VSEPR – molecular shape • Molecular Polarity • Bond ...
Transcript of VSEPR – molecular shape • Molecular Polarity • Bond ...
Lecture 29 Chapter 9 Sections 4-5
• VSEPR – molecular shape
• Molecular Polarity
• Bond Lengths & Energies
Molecular Shape• The molecular shape describes how the ligands (not the electron
groups) are arranged in space.• This group of molecules all have different shapes.• Electron groups all feel each other, but we can “see” only atoms
Example – H3O+
• What is the steric number?
• Electron group shape?
• How many ligands?
• Molecular shape?
Four
Three
Trigonal pyramid
Tetrahedral
Lone pairs are ‘larger’ than bonding pairs
Experiments show that sulfur tetrafluoride has bond angles of 86.8° and 101.5 °.
This is because the lone pair wants to take up more space – it distorts the positions of the other atoms.
Dipole moments
• Recall that polar bonds result from electronegativitydifferences
• Polar bonds can result in polar molecules, depending on the molecule’s geometry
• A polar molecule will align itself in an electric field• Polar things also interact with each other
– “like dissolves like”• The extent to which the molecules align in a field is
referred to as the dipole moment and has the Greek symbol mu, µ.
rqrr=µ
Bond Lengths and Energies
• Bond length – the nuclear separation distance where the molecule is most stable.1. The smaller the principle quantum numbers of the valence
orbitals, the shorter the bond.2. The higher the bond multiplicity, the shorter the bond.3. The higher the effective nuclear charge of the bonded atoms, the
shorter the bond.4. The larger the electronegativity difference, the shorter the bond.
• Bond energy – the stability of a chemical bond1. Bond strength increases as more electrons are shared between the
atoms2. Bond strength increases as the electronegativity difference (∆χ)
between bonded atoms increases.3. Bond strength decreases as bonds become longer.
CLOSE BOOKS!Which molecule has the shortest bond length?
54321
25%25%25%25% 1. F2
2. Cl2
3. Br2
4. I2