CHM 238 - Ch 2 McMurray - Chapter 2 Polar Covalent Bonds...

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Chapter 2: Polar Covalent Bonds; Acids and Bases
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Electronegativity Each element has a pull for the electrons that it shares with another atom. This pull is referred to as electronegativity. Electronegativity increases going up and to the right on the Periodic Table.
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Bond Polarity: Differences in the electronegativities of the atoms sharing a pair of electrons implies unequal sharing. The shared electrons will spent a greater proportion of time around the more electronegative atom.
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Bond Polarity and Molecular Polarity: A polar bond has a dipole-moment , which is a vector quantity, has magnitude and direction. μ = δ x d , where δ is the size of the charge at one end of the polar bond and d is the distance separating the two charges. The direction of the dipole-moment is from the positive end to the negative end of the polar bond.
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Bond Polarity: The more electronegative atom becomes more negative. The less electronegative atom becomes more positive. H Cl 2.2 3.2 H Cl δ + δ- δ - indicates a partial charge due to unequal sharing of electrons
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Dipoles: A polar bond is an instance of a dipole: Dipole: two opposite charges separated by a distance. Dipoles are often depicted by a dipole arrow. Starts with a plus (+) and extents to a point pointing to the negative end of the dipole. H Cl 2.2 3.2 H Cl δ + δ- H Cl
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More Examples Dipoles: Cl O O N C N H B O F P S
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Bond Dipoles: Dipole Moments of Some Common Covalent Bonds Bond μ (Debyes) Bond μ (Debyes) C—N 0.22 H—C 0.3 C—O 0.86 H—N 1.31 C—F 1.51 H—O 1.53 C—Cl 1.56 C=O 2.4 C—Br 1.48 C≡N 3.6 C—I 1.29
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Bond Polarity and Molecular Polarity: A molecule with polar bonds has a dipole-moment that is equal to the vector sum of the dipole-moments of the polar bonds in the molecule.
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Molecular Polarity:
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Molecular Polarity & Lone Pairs:
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Formal Charges: Electron “book-keeping” for an atom. If an atom loses an electron, it increases it charge by one. If an atom gains an electron, it deceases it charge by one.
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Formal Charges: If an atom uses an unpaired electron to form a covalent bond, no change in charge. (Why?) If an atom uses a lone pair of electrons to form a covalent bond, it increases it charges by one. (Why?)
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Formal Charge: If an atom accepts an electron pair to form a covalent bond, it decreases it charge by one. (Why?) All of the above changes can be accommodated by the formula: F.C.= # of Valence Electrons - #of Bonds - #of Unshared Electrons
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Formal Charges: H N H H H For N: 5 - 4 - 0 = 1 For H: 1 - 1 - 0 = 0 So the nitrogen has a positive formal charge while the hydrogens are formally neutral.
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Rules for Formal Charges: Formal charges are an atomic property. The sum of the formal charges of each of the atoms in a structure is equal to the charge of the structure.
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Common Charged Atoms in Lewis Structures.
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Charged Atoms in Structures: Learn what is “normal” for the bonding of each element family.
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