The Chemical Bond Chapter 9 This is our home as seen from far-out space. Its surface and atmosphere are composed of some free elements as well as ionic and molecular compounds. We look deeper into the nature of compounds in this chapter.
Valence e : The outer s and p e in the atoms of a representative element. These e! 1) are involved in bonding! 2) determine the chemical properties of an element! The Octet Rule: Representative elements form bonds so as to have access to eight outer e.! NOTE: There are some exceptions.!
Representative or Main Group Elements: [NG] ns x np y These are the elements shown in blue in the figure above.!
Chemical Bonds! Intramolecular Forces: Forces that hold atoms together.! Ionic Bonds: Electrostatic forces of attraction between ions (strongest bonds).! Ø Ions are formed by e transfer.!
Chemical Bonds! Intramolecular Forces: Forces that hold atoms together.! Covalent Bonds: The forces of attraction between two atoms that are sharing e.! Ø Typical for molecular substances.! Ø Molecules attracted to each other weakly (Intermolecular forces).! Ø Often found between nonmetal atoms, or nonmetal metalloid.!
I. Follow the recipe. II. No extraneous marks. Lewis (Electron-Dot) Structures! Gilbert N. Lewis (1875 1946)! An American physical chemist known for the discovery of the covalent bond and his concept of electron pairs.!
Resonance Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule. The resulting electron structure of the molecule is given by the average (or superposition) of these resonance structures.! Note: Resonance is NOT a dynamic process. That is, the e that form a double or triple bond DO NOT move from place to place amongst the outer atoms.! Example:! Draw the complete Lewis (electron-dot) structure for the carbonate polyatomic ion.! CO 3 2
Resonance Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule. The resulting electron structure of the molecule is given by the average (or superposition) of these resonance structures.! Note: Resonance is NOT a dynamic process. That is, the e that form a double or triple bond DO NOT move from place to place amongst the outer atoms.! Example:! Draw the complete Lewis (electron-dot) structure for sulfur dioxide.! SO 2
Disregard the following section:! Section 9-6, pages 294 298: Formal Charge
Electronegativity Electronegativity: The ability of an atom of an element to attract electrons to itself in a covalent bond.!
Electronegativity Electronegativity: The ability of an atom of an element to attract electrons to itself in a covalent bond.! Arrows indicate the directions of increasing electronegativity, excluding noble gases.
Polar Covalent Bond Ø! Ø! Ø! Ø!
Polar Covalent Bond Note: Chemists use one or the other, or both.! For the HF molecule:!
Polar Covalent Bond Note: We DO NOT draw Lewis (electron-dot) structures for ionic compounds. In a practical sense, it s pointless. When you see them drawn in your e-book, ignore them.!
The electronegativities of the atoms joined by the bond are the same. The bonding e are shared equally there is no separation of charge.! Example: O 2! Nonpolar Covalent Bond
Problem 53, page 315! Indicate the polarity for each of the following bonds.! (a) N H! (b) Li H! (c) F O! (d) O Cl! (e) S Se!
Geometries of Simple Molecules:! VSEPR Theory! V.S.E.P.R.! Main Idea:!!!!!
VSEPR Example 1: CO 2! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
Using the AB x E y Classification System Ø Ø Ø A represents the central atom.! x and y are integers (1, 2, 3, ).! B x represents x atoms bonded to the central atom, A.! Note: B x is used to represent all atoms bonded to the central atom, even if they are atoms of different elements. The identity of the element is irrelevant for geometric classification purposes.! Ø E y represents y lone pairs of e on the central atom, A.!
VSEPR Example 2: HCN! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
VSEPR Example 3: BF 3! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
VSEPR Example 4: SO 2! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
VSEPR Example 5: CCl 4! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
VSEPR Example 6: NH 3! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
VSEPR Example 7: H 2 O! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
VSEPR Example 8: NO 3! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
Geometries of Simple Molecules:! VSEPR Theory! A few things to keep in mind:! Ø Double and triple bonds are treated the same as single bonds when determining the geometry of a molecule.! Ø Molecular geometry is described by the bonded atoms and DOES NOT include the unshared pairs of electrons.!
VSEPR Example 9: CH 4! 1. Draw the Lewis (electron-dot) structure.! 2. What arrangement of the e pairs allows them to be as far apart as possible to minimize the repulsions?! Molecular Geometry:! Bond Angle:! Classification:!
For each of the following, draw the Lewis (electron-dot) structure and determine the molecular geometry.! 1. PO 3 3! 2. O 3! 3. CS 2! 4. Cl 2 O! 5. ClO 4!
Problem 65, page 316! When a molecule or ion has more than one central atom, the geometry is determined at each central atom site. What are the geometries around the two central atoms in each of the following?! (a) NCl 2 NO (2 central N atoms)! (b) CH 3 OCl (central C and O atoms)!
I. Nonpolar Molecules! Polarity of Molecules Ø In a molecule, the combined effects of the individual bond dipoles are known collectively as the net dipole moment.! Ø If the geometry of the molecule is such that equal dipoles cancel, the net dipole moment is zero, which means that the molecule is nonpolar.!
Molecular Polarity Example 1: Cl 2!
Polarity of Molecules II. Polar Molecules! Ø! Ø!
Polarity of Molecules 1 The molecule has a symmetrical geometry (i.e., linear, trigonal planar, or tetrahedral), but the terminal atoms (ligands) are not all the same.! Examples: HCN and CHCl 3!
Polarity of Molecules 2 The geometry of the molecule is V-shaped or trigonal pyramid. The geometry leads to dipole moments that do not cancel.! Example: H 2 O!
Molecular Polarity Example 2: HCl! Let ε electronegativity. ε Cl > ε H Cl will be δ and H will be δ +. HCl is a polar molecule.
Molecular Polarity Example 3: SO 3! Let ε electronegativity. ε O > ε S each O will be δ, and S will be δ +. The bond polarities are arranged symmetrically, and cancel. SO 3 is a nonpolar molecule.
Molecular Polarity Example 4: CH 4! Let ε electronegativity. ε C > ε H C will be δ, and each H will be δ +. Again, the bond polarities are arranged symmetrically, and cancel. CH 4 is a nonpolar molecule.
Molecular Polarity Example 5: H 2 S! Let ε electronegativity. ε S > ε H S will be δ, and each H will be δ +. The polar bonds result in a net dipole moment (analogous to the water molecule). H 2 S is a polar molecule.
Problem 67, page 316! Discuss the polarities of the following molecules.! (a) SF 2! (b) CS 2! (c) CCl 2 F 2 (C is the central atom.)! (d) NOCl (N is the central atom.)! (e) Cl 2 O!
Problem 72, page 316! Which molecule has a greater dipole moment, CHCl 3 or CHF 3? Explain using principles of molecular structure.!
Directions: Using principles of chemical bonding and molecular geometry, explain each of the following observations. Lewis structures (electron-dot diagrams) and sketches of molecules may be helpful as part of your explanations. For each observation, your answer must include references to both substances.!
1. The XeO 3 molecule is polar.!
2. The SeO 2 molecule is polar, whereas the SeO 3 molecule is not.!
3. The atoms in a C 2 H 4 molecule are located in a single plane, whereas those in a C 2 H 6 molecule are not.!
H H O H H C α C β C γ C σ H H O ε H H