Chapter 8: Drawing Alkanes

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Have you ever, you know, stared at a molecule drawn flat on a page and just wondered how that simple drawing really captures its constant bending and twisting in three dimensions?

Right.

Or how chemists manage to name these, like, incredibly complex structures in a way that actually makes sense.

What if we told you that unlocking the secrets of the simplest organic molecules, well, it can fundamentally change how you see the chemical world around you.

Today we're doing a deep dive into alkanes, the really the foundational carbon skeletons of organic chemistry.

The backbone ascension.

Exactly.

And our guide for this is the incredibly practical organic chemistry for dummies.

Second, our mission to really distill the core ideas,

understand their structure, see how they move and even uncover their surprisingly reactive side.

Yeah.

And what's fascinating here is that even these simplest molecules, alkanes, they introduce us to really profound concepts, things like how structure influences behavior, how atoms orient themselves in 3D space and, you know, the fundamental rules for naming organic compounds.

This deep dive is kind of your shortcut to truly getting it.

Okay, so let's start right at the beginning.

What are alkanes fundamentally?

Our source calls them the simplest organic molecules, basically just carbon single bonds.

And carbon hydrogen bonds, of course, saturated hydrocarbons.

Right.

But calling them simple.

Yeah.

It feels like it undersells them a bit.

They're really the gateway to understanding much bigger concepts, aren't they?

Absolutely.

I mean, this is where we first bump into ideas like isomerism, you know, same formula, say C4H10, but completely different structures like butane and isobutane and conformation, which is all about how these molecules can flex and bend around those single bonds.

Then there's stereochemistry, the whole 3D arrangement aspect, and of course nomenclature, the naming rules.

Okay, let's actually tackle nomenclature first.

Yeah.

That naming system, our source has this fantastic tip, it feels counterintuitive at first, but for converting a name to a structure,

read it right to left.

It works surprisingly well.

It really does.

So, for example, if you need to draw, let's say, 4T -butyl -2 -3 -some -5 -trimethylheptane.

Sounds complicated, right?

Definitely a mouthful.

But you don't start at the 4T -butyl part.

You go straight to the end.

Heptane, okay, heptamines seven.

So you draw a seven carbon chain.

That's step one.

Get the skeleton down first.

Exactly.

Then you number it.

One to seven.

And then you go back and add the substituents based on the numbers.

So a T -butyl group goes on carbon four.

Okay.

And then methyl groups on carbons two, three, and five.

Boom, boom, boom.

Building with Legos.

But you start with the base plate, the heptane.

That's a great analogy.

And it leads right into a common pitfall, which is,

why do two drawings sometimes look different on paper, but actually represent the exact same molecule?

Yeah, I've definitely been confused by that.

You draw it one way, the book draws it slightly differently.

Right.

And the core issue is, well, these 2D drawings are just representations of 3D things.

Those bonds, they aren't really at 90 degree angles like we sometimes draw them for convenience.

They're more spread out.

Exactly.

Around a typical carbon in an alkane, they're at about 109 .5 degrees.

It's a tetrahedral geometry.

So whether you draw the chain perfectly straight or kind of snaking up and down, it can still be the exact same molecule as long as connectivity is the same.

So the way it's drawn is just a convention sometimes.

Precisely.

And that's why the source really, really emphasizes using molecular models, those plastic ball and stick kits.

Oh yeah, those are essential.

They are crucial, especially when you're starting out.

It helps you build that intuition for the 3D reality behind the flat drawings.

You can twist them and see it.

Okay.

So speaking of twisting,

that leads us to conformations, right?

These carbon single bonds, they're not rigid rods.

They can rotate.

Exactly.

They spin.

And this constant rotation allows alkanes to exist in all these different shapes, or as chemists call them, conformations.

Shapes that don't involve breaking any bonds.

Correct.

It's just rotation around the single bonds.

Think about your own body.

You can bend and flex into different conformations, lying down to read, sitting up straight, stretching.

Different shapes.

Same person.

Exactly.

And just like your body might prefer a comfy chair over, say, doing a handstand for hours.

