Name: octet rule exceptions in lewis structures
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Exceptions to the Octet Rule

The octet rule explains chemical bonding in main group compounds by predicting that each atom reaches an 8-electron configuration. However, there are three major exceptions to this rule. 

The first exception is odd-electron species. Most molecules and ions have an even number of electrons. However, certain molecules, called radicals, have one or more unpaired electrons. Radicals with an odd number of unpaired electrons cannot achieve octets. 

The superoxide anion, a radical with one unpaired electron, has 13 valence electrons. It can be represented by two contributing structures where one oxygen has only seven electrons and thus cannot reach an octet. 

The second exception is atoms that form an incomplete octet. For instance, hydrogen, helium, and lithium tend to reach a duet, whereas elements of group 2 and 13, like beryllium and boron, often form molecules with four and six electrons around them, respectively. 

Consider aluminum chloride, which has 24 valence electrons. While all chlorine atoms reach the octet, aluminum gets only 6 valence electrons &#8212; an incomplete octet. 

Although aluminum chloride is stable, it reacts with molecules like ammonia that have an unshared pair of electrons. The nitrogen in ammonia donates its lone pair to aluminum, forming a special bond called a coordinate covalent or dative bond. 

The third exception is elements that can accommodate more than 8 valence electrons or an expanded octet. These elements are located in the third row of the periodic table and below. 

Elements, such as phosphorus, sulfur, or iodine, have access to d orbitals, allowing them to accommodate more than 8 valence electrons &#8212; often up to 12 or 14. 

Consider the tetrachloroiodide anion, which has 36 valence electrons. Even after assigning the bonding electron pairs and satisfying the octet for all atoms, 4 valence electrons remain unassigned. 

These electrons are placed on the central iodine atom, yielding an expanded octet with 12 electrons. Molecules with more than 8 valence electrons around the central atom are called hypervalent.

Remember, elements from the second row of the periodic table, such as carbon or oxygen, have only s and p orbitals and never form hypervalent compounds because collectively they can only hold up to 8 valence electrons.

exceptions 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:

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

<h4>Odd-electron Molecules</h4>

Molecules that contain an odd number of electrons are called 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 an odd-electron molecule like NO, the following steps are considered:

<ol>
	<li>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 indicates that it is a free radical, where not every atom has eight electrons in its valence shell.</li>
	<li>Draw a skeleton structure of the molecule. A skeleton structure with an N&#8211;O single bond can easily be drawn.</li>
	<li>Distribute the remaining electrons as lone pairs on the terminal atoms. In this case, there is no central atom, so the electrons are distributed around both atoms. Eight electrons are assigned to the more electronegative atom in these situations; thus, oxygen has the filled valence shell: 
	<img alt="Nitric oxide Lewis structure diagram; chemical bonding; electron dot representation." src="/files/core/CDNSource/lessons/1066/1066Image_1.svg" style="height: 40px; display: block; margin: 0 auto;" data-altupdatedat="1754904545000" data-alt="Nitric oxide (NO) Lewis structure, chemical bonding diagram, molecular chemistry study." /></li>
	<li>Place all remaining electrons on the central atom. Since there are no remaining electrons, this step does not apply.</li>
	<li>Rearrange the electrons to make multiple bonds with the central atom in order to obtain octets wherever possible. Although an odd-electron molecule cannot have an octet for every atom, each atom should get electrons 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, one of the lone pairs from oxygen is utilized to form a NO double bond. (Another lone pair of electrons cannot be taken from oxygen to form a triple bond because nitrogen would then have nine electrons:) 
	<img alt="Nitric oxide Lewis structure diagram showing nitrogen-oxygen bond with electrons and lone pairs." src="/files/core/CDNSource/lessons/1066/1066Image_2.svg" style="height: 40px; display: block; margin: 0 auto;" data-altupdatedat="1754904545000" data-alt="Nitric oxide (NO) Lewis structure, molecular formula diagram, illustrating chemical bonding." /></li>
</ol>

<h4>Electron-deficient Molecules</h4>

Some molecules, however, 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, the 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&#8211;F single bonds. This suggests the best Lewis structure has three B&#8211;F single bonds and electron-deficient boron. The reactivity of the compound is also consistent with electron-deficient boron. However, the B&#8211;F bonds are slightly shorter than what is actually expected for B&#8211;F single bonds, indicating that some double-bond character is found in the actual molecule.

