Lewis electron-dot structures

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1 CHAPTER 9 SUPPLEMENTAL Lewis electrondot structures shows how atoms are bonded by giving the location of bonded electron pairs and position of (nonbonding) lone pairs of electrons 1. Count the total number of valence electrons of the atoms in the species. Add one electron to the count for each negative charge, and subtract one electron for each positive charge. Group # valence electrons example 1 1 H 3 3 B, Ga 4 4 C, Si 5 5 P, As 6 6 S, O 7 7 Cl, I 8 8 Ne, Xe 1. Skeleton structures Identify the central atom(s) and the terminal atom(s) in your compound. a) H and F are ALWAYS terminal atoms. b) Halogen atoms are USUALLY terminal, unless allhalogen species. c) B, C, N, and atoms (usually the more electropositive) of later rows in the periodic table are USUALLY central atoms. 3. Draw a skeleton structure with single bonds join the terminal atoms to the central atom. Count the number of electrons used ( per single bond) and subtract this number from calculated value of valence electrons. 1

2 4. Terminal atoms Use the remaining valence electrons to complete the octets around the terminal atoms (but NOT for hydrogen!) by adding lone pairs. Start with the most electronegative atoms first. 5. Central atoms (a) If there are any valence electrons left over, add them as lone pairs around the central atom(s). (b) If there are fewer than 8 electrons on the central atom you will need to make a multiple bond. Move a pair of nonbonding electrons from a terminal atom to make a new bond connecting the central atom. Atoms that often form multiple bonds are C,N, O, and S. We now have a legitimate Lewis dot structure. 3 Is it the best possible model we can make? 6. Utilize Formal charge Assigning formal charges to atoms in a Lewis dot structure lets us know whether there are better ways to distribute electrons in the structure. Formal charge compares the number of valence electrons an atom brings to a species to how many electrons are assigned to the atom in the Lewis dot structure. 4

3 Formal charge and Lewis Structure (a) Both electrons of a lone pair are assigned to the atom they drawn on. (b) Assume all bonds in the structure are nonpolar covalent and the two bonding electrons are equally shared (one electron is assigned to each atom). formal charge = formal charge = 1 (# of valence e ) (# of lone pair e ) (# of bonding e ) OR (# of valence e ) x (# of lone pairs ) (# of bonds ) ***You may find it easier to assign formal charges by counting and comparing electrons in your structure, than by using these formulas. 5 Formal charge and Lewis Structure (c) Since all electrons are assigned when calculating formal charges then the sum of formal charges of ALL atoms in a species MUST ADD UP to the total charge of the molecule or ion. (d) Positive and negative formal charges in the same structure indicate the structure may not be the best possible structure. 6 3

4 what DO WE DO? Eliminate differences in formal charge (when possible) by: a) using lone pairs to form double or triple bonds b) use double or triple bonds to form lone pairs c) Move the electrons AWAY FROM the atom with the more negative formal charge If it is not possible to eliminate all formal charge differences, the best structure probably puts negative formal charges on the more electronegative atoms 7 General tips for Lewis structures 1. C, N, O, and F normally OBEY the octet rule.. Terminal atoms normally OBEY the octet rule. 3. Central atoms MAY OR MAY NOT obey the octet rule. a) Be, B, Al (Group # and 3) can have less than an octet (electron deficient) b) Period # 3 (P, S, Cl) and periods below can have more than an octet (electronrich). 8 4

5 Does your Lewis Structure have Resonance Structures? If your structure treats two equivalent atoms differently, then there is likely one or more similar RESONANCE structures that can also be considered the best possible structure. The real structure is better represented by an average of the RESONANCE structures ( ). The electrons are DELOCALIZED. For ozone (O 3 ) 9 5

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