Definitely prefer the comfy chair.

Molecules are the same.

They naturally prefer lower energy conformations.

They're always kind of wiggling around, seeking the most stable, the least strained arrangement.

Okay.

But picturing that constant rotation for every bond in a molecule sounds

dizzying.

How do chemists actually visualize and draw these specific rotational states?

Is there a shorthand?

There is.

It's called the Newman projection.

It's a really clever way to look directly down one specific carbon bond.

So you're kind of sighting along the bond axis.

Precisely.

Imagine that CC bond is pointing straight at your eye.

The front carbon you see is represented by a point where three lines meet like a Y shape.

Those lines are the other three bonds attached to that front carbon.

Okay.

The back carbon is hidden behind the front one, so we draw it as a circle.

And the three bonds attached to that back carbon are shown as lines coming out from the edge of the circle.

Ah, I see.

So you can clearly see how the groups on the front relate to the groups on the back.

Exactly.

And the convention usually is to keep the front carbon groups fixed in place.

And imagine rotating the back carbon, the circle, to see the different conformations.

So as that back carbon rotates,

there are infinite possible angles, infinite conformations.

But are there specific ones we focus on?

Yes.

While there are infinite possibilities, two main types are the most important energetically,

eclipsed and staggered.

Eclipsed and staggered, okay.

In an eclipsed conformation, if you look down that Newman projection,

the bonds on the front carbon line up perfectly with the bonds on the back carbon.

The dihedral angle, that's the angle between a bond on the front and a bond on the back, is zero degrees.

So they're hiding behind each other.

Exactly.

Although we usually draw them slightly offset in the Newman projection, just so you can see what's back there.

Now in a staggered conformation, the bonds on the back carbon are positioned exactly halfway between the bonds on the front carbon.

The dihedral angle is 60 degrees.

They're as far apart as they can possibly get.

And one is more stable than the other.

Oh, definitely.

When those bonds in the electron clouds around them are eclipsed, they repel each quite strongly.

We call this repulsion torsional strain.

Like they're bumping into each other.

Sort of, yeah.

Electron cloud repulsion.

Because the staggered conformation gets those groups further apart, it has much less torsional strain.

So generally speaking, staggered conformations are more stable, lower in energy than eclipsed conformations.

Okay, that makes sense for simple ethane.

But what about something slightly bigger, like butane, C4H10?

Does it get more complicated?

It does, because now not all staggered conformations are equal in energy, and not all eclipsed conformations are equal either.

It gets really interesting.

How so?

Well, let's look down the central bond, the C2 -C3 bond of butane.

Now you have methyl groups, CH3, attached to both C2 and C3, as well as hydrogens.

Right, bigger groups involved now.

Exactly.

So let's consider the staggered conformations first.

There's one called the anti -conformer.

Anti means the two big methyl groups are positioned 180 degrees apart from each other in the Newman projection, as far away as possible.

That sounds stable.

It is.

That's the lowest energy conformation for butane.

Minimal strain.

But there's another type of staggered conformation called gauche.

In a gauche conformer, the two methyl groups are staggered, yes, but they're only 60 degrees apart.

They're kind of flanking each other.

Closer than ambide.

Right.

And that closeness introduces a little bit of extra repulsion between the methyl groups themselves, not just the bonds.

We call this a gauche interaction, and it makes the gauche conformer slightly higher in energy than the anti -conformer.

The basic principle.

Big groups want to be as far apart as possible.

Okay.

And what about the eclipsed forms for butane?

Also, different energies.

The absolute worst case, the highest energy conformation is when the two methyl groups are directly eclipsing each other.

Dihedral angle of zero.

We sometimes call this totally eclipsed.

Maximum repulsion.

Maximum straight.

Yeah.

There are other eclipsed conformations, too, where maybe a methyl group is eclipsing a hydrogen.

Those are still high energy because they're eclipsed, but they're not as high energy as the totally eclipsed conformation, where the two biggest groups are clashing head on.

So there's a whole landscape of energy levels.

Precisely.

