CHM136H1 Lecture Notes - Lecture 4: Chemical Polarity, Valence Electron, Lone Pair

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1/14/2019 Lec 4 - Google Docs 1/2
Polar Covalent Bonds (Lecture 4)
Polar covalent bonds → electrons are not always shared equally in bonds, due to a difference
in electronegativity of the atoms (leads to partial charges )
- IMPORTANT because most chemical reactions would not occur without EN differences
between atoms in bonds
Electronegativity → the intrinsic ability of an atom to attract shared electrons in a covalent
bond (increases to the left and up the periodic table)
- ΔEN < 0.5 → symmetrical covalent bond
- 0.5 < ΔEN < 2 → polar covalent bond
- ΔEN > 2 → ionic bond
Electrostatic Potential Maps → show calculated charge distributions
- Red for electron-rich regions; blue for electron-poor regions
- Inductive effect → shifting of electrons in a bond in response to EN of nearby atoms
Dipole Moment (𝛍) → measure of charge separation in a molecule (ie. net molecular polarity)
- Whole molecules are often polar due to the vector sum of individual bond polarities and
contributions from lone-pair electrons
- ** in symmetrical molecules, the dipole moments of each bond cancel out (no net dipole
moment) **
Formal Charge comparing the bonding of an atom in a molecule to its valence electron
structure (as in the periodic table); ie. “electron bookkeeping
- Eg. if an atom has one more electron in the molecule that it should have, it as a “-1
charge; “+1charge if it has one less electron
- Formal charge = (number of valence electrons in free atom) - (number of bonding
electrons/2) - (number of nonbonding electrons, ie. lone pairs)
Resonance many molecules have structures that cannot be adequately shown with a single
representation; thus, we consider structures that contribute to the final structure but differ in the
position of pi bonds or lone pairs
- Such a structure has delocalized electrons and is represented by two resonance forms
- Notice how the resonance forms are connected by a double-headed arrow (↔)
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