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9.2: Lewis Symbols and the Octet Rule

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Lewis Symbols and the Octet Rule

9.2: Lewis Symbols and the Octet Rule

Chemical bonds are complex interactions between two or more atoms or ions, which reduce the potential energy of the molecule. Gilbert N. Lewis developed a model called the Lewis model that simplified the depiction of chemical bond formation and provided straightforward explanations for the chemical bonds seen in most common compounds.  

Lewis Model

The Lewis model depicts chemical bond formation by the sharing or transfer of valence electrons, which helps to attain a stable electron configuration. An ionic bond is formed when electrons are transferred between a metal and a nonmetal, whereas a covalent bond is formed when electrons are shared between two nonmetals.

The Lewis model is solely used to describe bond formations without taking into account the energy changes associated with the attractions and repulsions between electrons and nuclei on neighboring atoms. While these interactions are central to chemical bonding, the exact determination of the values is complex. Instead, Lewis designed special drawings to depict chemical bonds using valence electrons, called Lewis symbols. 

Lewis Symbols

Lewis symbols describe valence electron configurations of atoms and monatomic ions. A Lewis symbol consists of an elemental symbol surrounded by one dot for each of its valence electrons. For example, sodium has one valence electron; so one dot is drawn around the symbol Na. 


For main group elements, the number of valence electrons is indicated by a lettered group number in the periodic table. For example, lithium (Li) belongs to group IA and has one electron; beryllium (Be) is a group IIA element and has two valence electrons. 

There are exceptions to the Lewis model. In helium, the number valence electron is not the same as the group number. The transition metals, lanthanides, and actinides have incompletely filled inner shells; hence they cannot be written in simple Lewis dot symbols. 

The Octet Rule

The halogen molecules (F2, Br2, I2, and At2) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of s- or p-block atoms to form enough bonds to obtain eight valence electrons is known as the octet rule. The octet rule predicts the combinations of atoms that will have lower potential energy when they bond together.

The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons); this is especially true of the nonmetals of the second period of the periodic table (C, N, O, and F). 

  • Group 14 elements have four electrons in their outermost shell and therefore require four more electrons to reach an octet. These four electrons can be gained by forming four covalent bonds, like carbon in CH4 (methane) and silicon in SiH4 (silane). 
  • Group 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. These atoms form three covalent bonds, as in NH3 (ammonia). 
  • Group 16 elements such as oxygen and other atoms obtain an octet by forming two covalent bonds - like bonding with two hydrogen atoms in H2O (water).

There are exceptions to the octet rule. Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. In this case, hydrogen is said to have reached a duet. The transition elements and inner transition elements also do not follow the octet rule.

This text is adapted from Openstax, Chemistry 2e, Chapter 7.3: Lewis Symbols and Structures.

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