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Hi! I’m Deboki Chakravarti and welcome to Crash Course Organic Chemistry!
Before we get too deep into chemical reactions and equations and graphs, let's take a deep breath…
Did you smell anything?
Maybe the damp soil of a houseplant or musty curtains or whatever your roommate microwaved for lunch?
Almost everything you smell is an organic chemical.
Sometimes we can even smell whether or not a reaction has taken place.
Take lemons, for example.
A major component of that bright lemony smell is limonene.
Limonene has two double bonds in its structure, and enzymes can catalyze a reaction to make just one of them add a water molecule.
The product of this enzyme-catalyzed reaction, called alpha-terpineol, smells like lilacs instead.
So how do reactions like these happen and why don’t lemons smell like lilacs if you leave them out for a while?
To answer questions like these with organic chemistry, we need to think about concepts like reaction kinetics, thermodynamics, spontaneity, and free energy.
So in this episode, we’ll learn how different reactions can move at vastly different speeds and how not everything that could happen actually does.
[Theme Music]
In order to have a really useful chemical reaction, two things need to be true:
there should be more products than reactants when the reaction is done, and the reaction must take place at a reasonable speed.
We use two terms to describe these conditions:
thermodynamics for the energy and reaction progress, and kinetics for how fast a reaction happens.
You might remember from general chemistry that thermodynamics involved all these different terms like ΔH, ΔG, and ΔS.
We'll need these ideas to understand organic chemical reactions, so let's review.
First, ΔH is enthalpy, which means the change in heat of a reaction at constant pressure.
A good way to understand enthalpy is with chemical hot and cold packs that you might use for an injury.
A hot pack gets hot because the products of the reaction: the dissolved salt ions in solution
have lower energy than the reactants: the undissolved salt.
This is called an exothermic process because heat is given off by the reaction, warming the pack, and the sign of ΔH is negative.
On the other hand, a cold pack pulls in heat, usually from our bruised knee, so it feels cold to the touch.
The products have higher energy than the reactants.
This is called an endothermic process because heat moves into the pack, and the sign of ΔH is positive.
Second, ΔS is entropy, which is the idea that the world is always moving toward more chaos.
Entropy explains why a sugar cube dissolves in tea, or why a marching band needs to work hard to keep a sharp formation instead of dissolving into a random noise-making crowd.
For example, entropy increases and ΔS is positive when one reactant splits up into two products, or when a ring compound opens into a more wiggly chain.
The relationship of enthalpy, entropy, and temperature, which we represent with T in units of Kelvin, tells us about another important term:
ΔG or Gibbs free energy.
This allows us to predict whether a reaction occurs spontaneously or whether we have to use force to get the reaction to move along.
The sign of ΔG tells us the most important thing about our reaction: whether it happens spontaneously or not.
A negative ΔG means a reaction is spontaneous, and may have energy left after reacting to do work, like push a piston or explode.
And a positive ΔG means a reaction is nonspontaneous and we have to force it to happen somehow.
The relationship between these four terms can be expressed mathematically:
ΔG equals ΔH minus T times ΔS.
Finally, we need to be familiar with K, the equilibrium constant, which we talked about in episode 11 when we talked about acidity.
We called this term Ka or the acid dissociation constant, but all Ks are defined the same way:
the amount of products divided by the amount of reactants when the forward and reverse reaction rates are equal.
As a quick refresher, if K is large, the reaction mixture is mostly products.
And if K is small, the reaction mixture is mostly reactants.
To relate K to all the other terms, we have a second equation:
ΔG equals negative R times T times the natural log of K.
Basically, this equation tells us that a negative ΔG, which means a reaction is spontaneous, will have a large K.
So we know the opposite relationship too: a positive ΔG will have a small K, and the reaction is nonspontaneous.
It can be tough to understand all these relationships just from equations, but we can visualize them with energy diagrams.
The y-axis represents the energy of the reaction, and that's usually represented by ΔH or ΔG.
For this episode, we'll just think about the enthalpy of the reaction, ΔH, for simplicity.
The x-axis of an energy diagram represents the reaction progress, a line that starts with reactants on the left and moves towards products on the right.
The peak of the hill in between reactants and products is the transition state.
Not all hills are perfect cartoon parabolas that start and end at the same flat grassy level.
We have to pay attention to how high or low the reactants and products are in an energy diagram, because that tells us even more about the reaction.
If the products end up below the reactants, they are lower energy, so ΔH is negative.
This kind of chemical reaction can give off heat, like our hot pack, and the reaction is considered exothermic.
If the products end up higher than the reactants, they are higher energy and the reaction needs outside energy to happen, so ΔH is positive.
This kind of chemical reaction needs heat to flow in so it can make the products, like our cold pack, so it's considered endothermic.
Sometimes, just being at room temperature is enough extra energy to make it happen.
