Chem 110 Basics of Chemical Bonding

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Information about Chem 110 Basics of Chemical Bonding

Published on November 27, 2008

Author: cpesison


Slide 1: Chapter 03 Basics of Chemical Bonding Slide 2: Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that participate in chemical bonding. Slide 3: Lewis dot Symbol consists of the symbol of an element and one dot for each valence electron in an atom of the element. Slide 4: 1s22s1 1s22s22p5 1s2 1s22s22p6 [He] [Ne] Ionic Bond is the electrostatic force that holds ions together in an ionic compound. There is a complete transfer of electrons. Slide 6: A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Lewis Structure is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between atoms, and lone pairs are shown as pairs of dots Lone Pairs are pairs of valence electrons that are not involved in covalent bond formation Why should two atoms share electrons? Lewis structure of F2 Slide 7: Lewis structure of water If two atoms share two pairs of electrons, the covalent bond is called a double bond A triple bond arises when two atoms share three pairs of electrons. + + or or Slide 8: Bond length is the distance between the nuclei of two covalently bonded atoms in a molecule Some common bond lengths Triple bond length < Double Bond Length < Single Bond Length Slide 9: Comparison of Some General Properties of an Ionic Compound and a Covalent Compound Slide 10: When a bond is made up of equally shared electron pairs we call it a nonpolar covalent bond. Slide 11: electron rich region electron poor region e- rich e- poor d+ d- Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms Slide 12: Electron Affinity - measurable, Cl is highest Electronegativity - relative, F is highest Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Slide 13: The Electronegativities of Common Elements Slide 15: Classification of bonds by difference in electronegativity Nonpolar Covalent equal sharing of e- Difference Bond Type 0 Covalent  2 Ionic 0 < and <2 Polar Covalent Slide 16: Comparison of Bonds Slide 18: Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 Nonpolar Covalent Practice on Differences in Electronegativity Slide 33: Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. Complete an octet for all atoms except hydrogen If structure contains too many electrons, form double and triple bonds on central atom as needed. Writing Lewis Structures Slide 34: Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Write the Lewis structure of nitrogen trifluoride (NF3). Slide 35: Write the Lewis structure of the carbonate ion (CO32-). Step 1 – C is less electronegative than O, put C in center Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e- Slide 41: Two possible skeletal structures of formaldehyde (CH2O) The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. Slide 42: formal charge on C = 4 - 2 - ½ x 6 = -1 formal charge on O = 6 - 2 - ½ x 6 = +1 -1 +1 Slide 43: formal charge on C = 4 - 0 - ½ x 8 = 0 formal charge on O = 6 - 4 - ½ x 4 = 0 0 0 Slide 44: For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. Lewis structures with large formal charges are less plausible than those with small formal charges. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH2O? Formal Charges can predict the more stable and thus more plausible Lewis structure Slide 45: The Incomplete Octet BeH2 BF3 Exceptions to the Octet Rule Slide 46: Odd-Electron Molecules NO The Expanded Octet (central atom with principal quantum number n > 2) SF6 Slide 58: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. AB2 2 0 Valence shell electron pair repulsion (VSEPR) model Slide 60: AB2 2 0 linear linear AB3 3 0 Slide 62: AB2 2 0 linear linear AB4 4 0 Slide 64: AB2 2 0 linear linear AB4 4 0 tetrahedral tetrahedral AB5 5 0 Slide 66: AB2 2 0 linear linear AB4 4 0 tetrahedral tetrahedral AB6 6 0 Slide 69: Comparison of lone pair and bonding pair repulsions Slide 70: AB3 3 0 trigonal planar trigonal planar AB2E 2 1 Slide 71: AB3E 3 1 AB4 4 0 tetrahedral tetrahedral Slide 72: AB4 4 0 tetrahedral tetrahedral AB2E2 2 2 Slide 73: AB5 5 0 trigonal bipyramidal trigonal bipyramidal AB4E 4 1 Slide 74: AB5 5 0 trigonal bipyramidal trigonal bipyramidal AB3E2 3 2 Slide 75: AB5 5 0 trigonal bipyramidal trigonal bipyramidal AB2E3 2 3 Slide 76: AB5E 5 1 Slide 77: AB4E2 4 2 Slide 79: Predicting Molecular Geometry Draw Lewis structure for molecule. Count number of lone pairs on the central atom and number of atoms bonded to the central atom. Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO2 and SF4? AB2E bent AB4E distorted tetrahedron

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