Organic Chemistry I Course Pack

Organic Chemistry I Course Pack - Hanadi SLEIMAN CHEM 212...

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Hanadi SLEIMAN CHEM 212 Chapter 1 - 1 -
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- 2 - I. Lewis Structures 3 How to write Lewis structures 3 II. VSEPR: The Shape of Molecules 5 III. Bond Polarity & MoleculE Polarity 7 IV. Orbital Theory 8 Atomic Orbitals 8 Molecular Orbitals 9 Hybridization 12 Summary 14 Applications 15 V. Acids and Bases 17 Trends 18 Lewis Definition of Acids and Bases 20 VI. Exercises 20 VII. Solutions 22 VIII. Suggested Problems from Solomons, 9 th edition 26
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I. Lewis Structures G.N. Lewis: Theory of Bonding based on the octet rule: atoms of C, N, O, F attain stable configurations when they have 8 electrons in their outer = valence shell. How to write Lewis structures Step 1 Count the number of valence electrons available; do not forget to add 1 e - for each (-) charge, and subtract 1 e - for each (+) charge on the molecule. Examples CF 4 : C = 4 valence e - ; F = 7 valence e - Total valence electrons = - 2 e - 4 e 4174 3 ×+ × = 1 2 ×+ ×+= NO 3 - : N = 5 valence e - ; O = 6 valence e - ; (-) = 1 valence e - Total valence electrons = 51 63 Step 2 Connect the bonded atoms by a line, representing 2 shared electrons. C F F F F N O O O Step 3 Count the number of shared electrons, and subtract this from the total number of valence electrons; this gives the number of electrons to be added to each atom to complete the structure. Examples : CF 4 = 32 – 8 = 24 NO 3 - = 24 – 6 = 18 Step 4 Add electrons in pairs to each atom to complete its octet (add 2 electrons for Hydrogen). This is not always possible for all atoms. C F F F F N O O O In CF 4 all atoms have their octet, whereas in NO 3 - , the Nitrogen atom has only 6 e - . Step 5 If one or more atoms have fewer than 8 electrons, form double or triple bonds to complete their octet. N O O O An electron pair on Oxygen is shared with Nitrogen to form a double bond. - 3 -
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Step 6 Assigning Formal Charges to atoms; by definition: Formal Charge = # (valence e - ) - # (e - assigned to atom) Formal Charge = # (valence e - ) – [# (bonds) + #(unpaired e - )] Examples: - 4 - C F F F F F = 7 - (1+6) = 0 C = 4 - (4 + 0) = 0 N O O O O = 6 - (1+6) = -1 O = 6 - (2+4) = 0 O = 6 - (1+6) = -1 N = 5 - (4+0) = +1 Thus we write: N O O O Direct Application : Write the Lewis structures of NH 4 + , CO 3 2- , C 2 H 4 , CH 3 NO 2 , CO 2 , C 2 H 5 Cl. Solution: N H H H H C O O O CC H H H H C H H N H O O C O O H H H H H Cl
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II. The Shape of Molecules: VSEPR Method The Lewis structure of a molecule gives no information on the shape of this molecule. VSEPR stands for Valence-Shell Electron Pair Repulsion; the VSEPR theory is a model which helps to predict the geometry of molecules; its premise is that electron pairs on an atom will tend to be as far apart from each other as possible. Coordination Number CN = electron pairs on an atom = bonding pairs + lone pairs. Note that double and triple bonds are counted as 1 “bonding pair”. Basic shapes relevant to Organic Molecules: Coordination Number CN Shape Bond Angles CN = 4 Tetrahedral 109.5 ° CN = 3 Trigonal Planar 120 ° CN = 2 Linear 180 ° Examples: In Methane CH 4 , the central Carbon has CN = 4 (4 bonding pairs and 0 lone pairs); thus the molecule has a Tetrahedral geometry.
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This note was uploaded on 03/06/2012 for the course CHEM 302 taught by Professor Sleiman during the Winter '12 term at McGill.

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Organic Chemistry I Course Pack - Hanadi SLEIMAN CHEM 212...

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