which is similar to water. Here, sulfur obeys the octet rule by forming two lone pairs around the central sulfur atom and two bonds with hydrogen. But when combined with fluorine, sulfur has a completely different behavior with six bonds or 12 involved electrons.
Sulfur hexafluoride proves that compounds that some compounds do violate the octet rule. Irving Langmuir argued that this happens simply because of the large size of sulfur atoms which allows more fluorine atoms to pack around it. He stated that the bonding between the atoms of sulfur hexafluoride might be a combination of ionic and covalent. However, the definitive answer to the question of why this violation happens has not been resolved yet.