And you can actually plot this out on an energy diagram as you rotate that C2C3 bond through 360 degrees.

You see peaks for the eclipsed forms, valleys for the staggered forms, with the anti -conformation being the deepest valley.

Fascinating.

Okay, let's switch gears slightly.

Alkenes don't just form chains, they also form rings.

Cycloalkanes.

Right.

And naming them is pretty intuitive, actually.

You just count the carbons in the ring, use the regular alkane name for that number of carbons, like propane or hexane, and stick the prefix cyclo on the front.

So cyclopropane for three carbons, cyclomethane for four, cyclopentane, cyclohexane.

Got it.

Simple enough.

But rings immediately bring up another important concept.

Stereochemistry on the ring.

How do substituents, groups attached to the ring, orient them themselves?

Because the ring isn't just a flat line, it has sides.

Exactly.

Think of a ring like cyclopentane as having a top face and a bottom face.

If you have two substituents attached to the ring, say two methyl groups, and they're both pointing towards the same face, maybe both pointing up relative to the ring that's called a cis -stereoisomer.

Just for same side.

Correct.

If one substituent is pointing up and the other is pointing down off opposite faces, that's a trans -stereoisomer.

Transfer across or opposite?

Mm -hmm.

And cis and trans -isomers are types of stereoisomers.

Remember, that means they have the same atoms connected in the same order, but they differ in their specific 3D arrangement.

Like 1 -bar -2 -dimethylcyclo -pentane, it can exist as a cis -isomer or a trans -isomer, and they are different molecules with different properties.

Okay.

Now, you mentioned cyclohexane earlier.

Six -carbon ring.

You said it's particularly interesting.

Oh, absolutely.

Cyclohexane is super common, and here's the key thing.

Even though we often draw it as a flat hexagon on paper for simplicity, it's not flat at all.

It's puckered.

Very much so.

It predominantly adopts a specific 3D shape called the chair conformation.

It looks kind of like a lounge chair.

This chair shape is the most stable conformation for cyclohexane because it minimizes all types of angle strain, torsional strain.

Are there other shapes?

Yes.

There are others like the boat conformation, which looks like a boat, and twist boats and half chairs, but these are all higher in energy, less stable than the chair.

So at room temperature, most cyclohexane molecules are hanging out in the chair form.

How do you even draw that chair?

It sounds tricky.

It takes a little practice, but the source gives a good tip.

Start by drawing two parallel lines slightly offset and angled downwards.

Then from the right end of the top line and the left end of the bottom line, draw lines converging downwards to form a V that's like the tail or foot rest of the chair.

Then from the other ends of those parallel lines, draw lines converging upwards to form another V that's the nose or back rest of the chair.

With practice, you get the hang of making it look right.

Got it.

Parallel lines, downward V, upward V.

What about the hydrogens on this chair?

Are they all equivalent?

Nope.

That's another crucial point.

The hydrogens on a cyclohexane chair occupy two distinct types of positions.

There are axial hydrogens and equatorial hydrogens.

Axial and equatorial, like the earth.

Exactly like the earth.

Axial hydrogens stick straight up or straight down, parallel to an imaginary axis running through the center of the chair.

Equatorial hydrogens stick out sort of sideways around the equator of the chair ring.

Each carbon in the chair has one axial hydrogen and one equatorial hydrogen.

And drawing those correctly is important.

Very important.

A good tip is to draw the axial hydrogens first.

On the up points of your chair drawing, like the nose, the axial bond goes straight up.

On the down points, like the tail, the axial bond goes straight down.

Fill those in first.

Then the equatorial bonds basically point and slightly up or down, generally parallel to the ring bonds next door.

It sounds complicated, but drawing them systematically helps.

Axial first, then equatorial.

Makes sense.

Now here's something that blows my mind a bit.

You said cyclohexane isn't static.

It does something called a ring flip.

Yes.

At room temperature, cyclohexane chairs are constantly interconverting.

It's not a static structure.

What happens in a ring flip is essentially the nose carbon flicks down and the tail carbon flips up.

