Cyclopropane
replied to your post “n-Alkanes”
How come the three carbons can’t make a triangle?
They could, but it wouldn’t be an isomer of propane.
You’re jumping ahead of me a little, but say hello to cyclopropane, which probably deserves a post of its own.
Cyclopropane is the simplest example of a class of hydrocarbons called cycloalkanes. They’re also referred to as naphthenes, but this term only seems to have traction in the petroleum industry.
The difference between cyclopropane and propane lies in the number of hydrogen atoms. Propane is C3H8, which we could also write out as CH3-CH2-CH3. Notice how the end carbons have three hydrogen atoms apiece, but the middle one only has two. This is because carbon has to use all four of its bonds. The ones on the end only use one bond to form the carbon chain, so that means they have room for three hydrogen atoms. The one in the middle is bonded to both of the end carbons, so it has two bonds left over for hydrogen atoms.
But in cyclopropane, the carbon atoms don’t form a chain, they form a ring, so there are no ends. Each carbon atom is bonded to two of its closest neighbor carbons, leaving two bonds open for hydrogen atoms. Consequently, the formula for cyclopropane is
C3H6.
In other words, cyclopropane has two fewer hydrogen atoms than regular propane, so the two molecules cannot be isomers of one another.
All alkanes have the chemical formula
CnH2n+2,
while all cycloalkanes have the formula
CnH2n.
So while the two classes are very similar in many respects, they are different in a very fundamental way.
Having said all of that, there is an important obstacle to forming a triangle of carbon atoms like this. As I said in the original post, carbon atoms like to keep their valence electrons spaced apart evenly, so the bonds they form tend to be at 109.5 degree angles from one another. But in a triangular shape, that doesn’t work at all. Since the carbon atoms are essentially identical, they form an equilateral triangle together, with equal angles of sixty degrees in each corner.
This would be something of a strain on the molecule, as best illustrated by this plastic model kit.
A decent model kit is designed to bend a little in order to show where strain might exist in a structure. If chemical bonds were inflexible and rigid, we might expect that cyclopropane would be extremely unstable, and probably difficult to form in the first place.
But cyclopropane isn’t nearly as unstable as you might expect. Don’t get me wrong, it’s pretty flammable, but so is regular propane, and its bonds aren’t under any strain at all. As it turns out, the real world chemical bonds of cyclopropane bow out, just like the plastic connectors in this model. It’s not an ideal situation for a molecule, but under the right conditions, the atoms will go for it, and once they do, they can stay that way.
bee-squared-official
