Chapter 6: Chemical Reactivity and Mechanisms

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Welcome to the Deep Dive.

We sift through the sources to give you the key insights fast.

Today we're really plunging into the heart of chemical reactions,

reactivity mechanisms,

we're drawing heavily from Capture Six of David Klein's organic chemistry text.

Our goal is to give you a kind of shortcut, a high level view, to really get the why and how behind reactions.

Energy, electrons moving,

the skills you need.

And to kick things off, think about this, Alfred Nobel, famous for the Nobel Prize, well he made his fortune developing dynamite.

Yeah, it's a pretty explosive start, isn't it?

But it's a great connection because those fundamental principles behind dynamite, they're exactly the same principles driving every chemical reaction, whether it's super fast, like an explosion, or, you know, slow and steady in a biological system.

Exactly.

So let's unpack that.

Let's start with why reactions even happen.

That takes us into thermodynamics.

Right, thermodynamics.

It's all about energy levels, reactants versus products.

It tells you essentially what's the best possible outcome, the maximum yield you can expect.

And the first big player here is enthalpy, symbol AH.

Most people think of it as heat, basically.

The exchange of heat energy between the reaction, the system, and everything else, the surroundings, assuming constant pressure, of course.

And we remember the basics.

Breaking bonds costs energy.

You put energy in.

But forming bonds, that releases energy.

We also talk about how bonds break,

homolytic cleavage, that's even splitting, gives you radicals, little fishhook arrows show that.

But then there's heterolytic cleavage, uneven.

One atom takes both electrons, makes ions.

We use the normal two -headed arrows for that.

Precisely.

And the energy needed for that even split the homolytic kind, that's the bond association energy, or BDE.

We usually measure it under standard conditions, so we write 80 degrees.

For a whole reaction, the overall 80 degrees is just while you add up the energy for all the bonds you break and subtract the energy released when you form the new ones.

And the sign tells you if it's releasing heat or absorbing it, from the system's point of view, remember?

So negative H degrees means exothermic, energy released.

Positive H degrees means endothermic, energy absorbed.

Yeah, that's sign convention.

It sometimes tricks people up, especially for their thinking physics, where the focus might be on the surroundings.

But in chemistry, negative H means the system, the chemicals, lost energy.

Got hotter outside, sure, but the chemicals have less energy.

So calculating 80 degrees with BDEs, like for, say, tert -butane reacting with chlorine, it's just bookkeeping, really.

Cost of breaking minus payoff of forming.

Very practical.

But, and this is a big but, that always isn't the whole picture.

It doesn't explain everything.

You see lots of endothermic reactions that absorb heat happen just fine, spontaneously even.

Which leads us to the second factor, entropy, SS.

Ah, entropy, often called disorder, right?

Yeah, that's the common shorthand.

But it's maybe more accurate to think about it in terms of probability or the number of ways energy can be spread out in a system, the number of possible arrangements or microstates.

Like gas expanding.

If you open a valve between two chambers, the gas spreads out.

It doesn't just continuously squish back into one corner later.

Exactly, and why not?

Not because of energy change, really.

It's because there are just vastly more ways for the gas molecules to be spread out across both chambers than crammed into one.

More arrangements means higher entropy, higher probability.

Nature tends towards the most probable state.

So for something to be spontaneous, the total entropy has to go up.

That's the key.

The total entropy system plus surroundings.

AA total equals A system plus AS surroundings.

That has to increase.

And here's a neat point.

The system's entropy can decrease, like when something gets more ordered, as long as the surroundings entropy increases even more to compensate.

Okay, so what makes system entropy, HS system, go up?

Generally, two main things.

One,

making more molecules.

Like one reactant breaking into two products, more particles, more ways to arrange them.

Two, increasing freedom of movement.

Think about breaking open a ring structure, and a cyclic molecule can wiggle and rotate much more freely than a constrained ring.

More conformational freedom, higher entropy.

Right, which leads us neatly to the big one, the one that combines both.

It gives free energy A.

Yes, the ultimate decider for spontaneity.

The equation is Elegant HE,

AAAHD.

That H term, it actually relates to the entropy change of the surroundings.

And the T's term, that's directly about the entropy change of the system.

It puts it all together.

And the rule is simple.

The rule is simple.

For a reaction to be spontaneous, T must be negative.

We call that exergonic.

If day is positive, it's endergonic, non -spontaneous.

Needs energy input to go.

And that T in the equation, temperature, that's important too.

Hugely important.

