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Hi! I’m Deboki Chakravarti and welcome back to Crash Course Organic Chemistry!

What do Saturn, a diamond, and Giant’s Causeway in Ireland all have in common? Hexagons. From honeycombs to the shells of tortoises, hexagons appear all over the natural world.

Even when you push bubbles together, they’ll make hexagons. But the scientific reasons why hexagons seem to be everywhere are debatable. Some people consider hexagons a space-saving wonder.

For example, pencil companies realized that they could pack more hexagonal pencils in a box and save storage space. So that’s why every Number 2 is hexagonal. Hexagons are even in the arrangement of carbon atoms, like the close-packed that structure gives diamond its famous hardness.

And in graphite, hexagonal sheets of carbon conduct in one direction but not the other. In this episode, we’ll uncover the secrets of one of organic chemistry’s favorite manifestations of this shape: the cyclohexane molecule. [Theme Music]. As we learned last time, cycloalkanes are named like other alkanes, except they have the prefix cyclo- right before the root name.

A plain old five-carbon ring would be called cyclopentane. If we just have one substituent, we don’t need a number. Like methylcyclopentane.

But if there are two or more substituents, the carbons are numbered to give those substituents the lowest possible numbers. And the substituents are put in alphabetical order, like we normally do. For example, 1-ethyl-2-methylcyclopentane.

A cycloalkane can be named as a substituent too. Like when the acyclic (or no ring) part of the molecule has more carbons, or if the chain has an important functional group. For example, 4-cyclopropylhept-1-yne.

Cycloalkanes have two distinct faces, because they don’t have free rotation around their bonds. In a cycloalkane with two substituents pointing toward the same face of the ring, we add the prefix cis-. And if they point toward opposite faces of the ring, we add the prefix trans-.

The most common cycloalkanes are cyclopentanes and cyclohexanes, because of what we talked about last episode: they have low ring strain so they’re more stable. Ring strain comes from a combination of two things. There’s angular strain, or deviating from the ideal angle of 109.5 degrees for sp3 hybridized carbons.

And there’s torsional strain, or having bonds in an eclipsed conformation. Nowadays we know what causes ring strain, and we know that high strain means more energy and less stability. But organic chemists initially figured out ring strain (and lots of other things about chemical reactions) by setting compounds on fire.

To see how, let’s go to the Thought Bubble. Combustion reactions are when hydrocarbons burn in oxygen, releasing energy stored in their carbon-carbon bonds. One way to think about combustion is that plants incorporate energy from the sun into molecules.

So when we burn wood in a campfire, or ancient plants that became fossil fuels, it’s kind of like we’re undoing photosynthesis. Besides using heat to toast marshmallows, organic chemists can also measure the heat produced in a combustion reaction to learn about the formula of chemical compounds and the energy stored in their bonds. In the inner chamber of an instrument called a bomb calorimeter, we burn a compound, measure the temperature change, and use that measurement to calculate the energy of the reaction.

Don’t worry, it’s safe! But if there was a bunch of energy in a closed system like this... that’s a bomb. Hence the name.

Organic chemists used bomb calorimetry to quantify ring strain, by comparing the heat released from the combustion of an alkane with a cycloalkane. Any extra heat from the cycloalkane has to be from the energy of the ring strain being released. So we have to do a little math to get there.

For example, we can use a bomb calorimeter to burn a very long acyclic alkane, and find that the average energy released per CH2 unit is 658.6 kilojoules per mole. And when we burn cyclopropane, which we know is a very strained ring, it releases 697.1 kilojoules per mole per CH2 unit. So the ring strain energy contributed 38.5 kilojoules per mole.

We have to multiply that by three for the three CH2 units in cyclopropane, for a total of 115 kilojoules per mole. There's quite a bit of extra energy packed into that ring. Thanks, math!

And thanks to you, Thought Bubble! Through these sorts of combustion experiments, the cyclohexane ring was found to have no ring strain, and cyclopentane is pretty low too. Both of these cycloalkanes decrease ring strain because their angular strain is relatively low.

In a pentagon, the angles are 108 degrees and in a hexagon, they’re 120 degrees. These aren’t very far from the comfy 109.5 degrees for sp3 hybridized carbons. And even though flat cycloalkanes would have bonds in an eclipsed conformation, they can reduce torsional strain by puckering.

