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Chemistry 2131:
Organic Chemistry for the Life Sciences (3)

Conformations of Alkanes and Cycloalkanes

• last class we learned how to arrange atoms based on names and how to name different constitutional isomers of alkanes. Now we move on to a more subtle kind of difference between molecules.
• when we discussed orbital hybridization and the definition of sigma bonds (single bonds) we noted that there is free rotation about these bonds. Clearly we are still talking about the same molecule as we swing the atoms around in space. There is however a real difference in shape between the molecules when rotation occurs to different extents. We call the different forms conformational isomers or confomers. We refer to each possible position as a conformation
• the easiest way to start looking at this is to make ethane. Hold the molecule so that you are sighting along the C-C. You can see that there are an infinite number of possibilities as you spin a carbon around. But are they equivalent? The answer is no.
• to look at this in a little more detail, we can use a type of presentation of what we are doing with the models. It's called a Newman projection. In this kind of projection, you sight along a certain C-C bond. You depict the head on view of one of the carbons with a circle. Then you draw the three groups coming off of the front carbon at 120¡ angles. Remember this has nothing to do with the real bond angle, it's merely a two-dimensional representation of a three-dimensional object. The three groups on the front carbon are depicted going to the centre of the circle. The three groups coming off of the rear carbon extend only to the perimeter of the circle (that's how you tell what's on the front and what's on the back).
• the two most different forms of ethane differ from each other by a 60 degree rotation. When the groups are all right in front of each other, this is referred to as the eclipsed conformation. If you rotate the rear carbon 60 degree you get the other conformation, the staggered conformation, where the groups are as far apart as possible.
• what about the stability or potential energy of these different confomers? If you plot a potential energy graph versus dihedral angle, which is the angle of rotation from where you start. If we start fully eclipsed and start rotating the front carbon, what we find is that between the eclipsed and staggered forms there is about a 3 kcal/mol difference. We get staggered conformations at 60, 180 and 300 degrees. For ethane these staggered conformations are identical.
• why are the confomers of different energy? It is not fully understood, but probably arises from repulsion between the electron pairs of the C-H bonds on the adjacent carbon atoms
• that was a very simple example. Let's make life a little trickier, and look at butane. We'll sight along the C-C bond between carbons 2 and 3. Start with both methyl groups up. We call this the fully eclipsed conformation, because the rear methyl group is eclipsed by the front methyl group. If we rotate the back carbon by 60 degrees, we arrive at the first staggered conformation. We call this the gauche conformation. If we continue on for another 60 degrees we get to another eclipsed conformation, but now the methyl group is eclipsed by a hydrogen atom, not another methyl group. If we go another 60 degrees, we arrive at another staggered conformation. This one we call anti, because the methyl groups are as far as possible from each other. If we continue on to 360 degrees we go back through these forms.
• what about the energy of these confomers? There is about a 5 kcal/mol difference between the fully eclipsed and the anti staggered conformations. Furthermore, the gauche staggered confomer is about 0.9 kcal/mol higher in energy than the anti conformation. The partially ecliped confomer is approximately 1.5 kcal/mol lower in energy than the fully eclipsed confomer.
• why should there be such large differences in energy? There is a problem of fitting the methyl groups so close together. This is called steric strain. The groups are bulky and they simply bump into each other when forced so close together.
• one of the big take home messages from this section is that the most stable conformation for a longer alkane is the fully extended conformation, where all of the groups are in the anti staggered conformation. That is not to say that there isn't free movement, but the more extended the more stable.