Temperature directly scales the entropy contribution, T s.

So if H and T day are pulling in opposite directions, say, an endothermic reaction that increases entropy, then temperature can be the deciding factor whether D ends up positive or negative.

And this connects back beautifully to explosives, doesn't it?

Perfectly.

Explosives like nitroglycerin or TNT are engineered for a massively negative day.

Why?

Well, they're usually highly exothermic, so age is very negative.

Big energy release.

But also, they produce enormous volumes of gas from a small amount of liquid or solid.

That's a huge increase in entropy that's very positive.

So you get a favorable push from both enthalpy and entropy, making day extremely negative.

Boom.

Okay.

So Kate tells us if it can happen, but even spontaneous reactions don't just instantly finish, right?

They reach a balance point.

Exactly.

They reach equilibrium.

That's the point where the system has reached its lowest possible free energy.

It's stable.

The forward and reverse reactions are still happening, but their rates are equal, so there's no net change in the amounts of reactants and products.

And we quantify that balance with Keck, the equilibrium constant.

Right.

Keck is just the ratio of products to reactants at equilibrium.

Tells you which side is favored.

And D and Keck are directly linked.

Directly.

Through the equation, G equals RT ln Keck.

R is the gas constant.

T is temperature in Kelvin.

What this shows is powerful.

If Ag is negative, spontaneous, then ln Keck must be positive, meaning Keck is greater than one.

Products are savored.

If Ag is positive, non -demontaneous.

Keck is less than one.

Reactants are favored.

And because of that logarithm, even a small change in Keck can mean a huge change in Keck.

A big shift in the equilibrium position.

So thermodynamics tells us the destination, the maximum possible yield based on G and Keck.

That's it in a nutshell.

It's about the relative stability and the equilibrium balance.

Now you mentioned biological systems earlier.

People sometimes get confused about life and the second law of thermodynamics.

This idea that entropy always increases.

Life seems so ordered.

Ah, yes.

A classic point of confusion.

But life absolutely does not violate the second law.

Highly ordered biological processes like building proteins or DNA are decreasing entropy locally within the organism.

But they are always, coupled to other reactions that are massively favorable and increase entropy much more.

Like breaking down food,

metabolism.

Precisely.

Think about metabolizing glucose.

It's highly exergonic, releases heat, increasing surroundings entropy, and breaks down a complex sugar into simpler molecules like CO2 and water, increasing system entropy.

The overall entropy change for the whole universe is definitely positive.

Living things are incredibly efficient at generating entropy overall, even while creating local order.

We're basically entropy machines.

Entropy machines.

I like that.

Okay, so we've covered the why thermodynamics, but that doesn't tell us how fast things happen, which brings us to the other side of the coin.

Kinetics.

Speed.

Right.

And this distinction is absolutely crucial.

Thermodynamics asks, will it go?

Kinetics asks, how fast will it go?

And here's where it gets really interesting.

Think about diamonds.

Diamonds spontaneously turn into graphite.

Thermodynamically speaking, graphite is more stable under normal conditions.

G for that conversion is negative.

But my diamond ring isn't turning into pencil lead.

Exactly.

Because kinetically, that reaction is unbelievably slow.

The energy barrier to get it started is enormous.

So just because something is spontaneous, negative G, doesn't mean it happens at any noticeable rate.

Speed and spontaneity are different things.

Got it.

So kinetics studies the factors that affect that rate.

Yes.

And we often express the rate mathematically with a rate equation.

Usually something like rate T times the concentrations of reactants may be raised to some power.

Avoid X, B, Z, that kind of thing.

That's the one.

Senki is the rate constant.

It bundles up all the factors affecting the rate except concentration.

Things like temperature, catalysts.

We'll get to those.

The exponents, X and Y, tell you how sensitive the rate is to the concentration of each reactant.

You have to So the rate of reaction of each of those exponents gives the overall reaction order, first order, second order.

Correct.

Now what affects that weight constant and kinky?

Three main things.

Number one, and arguably the most important, the energy of activation.

Ele.

The energy barrier, the hump on the energy diagram.

Exactly.

It's the minimum amount of energy colliding molecules need to actually react and form products.

A low Ea means it's easy to get over the hump.

Lots of molecules have enough energy, so the reaction is fast.

Hyeta, slow reaction.

Makes sense.

Factor number two.

Temperature.

Turn up the heat, molecules move faster.

They collide more often, sure.

But more importantly, a higher temperature means a much larger fraction of those molecules have enough energy to overcome the accurate barrier.

