Practice Worksheet for Lewis Structures

Practice Worksheet for Lewis Structures 1. Main concepts - Lewis Structures a. Connectivity b. Bonds & Lone pairs c. Electron Geometry & Molecular Shape d. Resonance Structures Formal Charge (FC) is a charge an atom in the molecule or ion would have if all bonding electrons were shared equally between the bonded atoms. 1 FC= # of valence e - - # of shared e- - #nonbonding e-) 2 e. Assigning Oxidation number using Lewis Structures i. Oxidation number is 0 for atoms in an element. ii. The sum of all oxidation numbers in a molecule or ion must add up to the total charge. iii. In compounds, alkalis (group 1) have oxidation number +1; alkaline earths (group 2) have oxidation number +2. iv. In compounds, fluorine (F) always has oxidation number -1. Other halogens (Cl ext.) have oxidation number -1, except when bonded to elements that are more electronegative such as fluorine or oxygen, where they can have positive oxidation numbers. v. In compounds, hydrogen has oxidation number +1 (when bonded to elements with higher electronegativity). When bonded to metals hydrogen has oxidation number -1 (when bonded to elements with lower elecronegativity). vi. In compounds, oxygen has oxidation number -2, except when bonded to fluorine, where fluorine's rule takes precedence (in OF2 oxygen has oxidation number +2), and if there are O-O (peroxide) bonds. (in H2O2 or Na2O2 oxygen has oxidation number -1). 2. Useful Ideas to Consider When Drawing Lewis Structures a. Arrange atoms so that the central atom has the lowest electronegativity. b. Determine the number of valence electrons. c. Place one electron pair between bonded atoms to form a single bond. d. Use remaining pairs as lone pairs around terminal atoms (except H) so each atom is surrounded by 8 electrons e. If the central atom has fewer than 8 electrons, move one or more of the lone pairs on a terminal atom into a position intermediate between the center and terminal position to form a multiple bond. f. The octet rule is a useful guideline when drawing Lewis structures g. Carbon forms four bonds (four single bonds; two single bonds and one double bond; or one single bond and one triple bond.) In uncharged species, nitrogen forms three bonds and oxygen two. Hydrogen typically forms only one bond to another atom. h. When multiple bonds are formed, both of the atoms involved are usually one of the following: C, N, O, and S. Oxygen has the ability to form multiple bonds with a variety of elements. Carbon forms many compounds having multiple bonds to another carbon, or to nitrogen or oxygen. i. Nonmetals may form single, double, and triple bonds but never quadruple bonds. j. Always account for single bonds and lone pairs before forming multiple bonds. k. Be alert for the possibility that the molecule or ion you are working on is isoelectronic with a species you have seen before. 1 Example (for CO2) Valence =4(C) +2*6(O) =16 e-(total number of valence electrons you have) Octet = 8(C) +2*8(O) =24 e- (total number of electron needed for each atom to be satisfied by an octet rule (8 e- each)) Exceptions: H only needs 2 eBe sometimes only contribute 4 eB usually only contribute 6 eO(need)  V(have) 24  16 Bonds (total number of bonds) = = = 4 bonds 2 2 V  B  2 16  4  2 = 4 lone pairs or Lone pairs (total number of nonbonding electron pairs) = = 2 2 nonbonding electron pairs (or 8 nonbonding electrons) 1. Choose most stable resonance structure based on Formal Charge for SCN- and N2O 2. Consider the following different Lewis structure for OClO- [O Cl O ] [O Cl O ] [O Cl a. Use the boxes to indicate the formal charge for each atom. b. Circle the least stable Lewis structure. 3. Assign oxidation numbers to each of the atoms in the following molecules: a. PO43- P_____ O_____ b. NaH Na____ H_____ c. C2H6 C_____ H_____ 2 O] 4. In which of the following species is it necessary to employ an expanded valence shell to represent the Lewis structure: PO43-, PI3, ICl3, OSCl2, SF4, ClO4-? Explain your choices. 5. Only one of the following Lewis structures is correct. Select that one and indicate the errors in the others. a) cyanate ion b) carbide ion c) hypochlorite ion d) nitrogen (II) oxide 6. One of the following ions has a trigonal planar shape: SO32-, PO43-, PF6-, CO32-. Which ion is it? Explain. 7. Use VSEPR theory to predict the geometric shapes of the following molecules and ions: (a) N2 (b) HCN (c) NH4+ (d) NO3- (e) NSF. 8. Give the Lewis dot structures, electron pair geometry, and molecular shape of the following molecules. For each atom on each molecule, indicate the formal charge and oxidation number. a. b. c. d. e. f. g. SeCl4 I3 PSCl3 IF4PH2TeF42N33 h. i. j. k. l. m. n. o. p. q. r. s. t. SeOCl4 PH4+ NOICl2PCl3 PF5 ICl3 IF5 SnCl2 NO3BrF3 XeF2 BF3 9. Below are 6 possible Lewis structure. Not all the structures below are reasonable Lewis structures. In the box to the right of each Lewis structure, write the letter(s) corresponding to all the following that apply. (a) Atom(s) with less than an octet (b) Non-zero formal charge(s) 4 (c) Disallowed Expanded octet molecule #Ve- Lewis Structure (with Formal Charge on each atom) SO2 or NO2- CCl4 CH2Cl2 ClO3- NH2- CH2O 5 Steric # e--geometry Shape Dipole? PF5 H3PO4 C2N2 N3- O3 NOCl NH4+ 6 Valence Electron Pairs Steric Number = #of atoms + # of lone pairs attached to the center atom 2 3 Electron Pair Geometry Number of Bond Pairs (Number Number of Lone Pairs Molecular Shape Examples of atoms bonded to the center atom) Linear(<180°) Trigonal Planar (<120°) Trigonal Planar Tetrahedral (<109.5°) Tetrahedral 2 3 0 0 Linear Trigonal Planar BeCl2 BF3 2 4 1 0 Bent Tetrahedral SO2 CH4, BF4 - 3 1 NH3, PF3 2 5 2 0 H2O, SCl2 PF5 4 1 3 2 T-shaped ClF3 2 3 Linear XeF2 6 6 Tetrahedral Trigonal Bipyramidal Trigonal Bipyramidal Trigonal Bipyramidal Trigonal Bipyramidal Octahedral Octahedral Trigonal Pyramidal Bent Trigonal Bipyramidal Seesaw 6 5 0 1 SF6 ClF5 6 Octahedral 4 2 Octahedral Square Pyramidal Square Planar 3 4 4 4 5 5 5 5 7 SF4 XeF4