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Hi!
I’m Deboki Chakravarti and welcome back
to Crash Course Organic Chemistry!
Imagine, for a second, that you’ve never
seen a real life cat.
Not even a picture or a single YouTube video.
You’ve only seen simple 2D drawings of them, like a stick figure drawn out by your cousin or a misshapen fuzzy blob.
From those drawings, you’d get a sense of an average cat: two ears, some whiskers, four legs, and a tail.
But without imagining a 3D cat, you wouldn’t have a great idea of how they fit into the world -- including things like their fluffy fur, chonky bellies, or sharp retractable claws.
The organic molecules that make up cats (and everything living on Earth) aren’t the flat, lifeless structures we’ve been drawing, either.
From the simplest organic molecule, methane, to the most complex proteins, all of these compounds can be plotted on a 3D cartesian coordinate system with x, y, and z axes.
By understanding how molecules have 3D shapes, we can better understand how the structure of any molecule affects what it can do.
Some compounds fit together like puzzle pieces, and other combinations are like trying to force a square peg into a round hole.
[Theme Music]
If you’re super rusty on VSEPR (and no worries if you are!)
or if you haven’t heard of hybridization or valence bond theory, you may want to watch Crash Course General Chemistry episodes 24 and 25.
Hank did a really good job explaining those ideas, so I’m just going to do a quick refresher and build from there.
Since organic chemistry became a thing, there have been lots of improvements to theories about how atoms interact and form chemical bonds with each other to make molecules.
It’s like the quote “standing on the shoulders of giants.”
Each theory explains an observed phenomenon that the previous theory couldn’t quite nail down.
In 1916, Lewis structures helped us think about how atoms and electrons are arranged in a molecule.
They’re a powerful tool, which is why we still use them today for simple 2D drawings.
These structures use straight lines to represent covalent bonds and dots for unbonded valence electrons.
Valence Shell Electron Pair Repulsion Theory, or VSEPR, was first proposed in 1957 and started to explain the observed 3D shapes of these molecular structures.
VSEPR is the theory that the 3D shape of a molecule is determined by a central atom’s lone pairs of electrons and the other atoms it’s bonded to.
There are five generally accepted VSEPR electron-pair geometries, which are the 3D shapes that take lone pairs of electrons and bonds into account.
In organic chemistry, we use three of these five geometries: linear, trigonal planar, and tetrahedral.
On the other hand, there are lots of molecular shapes, which describe how atoms in a molecule relate to each other and pretend that lone pairs are invisible.
For example, we say that water has a bent molecular shape.
In a water molecule, the central oxygen atom has two lone pairs of electrons, and two hydrogens bonded to it.
The bond angles and the lone pairs are important-- they lock the molecule into its 3D shape.
But when we call it bent, that only describes the atoms.
As scientists started to widely accept VSEPR as a way to explain the molecular shapes that they saw experimentally, quantum theory became a thing.
And so did the idea of orbitals, places where we’re most likely to find electrons around atoms.
There are four distinct atomic orbital names and shapes: s, p, d, and f.
Because of how orbitals are positioned, even orbitals couldn’t completely explain the 3D shapes predicted by VSEPR and observed experimentally.
Something was still missing!
That something was an idea called orbital hybridization.
Basically, you can mix one s-orbital with its spherical shape, and one p-orbital with its 3D figure-eight shape to make two hybrid atomic orbitals that sort of look like both.
Kind of like mixing a donkey and a horse to get a mule.
It’s important to pay attention to the number of atomic orbitals we mix, because that’s how many hybrid orbitals are produced.
In other words, if we mix one s orbital and one p orbital we get two sp orbitals.
If we mix one s orbital and two p orbitals we get three sp2 orbitals.
And if we mix one s orbital and three p orbitals we get four sp3 orbitals.
Orbital hybridization helps explain the 3D geometries predicted by VSEPR.
For example, let’s look at the tetrahedral shape of a methane molecule.
A carbon atom has four valence electrons, which all need to be unpaired to bond.
To make sure they’re all unpaired, it makes four sp3 hybrid orbitals, and we can just imagine sticking one electron in each.
Each hydrogen atom has a 1s orbital with one unpaired electron.
