But the same molecular formula can be achieved by eliminating H2 another way.
Form one carbon-carbon double bond...this new species is an alkene.
Eliminate another H2 with a 2nd double bond, and the molecule is C2H2n-2.
But this formula also describes an alkyne with one triple bond.
Both alkenes and alkynes (with however many multiple bonds they may contain) are unsaturated hydrocarbons, e.g., they have been dehydrogenated with respect to their "parent" alkane. They have also suffered a change in hybridization at the point of the multiple bond.
While alkanes were sp3, alkenes are sp2, and alkynes are sp at the multiple bond sites. And their geometry changes accordingly.
In particular, H H H H
\ / \ /
H - C1 H H H C8- H
\ / \ \ /
2C = C3 H - C6- C7 , trans-2-cis-4-octadiene
/ \ / \
H 4C = C5 H (a diene has two double bonds
/ \ a triene three, etc. "Monoene"
H H isn't used; it's just "-ene")
has carbons 1 through 6 in a single plane because of the flat centers at carbons 2 through 5 due to their trigonal planar hybridization. Since the sigma bond between 5 and 6 is a (relatively) free rotor, the tail of the molecule can adopt many conformations, but its head is rigid.The nomenclature has added two new terms to our chemical vocabulary, cis and trans, which indicate that the chain or substituents continue on the "same side" or the "opposite side" of the double bond beginning at the given numbered carbon. This is necessary since, due to that planar rigidity, one cannot swap sides, and the molecule is locked into either the cis or the trans configuration. It should be obvious that cis and trans isomers have different proporties. For example:
Cl Cl
\ /
C = C cis-1,2-dichloroethene has a dipole moment due to the
/ \ electronegativity differences
H H
but
Cl H
\ /
C = C trans-1,2-dichloroethene has none by vector cancellation
/ \
H Cl
The pi bonds (one in a double bond and two in a triple) are exposed on the flanks of the sigma; so they are much more reactive. For example, while one can brominate an alkane by first dissociating the bromine diatom with ultraviolet light to make a highly reactive bromine (radical: species with an unpaired electron) atom, bromination of alkenes can occur in the dark. Furthermore, instead of a substitution of a bromine for a hydrogen as would occur in alkanes, diatomic bromine adds straight across the double bond. In fact, this reaction is characteristic of alkenes and is often used as a test for the presence of double bonds in straight- and branched-chain hydrocarbons.H Br-Br H Br Br \ / \ / C = C -----> H - C - C - H 1,2-dibromoethane / \ / \ H H H HHowever, when (even n) alkenes with alternating double bonds condense with themselves to produce arenes, the most famous of which is benzene, this halogen addition chemistry is lost for a very interesting reason:
H H H H
\ / \ /
C - C C = C ortho difluorobenzene
// \\ / \
H - C C - H or H - C C - H (where ortho means
\ / \\ // on adjacent carbons. Para
C = C C - C means on opposite carbons,
/ \ / \ and meta on carbons
F F F F one C removed.)
seems to suggest that there are two different species, one of which has a shorter, stronger bond on the carbons to which the fluorines are attached and one of which doesn't. But this isn't true.Instead, all the bonds in benzene are neither of "order 1" (single) nor "order 2" (double) but rather have a bond order of 1.5. In other words, those two pictures we can draw of that (disubstituted) benzene are a hoax. The 3 pi electrons can lower their wavelengths considerably by spreading out around the entire ring! And so they do.
Chemists may speak of those two pictures "resonating" in the sense that the molecule spends half its time in one configuration and half in the other, but this is truly unfortunate parlance. Instead benzene spends all of its time with all of its carbon-carbon bonds at 1.5 order. The only real advantage to being able to come up with "resonance structures" like these is that the more such not-unreasonable structures that can be written, the greater the relaxation freedom of the pi electrons and the more stable the compound.
This analysis isn't restricted to arenes but works for complex Lewis structures as well. Take, for example, the carbonate ion:
_.. _..
: O : : O : : O :
\ .. \\ .._ \ .._
C = O or C - O : or C - O :
/ ¨ / ¨ // ¨
: O : : O : : O :
- ¨ - ¨
implies an extra degree of stability since the (single) pi electron can spread itself out over all three CO bonds. Clearly the bond order for those is 1.33 as the proper average over the "resonance structures."
The extra stability inherent in the alternating double bonds of arenes would be lost if bromine addition occurred as it does in alkenes. But the stability is so strong that arenes refuse to add halogens, preferring instead to permit the substitution of one of its hydrogens for a single halogen atom, preserving all the double bonds and thus the "resonance structures."Arenes have other interesting substitution reactions as well. For example, when attacked by nitric acid, they will trade a hydrogen for a nitroso group if the powerful dessicant, conc. sulfuric acid, is around to tie up the water also produced.