Lewis Dot Structures: Exceptions: Videos & Practice Problems

In chemistry, the octet rule is a guideline that elements tend to combine in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. However, there are exceptions to this rule. Elements in Group 2A can be stable with just four valence electrons, while those in Group 3A can be stable with six. Transition elements can have expanded octets with more than eight electrons, such as 10, 12, 14, or 16 valence electrons, based on twice their group number.

Free radicals are molecules or ions with an unpaired electron, which can make them highly reactive and potentially damaging to cells. They have an odd number of total valence electrons. To determine the best Lewis dot structure for a radical, one should place the unpaired electron on the element that results in the least formal charges. The formal charge can be calculated using the formula:

group number−number of bonds−nonbonding electrons A Lewis structure with atoms having formal charges close to zero is considered the most stable representation of the molecule. When the total number of valence electrons is an odd number, it often indicates the presence of a radical.

Lewis Dot Structures: Exceptions

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 electronic 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 around them (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, group 4A can have 8 electrons, group 5A can have 10, group 6A can have 12, and this trend continues up to group 8A, which can accommodate 16 electrons.

These exceptions to the octet rule are important when drawing Lewis dot structures, as they illustrate that some elements can achieve stability with fewer or more than eight electrons in their valence shell. Understanding these exceptions is crucial for predicting the behavior of various elements in chemical bonding.

Lewis Dot Structures: Exceptions Example 1

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, contributes 8 valence electrons, while each bromine atom, from group 7A, contributes 7 valence electrons. With two bromine atoms, the total valence electrons sum up to 22 (8 from xenon + 7 × 2 from bromine).

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 will have three lone pairs (6 electrons) and one bond (2 electrons), totaling 8 electrons for each bromine. This arrangement uses up 16 of the 22 available valence electrons.

With 6 electrons remaining, we place these around the xenon atom as three lone pairs. This results in xenon having a total of 10 electrons (2 from each bond with bromine and 6 from the lone pairs). Although this configuration breaks the octet rule, it is permissible for xenon, as it can accommodate more than 8 electrons due to its position in the periodic table. In fact, xenon can ideally hold up to 12 or even 16 electrons in certain compounds, but in this case, it is stable with 10 electrons.

Thus, the Lewis dot structure for xenon dibromide illustrates xenon at the center with two bromine atoms bonded to it, each surrounded by three lone pairs, demonstrating that central atoms like xenon can exceed the octet rule while remaining stable.

Determine the Lewis Dot Structure for the following compound:SOCl₂

• A.
• B.
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• D.

Answer: (A)

Lewis Dot Structures: Exceptions

Free radicals are molecules or ions characterized by having an unpaired electron, which can lead to damage in healthy cells. These radicals typically possess an odd number of total valence electrons, making them unstable. For instance, consider a molecule containing nitrogen and oxygen. Nitrogen, found in group 5A of the periodic table, has 5 valence electrons, while oxygen, from group 6A, has 6 valence electrons. When combined, this molecule totals 11 valence electrons, indicating the presence of a free radical due to the odd count.

To visualize this, one can draw the Lewis dot structure, which helps in understanding the bonding and electron distribution within the molecule. The formal charge is a useful concept in this context, calculated using the formula:

Formal Charge = Group Number - (Bonds + Nonbonding Electrons)

For nitrogen, with 5 valence electrons, making 2 bonds and having 3 nonbonding electrons, the formal charge is:

Formal Charge (N) = 5 - (2 + 3) = 0

For oxygen, with 6 valence electrons, making 2 bonds and having 4 nonbonding electrons, the formal charge is:

Formal Charge (O) = 6 - (2 + 4) = 0

Both atoms achieve a formal charge of 0, which is ideal for stability in a molecule. The key takeaway is that when calculating the total number of valence electrons results in an odd number, it strongly suggests the presence of a free radical. Understanding these concepts is crucial for exploring the implications of free radicals in health and nutrition.

