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Lewis Electron-Dot Formulas

Lewis Electron-Dot Symbols

A Lewis electron-dot symbol is a symbol in which the valence shell electrons are represented by dots placed around the letter symbol of an element. 


 Lewis symbols are usually written only for the main group elements.  This is easy to do since the number of valence electrons is equal to the group number.  Lewis symbols are shown above for the elements in the second period.  The exact placement of the single dots does not matter, for example the single dot for fluorine can be written on any of the four sides.  Notice the pairing of dots does not always correspond to the pairing of electrons in the ground state.

The Octet Rule 

The ns2np6 electron configuration of the valence shell of all of the inert gas atoms except helium is commonly called an octet of electrons.  In forming bonds or ions, atoms often gain, lose, or share electrons until they have achieved an outer shell that contains an octet of electrons.  This is called the octet rule. 

Lewis Structures

The first thing we must do in writing a Lewis structure is to deduce the skeletal structure.  The skeletal structure of a polyatomic ion or molecule indicates the order in which the atoms are joined to one another.  This usually consists of one or more central atoms and at least two terminal atoms.  A central atom is bonded to two or more atoms in the structure, while a terminal atom is bonded to only one other atom.  We write the skeletal structure by writing the symbol for the central atom surrounded by the terminal atoms.  We write a single dash to connect each terminal atom to the central atom.  Up to this point we have made no attempt to account for the valence electrons.  The following considerations help us to write plausible skeletal structures.

After choosing a skeletal structure for a polyatomic molecule or ion, follow these steps.

  1. Determine the total number of valence electrons, this is the sum of the valence electrons for each atom.  For a polyatomic anion add one electron for each unit of negative charge. For a polyatomic cation subtract one electron for each unit of positive charge.
  2. Write the skeletal structure, and connect bonded atoms with an electron-pair bond.
  3. Place electrons around terminal atoms so that each has an octet (except hydrogen).
  4. Assign any remaining electrons as lone pairs around the central atom.
  5. If the central atom has fewer than eight electrons (an octet), a multiple bond(s) is likely.  Move one or more pairs of electrons from a terminal atom(s) to a region between it and the central atom to form a double or triple bond.

Molecules That Don't Follow the Octet Rule

The exceptions fall into three categories, each identified by some structural characteristic.

In Lewis structures with an odd number of electrons, it is not possible for all the valence electrons to be in pairs, nor is it possible for all the atoms to have an octet.  One example is nitrogen monoxide, NO, which has 11 electrons.

This consists mostly of molecules containing Group IIA or IIIA atoms.  Consider boron trichloride, BCl3. The molecule consists of boron surrounded by the more electronegative chlorine atoms.  If you connect boron and chlorine atoms by electron pairs and fill out the chlorine atoms with octets of electrons, you obtain:    

The second-period elements carbon, nitrogen, oxygen and fluorine almost always obey the octet rule, with the odd-electron exception.  The valence shell of these atoms holds a maximum of eight electrons.  The third-period elements can hold up up to 18 valence shell electrons.  Apparently third-period elements can depart from the octet rule by having more than eight electrons in their valence shells.  An example of the expanded octet is seen in the compound xenon pentafluoride, XeF5.

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