Lewis Dot Structure Example - Octet Rule Exception

Lewis dot structures are useful to predict the geometry of a molecule. Sometimes, one of the atoms in the molecule does not follow the octet rule for arranging electron pairs around an atom. This example uses the steps outlined in How to Draw A Lewis Structure to draw a Lewis structure of a molecule where one atom is an exception to the octet rule. Review of Electron Counting The total number of electrons shown in a Lewis structure is the sum of the valence electrons of each atom. Remember: non-valence electrons are not shown. Once the number of valence electrons has been determined, here is the list of steps normally followed to place the dots around the atoms: Connect the atoms by single chemical bonds.The number of electrons to be placed is t-2n, where t is the total number of electrons and n is the number of single bonds. Place these electrons as lone pairs, starting with outer electrons (besides hydrogen) until every outer electrons has 8 electrons. Place lone pairs on most electronegative atoms first.After lone pairs are placed, central atoms may lack an octet. These atoms form a double bond. Move a lone pair to form the second bond.Question:Draw the Lewis structure of the molecule with molecular formula ICl3.Solution:Step 1: Find the total number of valence electrons.Iodine has 7 valence electronsChlorine has 7 valence electronsTotal valence electrons 1 iodine (7) 3 chlorine (3 x 7)Total valence electrons 7 21Total valence electrons 28Step 2: Find the number of electrons needed to make the atoms happyIodine needs 8 valence electronsChlorine needs 8 valence electronsTotal valence electrons to be happy 1 iodine (8) 3 chlorine (3 x 8)Total valence electrons to be happy 8 24Total valence electrons to be happy 32Step 3: Determine the number of bonds in the molecule.number of bonds (Step 2 - Step 1)/2number of bonds (32 - 28)/2number of bonds 4/2number of bonds 2This is how to identify an exception to the octet rule. There are not enough bonds for the number of atoms in molecule. ICl3 should have three bonds to bond the four atoms together. Step 4: Choose a central atom.Halogens are often the outer atoms of a molecule. In this case, all the atoms are halogens. Iodine is the least electronegative of the the two elements. Use iodine as the center atom.Step 5: Draw a skeletal structure.Since we do not have enough bonds to connect all four atoms together, connect the central atom to the other three with three single bonds.Step 6: Place electrons around outside atoms.Complete the octets around the chlorine atoms. Each chlorine should get six electrons to complete their octets.Step 7: Place remaining electrons around the central atom.Place the remaining four electrons around the iodine atom to complete the structure. The completed structure appears at the beginning of the example. Limitations of Lewis Structures Lewis structures first came into use early in the twentieth century when chemical bonding was poorly understood. Electron dot diagrams help illustrate electronic structure of molecules and chemical reactivity. Their use remains popular with chemistry educators introducing the valence-bond model of chemical bonds and they are often used in organic chemistry, where the valence-bond model is largely appropriate. However, in the fields of inorganic chemistry and organometallic chemistry, delocalized molecular orbitals are common and Lewis structures dont accurately predict behavior. While its possible to draw a Lewis structure for a molecule known empirically to contain unpaired electrons, use of such structures leads to errors in estimating bond length, magnetic properties, and aromaticity. Examples of these molecules include molecular oxygen (O2), nitric oxide (NO), and chlorine dioxide (ClO2). While Lewis structures have some value, the reader is advised valence bond theory and molecular orbital theory do a better job describing the behavior of valence shell electrons. Sources Lever, A. B. P. (1972). Lewis Structures and the Octet Rule. An automatic procedure for writing canonical forms. J. Chem. Educ. 49 (12): 819.  doi:10.1021/ed049p819Lewis, G. N. (1916). The Atom and the Molecule. J. Am. Chem. Soc. 38 (4): 762–85. doi:10.1021/ja02261a002Miessler, G.L.; Tarr, D.A. (2003). Inorganic Chemistry (2nd ed.). Pearson Prentice–Hall. ISBN 0-13-035471-6.Zumdahl, S. (2005). Chemical Principles. Houghton-Mifflin. ISBN 0-618-37206-7.