The two flourines that share single bonds with boron have seven electrons around them six from their three lone pairs and one from their single bonds with boron. This is the same amount as the number of valence electrons they would have on their own, so they both have a formal charge of zero.
Finally, boron has four electrons around it one from each of its four bonds shared with fluorine. This is one more electron than the number of valence electrons that boron would have on its own, and as such boron has a formal charge of This structure is supported by the fact that the experimentally determined bond length of the boron to fluorine bonds in BF 3 is less than what would be typical for a single bond see Bond Order and Lengths.
Boron on the other hand, with the much lower electronegativity of 2. This formal charge-electronegativity disagreement makes this double-bonded structure impossible. However the large electronegativity difference here, as opposed to in BH 3 , signifies significant polar bonds between boron and fluorine, which means there is a high ionic character to this molecule. This suggests the possibility of a semi-ionic structure such as seen in Figure None of these three structures is the "correct" structure in this instance.
The most "correct" structure is most likely a resonance of all three structures: the one with the incomplete octet Figure 4 , the one with the double bond Figure 5 , and the one with the ionic bond Figure 6. The most contributing structure is probably the incomplete octet structure due to Figure 5 being basically impossible and Figure 6 not matching up with the behavior and properties of BF 3. As you can see even when other possibilities exist, incomplete octets may best portray a molecular structure.
This structure completes boron's octet and it is more common in nature. This exemplifies the fact that incomplete octets are rare, and other configurations are typically more favorable, including bonding with additional ions as in the case of BF 3.
More common than incomplete octets are expanded octets where the central atom in a Lewis structure has more than eight electrons in its valence shell. In expanded octets, the central atom can have ten electrons, or even twelve.
Molecules with expanded octets involve highly electronegative terminal atoms, and a nonmetal central atom found in the third period or below , which those terminal atoms bond to. Expanded valence shells are observed only for elements in period 3 i. The 'octet' rule is based upon available n s and n p orbitals for valence electrons 2 electrons in the s orbitals, and 6 in the p orbitals.
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:. 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 A strict adherence to the octet rule forms the following Lewis structure:. If we look at the formal charges on this molecule, we can see that all of the oxygen atoms have seven electrons around them six from the three lone pairs and one from the bond with sulfur. This is one more electron than the number of valence electrons then they would have normally, and as such each of the oxygen atoms in this structure has a formal charge of If instead we made a structure for the sulfate ion with an expanded octet, it would look like this:.
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.
This is the same amount of electrons as the number of valence electrons that oxygen atoms have on their own, and as such both of these oxygen atoms have a formal charge of zero. The two oxygens with the single bonds to sulfur have seven electrons around them in this structure six from the three lone pairs and one from the bond to sulfur. That is one electron more than the number of valence electrons that oxygen would have on its own, and as such those two oxygens carry a formal charge of Remember that with formal charges, the goal is to keep the formal charges or the difference between the formal charges of each atom as small as possible.
As a result, the second period elements more specifically, the nonmetals C, N, O, F obey the octet rule without exceptions. Phosphorus pentachloride : In the PCl 5 molecule, the central phosphorus atom is bonded to five Cl atoms, thus having 10 bonding electrons and violating the octet rule.
The overall geometry of the molecule is depicted trigonal bipyramidal , and bond angles and lengths are highlighted. However, some of the third-period elements Si, P, S, and Cl have been observed to bond to more than four other atoms, and thus need to involve more than the four pairs of electrons available in an s 2 p 6 octet.
Although the energy of empty 3d-orbitals is ordinarily higher than that of the 4s orbital, that difference is small and the additional d orbitals can accommodate more electrons. Therefore, the d orbitals participate in bonding with other atoms and an expanded octet is produced. Examples of molecules in which a third period central atom contains an expanded octet are the phosphorus pentahalides and sulfur hexafluoride.
Sulfur hexafluoride : In the SF 6 molecule, the central sulfur atom is bonded to six fluorine atoms, so sulfur has 12 bonding electrons around it. The overall geometry of the molecule is depicted tetragonal bipyramidal, or octahedral , and bond angles and lengths are highlighted. For atoms in the fourth period and beyond, higher d orbitals can be used to accommodate additional shared pairs beyond the octet.
The relative energies of the different kinds of atomic orbital reveal that energy gaps become smaller as the principal energy level quantum number n increases, and the energetic cost of using these higher orbitals to accommodate bonding electrons becomes smaller.
Privacy Policy. Skip to main content. Basic Concepts of Chemical Bonding. Search for:. Exceptions to the Octet Rule The Incomplete Octet While most elements below atomic number 20 follow the octet rule, several exceptions exist, including compounds of boron and aluminum. Key Takeaways Key Points The octet rule states that atoms with an atomic number below 20 tend to combine so that they each have eight electrons in their valence shells, which gives them the same electronic configuration as a noble gas.
The two elements that most commonly fail to complete an octet are boron and aluminum; they both readily form compounds in which they have six valence electrons, rather than the usual eight predicted by the octet rule. While molecules exist that contain atoms with fewer than eight valence electrons, these compounds are often reactive and can react to form species with eight valence electrons. For example, BF 3 will readily bind a fluoride anion to form the BF 4 — anion, in which boron follows the octet rule.
Key Terms atomic number : The number, equal to the number of protons in an atom, that determines its chemical properties. Symbol: Z. Odd-Electron Molecules Molecules with an odd number of electrons disobey the octet rule. Learning Objectives Describe the deviation from the octet rule by free radicals.
Key Takeaways Key Points While the majority of compounds formed from atoms below atomic number 20 follow the octet rule, there are many examples of compounds that do not. Having an odd number of electrons in a molecule guarantees that it does not follow the octet rule, because the rule requires eight electrons or two for hydrogen around each atom.
The most commonly encountered stable species that exist with an odd number of electrons are nitrogen oxides, such as nitric oxide NO and nitrogen dioxide NO 2 , both of which are free radicals and disobey the octet rule. Key Terms metastable : Of or pertaining to a physical or chemical state that is relatively long-lived, but may decay to a lower energy state when perturbed.
They vary in reactivity and stability from highly reactive, occurring as transient short-lived species, to metastable. However, some of the third-period elements Si, P, S, and Cl have been observed to bond to more than four other atoms, and thus need to involve more than the four pairs of electrons available in an s 2 p 6 octet.
Although the energy of empty 3d-orbitals is ordinarily higher than that of the 4s orbital, that difference is small and the additional d orbitals can accommodate more electrons. Therefore, the d orbitals participate in bonding with other atoms and an expanded octet is produced.
Examples of molecules in which a third period central atom contains an expanded octet are the phosphorus pentahalides and sulfur hexafluoride. For atoms in the fourth period and beyond, higher d orbitals can be used to accommodate additional shared pairs beyond the octet.
The relative energies of the different kinds of atomic orbital reveal that energy gaps become smaller as the principal energy level quantum number n increases, and the energetic cost of using these higher orbitals to accommodate bonding electrons becomes smaller.
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