Transcript Chapter 9, Part 1

December 1, 2009
•T H E “ M A K E U P ” L E C T U R E
•T O P I C S
Molecular Polarity (8.7)
Introduction to Bonding Theories (9.1)
Valence Bond Theory (9.2)
Molecular Polarity
 Polarity = uneven distribution of charge
 Bond polarity = Electrons drawn closer to the more
electronegative atom
 Molecular polarity = Molecule as a whole has a net separation
of charge
A polar molecule must have polar bonds
 A molecule is polar if the directions of the polar bonds don’t
cancel/offset eachother

Nonpolar Examples: CH4, CO2
 Polar Examples: H2O, NH3

Nonpolar Examples: CH4, CO2
Polar Examples: H2O, NH3
Why is Polarity Important?
 Polarity dictates many molecular properties

Physical state (solid, liquid, gas)


CO2 (44 g/mol) vs. H2O (18 g/mol)
Solubility
Chapter 9- Chemical Bonding Theories
 Valence Bond Theory: Uses Lewis Structures
 Bonds form using shared electrons between overlapping
orbitals on adjacent atoms.
 Orbitals arrange around central atom to avoid each other.
 Two types of bonds: sigma () and pi ().
 Qualitative, visual- good for many atom systems in ground
state
 Molecular Orbital Theory: Uses MO Diagrams
 Orbitals on atoms “mix” to make molecular orbitals, which go
over 2 or more atoms.
 Two electrons can be in an orbital.
 Quantitative- needed to describe excited states
Sigma () Bonding
 Orbitals on bonding atoms overlap directly between
bonding atoms
Sigma () Bonding
Consider VSEPR Shapes and bonding:
What’s wrong with this picture?
 Atoms bond by having their valence orbitals overlap
Bonding orbitals are not the same shape as atomic orbitals
2pz
2px
2py
2s
Orbitals in CH4
 Electron configurations:
 H = 1s1
 C = 1s22s22p2
Atomic orbitals change
shape when they make
molecules
Hybrid Orbitals