<img alt="Beryllium hydride and boron trifluoride Lewis structures, molecular geometry illustration." src="/files/core/CDNSource/lessons/1066/1066Image_3.svg" style="height: 80px;" data-altupdatedat="1754904545000" data-alt="Beryllium hydride and boron trifluoride Lewis structures; molecular geometry diagram." />

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:

<img alt="Lewis acid-base reaction equation; BF3 and NH3 forming BF3NH3 complex; molecular diagram." src="/files/core/CDNSource/lessons/1066/1066Image_4.svg" style="height: 80px;" data-altupdatedat="1754904545000" data-alt="Boron trifluoride ammonia reaction diagram; Lewis structures showing bond formation; chemical reaction." />

<h4>Hypervalent Molecules</h4>

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 &#8805; 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, such as PCl5, and SF6. In PCl5, the central atom, phosphorus, shares five pairs of electrons. In SF6, sulfur shares six pairs of electrons.

<img alt="Molecular geometry diagram; phosphorus pentachloride and sulfur hexafluoride Lewis structures." src="/files/core/CDNSource/lessons/1066/1066Image_5.svg" style="height: 110px;" data-altupdatedat="1754904545000" data-alt="Phosphorus pentachloride, sulfur hexafluoride Lewis structures, chemical bond diagrams." />

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

<img alt="Lewis structures of SF6 and XeF6 molecules, depicting electron configurations in diagram form." src="/files/core/CDNSource/lessons/1066/1066image_6.svg" style="height: 110px;" data-altupdatedat="1754904545000" data-alt="Lewis structures of SF6 and XeF4 molecules with electron pair configurations." />

In the Lewis structures for these molecules, there are 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.

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chemistry 2e, Section 7.3: Lewis Symbols and Structures.</a>

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


	...

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octet rule exceptions in lewis structures

Some covalent molecules do not follow the octet rule in their Lewis structures. Their central atoms can fall into three groups: odd-electron molecules, electron-deficient molecules, and hypervalent molecules. These cases show why a Lewis structure may not always give every atom eight valence electrons.

Odd-electron molecules have an odd number of valence electrons, so one electron remains unpaired. Molecules with an odd number of electrons are called radicals. Nitric oxide, NO, is one example, and it forms in internal combustion engines when oxygen and nitrogen react at high temperature. To draw the Lewis structure for NO, first count the valence electrons. Nitrogen contributes 5 and oxygen contributes 6, for a total of 11.

The odd electron means not every atom can reach an octet. A single N–O bond can be drawn first, and the remaining electrons are placed as lone pairs on the terminal atom. In NO, oxygen is the more electronegative atom, so it gets a full valence shell first. Because no electrons remain for a central atom, the next step is to rearrange electrons to make multiple bonds where possible. In NO, one lone pair from oxygen forms an N=O double bond, which brings nitrogen closer to an octet. A triple bond is not used, because that would give nitrogen nine electrons.

Electron-deficient molecules are another exception. In these compounds, the central atom has fewer electrons than needed for a noble-gas configuration. This often happens with central atoms from groups 2 and 13, especially when the outer atoms are hydrogen or other atoms that do not form multiple bonds. Beryllium dihydride, BeH2, and boron trifluoride, BF3, are examples. In their Lewis structures, beryllium has only four electrons and boron has only six.

BF3 can be drawn with a boron-fluorine double bond, but experimental bond lengths are closer to B–F single bonds. That evidence supports the best Lewis structure as three B–F single bonds with electron-deficient boron. The compound is also very reactive for the same reason. Boron in BF3 readily combines with a molecule that has a lone pair, such as NH3, because the lone pair on nitrogen can be shared with boron.

Hypervalent molecules are the third group. These molecules have a central atom with more than eight electrons in its valence shell. Second-period elements, with only four valence orbitals, can hold only eight electrons. Third-period and higher elements can share more than four pairs of electrons because they have empty d orbitals in the same shell. Examples include PCl5 and SF6, where phosphorus shares five pairs of electrons and sulfur shares six pairs.

Some hypervalent molecules, such as IF5 and XeF4, also have lone pairs on the central atom. In these Lewis structures, electrons remain after the outer atoms have filled octets, and those extra electrons are placed on the central atom. Together, odd-electron, electron-deficient, and hypervalent molecules show the main ways Lewis structures can depart from the octet rule.

Octet Rule Exceptions in Lewis Structures

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Learn why some Lewis structures break the octet rule with odd-electron, electron-deficient, and hypervalent molecules like NO, BF3, and SF6.

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Octet Rule Exceptions in Lewis Structures

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