Limitations of The Octet Rule

Octet rule becomes relevant when considering main-group elements of the second and third periods in the periodic table, viz.: Carbon, nitrogen, oxygen, and fluorine. While the octet rule is very useful in terms of what might be expected from the behavior of many elements, there are limitations. Thus, there exist exceptions or cases when the octet rule just fails to work, making chemical bonding much more complex. Such examples are hydrogen and helium, known to be stable with fewer than eight electrons. Some atoms also can have more than the octet in so-called hypervalent molecules, which is an antipode of a simplistic view of electron configuration.

Understanding of the Octet Rule

The octet rule suggests that atoms bond in a way that enables them to achieve a full outer shell of eight electrons for each atom hence offering them stability. This is best applied to main group elements, particularly to the second and third periods of the periodic table. These are done by three main mechanisms: ionic bonding, wherein electrons are transferred from atom to atom; covalent bonding, where they will be shared; and coordinate covalent bonding, in which both electrons come from one of the participating atoms.
While the octet rule can be applied to explain a good number of chemical reactions, it is not universal. For example, certain elements are stable with fewer than eight electrons. Other atoms can hold more than an octet. Molecules such as these are called hypervalent. Examples include SF₆ and PF₅. The octet rule is a guiding principle but one that falls way short of describing the full complexity of chemical bonding and molecular stability.

Exceptions to the Octet Rule

The octet rule does have some notable exceptions because of its limitations. First of these are the molecules that have incomplete octets wherein the central atom has less than eight electrons. Examples of this are boron trifluoride, BF₃, and aluminum chloride, AlCl₃, wherein the boron and aluminum atoms achieve stability even with just six electrons in their valence shells.
The second category refers to those molecules that contain an odd number of electrons. For example, NO₂ and ClO₂ do not follow the octet rule as all the electrons therein their structure is not paired.
The rule is also not applicable to hypervalent molecules in cases when more than eight electrons are allowed on center atoms. Starting from the third-period elements, like phosphorus and sulfur, d-orbitals can participate in the valence shell of the abovementioned elements, thus forming compounds that violate the octet rule but remain stable.

Real-Life Applications and Importance

In real-life applications, the constraints of the octet rule realize themselves in several areas: from materials science to biochemistry and pharmacology. As one simple example, the development of pharmaceuticals is based on the idea of intermolecular interactions at the atomic level. For many biologically active compounds, the octet rule is not observed strictly; therefore, advanced theories like molecular orbital theory are needed to predict their behavior.
On the other side, new material creation in material science depends on manipulating atomic interactions. Properties in most modern materials, for example, semiconductors or catalysts, boil down indeed to electron-deficient/hypervalent species, which cannot be accounted for by the octet rule alone.

Limitations to the Octet Rule

Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:

• Odd-electron molecules have an odd number of valence electrons and therefore have an unpaired electron.
• Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration.
• Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.

Odd-electron Molecules

We call molecules that contain an odd number of electrons free radicals. Nitric oxide, NO, is an example of an odd-electron molecule; it is produced in internal combustion engines when oxygen and nitrogen react at high temperatures.

To draw the Lewis structure for odd-electron molecules like NO, we follow the same five steps we would for other molecules, but with a few minor changes:

1. Determine the total number of valence (outer shell) electrons. The sum of the valence electrons is 5 (from N) + 6 (from O) = 11. The odd number immediately tells us that we have a free radical, so we know that not every atom can have eight electrons in its valence shell.

2. Draw a skeleton structure of the molecule. We can easily draw a skeleton with an N–O single bond:
   N–O

3. Distribute the remaining electrons as lone pairs on the terminal atoms. In this case, there is no central atom, so we distribute the electrons around both atoms. We give eight electrons to the more electronegative atom in these situations; thus oxygen has the filled valence shell:

4. Place all remaining electrons on the central atom. Since there are no remaining electrons, this step does not apply.

5. Rearrange the electrons to make multiple bonds with the central atom in order to obtain octets wherever possible. We know that an odd-electron molecule cannot have an octet for every atom, but we want to get each atom as close to an octet as possible. In this case, nitrogen has only five electrons around it. To move closer to an octet for nitrogen, we take one of the lone pairs from oxygen and use it to form a NO double bond. (We cannot take another lone pair of electrons on oxygen and form a triple bond because nitrogen would then have nine electrons:)

Electron-deficient Molecules

These are a few molecules that contain central atoms that do not have a filled valence shell. Generally, these are molecules with central atoms from groups 2 and 13, outer atoms that are hydrogen, or other atoms that do not form multiple bonds. For example, in the Lewis structures of beryllium dihydride, BeH₂, and boron trifluoride, BF₃, beryllium, and boron atoms each have only four and six electrons, respectively. It is possible to draw a structure with a double bond between a boron atom and a fluorine atom in BF₃, satisfying the octet rule, but experimental evidence indicates the bond lengths are closer to that expected for B–F single bonds. This suggests the best Lewis structure has three B–F single bonds and an electron-deficient boron. The reactivity of the compound is also consistent with an electron-deficient boron. However, the B–F bonds are slightly shorter than what is expected for B–F single bonds, indicating that some double bond character is found in the actual molecule.

