AP Chemistry

Chapters 9

Vocab (Ch 9) VSEPR- Valence Shell e- Pair Repulsion bonding pair non bonding pair – lone pair of electrons electron domain – regions around a central atom

where e- are likely to be found. molecular geometry- the arrangement of atoms in

space electron domain geometry- the arrangement of e-

domains about the central atom of a molecule The Molecular geometry is a derivation of the

Electron-Domain geometry See Table 9.2 (page 309)

Electron Domains

The e- in a multiple bond constitute a single e- domain.

# of e- domains = (# of atoms bonded to the central

atom) + (# of non bonding pairs on the central atom)

Page 306

Molecular Shapes Website

See VSEPR table handout for molecular shapes

http://www.molecules.org/VSEPR_table.html

See B& L page 307-309

Effect of Non bonding e- and multiple bonds on Bond Angles

e- domains for non-bonding e- pairs exert greater repulsive forces on adjacent e- domains and thus tend to compress the bond angles

e- domains for multiple bonds exert a greater repulsive force on adjacent e- domains than do single bonds.

Page 310

Molecules with Expanded Valence Shells These shapes generally contain axial and

equatorial positions See B&L pg. 312 Variations of the trigonal bipyramidal shape

show lone electron pairs in the equatorial position

Variations of the octahedral shape show lone electron pairs in the axial positions

Page 311

Molecules With More Than One Central Atom

You can use the VSEPR theory for molecules with more than one central atom, such as, CH3NH2.

Pages 313-314

Bond Polarity Bond Polarity is a measure of how equally the

e- in a bond are shared between the 2 atoms of the bond.

Polarity is used when talking about covalently bonded molecules.

If the molecule has only 2 different atoms, such as, HF or CCl4 you can calculate the electronegativity difference and determine the type of covalent bond (polar or non-polar).

Polarity and Bond Type

Electronegativity Difference

Bonding Type

<0.5 Non-polar covalent

0.5 – 1.9 Polar covalent

> 2.0 ionic

Dipole Moment Dipole Moment – the measure of the

amount of charge separation in the molecule.

For a molecule with more than 2 atoms, the dipole moment depends on both the polarities of the individual bonds and the geometry of the molecule.

The overall dipole moment of a polyatomic molecule is the sum of its bond dipoles.

See B&L page 315 figure 9.9

Dipole Moment For each bond in the molecule, consider

the bond dipole (the dipole moment due only to the 2 atoms in that bond)

The dipole “arrow” should point toward the more electronegative atom in the bond

The overall dipole moment of a polyatomic molecule is the sum of its bond dipoles. (Consider the magnitude and direction of the bond dipoles)

Different Theories VSEPR Theory (using Lewis Dot Structures) Valence Bond Theory (using hybridization) Molecular Orbital Theory (shows allowed states

for e- in molecules)

Go to the following web-site for a compare and contrasting of the 3 different theories

http://www.chem.ufl.edu/~chm2040/Notes/Chapter_12/theory.html

sp Hybrid Orbitals

See section 9.5 pages 318-320

sp2 and sp3 Hybrid Orbitals

See section 9.5 pages 320-322

d Orbital Hybridization

See section 9.5 pages 322

Multiple Bonds and Hybridization

See section 9.6 pages 324-326

Delocalized π Bonding

See section 9.6 pages 327-330

Sigma Bonds ( σ )

Sigma bonds occur when the e- density is concentrated between the 2 nuclei.

These are single covalent bonds. Sigma bonds can form from the

overlap of an s orbital with another s orbital, an s orbital with a p orbital, or a p orbital with a p orbital.

Pi Bonds ( π )

Overlap of 2 “p” orbitals oriented perpendicularly to the inter-nuclear axis

This overlap results in the sharing of electrons.

The shared electron pair of a pi bond occupies the space above and below the line that represents where the two atoms are joined together.

Hybridization

An atom in a molecule may adopt a different set of atomic orbitals (called hybrid orbitals) than those it has in the free state.

See B&L pages 319-322 for explanation and diagrams of electron promotion

Multiple Bond A multiple bond consists of one sigma

bond and at least one pi bond. A double bond consists of one sigma

bond and one pi bond. A triple bond consists of one sigma

bond and two pi bonds. A pi bond always accompanies a sigma

bond when forming double and triple bonds.

Sigma and Pi Most of the bonding that you have seen

so far has bonding e- that are localized. σ and π bonds are associated with the

2 atoms that form the bond (and NO other atoms)

Delocalized bonding can occur in molecules that have π bonds and more than one resonance structure.

Molecular Orbital Diagram Energy Level Diagram (Molecular Orbital

Diagram) The H2 molecule is the easiest molecule

to plot on the molecular orbital diagram Whenever 2 atomic orbitals overlap,

2 molecular orbitals form (one is a bonding orbital and one is an anti-bonding orbital).

This is not on the AP exam

Paramagnetism

Molecules with one or more unpaired electrons are attracted into a magnetic field

The more unpaired electrons in species, the stronger the force of attraction

This behavior is called paramagnetism

Diamagnetism

Substances with no unpaired electrons are weakly repelled from a magnetic field

This property is called diamagnetism

Diamagnetism is much weaker than paramagnetism

Problems to Try Ch 9 # 5-13, 15, 23, 27, 31, 32, 34, 40, 43, 44 (a

and c), 63 AP Exam Problems to Try

1999 # 82000 # 7 (last section)2002 # 62003 # 82004 # 7 & # 8