Chapter 2 Lecture

Chapter 2 Lecture - Polar Covalent Bonds: Electronegativity...

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Polar Covalent Bonds: Electronegativity Covalent bonds can have ionic character These are polar covalent bonds Bonding electrons attracted more strongly by one atom than by the other Electron distribution between atoms is not symmetrical
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Bond Polarity and Electronegativity Electronegativity (EN): intrinsic ability of an atom to attract the shared electrons in a covalent bond Inductive Effect: shifting of sigma bonded electrons in response to nearby electronegative atom Symmetrical Covalent Bonds C – C C – H (non-polar) Polar Covalent Bonds C – O (polar) δ + δ -
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The Periodic Table and Electronegativity
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Bond Polarity and Inductive Effect Nonpolar Covalent Bonds : atoms with similar EN Polar Covalent Bonds : Difference in EN of atoms < 2 Ionic Bonds: Difference in EN > 2 C–H bonds, relatively nonpolar C-O, C-X bonds ( more electronegative elements) are polar Bonding electrons toward electronegative atom C acquires partial positive charge, δ + Electronegative atom acquires partial negative charge, δ - Inductive effect: shifting of electrons in a bond in response to EN of nearby atoms
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Electrostatic Potential Maps Electrostatic potential maps show calculated charge distributions Colors indicate electron- rich (red) and electron- poor (blue) regions Arrows indicate direction of bond polarity
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Polar Covalent Bonds: Dipole Moments Molecules as a whole are often polar from vector summation of individual bond polarities and lone-pair contributions Strongly polar substances soluble in polar solvents like water; nonpolar substances are insoluble in water. Dipole moment ( μ ) - Net molecular polarity, due to difference in summed charges μ - magnitude of charge Q at end of molecular dipole times distance r between charges μ = Q × r, in debyes (D), 1 D = 3.336 × 10 - 30 coulomb meter length of an average covalent bond, the dipole moment would be 1.60 × 10 - 29 C m, or 4.80 D.
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Absence of Dipole Moments In symmetrical molecules, the dipole moments of each bond has one in the opposite direction The effects of the local dipoles cancel each other
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Drawing Lewis Structures Draw molecular skeleton: this will come with practice, remember atom valence Determine number of available valence electrons: add an electron for each negative charge, remove an electron for each positive charge Draw all single covalent bonds and lone pairs: give as many atoms as possible full octets, assigning lone pairs to most electronegative atoms Convert lone pairs to multiple bonds if needed: to satisfy octet rule for as many atoms as possible Assign formal charges to all atoms: the sum of all charges must equal total charge of the molecule
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Nitromethane: CH 3 NO 2 H H H C N O O Valence E = 3(1) + 1(4) + 1(5) + 2(6) = 24 e - H H H C N O O •• •• •• •• •• Step 5……………
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Chapter 2 Lecture - Polar Covalent Bonds: Electronegativity...

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