Lewis Structures ©2011 University of Illinois Board of Trustees

Lewis Structures Are models The representations of the electron arrangements in atoms, ions, or molecules by showing the valence electrons as dots placed around the symbols for the elements Also called Lewis Dot Diagrams or Electron Dot Diagrams Can be drawn for atoms, molecules, anions, cations or ionic compounds Useful when determining the geometry or shape of a molecule ©2011 University of Illinois Board of Trustees

Lewis Structures Why are they important? You can visualize the electrons involved in chemical bonds You can gain a greater understanding of how chemical bonds form ©2011 University of Illinois Board of Trustees

Lewis Structures 1) Count up total number of valence electrons 2) Decide on arrangement of atoms and connect all atoms with single bonds - “least electronegative atoms usually in the middle - “single” atoms usually in center; C always in center, H always on outside. 3) Complete octets on exterior atoms, lone pair electrons (not H, though) 4) Check: - valence electrons math with Step 1 - all atoms (except H) have an octet;if not, try multiple bonds - any extra electrons?Put on central atom ©2011 University of Illinois Board of Trustees

Lewis Structures How Are They Drawn? 1)Count up all of the valence electrons for each atom in the formula for anions, add the charge of the ion to the number of valence electrons for cations, subtract the charge of the ion from the number of electrons 2) Determine the number of octet electrons the formula should have 3) Determine the Number of Bonding Electrons Subtract valence electrons from octet electrons 4) Determine the Number of Bonds Divide # Bonding electrons by 2 5) Draw the Structure with the correct Number of Bonds, least electronegative element is usually the central atom Bond all atoms together by single bonds, then add in the multiple bonds until the rules in the notes are followed 6) Determine number of Lone Pair Electrons and arrange them around the atoms until the octet rule is satisfied for all atoms (except Hydrogen) ©2011 University of Illinois Board of Trustees

Some Examples CO 2 1) Valence Electrons = 16  Carbon is in group 4, 4 valence  Oxygen is in group 6, 6 valence X 2 atoms = 12 total  = 16 2) Octet Electrons = 8 ea. X 3 = 24 3) Bonding Electrons = 24 – 16 = 8 4) Number of Bonds = (8/2) = 4 5) O = C = O 6) Non bonding electrons = 16 – 8 = 8 ©2011 University of Illinois Board of Trustees

Carbon Dioxide This Lewis Structure uses the correct number of electrons, but does not obey the octet rule. This Lewis Structure uses the correct number of electrons and obeys the octet rule. O = C = O ©2011 University of Illinois Board of Trustees

Hydrogen Cyanide HCN 1) Valence electrons = 10 2) Octet electrons = 18 3) Bonding electrons = 8 4) Number of Bonds = 4 5) H – C N 6) Number of Nonbonding electrons = 2 H – C N: ©2011 University of Illinois Board of Trustees

Exceptions to the Octet Rule (there are always exceptions) BH 3 Each hydrogen accommodates 2 electrons, or one bond. The boron atom in BH 3, on the other hand, has only 6 total electrons. Because boron is a smaller atom, it does not have enough space to accommodate a full octet of 8 electrons. ©2011 University of Illinois Board of Trustees

BF 3 BF 3 has two potential Lewis structures shown below. Left structure (Structure I) is better because it minimizes interaction between molecules. Structure II shows B and F with formals charges. F is a more electronegative atom (attracts more electrons) and will have more 3 lone pairs, shown in Structure I. ©2011 University of Illinois Board of Trustees

Formal Charge Determines which Lewis structure is correct if many structures are possible Compares number of electrons around a bonded atom to the number of electrons a lone atom possesses Formal charge = #valence electrons – (#nonbonding + ½ bonding) The best Lewis Structure is the one where the formal charges are as close as possible to zero ©2011 University of Illinois Board of Trustees

Figure I: Formal charge on top Fluorine is 0. Formal charge on right Fluorine is 0. Formal charge on left fluorine is 0. Formal charge on Boron is 0. Figure II: Formal charge on top Fluorine is 0. Formal charge on right Fluorine is 1. Formal charge on left fluorine is 0. Formal charge on Boron is -1. ©2011 University of Illinois Board of Trustees

NO – Nitric Oxide Exception: odd number of electrons Consider NO – 11 valence electrons Best course of action:  Maximize number of bond  Make sure neither atom in the 2 nd period exceeds an octet One atom will have an odd electron count ©2011 University of Illinois Board of Trustees

1 st 2 nd or 3 rd period elements as central atoms Expanded octets: an exception Atoms in the 3 rd period or higher can old more than 8 electrons They can hold 8, 10, or 12 electrons around the central atom Examples: XeF 4, SF 6, PCl 5 ©2011 University of Illinois Board of Trustees

Bonding and Shapes of Molecules Number of Bonds Number of Unshared Pairs on Central Atom ShapeExamples Linear Trigonal planar Tetrahedral Trigonal Pyramidal Bent BeCl 2 CO 2 BF 3 CH 4, SiCl 4, CCl 4 NH 3, PCl 3 H 2 O, -Be- =C= B C N : O : : Covalent Structure ©2011 University of Illinois Board of Trustees

Linear BeCl 2 ©2011 University of Illinois Board of Trustees

Carbon dioxide Linear geometry Linear ©2011 University of Illinois Board of Trustees

Trigonal Planar BF3 ©2011 University of Illinois Board of Trustees

Tetrahedral Methane –The first member of the paraffin (alkane) hydrocarbons series. a.k.a. (marsh gas, CH 4 ). Methane ©2011 University of Illinois Board of Trustees

Tetrahedral SiCl4 Silicon tetrachloride ©2011 University of Illinois Board of Trustees

Tetrahedral Carbon tetrachloride – “carbon tet” had been used as dry cleaning solvent because it is extremely non-polar. Carbon tetrachloride – CCl 4 ©2011 University of Illinois Board of Trustees

Pyramidal NH 3.. H H H N Ammonia ©2011 University of Illinois Board of Trustees

Trigonal Pyramidal Phosphorus trichloride ©2011 University of Illinois Board of Trustees

Bent.. HH O Water SO 2  (-)  (+) Polar molecule ©2011 University of Illinois Board of Trustees