Write the Lewis structure for each molecule. These molecules do not follow the octet rule.
Some conditions are there which violate the octet rule they are :
Three cases can be constructed that do not follow the Octet Rule, and as such, they are known as the exceptions to the Octet Rule. 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
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 (NONO). 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? The unpaired electron is usually placed in the Lewis Dot Structure so that each element in the structure will have the lowest formal charge possible. The formal charge is the perceived charge on an individual atom in a molecule when atoms do not contribute equal numbers of electrons to the bonds they participate in. The formula to find a formal charge is:
Formal Charge= [# of valence e- the atom would have on its own] - [# of lone pair electrons on that atom] - [# of bonds that atom participates in]
No formal charge at all is the most ideal situation. An example of a stable molecule with an odd number of valence electrons would be nitrogen monoxide. Nitrogen monoxide has 11 valence electrons (Figure 1). If you need more information about formal charges, see Lewis Structures. If we were to consider the nitrogen monoxide cation (NO+NO+ with ten valence electrons, then the following Lewis structure would be constructed:
Figure 1. Lewis dot structure for the NO+NO+ ion with ten valence electrons.
Nitrogen normally has five valence electrons. In Figure 1, it has two lone pair electrons and it participates in two bonds (a double bond) with oxygen. This results in nitrogen having a formal charge of +1. Oxygen normally has six valence electrons. In Figure 1, oxygen has four lone pair electrons and it participates in two bonds with nitrogen. Oxygen therefore has a formal charge of 0. The overall molecule here has a formal charge of +1 (+1 for nitrogen, 0 for oxygen. +1 + 0 = +1). However, if we add the eleventh electron to nitrogen (because we want the molecule to have the lowest total formal charge), it will bring both the nitrogen and the molecule's overall charges to zero, the most ideal formal charge situation. That is exactly what is done to get the correct Lewis structure for nitrogen monoxide:
Figure 2. The proper Lewis structure for NONO molecule with 11 valence electrons
The Lewis structure for boron tribromide is drawn with a B in the center with three lines connecting to three Brs, each of which contains three pairs of dots around it. The B is the symbol for the single boron atom, and the Brs are the symbol for the three bromine atoms in boron tribromide.
Each of the dots represents a valence electron. Each line represents a bonded pair of electrons that allow the different atoms to combine together to form a single molecule. Separated, each bromine atom contains seven valence electrons, and each boron atom contains three valence electrons, leading to a total of 24 valence electrons. This molecule's configuration of electrons is stable, even though it leaves boron with only six electrons in its outer shell.
The Lewis dot structure illustrates these 24 electrons through the dots and lines. Instead of dots, each pair of electrons can also be represented by a line. Conversely, all bonding lines can be represented by dots. The formal charge of an atom can be determined using the Lewis dot structure by taking the total number of valence electrons and subtracting from that the non-bonding valence electrons. Subtracting half the number of bonding electrons from that number gives the formal charge.