The whole ring sort of inverts itself into an alternative chair conformation.

Wow.

And what happens to the hydrogens during this flip?

This is a really neat part.

Every single axial hydrogen becomes equatorial and every single equatorial hydrogen becomes axial.

They completely swap positions.

Oh.

Yeah.

And again, the source really hammers this home.

Get out those molecular models.

You absolutely have to build a cyclohexane chair and physically perform the ring flip yourself to truly understand how those positions interchange.

Seeing it in 3D makes all the difference.

I believe it.

Okay.

So simple cyclohexane flips between two identical chairs.

But what if you have substituents on the ring, like a methyl group or something bigger?

Ah, now it gets even more interesting because with substituted cyclohexanes, the two chair conformers, the one before the flip and the one after the flip, are often not identical in energy.

One is usually more stable than the other.

Why is that?

Let's take isopropylcyclohexane.

An isopropyl group is pretty bulky.

Now imagine the chair conformation where that big isopropyl group is sticking straight up or down in an axial position.

Okay.

I'm picturing it crowded.

Very crowded.

When a large group is axial, it bumps into the axial hydrogens that are on the same side of the ring, two carbons away.

These hydrogens are called the di -axial hydrogens relative to the group.

This bumping causes steric strain,

basically.

Things trying to occupy the same space.

It's specifically called 1 -meth -3 -diaxial strain or 153 -diaxial interactions.

That's unfavorable.

Highly unfavorable.

It raises the energy of that conformer significantly.

Now consider what happens after the ring flip.

The isopropyl group, which was axial, becomes equatorial.

Exactly.

And when it's equatorial, it's sticking out away from the ring, not bumping into those axial hydrogens nearly as much.

Much happier there.

Much happier, much lower energy, much more stable.

So the fundamental principle is large substituents strongly prefer to be in the equatorial position on a cyclohexane ring to avoid 1 -meth -3 -diaxial strain.

The equilibrium will heavily favor the conformer where the big group is equatorial.

That's a really important concept for stability.

Now you mentioned cis and trans earlier.

Does the flip change whether something is cis or trans?

Excellent question.

And the answer is absolutely not.

This is a crucial distinction people sometimes mix up.

Confirmation versus configuration.

Okay, break that down.

A ring flip changes the confirmation how the molecule is bent or twisted in space, like changing between chair forms or going axial to equatorial.

But it does not change the configuration.

Configuration refers to the fixed 3 -D arrangement that can only be changed breaking and forming bonds.

So if two groups are cis on the ring, on the same face, they will still be cis after the ring flip, even though one might have gone from axial to equatorial and the other from equatorial to axial.

Cis stays cis, trans stays trans during a ring flip.

Got it.

Flipping is just flexing, not rearranging the fundamental connections or stereochemistry.

Precisely.

Changing from cis to trans or vice versa would require a chemical reaction.

Okay, so putting this together.

If a question asks you to draw the most stable chair confirmation of, say, a dis -substituted cyclohexane,

how do you approach that?

Good question.

First, you need to figure out the relative positions based on cis -trans.

Sometimes drawing a simple Hayworth projection first helps.

That's the flat hexagon view where you show bonds as going up or down, wedges dashes to represent the faces.

This helps you see if, for example, a cis 1 -3 relationship means both groups are up or both down.

Okay.

Then you translate that to the chair.

For cis 1 -3 dimethyl cycle hexane, for instance, being cis, means both methyl groups could be axial or both could be equatorial after you consider the chair structure, which is more stable.

And one where they're both equatorial to avoid that 1 -philon -3 -diaxial strain.

Exactly.

The diaquatorial conformer is much lower in energy.

The diaxial conformer would have lots of strain.

Now, what if you have a molecule where, because of the cis -trans relationship, one group must be axial and one must be equatorial, like in many trans -wannable 2 -dissubstituted cases?

Then you can't have both equatorial.

Right.

In that case, the most stable conformer will be the one where the larger of the two groups occupies the equatorial position, leaving the smaller group to take the less favorable axial position.

Minimize the strain as much as possible.

