Lewis Dot Structures: Exceptions (Simplified): Videos & Practice Problems

In chemistry, the octet rule is a fundamental principle that states atoms tend to bond in such a way that they have eight electrons in their valence shell, achieving a stable electron configuration. However, certain elements can deviate from this rule and still maintain stability, particularly those in groups 2A through 6A of the periodic table.

For elements with incomplete octets, the number of electrons they can accommodate is determined by the formula:

Non-octet electrons = 2 × Group number

For example, elements in group 2A can have 4 electrons (2 × 2), while those in group 3A can have 6 electrons (2 × 3). This pattern continues, allowing group 5A elements to stabilize with 10 electrons, group 6A with 12, and so forth. Thus, elements can possess more than 8 electrons around them and still be stable, as seen in their Lewis dot structures.

Understanding these exceptions to the octet rule is crucial for predicting the behavior of certain elements during chemical bonding. It highlights the diversity of electron configurations and the complexity of molecular stability beyond the traditional octet framework.

To draw the Lewis dot structure for xenon dibromide (XeBr₂), we start by determining the total number of valence electrons available. Xenon, a noble gas in group 8A, has 8 valence electrons, while bromine, located in group 7A, has 7 valence electrons. Since there are two bromine atoms, the total number of valence electrons is:

Total Valence Electrons = 8 (from Xe) + 2 × 7 (from Br) = 22

In the Lewis structure, xenon is placed at the center, bonded to the two bromine atoms. To satisfy the octet rule for the surrounding bromine atoms, we distribute electrons around them. Each bromine atom will have three lone pairs (6 electrons) and one shared pair (2 electrons from the bond), totaling 8 electrons for each bromine:

Electrons around each Br = 6 (lone pairs) + 2 (shared) = 8

This arrangement uses up 16 of the 22 valence electrons. The remaining 6 electrons are placed around the xenon atom as lone pairs. These 6 electrons are distributed evenly, resulting in three lone pairs around xenon:

Electrons around Xe = 6 (lone pairs) + 4 (from bonds) = 10

Thus, the final Lewis structure shows xenon with 10 electrons surrounding it, which breaks the octet rule. While the ideal non-octet number for xenon would be 16, in this case, it is stable with 10 electrons. This illustrates that central atoms like xenon can exceed the octet rule, making it an exception in molecular structures.

In the study of electron molecules, free radicals play a significant role. Free radicals are defined as molecules or ions that possess a single unpaired electron, which results in an odd number of total valence electrons. This characteristic is crucial for identifying radical compounds. When drawing these structures, the unpaired electron is typically placed on the less electronegative element, with the exception of hydrogen. This is due to hydrogen's adherence to the duet rule, which states that it can only accommodate two electrons.

For instance, consider nitrogen monoxide (NO). In this molecule, the unpaired electron is located on the nitrogen atom, which is less electronegative than oxygen. To determine if nitrogen monoxide is a radical compound, we can calculate the total number of valence electrons. Nitrogen, being in group 5A, contributes 5 valence electrons, while oxygen, from group 6A, contributes 6 valence electrons. Adding these together gives a total of 11 valence electrons, which is indeed an odd number. This confirms that nitrogen monoxide is classified as a radical molecule.

When analyzing radical or free radical compounds, it is essential to check for the odd number of valence electrons, as this is a key indicator of their radical nature.

The Lewis structure for nitrogen dioxide (NO₂) illustrates the concept of radicals in chemistry. To begin, we calculate the total number of valence electrons available for the molecule. Nitrogen, located in group 5A of the periodic table, contributes 5 valence electrons, while each oxygen atom, found in group 6A, contributes 6 electrons. With two oxygen atoms, the total valence electrons sum up to:

5 (from N) + 6 × 2 (from O) = 17 valence electrons.

Since 17 is an odd number, this indicates the presence of a radical, which is a species with an unpaired electron. In constructing the Lewis structure, nitrogen is placed at the center, forming single bonds with each oxygen atom. Initially, this configuration uses 16 of the available electrons, leaving one unpaired electron. At this stage, nitrogen does not satisfy the octet rule, as it only has 5 electrons surrounding it.

To address this, we can create a double bond between nitrogen and one of the oxygen atoms. This is achieved by using one of the lone pairs from oxygen to form a double bond, allowing nitrogen to have 7 electrons around it. The resulting structure shows nitrogen with a double bond to one oxygen and a single bond to the other, with the unpaired electron remaining on the nitrogen atom, confirming that nitrogen dioxide is indeed a radical.

This structure highlights the importance of the octet rule, which states that atoms tend to form bonds until they are surrounded by eight valence electrons, and illustrates how radicals can exist when there is an odd number of valence electrons.