sternman318
Hazard to Others
Posts: 121
Registered: 21-4-2011
Member Is Offline
Mood: No Mood
|
|
Understanding Bonding/Antibonding
I am beginning my Organic Chemistry course, and right now its just covering basics, especially bonding. One thing that confuses me is molecular
orbitals, more specifically bonding vs antibonding. I understand that bonding orbitals occur when the bonding electron's wave functions
'constrcutively add'(same phase) together, resulting in a low energy, stable bond. However if they add together 'destructively' ( out of phase), an
antibonding orbital forms, which is higher in energy and less stable. I have a feeling those are bare-bones, simplistic explanations for what is
really going on.
My question is, when do these orbitals occur? Can MO's ( molecular orbitals) alternate between bonding and antibonding? Do molecules change between
them based on their energy level ( i.e. excited molecule will form an antibonding MO)?
I understand the basic concept of them, just not the application of them
|
|
|
DDTea
National Hazard
Posts: 940
Registered: 25-2-2003
Location: Freedomland
Member Is Offline
Mood: Degenerate
|
|
A simple rule to remember is that the number of atomic orbitals "in" MUST equal the number of molecular orbitals "out." So suppose two carbon s
orbitals form a sigma bond between one another (i.e., just a single bond). TWO MO's must come out: a higher-energy, sigma bonding orbital and a
lower-energy, sigma antibonding. The two electrons that formed the sigma bond between the two carbon atoms will populate the lower-energy, sigma
bonding orbital (because it's vacant). You're cool with that so far though.
These orbitals occur any time that a bond is formed between the orbitals of two (or more) atoms, and MO diagrams can get complicated very quickly.
They can be simplified by considering molecular symmetry and by focusing on only the frontier orbitals (Highest Occupied Molecular Orbital and Lowest
Unoccupied Molecular Orbital). The HOMO and the LUMO will become your friends as you study O. chem further; entire books can and have been written
about them (see, for example, "Frontier Orbitals in Organic Chemical Reactions" or its updated version, "Molecular Orbitals and Organic Reactions").
One important example of HOMO/LUMO transitions is the pi-->pi* electronic transition observed in UV/Vis spectroscopy of conjugated molecules. An
electron absorbs a quantum of light equivalent to the energy difference between the HOMO and LUMO and is excited to the higher energy MO. From the
wavelength of light absorbed, the energy difference between the pi and pi* orbitals can be calculated. Based on this energy difference, it is
possible to, for example, calculate the length of a conjugated chain. Although several other electronic transitions occur in UV/Vis spectroscopy of
organic molecules, the pi-->pi* transitions are the most obvious and distinguishable peaks.
MO's are a lot like atomic orbitals. For example, they obey Hund's rule as they are populated with electrons. Electrons will preferentially fill
lower energy, bonding MO's before moving up to anti-bonding MO's. If an electron finds its way into a higher energy, anti-bonding MO while a
lower-energy, bonding MO is unfilled, the electron will emit energy until it can "drop" to the bonding MO. Similarly, if a high-energy X-Ray knocks
off an inner-shell electron from an atom, electrons from higher shells will drop to fill it; this gives rise to X-Ray Fluorescence spectra.
In an introductory course, the main application of MO's is spotting regions of high and low electron density. This is important for predicting the
reactivity of molecules under different conditions.
[Edited on 9-12-11 by DDTea]
"In the end the proud scientist or philosopher who cannot be bothered to make his thought accessible has no choice but to retire to the heights in
which dwell the Great Misunderstood and the Great Ignored, there to rail in Olympic superiority at the folly of mankind." - Reginald Kapp.
|
|
|
Magpie
lab constructor
Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline
Mood: Chemistry: the subtle science.
|
|
Does an empty orbital have any other properties other than being a defined region in space?
The single most important condition for a successful synthesis is good mixing - Nicodem
|
|
|
fledarmus
Hazard to Others
Posts: 187
Registered: 23-6-2011
Member Is Offline
Mood: No Mood
|
|
It also has a defined wave form and a defined number of nodes. These can be important in considering the mechanism of reactions. The shape and
wavefunction of an empty orbital may allow it to interact with filled orbitals of other molecules or other atoms on the same molecule; in the first
case defining the specificity of the reaction, and in the second leading to things like back-bonding and resonance stabilization.
|
|
|
sternman318
Hazard to Others
Posts: 121
Registered: 21-4-2011
Member Is Offline
Mood: No Mood
|
|
Thank you!! That helped a lot in clearing it up!
|
|
|
Magpie
lab constructor
Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline
Mood: Chemistry: the subtle science.
|
|
How can an orbital have a defined wave form and nodes when there are no electrons present? Is it like an empty room that affects its neighbors even
though no one is present?
The single most important condition for a successful synthesis is good mixing - Nicodem
|
|
|
watson.fawkes
International Hazard
Posts: 2793
Registered: 16-8-2008
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by Magpie  | | How can an orbital have a defined wave form and nodes when there are no electrons present? Is it like an empty room that affects its neighbors even
though no one is present? |
To continue your analogy, an empty orbital is like a guest room. Electrons can
stay there temporarily, but they can't stay indefinitely. It's as if they have to pay rent for their high-class accommodations (i.e. higher energy
state) and they always eventually run out of money. In the model of the Schrodinger equation, an orbital corresponds to an eigenvector for some
Hamiltonian. The corresponding eigenvalue is the total energy associated with the eigenvector. The eigenvector also has other quantum numbers
associated with it, involving both geometry and basic physical quantities such as energy or angular momentum. That's the mathematical picture.
Here are two example of how these "empty rooms" manifest. The first is electronic excitation states. The energy difference between the HOMO, highest
occupied molecular orbital, and any UMO, unoccupied molecular orbital, gives the strongest emission / absorption spectrum of a molecule. Wavelengths
of the spectrum correspond to energy differences between HOMO and the UMO.
The other is bonding possibilities. When a bonding event happens, electrons, being conserved, have to go somewhere, even if only for an attosecond or
less. They need guest rooms, basically, and a guest room is cheaper than being evicted (ionization). I don't know that the theory is fully developed,
but I'd have to believe that, say, ortho-, para-, meta-directing reaction ratios have a lot to do with the energies of the LUMO's (lowest UMO)
associated with the transient species. The greater the difference of energies of the LUMO over the HOMO, the less likely the reaction is to occur.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