So when these five atoms unite to form methane, each hydrogen’s 1s orbital overlaps with the carbon’s sp3 hybrid orbitals to form chemical bonds, called sigma bonds.
Sigma bonds are sort of like a handshake, made by the direct overlap of two orbitals that point at each other.
This idea of overlapping orbitals is the foundation of valence bond theory.
These bonds give methane a tetrahedral molecular shape, so it’s basically the same as these four balloons, tied together at the base.
These aren’t Hank’s balloons from 7 years ago, but they’re just as good at showing the shape and the science behind it.
Now, methane is a relatively simple way to think about orbital hybridization and valence bond theory.
But we can use these same ideas to explain the 3D shapes of molecules with double and triple bonds too.
For double bonds, let’s look at ethene.
If you’ve been paying attention to all this nomenclature we’ve been doing, that’s C2H4.
The carbon atoms are double-bonded to each other.
Here’s the Lewis Structure, which shows that each carbon has three things connected to it.
So each carbon needs to hybridize three atomic orbitals to make three sp2 hybrid orbitals.
Two sp2 orbitals overlap with two hydrogen 1s orbitals.
And one sp2 orbital overlaps with one of the other carbon’s sp2 orbitals.
That makes three sigma bonds!
Because each carbon made three sp2 hybrid orbitals, each carbon still has one p orbital left.
The leftover p orbitals on the two carbons are close enough to say,
“Hey! What’s up? If we share our electrons, we can make a bond too!”
This is called a pi bond, where the orbitals line up next to each other and sort of overlap sideways.
More valence bond theory!
Every double bond we’ll meet in this series has a sigma bond with orbitals that overlap end-to-end, and a pi bond with orbitals that overlap sideways.
This leads to the 3D molecular geometry of ethene: a trigonal planar arrangement around each of the carbon atoms.
Now, for triple bonds, let’s take a look at ethyne.
So we’re all on the same page with names, this one’s C2H2, with the carbon atoms triple-bonded to each other.
In this Lewis Structure, we can see that each carbon has two things connected to it, one hydrogen and the other carbon.
That’s how we know they’ll combine two atomic orbitals to make two sp hybrid orbitals.
Each carbon has two p orbitals left, which all lean over and share their electrons in two pi bonds!
So the triple bonds that we’ll meet in this course will always have one sigma bond and two pi bonds.
This makes the 3D molecular geometry linear around the carbon atoms.
When we draw alkynes, the groups they’re
bonded to should always be in a straight line, at 180 degree angles.
So like this.
NOT this.
Now, organic chemistry is all about carbon, so we’ve been focusing on orbital hybridization and valence bond theory in carbon-containing compounds.
But other elements have their own electron-pair geometries and hybrid orbitals going on too.
For example, oxygen in a water molecule is kind of similar to methane.
In water, the central oxygen atom is sp3 hybridized.
It has sigma bonds with two hydrogen atoms and has two lone pairs hanging out.
This gives water a tetrahedral electron-pair geometry and a bent molecular shape, remember?
Oxygen is often sp3 hybridized and forms single bonds in organic compounds like alcohols and ethers too.
But it can also be sp2 hybridized and form double bonds as a carbonyl group, which we see in aldehydes, ketones, and carboxylic acids.
In fact, carbonyl groups and valence bond theory were super important to figuring out the structure of a little molecule known as DNA.
You know, just the thing that holds our genetic information and provides the blueprint for our cells to grow and stay alive.
To see how, let’s go to the Thought Bubble.
Our story begins in 1869, when Swiss physiological chemist Friedrich Miescher isolated a novel substance from the nucleus of a cell.
In 1919, Russian biochemist Phoebus Levene was able to prove that DNA had three main pieces: five-carbon sugars, phosphate groups, and organic ring compounds called nitrogenous bases.
By 1944, we knew that our genetic information was held in DNA molecules, and Austrian biochemist Erwin Chargaff found a relationship in the ratios of the nitrogenous bases: A, T, G, and C.
But researchers struggled to figure out the 3D shape of DNA.
They worked with the idea that nitrogenous bases needed to be in the center of a double helix and bond with each other.
Any bond that formed between the bases had to be strong enough to hold the double helix together, but weak enough that the helix could open up for things like reading the genetic code or copying DNA strands.