Right.

That's why reactions generally speed up when heated.

There's that rule of thumb, isn't there?

Like 10 degrees Celsius increase doubles the rate.

Roughly, yes.

It depends on the Ea, but it's a useful guideline.

Higher T, faster rate.

And the third factor.

Steric considerations.

Basically, the shape and orientation of the molecules when they collide, even if they hit with enough energy, they might just bounce off if they aren't oriented correctly for bonds to break and form.

Geometry matters.

Okay.

Ea, temperature, sterics.

Which brings us to things that help reactions along.

Catalysts and enzymes.

Yes.

Speed demons.

A catalyst speeds up a reaction without being used up itself.

How?

Critically, it provides a different reaction pathway, an alternative route with a lower energy of activation, a lower Ea.

So it doesn't change the start or end points, doesn't change the Ea.

Absolutely not.

A catalyst has no effect on the overall thermodynamics or the equilibrium position, Kevin Keck.

It just helps the reaction reach equilibrium faster.

And enzymes are just biological catalysts, right?

Nature's highly specialized helpers.

Exactly.

Proteins, mostly, folded into specific shapes to bind reactants, we call them substrates, and lowered the Ea for specific biochemical reactions.

Life depends on them.

You know, a practical example could be brewing, making beer or wine.

Perfect example.

Fermentation turning sugars into ethanol is thermodynamically favorable.

B is negative.

But without yeast, it's incredibly slow.

The Ea is high.

Yeast contains enzymes that provide that low Ea pathway.

So you add yeast and boom, fermentation takes off.

The enzymes are the catalysts.

It's amazing how these concepts play out everywhere.

And circling back to nitroglycerin?

Yes, Nobel's challenge.

Nitroglycerin itself is super unstable, very low Ea, a primary explosive goes off easily, too dangerous.

Nobel's big innovation with dynamite was mixing nitroglycerin with diatomaceous earth, a stabilizer.

This essentially raised the Ea needed to initiate the explosion, making it a safer secondary explosive.

You needed a blasting cap.

But the medical connection is just wild.

Isn't it?

Factory workers handling it reported relief from angina, chest pain.

Decades later, Louis Ignarro and others figured out why.

The body metabolizes nitroglycerin to produce nitric oxide, NO.

And NO is a signaling molecule that relaxes blood vessels, easing chest pain.

Ignarro won a Nobel Prize for this discovery.

Which led to drugs like Viagra also working via related pathways.

Right.

And the supreme irony, Alfred Nobel himself apparently refused to take nitroglycerin for his own heart condition.

Wow.

And the prize Ignarro won was funded by Nobel's dynamite fortune.

It really ties together thermodynamics, kinetics, stabilization,

medicine.

Amazing.

It really does.

Okay.

Let's shift gears slightly.

Let's talk about how we actually visualize and describe these reaction steps.

Energy diagrams and arrow pushing.

The language of mechanisms.

Energy diagrams, as we touched on, give you that visual map.

The peaks are energy barriers.

EA, that's kinetics.

How high the peak is tells you how fast that step is.

The overall energy difference between the start reactants and the end products, that's A.

Thermodynamics tells you if it's favorable overall.

And on these diagrams, we see peaks and valleys.

Transition states and intermediates.

What's the difference again?

Good question.

Transition states are the absolute energy maxima, the very top of the humps.

They're fleeting, unstable, can't be isolated.

Bonds are in the process of breaking and forming simultaneously.

Think of it like the peak of a jump.

You can't just hover there.

Okay.

So transition states are peaks.

What about the dips, the valleys between peaks?

Those are intermediates.

They are actual chemical species, often reactive ones like carbocations or carbanions that exist for a finite, though maybe very short, lifetime.

They're local energy minima.

You can, in principle, sometimes isolate or detect them.

Think of standing on a step during a climb.

You can pause there for a bit.

And there's a way to predict what the transition state looks like.

Yes, the Hammond postulate.

It's a really useful guideline.

It says that the structure of a transition state resembles the species, reactants, or products intermediate that it's closer to in So for an exothermic step where the products are lower energy.

The transition state will be closer in energy to the reactants, so it will look more like the reactants.

We call it an early transition state.

And for an endothermic step, products higher energy.

The transition state will be closer in energy to the products, so it will resemble the products more.

A late transition state.

It helps us visualize that fleeting moment.

Okay, cool.

Now, the actual players in most organic reactions.

Nucleophiles and electrophiles.