An atom like the boron atom in BF₃, which does not have eight electrons, is very reactive. It readily combines with a molecule containing an atom with a lone pair of electrons. For example, NH₃ reacts with BF₃ because the lone pair on nitrogen can be shared with the boron atom:

Hypervalent Molecules

Elements in the second period of the periodic table (n = 2) can accommodate only eight electrons in their valence shell orbitals because they have only four valence orbitals (one 2s and three 2p orbitals). Elements in the third and higher periods (n ≥ 3) have more than four valence orbitals and can share more than four pairs of electrons with other atoms because they have empty d-orbitals in the same shell. Molecules formed from these elements are sometimes called hypervalent molecules. Figures given below show the Lewis structures for two hypervalent molecules, PCl₅ and SF₆.

In some of the hypervalent molecules, like IF₅ and XeF₄, some of the electrons in the outer shell of the central atom are lone pairs:

When we write the Lewis structures for these molecules, we find that we have electrons left over after filling the valence shells of the outer atoms with eight electrons. These additional electrons must be assigned to the central atom.

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Some Solved Examples

Example 1

Question: Octet rule cannot be applied to the non-metals after:

1) Carbon
2) Silicon
3) Oxygen
4) Nitrogen

Solution: The octet rule cannot be applied to the non-metals after silicon. These elements can “expand their octet” and have more than eight valence electrons around the central atom. Hence, the answer is the option (2) Silicon.

Example 2

Question: Octet rule is based upon:

1) Shape of the molecules
2) Chemical inertness of noble gases
3) Energy of the molecule
4) None

Solution: The Octet rule is based on the chemical inertness of noble gases. This rule is derived from the fact that noble gases have a stable electron configuration with a full outer shell, influencing the bonding behavior of other elements. Hence, the answer is the option (2) Chemical inertness of noble gases.

Example 3

Question: The group having species with complete octets in their ionic form is:

1)$\mathrm{O}^{2-}, \mathrm{F}^{-}, \mathrm{Na}, \mathrm{Mg}^{2+}$

2)$\mathrm{O}^{-}, \mathrm{F}^{-}, \mathrm{Na}^{+}, \mathrm{Mg}^{2+}$

3)(Correct)$\mathrm{O}^{2-}, \mathrm{F}^{-}, \mathrm{Na}^{+}, \mathrm{Mg}^{2+}$

4)$\mathrm{O}^{-}, \mathrm{F}^{-}, \mathrm{Na}, \mathrm{Mg}^{+}$

1) O2−,F−,Na,Mg2+ 2) O−,F−,Na+,Mg2+ 3) (correct) O2−,F−,Na+,Mg2+ 4) O−,F−,Na,Mg+

Solution: The group having species with a complete octet in their ionic form includes O^2-, F^-, Na⁺, and Mg²⁺. These ions achieve a noble gas configuration with a complete octet. Hence, the answer is option (3) O^2-, F^-}, Na⁺, Mg²⁺

Example 4

Question: Which of the following is an example of an odd-electron molecule?

1)NO₂
2)CO₂
3)N₂
4)O₂

Solution: Nitric oxide (NO) is an example of an odd-electron molecule because it has an odd number of valence electrons (11 in total). Hence, the answer is the option (1)(NO₂) .

Example 5

Question: The central atom in which of the following compounds is electron-deficient?

1)BeH₂
2)BF₃
3)CH₄
4) NH₃

Solution: In BF₃, the boron atom is electron-deficient because it only has six electrons around it instead of the full octet. Hence, the answer is the option (2)BF₃

Summary

The octet rule, in layman's terms, is one of the basic tenets of chemistry, guiding the formation of bonds between atoms to attain stability with a full outer shell of eight electrons. This paper has managed to point out some major limitations of the universal application of the octet rule. We have covered the concept, defining its tenets and mechanisms of bonding, but also its exceptions among incomplete octets, odd electron molecules, and hypervalent compounds.