Chemistry Week 9

Chemistry Week 9 - (2) Molecular Geometry TEXTBOOK READING:...

Info iconThis preview shows pages 1–3. Sign up to view the full content.

View Full Document Right Arrow Icon
(2) Molecular Geometry TEXTBOOK READING : BLB-10 , Chapters 8.5-7; 9.1-3, 5, pp. 317-328; 344-360, 361-367. Practice PROBLEMS: (Ch 8) 34, 46, 50, 54, 56, 81, 86; (Ch 9) 21, 22, 24, 26, 36 (a) Lewis Structures allow chemists to write molecular structures in a way that identifies areas of chemical reactivity and electronic behavior. They are useful for main-group covalent molecules and solids, but are not viable for metallic systems. Such structures can also give predictions of bond distances and bond angles. Rules for Writing Lewis Structures : 1. Create the skeleton showing how the atoms will be connected to give the final molecular structure. As a rule of thumb, the most electronegative elements are found on the extremities of molecules, the more electropositive elements are found on the interior. Thus, F, Cl, O are typically on the outer regions, B, C, N, Si, P, etc. are in the interior. H (with just one valence AO) is also found on the exterior. 2. Count the total number of valence electron pairs , taking into account the overall net charge of the molecule. The number of valence electrons contributed by an atom corresponds to its group number, e.g., B has 3 electrons, C has 4 electrons, N has 5 electrons, O has 6 electrons, etc. The number of pairs is the total number of valence electrons divided by two. 3. Draw a single covalent bond between each pair of atoms that are connected; this utilizes one electron pair for each line drawn. 4. Complete an octet of electrons around the “terminal atoms” (those at the extremities of the molecule) by placing lone pairs (nonbonding orbitals assigned to a single atom); utilize multiple bonds if the inner atoms do not establish an octet of electrons by shifting lone pairs or using the remaining pair of electrons; if there are several possible Lewis structures, then see if resonance can be established. 5. Calculate formal charges at each element: Formal Charge = # valence e (neutral atom) # valence e at atom in Lewis structure. To count the # of valence electrons around the element in the Lewis structure, a lone pair counts as two ; a bond pair counts as one (the electrons in a bond pair are equally shared by the two atoms). NOTE: the sum of all formal charges in the Lewis structure equals the total charge on the molecule. Examples: (5) The isocyanate ion, OCN , is an important ion in polymer chemistry. What is its Lewis structure? The number of valence electrons is (6 + 4 + 5 + 1) = 16 valence electrons or 8 electron pairs. The following Lewis structures are possible, each of which satisfies the octet rule at every atom: OCN +1 0 0 0 –2 –1 00 –1 We can calculate the formal charges for each Lewis structure:
Background image of page 1

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full DocumentRight Arrow Icon
(Left) O: 6 5 = +1 (3 bond pairs + 1 lone pair) C : 4 4 = 0 (4 bond pairs) N : 5 7 = 2 (1 bond pair + 3 lone pairs) (Middle) O: 6 6 = 0 (2 bond pairs + 2 lone pairs) C : 4 4 = 0 (4 bond pairs) N : 5 6 = 1 (2 bond pairs + 2 lone pairs) (Right) O: 6 7 = 1 (1 bond pair + 3 lone pairs) C : 4 4 = 0 (4 bond pairs) N : 5 5 = 0 (3 bond pairs + 1 lone pair) Although all three Lewis structures satisfy the octet rule, the one on the left gives a formal charge of +1 to oxygen.
Background image of page 2
Image of page 3
This is the end of the preview. Sign up to access the rest of the document.

Page1 / 13

Chemistry Week 9 - (2) Molecular Geometry TEXTBOOK READING:...

This preview shows document pages 1 - 3. Sign up to view the full document.

View Full Document Right Arrow Icon
Ask a homework question - tutors are online