The stars of the show.

About 95 % of the reactions you'll see in intro organic chemistry involve these guys.

It's all about electron -rich things meeting electron -poor things.

So nucleophiles first.

Nucleous lovers.

They're electron -rich.

Exactly.

They have electrons they want to donate.

They act as Lewis bases.

Look for atoms with lone pairs like the oxygen in an alcohol or an alkoxide or pi bonds like in alkanes or alkynes.

Those are your electron sources.

And some are stronger than others.

Definitely.

Factors like charge matter but also polarizability.

Bigger atoms like iodine or sulfur, HS, hold their outer electrons more loosely.

Their electron clouds are more easily distorted, making them really good nucleophiles.

Got it.

And the counterpart.

Electrophiles.

Electron lovers.

They're electron -poor.

Yep.

They want electrons.

They act as Lewis acids.

Look for atoms with a partial positive charge.

Plus, often due to an electronegative atom pulling electron density away like the carbon in methyl chloride, CH3Cl.

The chlorine pulls electrons, making the carbon electrophilic.

Or look for atoms with a full positive charge and often in an empty orbital like carcations.

That positive carbon with its empty p -orbital is practically screaming for electrons.

So identifying these centers.

That's really the key skill.

I would argue it's the single most important skill for understanding mechanisms.

If you can spot the nucleophile, electron source, the electrophile, electron sink, you know where the reaction is likely to happen.

You know which way the electrons will flow.

How do you spot them quickly?

Practice.

But you're scanning the molecule for clues.

Lone pairs.

Pi bonds.

Those potential nucleophiles.

Partial positive charges from induction.

Full positive charges.

Empty orbitals.

Those are potential electrophiles.

Okay, so once you've found the nucleophile electrophile, you show the reaction using arrow pushing.

The language itself.

Curved arrows show the movement of electron pairs.

The tail of the arrow always starts where the electrons are coming from, either a lone pair or a bond.

Usually a pi bond or a bond breaking.

The head of the arrow points exactly where those electrons are going, usually to form a new bond or sometimes onto an atom to become a lone pair.

And a really common mistake to avoid.

Oh, absolutely.

Never ever start the of a curved arrow on a positive charge.

Electrons flow from areas of high density to low density.

Positive charges are low density.

They accept electrons.

They don't donate them.

Okay.

And you mentioned there are only a few fundamental patterns for these arrows.

Just four main patterns build almost everything in ionic mechanisms.

Master these and you're set.

Pattern one.

Nucleophilic attack.

The nucleophile uses its electrons to form a new bond with the electrophile.

This might be just one arrow or sometimes a second arrow is needed if the electrophile needs to break a bond to make room.

Like pushing pi electrons onto oxygen in a carbonyl attack.

Okay.

Pattern two.

Loss of a leaving group.

An atom or group takes its bonding electrons and departs.

Often happens after or during another step like nucleophilic attack.

The group has to be stable on its own with those extra electrons.

A good leaving group.

Pattern three.

Proton transfers.

Moving H plus sign.

We saw this in acid -base chemistry.

Remember, it always involves two arrows.

One arrow from the base, nucleophile, grabbing the proton, H plus.

A second arrow from the bond holding the proton showing those electrons going back onto the atom the proton was attached to.

Right.

Don't just show the H plus falling off.

A base has to take it.

Precisely.

That's another common error for getting the base or the second arrow.

And the final pattern, number four.

Rearrangements.

Specifically carbocation rearrangements.

These deserve special attention.

Because carbocations aren't all equally stable, right?

Not at all.

Stability increases with more alkyl groups attached to the positive carbon.

Tertiary, three alkyl groups, is more stable than secondary.

Two groups.

Which is way more stable than primary.

One group.

Methyl is the least stable.

Why are more alkyl groups better?

Hyperconjugation.

It's a stabilizing effect where the electrons in adjacent CH or CC sigma bonds can overlap slightly with the MTP orbital of the carbocation.

More adjacent bonds, more overlap, more stabilization.

So carbocations rearrange to become more stable.

Exactly.

If a carbocation can rearrange to become more stable, like a secondary becoming a tertiary, it usually will.

The two common ways this happens are a hydride shift, an H atom with its two bonding electrons moves over, or a methyl shift, a CH3 group with its electrons moves over.

The driving force is always getting to that lower energy, more stable carbocation.

Always.

So even really complicated looking mechanisms are just combinations of these four patterns.

Nucleophilic attack, loss of leaving group, proton transfer, rearrangement.

