Chapter%2013%20Student%20Notes%20Part%202%20PHW

Chapter%2013%20Student%20Notes%20Part%202%20PHW - Lewis...

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Unformatted text preview: Lewis Structures to Represent Covalent Bonding Lewis structures (also called Lewis dot diagrams) are diagrams diagrams) that show the bonding of a molecule and the lone pairs of electrons that may exist in the molecule. A Lewis structure can be drawn for any covalently-bonded molecule. covalentlyThe Lewis structure is named after Gilbert N. Lewis, who Lewis, introduced it in his 1916 article entitled The Atom and the Molecule [Journal of the American Chemical Society 1916, Vol. 38, p. 762]. 1916, 38, 762]. 9/16/2008 Zumdahl Chapter 13 1 Drawing Lewis Structures 1. Count up the total number of valence electrons available (A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it. ion, 2. Calculate the total number of electrons needed (N) for each atom to have its own noble-gas set of electrons around it (two for hydrogen, eight for the nobleelements from carbon on in the periodic table). 3. Subtract the number in step 1 from the number in step 2. This is the number is of shared (or bonding) electrons present (S = ). 4. Assign bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion. 9/16/2008 Zumdahl Chapter 13 2 1 Drawing Lewis Structures (con't) 5. If any of the electrons earmarked for sharing remain, assign them in pairs them by making some of the bonds double or triple bonds. In some cases, there bonds. may be more than one way to do this. Typically, double or triple bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur. carbon, nitrogen, oxygen, sulfur. 6. Assign the remaining electrons ( octets to all atoms except hydrogen. ) as lone pairs to the atoms, giving 7. Determine the (C) on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion. C group no. - no. of lone pair electrons - (no. of electrons in bonding pairs) When choosing between more than one possible Lewis structure, the one with the is usually preferred. preferred. 9/16/2008 Zumdahl Chapter 13 3 Lewis Structure of Methane (CH4) 1. Count up the total number of valence electrons available (A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it. H H C H H H H C H H 9/16/2008 Zumdahl Chapter 13 4 2 2. Calculate the total number of electrons needed (N) for each atom to have its own noble-gas set of electrons around it (two for noblehydrogen, eight for the elements from carbon on in the periodic table). H H C H H H H C H H 9/16/2008 Zumdahl Chapter 13 5 3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S). (S H H C H H N-A = H H C H H 9/16/2008 Zumdahl Chapter 13 6 3 4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion. H H C H H 2 2 2 2 All bonding electrons assigned 9/16/2008 Zumdahl Chapter 13 7 Lewis Structure of Ethylene (C2H4) 1. Count up the total number of valence electrons available (A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it. H C H H H C H H C H C H 9/16/2008 Zumdahl Chapter 13 8 4 2. Calculate the total number of electrons needed (N) for each atom to have its own noble-gas set of electrons around it (two for noblehydrogen, eight for the elements from carbon on in the periodic table). H C H H H C H H C H C H 9/16/2008 Zumdahl Chapter 13 9 3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S). (S H C H H H C H H C H C H N-A = 9/16/2008 Zumdahl Chapter 13 10 5 4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion. H C H C H H H 2 C 2 C 2 2 H 2 H H 9/16/2008 Zumdahl Chapter 13 11 5. If any of the electrons earmarked for sharing remain, assign them in pairs by making some of the bonds double or triple bonds. In some cases, there may be more than one way to do bonds. this. Typically, double or triple bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur. sulfur. H H 2 2 C 2 C 2 2 H H 9/16/2008 Zumdahl Chapter 13 12 6 Lewis Structure of Carbon Monoxide (CO) 1. Count up the total number of valence electrons available (A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it. C O: 9/16/2008 Zumdahl Chapter 13 13 2. Calculate the total number of electrons needed (N) for each atom to have its own noble-gas set of electrons around it (two for noblehydrogen, eight for the elements from carbon on in the periodic table). 9/16/2008 Zumdahl Chapter 13 14 7 3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S). 9/16/2008 Zumdahl Chapter 13 15 4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion. 2 CO 9/16/2008 Zumdahl Chapter 13 16 8 5. If any of the electrons earmarked for sharing remain, assign them in pairs by making some of the bonds double or triple bonds. In some cases, there may be more than one way to do bonds. this. Typically, double or triple bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur. carbon, nitrogen, oxygen, sulfur. C 9/16/2008 2 O Zumdahl Chapter 13 17 6. Assign the remaining electrons (A - S) as lone pairs to the (A atoms, giving octets to all atoms except hydrogen. (2 lone pairs) C O O O 18 2 C O 2 C C 9/16/2008 Zumdahl Chapter 13 9 7. Determine the formal charge (C) on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion. What is the formal charge on the carbon atom in methane (CH4)? C group no. - no. of lone pair electrons - (no. of electrons in bonding pairs) 9/16/2008 Zumdahl Chapter 13 19 What is the formal charge on each hydrogen atom in methane (CH4)? C group no. - no. of lone pair electrons - (no. of electrons in bonding pairs) As should be the case for