AP Chemistry
th
Zumdahl Notes, 9 ed.
A Brief Collection of notes, Chapter 9
Feel free to open these files and annotate as you feel the
need…this is for your success.
Covalent bonding: orbitals: hybrids and the localized
electron model (building upon our knowledge of VSEPR, Lewis structures,
etc.)
• sp3 hybridization: figures on 417-418 for illustration…main thing: note in figure
9.5…energy level of hybrids is in between s and p orbitals!
• Biggest take: atoms hybridize so as to achieve the lowest/minimal energy for
the molecule
• Localized electron model…sounds to me like the hybrid orbitals that we would
predict from VSEPR (different name, perhaps?)
• http://chemwiki.ucdavis.edu/Textbook_Maps/General_Chemistry_Textbook_M
aps/Map%3A_Zumdahl's_%22Chemistry%22/08%3A_Bonding_General_Conce
pts/8.09_The_Localized_Electron_Bonding_Model
more
• sp2 hybridization: same thing, only one s and two p, leaving a p orbital as it was,
which then is involved in a pi bond, perpendicular to the planar arrangement of
the sigma bonded sp2 orbitals
•
Example: ethylene (properly, ethane) see page 420-421 for pictures
• sp hybridization: same thing, only different, this time one s and one p, leaving
two p orbitals for pi bonding…triple bond time, or two double bonds…see
example, 422, for carbon dioxide, 423 for nitrogen
• dsp3 hybridization: again, same thing, only different: one s, three p, one d, for
five hybrid orbitals of equivalent energy, all for sigma bonds around the central
atom…may have lone pairs occupying some of these orbitals
• d2sp3 hybridization: wow…one s, three p, two d, for six equivalent orbitals;
similarly, six locations for electron pairs, bonding or lone; illustrations on 424428
Molecular Orbital Model
• Idea: overlap of atomic orbitals can form either bonding or anti-bonding
molecular orbitals
•
•
Intuitively, yes, the bonding molecular orbital is at a lower energy point (favoring bonded
atoms), and the anti-bonding orbital is at a higher energy point (favoring separated atoms)
Anti-bonding orbitals are labelled with an asterisk
• At its simplest: hydrogen molecule: both of these are referred to as sigma
molecular orbitals:
•
•
Bonding: MO1 = 1sa + 1sb
Anti-bonding: MO2 = 1sa - 1sb
Molecular Orbital Model, continued
• Bond Order
•
•
Bond order = (#bonding electrons -#anti-bonding electrons)/2
Bond order is an indication of bond strength, as it reflects the difference between bonding
and antibonding electrons
•
•
Large bond order number means greater bond strength
Visit:
http://chemwiki.ucdavis.edu/Theoretical_Chemistry/Chemical_Bonding/Pictorial_Molecular
_Orbital_Theory/How_to_Build_Molecular_Orbitals
Bonding in Homonuclear Diatomic Molecules
• Bond order = 0
• Consideration: bonding orbitals are establishing constructive interference, and
antibonding orbitals are establishing destructive
•
•
•
Think of the need for opposite spin designations for the electrons if sigma, same spin if
pi…yes, this is something to remember
If two overlapping orbitals are axial to the two atoms…sigma bond
If two overlapping orbitals are parallel/above/below the two atoms…pi bond
Paramagnetism and diamagnetism
• Paramagnetic: a material attracted to a magnetic field
•
Associated with unpaired electrons
• Diamagnetic: a material repelled from a magnetic field
•
Associated with paired electrons
• If both present, paramagnetism will override, as its effect is stronger
• See figure 9.38, pg. 435…note filling of orbitals, magnetic effect, bond lengths
• Visit:
http://chemwiki.ucdavis.edu/Theoretical_Chemistry/Chemical_Bonding/Pictori
al_Molecular_Orbital_Theory
Bonding in Heteronuclear Diatomic Molecules
• Using molecular orbital diagram, fill in electrons, calculate bond order
•
•
If they are contiguous on the PT, consider drawing as if homonuclear instead
When the elements are further apart (very different), note illustration 9.42 on pg. 438 for
thought: using only orbitals involved in bonding…interesting note is that fluorine’s 2p
orbital is lower in energy than hydrogen’s 1s, showing tendency for fluorine to hold the
electrons closer.
Combining the localized electron and MO
models
• Reality check time: sigma bonds seem to be described reasonably well with
the models, while pi bonds are more delocalized
•
•
This allows us to feel more comfortable with what we are trying to describe
Note figures 9.45-48, describing resonance structures
•
•
We make this work out well
Delocalization explains the behavior and bond lengths we observe in the molecules!
Photoelectron Spectroscopy (PES)
• Technique: high energy photons are directed at a sample, and a detector
measures the energy of ejected electrons, which can be used to calculate the
energy of the electrons in the sample
•
Eelectron = hν – KE
• Realization: the peaks will correspond to the number of electrons present in the
atom (height is directly proportional to #)
•
See figures 9.49-9.50
• Visit:
http://chemwiki.ucdavis.edu/Physical_Chemistry/Spectroscopy/Photoelectron_
Spectroscopy/Photoelectron_Spectroscopy%3A_Theory
• Demo site: http://cbc.arizona.edu/chemt/Flash/photoelectron.html
More sites for help
• https://search.yahoo.com/yhs/search?p=molecular+orbitals+explained&ei=
UTF-8&hspart=mozilla&hsimp=yhs-001
• http://iverson.cm.utexas.edu/courses/310M/Handouts/Handoutsfl05/MOVB
Theory.pdf
• http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/mo.html
• http://www.chem.latech.edu/~upali/chem281/notes/Ch2-MO-Theory.pdf
• http://www.chm.davidson.edu/vce/MolecularOrbitals/index.html
• http://chemguide.co.uk/atoms/properties/atomorbs.html