2. Cycloalkanes:

• we turn now to another major group of alkanes, those that have a ring of carbon atoms in them. We call these cycloalkanes, because they have cyclic rings. Another way to look at it is that the last carbon is bonded to the first.
• the general formula of cycloalkanes is C_nH_2n
• to name simple cycloalkanes you simply add the prefix cyclo- to the name. For example cyclopentane. If there are groups hanging off of the ring you name them as you normally would. If there is only one group, there is no need to number. If more than one group is present you assigne the lower number to the group of lowest alphabetical order, and number such that the sum is minimal.
• cycloalkanes can have very different geometry than straight chain alkanes. As you make these beasts you will notice that they have a lot less rotational freedom.
• let's start with cyclopropane, the smallest possible cycloalkane. To make this one use the silvery pieces in your kit, they are more bendy than the black ones. The orbital hybridization of all of the carbon atoms is sp³, but the bond angles are 60 degrees not 109.5 as would be predicted. There are some serious strains on this molecule, this is another example of steric strain. The first is angle strain. Angle strain occurs when abnormal bond angles occur. The orbitals do not overlap head to head, instead they overlap at an angle. The resulting bond is not nearly as strong.
• the other steric strain experienced by cyclopropane is called nonbonded interaction strain, and that is what we talked about for ethane and butane. You will notice that the hydrogens of cyclopropane are forced into a fully eclipsed conformation.
• these strains make cyclopropane much less stable, or in other words more reactive than other alkanes
• another point that might have occured to you about this molecule is that it is planar. That means that it has a top and an bottom side. If you add more than one group onto the parent ring, the sides they are on become important. Let's add a couple of methyl groups two the ring. You can do this two ways, you can put them both on the same side or on opposite sides. These two different molecules are called cis-trans isomers or stereoisomers.
• if both groups are on the same face we call this the cis isomer. If they are on opposite sides they are trans.
• let's move on to cyclobutane. We can make this with normal black pieces. There are two ways you can arrange this molecule. If you make it planar, the bond angles are 90 degrees, but all of the hydrogen atoms are eclipsed. If you pucker the ring a little, the bond angle decreases and the hydrogens aren't as eclipsed. So in terms of the potential energy of the molecule, puckering introduces angles strain and reduces nonbonded interaction strain. It is a balance of these two that gives the actual structure of cyclobutane, which has a measured bond angle of 88 degrees. So we can conclude that there is a slight pucker.
• of course the same considerations for cis and trans isomers
• let's look at cyclopentane now. Let's start by making it in a planar conformation. IN the planar form the bond angle would be 108 degrees, which is very close to the ideal value of 109.5. So, there is little angle strain in a planar form, but all of the hydrogen atoms are fully eclipsed. The more favoured conformation is puckered, and is called the envelope conformation. Four of the carbons are in a plane and the fifth is popped out of the plane. There are five of these conformations in equilibrium with each other. The puckering reduces the bond angle somewhat, but also reduces the number of eclipsed hydrogens. The actual bond angles are around 105 degrees.

3. Conformations of cyclohexane:

• let's spend the rest of the time looking at the most important of the cycloalkanes, cyclohexane. You run into a lot of related compounds in biology, so it is important to understand the conformations these rings adopt
• if the molecule were planar, the bond angles would be 120 degrees. This is considerably greater than the 109.5 predicted for tetrahedral carbon.
• there are a number of puckered conformations that are possible. The most stable is called a chair. Another one that one hears about is the boat. Let's start with a boat conformation. In this conformation there are a number of steric problems. Number one is that there are two sets of hydrogens that are eclipsed. Another is what are called flagpole interactions. The hydrgeon atoms on the top carbons actually bash into each other.
• the most stable confomer is called a chair, not surprisingly because it looks like a chair. In the chair, the C-C bond angles are the ideal 109.5 degrees, so there is no angle strain, and all of the hydrogens are fully staggered.
• if you tilt the chair a little, you will notice that each carbon has a hydrogen going either straight up or down and one going out. We call the one going up or down axial and the one going out equatorial. This will become important later in the course, so figure them out now.
• the axial and equatorial bonds alternate, one up one down...
• chairs can flip from one chair to another. Mark a hydrogen in your ring with a coloured ball. Flip the ring and you will notice that the axial hydrogens become equatorial and vice versa.
• that doesn't make mush difference if you're talking about cyclohexane, but what if we have a methyl group on one of the carbons? Then the two possible chairs are different. In one chair the group is equatorial and in the other it is axial. In the axial conformation it experiences 1,3 diaxial interactions. It bumbs into the hydrogens on carbon 3 and 5.
• because of these interactions, the confomer with the group in the equatorial position is much more highly favoured. The equilibrium lies 95% to the equatorial side.
• as the size of the group increases, the preference for the equatorial position increases. For example, with a tert-butyl group, there is 10 000 times more in the equatorial position. In effect the ring is locked into this chair.