the neutral molecule CH4, 9/16/2008 Zumdahl Chapter 13 20 10 7. Determine the formal charge (C) on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion. What is the formal charge on each carbon atom in ethylene (C2H4)? C group no. - no. of lone pair electrons - (no. of electrons in bonding pairs) 9/16/2008 Zumdahl Chapter 13 21 What is the formal charge on each hydrogen atom in ethylene (C2H4)? C group no. - no. of lone pair electrons - (no. of electrons in bonding pairs) As should be the case for the neutral molecule C2H4, 9/16/2008 Zumdahl Chapter 13 22 11 7. Determine the formal charge (C) on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion. What are the formal charges on the atoms in carbon monoxide (CO)? C group no. - no. of lone pair electrons - (no. of electrons in bonding pairs) C O As should be the case for the neutral molecule CO 9/16/2008 Zumdahl Chapter 13 23 Resonance Structures There are cases where more than one equivalent Lewis structure .. .. can be drawn for a molecule: O O O O O O Typical O-O single bond length: 1.49 O- Typical O=O double bond length: 1.21 In O3 the experimental bond lengths are the same: 1.28 The actual structure can be described as a combination of the two Lewis structures, i.e., a .. .. O O O 9/16/2008 O O O Zumdahl Chapter 13 24 12 More Examples of Resonance Structures Carbonate anion, CO32- O C O O O -2 2- O C O -2 2- O C O O -2 2- Thiocyanate anion, NCS- N C S -1 - N C S -1 - N2O (nitrous oxide, laughing gas) N N O N N O N N O 9/16/2008 Zumdahl Chapter 13 25 Favored Resonance Structures When comparing resonance structures for the same molecule, usually those with the contribute more to the overall resonance hybrid. When formal charges are necessary, resonance structures that have and positive charges on the less electronegative elements typically typically contribute more to the overall resonance hybrid. 9/16/2008 Zumdahl Chapter 13 26 13 Breakdown of the Octet Rule Case 1: (free radicals) such as NO: radicals) NO: N O # of valence e- = 5 (N) + 6 (O) = 11 Lewis structure does not satisfy the octet rule but the molecule is stable, although somewhat reactive. Case 2: (mainly Be, B, and Al compounds): Be, -1 B F +1 (A) F B F F F F (B) Even though structure (B) obeys the octet rule, structure (A) is favored (Formal charges are lower & experimental evidence shows that there is no there double bond in BF3 ). 9/16/2008 Zumdahl Chapter 13 27 Breakdown of the Octet Rule (con't.) Case 3: Third and higher period elements can exhibit bonding where an octet on the central atom is exceeded -- can expand up 12 e-! SF6 F F F S F F F SF4 34 total valence e- 8 e- in bonds 24 e- in lone pairs 2 e- left over F F S F F Rule for Lewis structures: 9/16/2008 If electrons remain after satisfying the octet rule, Zumdahl central add lone pairs to the Chapter 13 atom! 28 14 The Shapes of Molecules: The VSEPR Theory VSEPR Valence Shell Electron-Pair Repulsion Theory 9/16/2008 Zumdahl Chapter 13 29 VSEPR Theory: The Basic Idea Electron pairs in the valence shell of an atom repel each other on a spherical surface formed by the underlying core of the atom. The geometry which applies to a particular arrangement is determined by the determined steric number (SN) of the central atom. The favored geometry is the one that SN) minimizes electron - electron repulsions. repulsions. "Steric" means "having to do with space." The steric number of an atom in a Steric" space. molecule can be determined by drawing the Lewis structure of the molecule and adding the number of atoms that are bonded to it and the number of lone pairs that it has. has. 9/16/2008 Zumdahl Chapter 13 30 15 Geometry and Steric Number SN = 2 SN = 3 SN = 4 SN = 5 Linear, 180 180 Trigonal planar, 120 120 Tetrahedral, 109.5 109.5 Trigonal bipyramidal 90 (equatorial - axial) 90 120 (equatorial equatorial) 120 Lone pairs occupy equatorial positions in preference to axial positions. SN = 6 Octahedral, 90 90 SN for Double and Triple bonds count the same as single bonded atoms When lone pairs are present, the situation is more complicated due due to repulsive forces Lone pair vs. lone pair > Lone pair vs. bonding pair > Bonding pair vs. bonding pair 9/16/2008 Zumdahl Chapter 13 31 Examples with no Lone Pairs on the Central Atom 9/16/2008 Zumdahl Chapter 13 32 16 Geometry is 9/16/2008 Zumdahl Chapter 13 33 Geometry is 9/16/2008 Zumdahl Chapter 13 34 17 Geometry is 9/16/2008 Zumdahl Chapter 13 35 Geometry is 9/16/2008 Zumdahl Chapter 13 36 18 Geometry is 9/16/2008 Zumdahl Chapter 13 37 Example: NH3 Geometry is 9/16/2008 Zumdahl Chapter 13 38 19 Example: H2O Geometry is 9/16/2008 Zumdahl Chapter 13 39 9/16/2008 Zumdahl Chapter 13 40 20 Chapter 13 Bonding; General Concepts 13.1 Types of Chemical Bonds 13.2 Electronegativity 13.3 Bond Polarity and Dipole Moments 13.4 Ions: Electron Configurations and Sizes 13.5 Formation of Binary Ionic Compounds 13.6 Partial Ionic Character of Covalent Bonds 13.7 The Covalent Chemical Bond: A Model 13.8 Covalent Bond Energies and Chemical Reactions 13.9 The Localized Electron Bonding Model 13.10 Lewis Structure 13.11 Resonance 13.12 Exceptions to the Octet Rule 13.13 Molecular Structure: The VSEPR Model 9/16/2008 Zumdahl Chapter 13 41 21 ...
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This note was uploaded on 09/04/2009 for the course CHEM 1310 taught by Professor Cox during the Spring '08 term at Georgia Tech.

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