They poke certain atoms above and below the molecular plane, so they’re not as lined up, and are slightly less strained. A puckered cyclopentane molecule pokes one carbon atom above the molecular plane, forming an envelope conformation. We can see this a little better by looking at a Newman projection along one of the bonds.

The thing about any molecular drawing or model is that they’re static, instead of the vibrating, moving molecule of reality. So the pucker of cyclopentane isn’t stuck on one carbon. Each carbon takes a turn poking above the molecular plane, so the molecule is kind of constantly wobbling around.

But cyclohexane is even more stable and free of ring strain because it can achieve a puckered structure that has all of the hydrogens staggered and all of the bond angles at 109.5 degrees. This basically magical structure is called the chair conformation of cyclohexane. Except instead of a plush armchair that you’d wanna take a nap in, it’s more like a crooked beach chair on the deck of a cruise ship.

Imagine a friend sitting with their head up and feet down. The two parallel lines in the middle are sort of like the “seat” of the chair. And if we look along those carbon-carbon bonds, we can draw a Newman projection and see that the hydrogens are staggered.

Cyclohexane can pucker in other ways too with a little bit of ring strain, like in the boat conformation. Which is also, kind of honestly like a goofy-looking boat. In the boat conformation, the footrest of the “chair” lines up with the headrest to make a little nook.

The angles between carbon atoms are still 109.5 degrees, but now there’s a little torsional strain. By drawing a Newman projection, we can see that the hydrogens are all eclipsed. But there are also two flagpole hydrogens that are sticking towards each other.

To relieve strain from this interaction, the boat twists a bit so the flagpole hydrogens don’t get too close. Sort of like turning your head for a kiss to avoid smashing noses. Most of the cyclohexane molecules in a sample will be in the chair conformation.

So it’s really important for us to learn how to draw them. We’ll practice here, starting with two parallel lines. Next, draw another set of parallel lines to make two sideways Vs.

Then, connect those two Vs with a third group of parallel lines. That’s the base structure of the chair! Now, we have to stick in the hydrogens.

Three on each face will be straight vertical up or down and they’re all parallel to each other. These are called the axial positions. And here’s a hint: the V-shaped points made by the carbon atoms can help you figure out if the axial bond goes up or down.

The Vs pointing up have the axial hydrogens above the plane, and the downward Vs have the axial hydrogens below the plane. And finally, we add in three hydrogens that are parallel to the bonds one step away, which are what we call the equatorial positions. The equatorial hydrogens form a ring around the axial hydrogens, sort of like Earth’s equator.

This can be tricky to wrap your head around with just words, so we’ve color-coded the equatorial bonds to the bonds in the chair structure that they match! So that’s one chair conformation. But because even low energy molecules are constantly moving, cyclohexane molecules are switching between conformations many times every second.

There’s one chair position where our friend is sitting comfortably with their head up and feet down, then there’s the boat, and finally there’s another chair position where their head is hanging down and their feet are sticking up. Plus lots of in-between-forms. When the chair converts from one position to the other, it’s called a chair flip.

But that name can be a little misleading! A chair flip does NOT mean we flip the chair upside down, and our friend falls out. We just move their head down and feet up.

Also, all the axial hydrogens become equatorial hydrogens. A chair flip also DOES NOT require bonds to be broken, they just rotate to shift the bonded groups around. In our all-hydrogen example, both chairs are equivalent.

They’re the same energetically, and we might wonder why we’d need to think about chair flips at all. But let’s say we swap out one of the hydrogens for a methyl group to create methylcyclohexane. If we do a chair flip with this molecule, you can see how the methyl group can switch between an axial position or an equatorial position.

And this has chemical consequences. For methylcyclohexane, the chair with the methyl group in the equatorial position has slightly lower energy. So most methylcyclohexane molecules in a sample will be in that chair conformation as they’re constantly converting back and forth.

To understand why the equatorial methyl group makes it slightly lower energy, we have to go back to our good ol’ tools: Newman projections. Let’s look along the “seat” carbon-carbon bonds in the middle. Now, we can see that when the methyl group is in the axial position, it’s a gauche conformation with respect to the “head” carbon.

But in the equatorial position, the methyl group is in the anti conformation. There’s another important effect here too. The axial methyl group has steric interactions with the other two axial hydrogens above the ring.