The difference in energy between where the reactants start and the top of the hill is called the activation energy, which is the energy required to get the reaction started.
But here's the thing: you can climb a hill from both directions, not just one.
So we can consider the height of the hill from the forward direction, reactants to products, and the reverse direction, products to reactants.
And activation energy can help us think about equilibrium.
Climbing a hill isn't easy, no matter how much you go hiking!
So a big reverse activation energy favors the products.
Once the molecules get to the other side, they want to chill in that low point instead of putting in all the effort to climb to the top again.
And a big forward activation energy favors the reactants because not all of them will have enough oomph to make it up and over the top.
Using the same hiking analogy, bigger hills with larger activation energies generally mean a slower chemical reaction.
I remember this because a big hill takes a lot more effort to hike up and over, while smaller hills are faster to climb.
No matter what, though, it takes energy!
For example, let's take a look at the reaction between but-1-ene and hydrogen bromide, which can form a primary or secondary carbocation.
We know that high-energy double bonds break to form the carbocations, and we get a new bond to hydrogen.
Both of these reactions are endothermic,
and it’s important to remember that it's the balance of making and breaking bonds that determines the overall enthalpy of the reaction.
But the energy diagram to form the primary and secondary carbocations are slightly different.
We learned in episode 14 that induction and hyperconjugation make the secondary carbocation more stable than the primary carbocation.
And from this energy diagram, we can conclude that the more stable secondary carbocation takes less energy to form because the hill is easier to climb!
So far, we've been exploring energy diagrams as one lonely hill with reactants on the left and products on the right.
But, as we've started to see in previous episodes, reactions in organic chemistry can be whole journeys with mountains and valleys and multi-step reactions.
So let's try working our way through a big reaction, for practice, and see what its energy diagram looks like.
Just for fun, we'll do a reaction between 2-methylprop-2-ene and methanol, which forms MTBE or methyl tert-butyl ether.
You might've heard of it as the controversial anti-knock gasoline additive, which basically means it adjusts how combustion works in engines.
If your brain's like mine, you might notice the sulfuric acid above the reaction arrow and go,
"Oh! That's a super strong acid, which means it completely dissociates in water"
but we don't have water here, we have methanol instead.
So our first instinct to do something with sulfuric acid is a good one, and it can transfer a proton to methanol.
With that in mind, let's rely on our tried-and-true logic for addition reactions:
our nucleophilic alkene is gonna attack the most acidic proton in the solution.
That's the proton on methanol!
Without the sulfuric acid to get the party started, the reaction between 2-methylprop-2-ene and methanol would be super slow,
almost to the point we might say they don’t react.
Anyway, with that attack, 2-methylprop-2-ene forms the more stable tertiary carbocation, and we’ve climbed over the first hill on our journey.
This carbocation is now a prime electrophilic target, and it’s still reactive even though we're sitting in a mini-valley and taking a break.
So next, the methanol comes in with its lone pairs of electrons, does a nucleophilic attack, and makes a new carbon-oxygen bond.
This forms an oxonium ion, we've climbed over the second hill, and we can take another trail mix break.
Finally, methanol swoops back in, deprotonates the hydrogen on the positive oxygen,
the electrons from the former O-H bond neutralize the positive charge on oxygen, and we tumble downhill as our final products form.
Our trek is done!
Looking at the full reaction mechanism from beginning to end, the snack breaks as we hike between the reactants and the products are called intermediates.
And the energy of the intermediates influences the kinetics and the thermodynamics of a reaction.
Intermediates have full charges and full bonds, while transition states on the hilltops have partial charges (drawn as lower case Greek deltas)
and partial bonds between reactants (drawn as dotted lines).
Now that we've made it through that beast of a reaction,
I have one last thing to add about the protonated methanol that we made using sulfuric acid that was so handy at the beginning.
We remade it at the end of the reaction, so it's a catalyst, something that speeds up a reaction but isn’t used up.
In an energy diagram, a catalyst flattens the activation energy hill.
We’ll be talking about catalysis a lot throughout this course, so consider this a sneak peek!
The reaction I mentioned at the beginning needs a catalyst to make the energy hill climbable,
which is why lemons don't just spontaneously start smelling like lavender.
In this episode, we did a lot of deep thinking about energy and equilibrium, which will help us predict even more chemical reactions.
Specifically, we:
Reviewed thermodynamics and enthalpy, entropy, and Gibbs free energy
Drew free energy diagrams of chemical reactions
Learned the difference between an intermediate and a transition state
And explained how catalysts lower the activation energy of a reaction
In the next few episodes, we’ll look more at the types of addition reactions that involve alkenes.
We’ll do our best to group these reactions by mechanism, so we can continue to puzzle through, and NOT memorize, the products that form.
Until then… thanks for watching this episode of Crash Course Organic Chemistry.
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