Why exceptions to the octet rule

In sulfur hexafluoride, the central sulfur atom has twelve electrons from its six bonds to fluorine atoms see Figure 5. Figure 4. Phosphorus pentachloride.

Left image: Lewis structure Right image: molecular model. Figure 5. Sulfur hexafluoride. Use the link below to answer the following questions:. Skip to main content.

Covalent Bonding. Search for:. Are rules always followed? Figure 1. Beryllium hydride. Figure 2. Boron trichloride. Summary Exceptions exist to the rules for covalent bonding. These exceptions apply to atoms whose electrons will not accommodate the normal octet rule. What is the next step? The overall geometry of the molecule is depicted tetragonal bipyramidal, or octahedral , and bond angles and lengths are highlighted.

For atoms in the fourth period and beyond, higher d orbitals can be used to accommodate additional shared pairs beyond the octet. The relative energies of the different kinds of atomic orbital reveal that energy gaps become smaller as the principal energy level quantum number n increases, and the energetic cost of using these higher orbitals to accommodate bonding electrons becomes smaller.

Privacy Policy. Skip to main content. Basic Concepts of Chemical Bonding. Search for:. Exceptions to the Octet Rule The Incomplete Octet While most elements below atomic number 20 follow the octet rule, several exceptions exist, including compounds of boron and aluminum. Key Takeaways Key Points The octet rule states that atoms with an atomic number below 20 tend to combine so that they each have eight electrons in their valence shells, which gives them the same electronic configuration as a noble gas.

The two elements that most commonly fail to complete an octet are boron and aluminum; they both readily form compounds in which they have six valence electrons, rather than the usual eight predicted by the octet rule.

While molecules exist that contain atoms with fewer than eight valence electrons, these compounds are often reactive and can react to form species with eight valence electrons. For example, BF 3 will readily bind a fluoride anion to form the BF 4 — anion, in which boron follows the octet rule. Key Terms atomic number : The number, equal to the number of protons in an atom, that determines its chemical properties.

Symbol: Z. Odd-Electron Molecules Molecules with an odd number of electrons disobey the octet rule. Learning Objectives Describe the deviation from the octet rule by free radicals. Key Takeaways Key Points While the majority of compounds formed from atoms below atomic number 20 follow the octet rule, there are many examples of compounds that do not. Having an odd number of electrons in a molecule guarantees that it does not follow the octet rule, because the rule requires eight electrons or two for hydrogen around each atom.

The most commonly encountered stable species that exist with an odd number of electrons are nitrogen oxides, such as nitric oxide NO and nitrogen dioxide NO 2 , both of which are free radicals and disobey the octet rule. Key Terms metastable : Of or pertaining to a physical or chemical state that is relatively long-lived, but may decay to a lower energy state when perturbed.

They vary in reactivity and stability from highly reactive, occurring as transient short-lived species, to metastable. Hydrogen is an exception because it can hold a maximum of two electrons in its valence level. Examples of Free Radical Molecules Recall that the Lewis structure of a molecule must depict the total number of valence electrons from all the atoms which are bonded together.

The Expanded Octet Main group elements in the third period and below form compounds that deviate from the octet rule by having more than 8 valence electrons. Learning Objectives Explain why some elements can form an expanded octet. Sulfur, phosphorus, silicon, and chlorine are common examples of elements that form an expanded octet. Phosphorus pentachloride PCl 5 and sulfur hexafluoride SF 6 are examples of molecules that deviate from the octet rule by having more than 8 electrons around the central atom.

Key Terms main group element : Elements that are not part of the transition metal block in the periodic table. Licenses and Attributions. 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 — 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 — often up to 12 or 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.

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

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:. 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, BeH 2 , and boron trifluoride, BF 3 , 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 BF 3 , 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 electron-deficient boron. The reactivity of the compound is also consistent with electron-deficient boron. However, the B—F bonds are slightly shorter than what is actually 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 BF 3 , 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, NH 3 reacts with BF 3 because the lone pair on nitrogen can be shared with the boron atom:.