n=3) and beyond The 'octet' rule is based upon available ns and np orbitals for valence electrons (2 electrons in the s orbitals, and 6 in the p orbitals). Beginning with the n=3 principle quantum number, the d orbitals become available (l=2). The orbital diagram for the valence shell of phosphorous is: Hence, the third period elements occasionally exceed the octet rule by using their empty d orbitals to accommodate additional electrons. Size is also an important consideration: - The larger the central atom, the larger the number of electrons which can surround it - Expanded valence shells occur most often when the central atom is bonded to small electronegative atoms, such as F, Cl and O. There is currently much scientific exploration and inquiry into the reason why expanded valence shells are found. The top area of interest is figuring out where the extra pair(s) of electrons are found. Many chemists think that there is not a very large energy difference between the 3p and 3d orbitals, and as such it is plausible for extra electrons to easily fill the 3d orbital when an expanded octet is more favorable than having a complete octet. Violations of the Octet Rule - Chemistry LibreTexts Skip to main content This action is not available. Following the Octet Rule for Lewis Dot Structures leads to the most accurate depictions of stable molecular and atomic structures and because of this we always want to use the octet rule when drawing Lewis Dot Structures. However, it is hard to imagine that one rule could be followed by all molecules. There is always an exception, and in this case, three exceptions. The Octet Rule is violated in these three scenarios: When there are an odd number of valence electrons When there are too few valence electrons When there are too many valence electrons Reminder: Always use the Octet Rule when drawing Lewis Dot Structures, these exceptions will only occur when necessary. Exception 1: Species with Odd Numbers of Electrons The first exception to the Octet Rule is when there are an odd number of valence electrons. An example of this would be the nitrogen (II) oxide molecule (\(NO\)). Nitrogen atom has 5 valence electrons while the oxygen atom has 6 electrons. The total would be 11 valence electrons to be used. The Octet Rule for this molecule is fulfilled in the above example, however that is with 10 valence electrons. The last one does not know where to go. The lone electron is called an unpaired electron. But where should the unpaired electron go? Species with incomplete octets are pretty rare and generally are only found in some beryllium, aluminum, and boron compounds including the boron hydrides. Let's take a look at one such hydride, \(BH_3\) (Borane). If one was to make a Lewis structure for \(BH_3\) following the basic strategies for drawing Lewis structures, one would probably come up with this structure (Figure 3): Figure 3: The structure of the Borane molecule. The problem with this structure is that boron has an incomplete octet; it only has six electrons around it. The central Boron now has an octet (there would be three resonance Lewis structures) However... In this structure with a double bond the fluorine atom is sharing extra electrons with the boron. The fluorine would have a '+' partial charge, and the boron a '-' partial charge, this is inconsistent with the electronegativities of fluorine and boron. Size is also an important consideration: The larger the central atom, the larger the number of electrons which can surround it Expanded valence shells occur most often when the central atom is bonded to small electronegative atoms, such as F, Cl and O. There is currently much scientific exploration and inquiry into the reason why expanded valence shells are found. The top area of interest is figuring out where the extra pair(s) of electrons are found. Many chemists think that there is not a very large energy difference between the 3p and 3d orbitals, and as such it is plausible for extra electrons to easily fill the 3d orbital when an expanded octet is more favorable than having a complete octet. This matter is still under hot debate, however and there is even debate as to what makes an expanded octet more favorable than a configuration that follows the octet rule. One of the situations where expanded octet structures are treated as more favorable than Lewis structures that follow the octet rule is when the formal charges in the expanded octet structure are smaller than in a structure that adheres to the octet rule, or when there are less formal charges in the expanded octet than in the structure a structure that adheres to the octet rule. The sulfate ion, SO 4 -2 . is an ion that prefers an expanded octet structure. If instead we made a structure for the sulfate ion with an expanded octet, it would look like this: Figure 13 Looking at the formal charges for this structure, the sulfur ion has six electrons around it (one from each of its bonds). This is the same amount as the number of valence electrons it would have naturally. This leaves sulfur with a formal charge of zero. The two oxygens that have double bonds to sulfur have six electrons each around them (four from the two lone pairs and one each from the two bonds with sulfur). The ICl 4 - ion thus has 12 valence electrons around the central Iodine (in the 5 d orbitals) Expanded Lewis structures are also plausible depictions of molecules when experimentally determined bond lengths suggest partial double bond characters even when single bonds would already fully fill the octet of the central atom. Despite the cases for expanded octets, as mentioned for incomplete octets, it is important to keep in mind that, in general, the octet rule applies. Practice Problems Draw the Lewis structure for the molecule I 3 - . Draw the molecule ClF 3 . The central atom for an expanded octet must have an atomic number larger than what? Draw the Lewis structure for the molecule NO 2 .