The crowding introduced by pushing these groups together is called diaxial strain. This makes an energy difference between the conformations, as we learned in episode 6. And the equatorial position is slightly preferred.

As more substituents are added to cyclohexane, the interactions become more noticeable. For example, let’s consider cis-1,3-dimethylcyclohexane. If both of the methyl groups are sticking up in the axial position, it’s strongly unfavorable.

Steric hindrance kicks in and they feel crowded and uncomfortable. So most molecules in the sample will be in the chair conformation where the methyl groups are equatorial. But now let’s consider trans-1,3 dimethylcyclohexane.

In both chair conformations, there’s one axial methyl group and one equatorial methyl group. So chair flips happen, but neither structure is energetically preferred over the other. There’s an axial methyl group in each one.

As the groups on the chair get even bigger and chunkier, the chair conformation where the substituent is in the equatorial position becomes even more favored by molecules in a sample. Neighboring atoms just get too crowded with too much diaxial strain in the axial position. In fact, more than 99.9% of this compound has the tert-butyl group in the equatorial position!

Now, we’ve been drawing a lot of chairs at this point, and, let’s be honest, you might want to flip a real chair. I know that I do! But take a breath, and when you’re ready, there’s one more thing to learn.

Remember skeletal structures of cyclohexane? They look like flat hexagons. So we also have to think about how to convert a flat structure to a chair drawing, especially when there are two or more substituents.

The first step is to see what the flat structure tells you. In this molecule, for instance, we know the solid wedge means the methyl group is coming out toward us, and the dashed wedge means the isopropyl group is pointing into the screen. So we have to make sure these groups stay on the correct side of the cyclohexane when we draw a chair.

The methyl needs to be above the plane, and the isopropyl needs to be below the plane. Next, we should number the carbons with the substituents just to help us keep track of them. These numbers don’t necessarily have to be IUPAC-official, it’s just helpful when we switch from a hexagon to a weird beach chair.

And speaking of weird beach chairs… it’s time to draw the base of the chair and number our carbon atoms. It’s really important that we number the same way we did on the flat cyclohexane. So if we numbered clockwise on the flat drawing, our numbers need to be clockwise on the chair too.

Then, we can stick in placeholder axial and equatorial bonds on the two carbons that have substituents. After that, we can use our reference numbers to know we should put the methyl on carbon 1 so it’s pointing above the plane (and it’s higher than the hydrogen on the same carbon). And we put the isopropyl on carbon 3 so it’s below the plane (and it’s lower than the hydrogen on the same carbon).

The isopropyl group is in an axial position, but it’s way bigger than the methyl group. When we flip a chair, all of the axial substituents become equatorial and vice-versa. So we need to do a chair flip here to get that chunky group in an equatorial position, which is what’s energetically preferred.

Let’s keep our eye on carbon 2 and pull it down. That forces carbon 3 up, and the isopropyl group moves into the equatorial position, but it’s still below the plane. At the same time, the methyl group moves into the axial position as carbon 1 pushes up, and it stays above the plane.

The axial methyl group has a lot less diaxial strain than the chunky, branched isopropyl. So this is our more stable chair! So the big takeaway in today’s episode is to practice!

It helps to practice drawing cyclohexanes, but if you have a model, that’s even better because you can actually change the positions of all the hydrogens and see them go from equatorial to axial, and you also get to play with a toy, and see chemistry, like, more visibly! And it’s okay if all this chair stuff doesn’t make any sense immediately… it’s hard to wrap your head around and takes a lot of practice. But it’s really important to know our cyclohexanes.

These organic molecules show up in carbohydrates, steroids, plant molecules, pesticides, and so many other things. And the conformation is so important to the function of these molecules. Today we talked about:.

How rings pucker to relieve angular and torsional strain,. How ring flips make axial substituents become equatorial,. And how more bulky substituents should be equatorial when drawing cyclohexanes in chair conformations, because that’s more energetically favorable.

Next episode we’ll be moving on to stereochemistry, which can also be tough, and dreaming about eating all the sugar we could want, which is pretty easy! Thanks for watching this episode of Crash Course Organic Chemistry. If you want to help keep all Crash Course free for everybody, forever, you can join our community on Patreon.