Some scientists presented evidence that the bonds between bases were hydrogen bonds, a relatively weak intermolecular force.
Other scientists, like Rosalind Franklin, made 3D representations that suggested the same.
But there was a problem: at that time, the agreed-upon structure for nitrogenous bases meant some oxygen and nitrogen atoms wouldn’t have the correct orbital hybridization and geometry to form hydrogen bonds.
American crystallographer Jerry Donohue eventually suggested that the textbooks were wrong and proposed a different structure: an oxygen atom that was sp2 hybridized.
Instead of an alcohol group, it was a carbonyl group.
With this key structural change, DNA’s nitrogenous bases could hydrogen bond to each other!
After this, the now-famous 1953 report on the structure of DNA was published, helping change the way we thought about genetics.
Talk about standing on the shoulders of giants.
Thanks, Thought Bubble!
So it’s pretty clear that the bonds of a molecule affect its structure.
We use the word isomers to describe molecules that have the same molecular formula, but different arrangements of atoms.
To help remember this, the word isomer comes from the Greek root isos-, meaning same.
From general chemistry we have words like isotope, meaning same number of protons, or isoelectronic, meaning same number of electrons.
Then, we add the Greek root -mer meaning part.
It shows up in words like polymer, meaning many parts, and monomer, meaning a single part.
So the term isomer means the same parts.
And two major kinds of isomers in organic chemistry are constitutional isomers, which are more common, and geometric isomers.
Constitutional isomers, also called structural isomers, are where two molecules have the same number and types of atoms as each other.
But the connections between the atoms can be super different, like we just saw with the oxygen atom in DNA.
Let’s use a more straightforward example, though.
In the first episode of this series, we mentioned octane and iso-octane, two components of gasoline.
They’re constitutional isomers, because they both have 8 carbons and 18 hydrogens.
In fact, that’s how iso-octane got its name… past chemists just stuck on the prefix “iso” to mean “same as octane.”
But their structures are pretty different.
In octane, the carbons are attached in a long chain without any branches, so it checks out as an IUPAC systematic name.
Iso-octane is branched, though, and if we’re going by IUPAC rules, it isn’t octane at all!
It has 3 carbon-chain substituents and we rule-followers would call it 2,2,4 trimethylpentane.
On the other hand, geometric isomers have the same number and types of atoms and the same connections between them.
But these compounds differ in how the bonds are spatially arranged.
To sort of visualize this, I can put the tips of my index fingers together and rotate one hand as far as I can.
See, I can twist one hand about 180 degrees without breaking contact.
And if my fingers were a single bond and my hand-atoms weren't held back by my arms, I could rotate it a full 360 degrees.
This is known as free rotation, when atoms can completely rotate around the axis of a bond.
The carbons on an ethane molecule can do this!
So ethane doesn’t have any geometric isomers.
However, if I touch two fingertips from each hand, I can’t rotate one hand without breaking a connection.
So a double bond between atoms doesn’t have free rotation, and molecules with double bonds can have geometric isomers that are spatially different.
Take, for example, the simple alkane pent 2-ene.
In one geometric isomer, the ethyl group and the methyl group are on the same side of the double bond, which we use the prefix cis to describe.
But in the other, these groups are opposite each other, which we use the prefix trans to describe.
For more complex alkenes and to stick with our trusty IUPAC rules, we use the prefixes E and Z.
But don’t worry about that for now, we’ll learn about E and Z nomenclature in another episode.
All this to say, electron orbitals and atomic bonds determine the shape of molecules, which determines what they can do, which determines basically every chemical reaction that keeps us and our universe going.
Even though it can be a little brain-bendy to imagine tiny molecules in 3D like you would a fluffy cat, without them, we wouldn’t exist to pet cats or watch YouTube videos
or take organic chemistry tests!
In this episode, we learned that:
Orbital hybridization and valence bond theory can help us explain 3D molecular structures,
Constitutional isomers have the same atoms but different atom-to-atom connections,
And geometric isomers have different spatial arrangements.
Next time, we’ll learn some techniques to understand isomers and atom connections even better.
Thanks for watching this episode of Crash Course Organic Chemistry.
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