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Chapter 9.6

CHEM 1A Chapter Notes - Chapter 9.6: Shap, Linus Pauling, Chemical Polarity


Department
Chemistry
Course Code
CHEM 1A
Professor
Arnold Yuan
Chapter
9.6

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Lecture 4: Shape Matters
PRE-READING NOTES
9.6 pg. 372-376
9.6 Electronegativity and Bond Polarity
shared electrons not always equally shared
polar covalent bond – intermediate in nature between a pure covalent bond and an ionic
bond
electron density greater on the more electronegative atom
electronegativity – ability of an atom to attract electrons to itself in a chemical bond
- Linus Pauling
- increases across a period in the periodic table
- decreases down a column in the periodic table
electronegativity inversely related to atomic size – larger atom, attracts electrons less
greater electronegativity difference = more polar bond
nonpolar– elements with identical electronegativity share atoms equally(0-0.4 ΔEN)
ionic – metal and nonmetal – large electronegativity difference (2.0+ ΔEN)
polar covalent – intermediate electronegativity difference (0.4-2.0 ΔEN)
dipole moment (μ) – anytime there is a separation of positive and negative charge
μ = qr where q is magnitude of charges and r is distance
for an electron and proton at 130 pm apart, μ = 6.3 D (debye = C m)
percent ionic character – ratio of a bond’s actual dipole moment to the dipole moment it
would have if the electron were completely transferred, multiplied by 100%
bond with >50% ionic character are ionic bonds
10.1-10.5 pg. 399-416
10.1 Artificial Sweeteners: Fooled by Molecular Shape
taste of food independent of metabolism
taste receptors active site binds tastant
artificial sugars bind same receptor as sugar
valence shell electron pair repulsion (VSEPR) theory – examine molecular shape
10.2 VSEPR Theory: The Five Basic Shapes
VSEPR theory – electron groups on interior atoms repel one another through coulombic
forces and determine the geometry of the molecule
Electron groups – lone pairs, single bonds, multiple bonds, single electrons
Two electron groups – linear geometry - 180° bond angle
Three electron groups – trigonal planar geometry - 120° bond angle, although different
types of electron groups exert slightly different repulsions (e.g. double vs. single bond)
Four electron groups: tetrahedral geometry – 109.5° angles
Five electron groups: trigonal bipyramidal geometry - 120° and 90°
- equatorial (3 bonds in the plane) and axial (2 bonds from the plane)
Six electron groups: octahedral geometry – 4 in single plane, 2 aove and below the
plane - 90°
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