Lewis Dot Structures: Exceptions Example 2

The Lewis structure for nitrogen dioxide (NO₂) illustrates the concept of radicals in chemistry. To begin, we calculate the total number of valence electrons: nitrogen contributes 5 electrons (as it is in group 5A), and each oxygen contributes 6 electrons (from group 6A), resulting in a total of 17 valence electrons (5 + 6 × 2 = 17). The presence of an odd number of valence electrons indicates that nitrogen dioxide is a radical.

In constructing the Lewis structure, nitrogen is placed at the center, forming single bonds with the two oxygen atoms. Initially, this configuration accounts for 16 electrons, leaving one unpaired electron. At this stage, nitrogen does not satisfy the octet rule, as it only has 5 electrons surrounding it (2 from one bond and 3 from the other).

To address this, we can create a double bond between nitrogen and one of the oxygen atoms by utilizing one of the lone pairs from oxygen. This adjustment allows nitrogen to achieve a total of 7 electrons around it, while still maintaining the radical nature of the molecule due to the presence of the unpaired electron. The final structure of nitrogen dioxide thus features a nitrogen atom bonded to two oxygen atoms, with one of those bonds being a double bond, and a lone electron remaining, confirming its status as a radical.

Draw the Lewis Dot Structure for the radical hydroxide, OH.

• A.
• B.
• C.
• D.

Answer: (B)

Draw the Lewis Dot Structure for POCl₃.

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• B.
• C.
• D.

Answer: (A)

Metalloids can sometimes adopt the bonding preferences of similar nonmetals. Based on your knowledge of expanded octets, draw the Lewis Dot Structure for the following ion, SiF₆^2–.

• A.
• B.
• C.
• D.

Answer: (B)

Here’s what students ask on this topic:

What are the exceptions to the octet rule in Lewis dot structures?

In Lewis dot structures, the octet rule states that atoms tend to have eight electrons in their valence shell. However, there are exceptions. Elements in Group 2A can be stable with just four valence electrons, while those in Group 3A can be stable with six. Transition elements can have expanded octets with more than eight electrons, such as 10, 12, 14, or 16 valence electrons, based on twice their group number. Additionally, free radicals, which have an odd number of total valence electrons, also break the octet rule.

How do you determine the best Lewis dot structure for a free radical?

To determine the best Lewis dot structure for a free radical, you should place the unpaired electron on the element that results in the least formal charges. The formal charge can be calculated using the formula: $\mathrm{Formal} \mathrm{charge}=\mathrm{Group} \mathrm{number}-\mathrm{number} \mathrm{of} \mathrm{bonds}-\mathrm{nonbonding} \mathrm{electrons}$. A Lewis structure with atoms having formal charges close to zero is considered the most stable representation of the molecule.

What is a free radical in chemistry?

A free radical in chemistry is a molecule or ion with an unpaired electron, making it highly reactive. Free radicals have an odd number of total valence electrons. They can cause damage to cells and are often discussed in the context of health and nutrition. In Lewis dot structures, the presence of an odd number of valence electrons usually indicates a free radical.

Why can some elements have more than eight electrons in their valence shell?

Some elements can have more than eight electrons in their valence shell due to the availability of d-orbitals in their electron configuration. This is common in transition elements and elements in periods 3 and beyond. These elements can expand their octet to accommodate 10, 12, 14, or even 16 valence electrons, based on twice their group number. This expanded octet allows them to form more bonds and achieve greater stability in certain compounds.

How do you calculate the formal charge of an atom in a Lewis structure?

The formal charge of an atom in a Lewis structure can be calculated using the formula: $\mathrm{Formal} \mathrm{charge}=\mathrm{Group} \mathrm{number}-\mathrm{number} \mathrm{of} \mathrm{bonds}-\mathrm{nonbonding} \mathrm{electrons}$. This formula helps determine the most stable Lewis structure by ensuring that the formal charges on all atoms are as close to zero as possible.