Put the biggest group in the comfiest equatorial seat.

You got it.

Okay, let's shift one last time.

Reactions.

Alkanes.

They're generally pretty unreactive, right?

Like paraffin wax.

Yeah, they're often called paraffins from the Latin paromethenes, meaning little affinity.

They don't have double bonds or functional groups that attract most reagents.

They're quite happy as they are under many conditions.

But there must be some reaction they do, otherwise they wouldn't be much use in synthesis.

There is one major reaction type that's really important, especially early in organic chemistry.

Free radical halogenation.

Free radical halogenation.

Okay.

This is often the first reaction you learn for alkanes, because it's a way to take an inert alkane and actually put a functional group on it, specifically a halogen atom, like chlorine, Cl, or bromine, Br.

So you're swapping a hydrogen for a halogen.

Exactly.

Let's use the chlorination of methane.

CH4 is the classic example.

You react methane with chlorine gas, Cl2.

Just mix them together.

Not quite.

The interesting thing here, as the source points out, is that this reaction is often photochemical.

It doesn't usually happen just by mixing in the dark at room temp.

You need energy input, typically in the form of light, ultraviolet light, often works well.

We represent light energy as high in reaction schemes.

So light triggers the reaction.

How?

It happens in three main stages.

Initiation, propagation, and termination.

Okay, stage one.

Initiation.

Getting things started.

Right.

The light energy is absorbed by a chlorine molecule, Cl2.

This energy is enough to break the relatively weak chlorine -chlorine bond, but it doesn't break unevenly.

It breaks symmetrically what we call homolytic cleavage.

Homolytic, meaning same splitting.

Exactly.

Each chlorine atom gets one electron from the bond pair that was holding them together.

Remember, a normal bond is two electrons.

So this forms two chlorine atoms, but because they have an unpaired electron, we call them chlorine radicals.

Radicals are a species with unpaired electrons, and they are typically very reactive.

And you draw the electron movement with special arrows.

Yes.

The source mentions using single -headed fishhook arrows to show the movement of just one electron, which is characteristic of radical reactions.

Okay, initiation creates chlorine radicals using light.

What's next?

Propagation.

Propagation.

Yeah.

This is the stage that actually forms the product and keeps the reaction going in a chain.

That highly reactive chlorine radical needs an electron to complete its octa.

So it bumps into a methane molecule.

And does what?

It rips off a hydrogen atom, not a proton, H +, but the whole H atom, including its one electron.

So the Cl radical reacts with HCH3 to form HCl, hydrochloric acid, and what's left from the methane.

Ah, CH3 with one electron missing from where the H was.

Precisely.

You're left with a methyl radical CH3.

Now this methyl radical is also unstable.

It has an unpaired electron.

So it needs to react too.

Right.

It bumps into another unreacted chlorine molecule, Cl2.

The methyl radical grabs one chlorine atom and one electron from Cl2 to form chloromethane, CH3Cl.

That's our desired product.

Okay.

We made the product.

But what's left over from the Cl2 molecule after the methyl radical took one Cl atom?

Another chlorine radical.

Bingo.

Another chlorine radical is regenerated.

And what can that chlorine radical do?

It can go back and react with another methane molecule.

Exactly.

See, it's a chain reaction.

The radical produced in the second propagation step is a reactant in the first propagation step.

So once initiated, these two propagation steps can theoretically cycle over and over again, consuming methane and Cl2 and producing HCl and chloromethane as long as the radicals persist.

Wow.

Okay.

So why does it ever stop?

What's termination?

Termination steps are what eventually break the chain.

They happen when two radical species bump into each other and combine, removing radicals from the cycle without generating new ones.

Like two chlorine radicals finding each other.

Yep.

Two Cl radicals can form Cl2.

Or two methyl radicals could combine to form ethane, CH3CH3, a side product.

Or a chlorine radical and CH3Cl.

Any reaction that consumes radicals without producing new ones is a termination step.

These become more likely when the concentration of radicals gets high or the reactants run low.

Initiation, propagation, termination.

No.

Now methane is simple.

All hydrogens are the same.

What about a larger alkane like butane, which we discussed earlier?

It has different types of hydrogen.

Excellent point.

Butane, CH3CH2CH3, has primary hydrogens on the NCH3 groups and secondary hydrogens on the middle CH2 groups.

And we classify hydrogens based on the carbon they're attached to.

Primary, secondary, tertiary.

Right.

Primary one degree carbon is attached to one other carbon, secondary two degrees to two, tertiary three degrees to three.

The hydrogens attached follow the same classification.

And this matters because the stability of the carbon radical formed during propagation depends on this.

How so?

Carbon radicals, like carbocations, are stabilized by having more alkyl groups attached to the radical center.

So tertiary radicals are more stable than secondary radicals, which are more stable than primary radicals.

More substituted, more stable radical.

Exactly.

So when that chlorine radical attacks butane, it could abstract a primary hydrogen to form a primary radical, or it could abstract a secondary hydrogen to form a secondary radical.

Which path is easier leads to a more stable intermediate.

Abstracting the secondary hydrogen because the secondary radical is more stable.

Correct.

So while you might get some one chlorobutane from attacking a primary H, you'll get more two chlorobutane from attacking a secondary H because the pathway via the more stable secondary radical is favored.

Chlorination shows some preference for attacking more substituted positions.

Some preference.

That leads into the last point.

Chlorination versus bromination.

You mentioned bromine.

Br2 can be used too.

Is it just the same?

It follows the same mechanism, initiation, propagation, termination.

Light or heat initiates it, forming bromine radicals Br2.

But there's a crucial difference in selectivity.

Selectivity, meaning which hydrogen it chooses to attack.

Precisely.

Remember how we said the chlorine radical attacks the secondary hydrogen and butane preferentially, but you still get some primary attack?

Chlorination is moderately selective.

Bromination, using Br2, is highly selective.

Why the difference?

It comes down to the stability and reactivity of the radicals involved.

A chlorine radical is very unstable, very high energy, very reactive.

It's like desperate to react, so it's not very picky.

It will grab almost any hydrogen it bumps into, although it has a slight preference for the one leading to a more stable radical.

Okay, reactive and less picky.

A bromine radical, on the other hand, is actually more stable, lower in energy, and therefore less reactive than a chlorine radical.

Because it's less desperate, it's much more choosy.

It will preferentially wait until it finds the hydrogen that leads to the most stable possible radical intermediate before it reacts.

So, for butane, the bromine radical would almost exclusively attack the secondary hydrogens?

Pretty much.

With bromination of butane, you get almost entirely to bromobutane.

It's highly selective for the more substituted position because the stability difference between the primary and secondary radicals matters much more to the less reactive bromine radical.

It's classic example of the reactivity selectivity principle.

Less reactive region, more selective outcome.

You nailed it.

It's a really powerful concept in synthesis.

Wow, okay.

That was quite the journey through alkanes.

We went from just drawing simple chains to understanding their names, seeing them twist and turn in 3D with confirmations like the cyclohexane chair.

And the ring flips.

And the ring flips.

All the way to their specific light -driven reactions with halogens and even the subtle differences between chlorine and bromine.

It really makes you appreciate just how much complexity and, well, personality exists even in these supposedly simple molecules.

Absolutely.

And if you connect this to the bigger picture, really mastering alkanes isn't just about, you know, memorizing some names and reactions.

It's about building that chemical intuition, understanding how tiny changes in 3D shape affect energy instability, how structure dictates reactivity.

That foundation serves you throughout all of organic chemistry.

It really sets the stage.

So here's a thought for you, our listener.

We talked a lot about how crucial molecular models are for grasping the 3D reality behind flat drawings of molecules.

Thinking beyond chemistry now, what other complex 2D representations that you encounter in science, or maybe even everyday life, could be radically clarified if you had a simple 3D model to play with.

Something to ponder.

Thank you so much for joining us on this deep dive into the world of alkanes.

We hope you feel a little more informed, maybe a bit more confident, and definitely curious to explore more organic chemistry.