That's the beauty of it.

They're the building blocks.

Sometimes two patterns happen at the exact same time.

We call that a concerted process.

But the fundamental electron movements are described by those four patterns.

Okay.

Drawing the arrows correctly seems critical.

Any other pitfalls?

The big one.

Violating the octet rule for second row elements.

Carbon, nitrogen, oxygen, fluorine.

These elements cannot have more than eight electrons in their valence shell.

That means no more than four bonds or combinations of bonds and lone pairs adding up to eight electrons.

Never draw an arrow that forces one of these atoms to exceed the octet.

It's probably the most common mistake students make when first learning mechanisms.

Double check your resulting structures.

Good warning.

So every arrow must achieve one of the four patterns and don't break the octet rule.

That's the mantra.

Now, predicting those carbocation rearrangements, you said they happen if they lead to more stability.

Secondary to tertiary, for example.

Yes, that's the main driver.

Do tertiary ones ever rearrange?

Generally, no, because they're already quite stable.

The main exception is if a rearrangement, even from tertiary, could lead to a resonance stabilized carbocation, like forming an allylic carbocation, positive charge next to a double bond, or a benzylic one next to a benzene ring.

Resonance offers significant extra stability, so that can be a driving force even for rearranging a tertiary carbocation.

Okay, one last detail on mechanisms.

The arrows we use in the overall reaction scheme.

Sometimes they're one way, sometimes equilibrium arrows.

How do we choose?

Ah, reversible versus irreversible arrows.

They aren't just stylistic choices.

They tell you about the thermodynamics of that specific step.

So, for nuclear philic attack?

It depends.

If the nucleophile, after attacking, could also function as a good leaving group, like water or halate, the step is likely reversible.

Use equilibrium arrows.

But if the nucleophile is very basic and forms a strong bond, making it a poor leaving group, like a green yard regent, RMGX acting as R, or a hydride ion H, then the attack is usually considered irreversible.

Single arrow.

How about loss of a leaving group?

Often reversible.

Good leaving groups are typically stable anions or neutral molecules, which can often act as nucleophiles themselves to reverse the reaction.

Equilibrium arrows are common here.

Proton transfers.

Technically, all proton transfers are reversible.

But practically speaking, if the p -CoI difference between the acid and the conjugate acid is really large, say, more than 10 p -CoI units, the equilibrium lies so far to one side that we often draw it as irreversible for simplicity.

If the p -CoV values are closer, definitely use equilibrium arrows.

And carbocation rearrangements.

Irreversible.

Generally, yes.

We usually draw them with a single arrow.

The stability difference between, say, a secondary and a tertiary carbocation is significant enough that the reverse rearrangement, going from more stable to less stable, is highly unfavorable and usually negligible.

Are there exceptions?

Like, things that drive reactions even if the step isn't super favorable?

Sure.

A classic is if a reaction step produces a gas, like CO2 or N2, that bubbles out of solution.

Le Chatelier's principle tells us removing a product shifts the equilibrium to the right.

So even if a step isn't inherently strongly favored, gas evolution can make it effectively irreversible.

Okay.

Wow.

That's a lot, but it feels like a complete picture.

So wrapping this up, what does this all mean for you, the listener?

We've really dug into the core principles, the why thermodynamics,

thinking about energy changes, probability and disorder, the ultimate decider, and where the reaction settles, equilibrium, Keck.

Then the how fast kinetics, understanding the energy barrier, EA, the role of temperature, and how catalysts, including enzymes, provide shortcuts by lowering that barrier.

And maybe most crucially for organic chemistry, the actual language of reactions, spotting the electron -rich nucleophiles and electron pore electrophiles.

And using curved arrows precisely to show electron flow according to those four fundamental patterns, nucleophilic attack, loss of leaving group, proton transfer, and those important carbocation rearrangements.

Mastering these ideas, these patterns,

it genuinely unlocks the ability to understand and predict a huge range of organic reactions, like learning the rules of the game.

It really is.

It stops being about memorizing reactions and starts being about understanding why they happen the way they do.

Exactly.

And hopefully you've seen how concepts connect everything from, you know, Nobel's dynamite to life -saving drugs like nitroglycerin, even to making beer.

It's all the same fundamental dance of electrons and energy.

So maybe as you go about your day, you'll start seeing these principles at play.

Think about how tiny changes at the molecular level drive everything we see and do.

Perhaps dig a bit deeper into one of the examples we talked about.

Yeah, keep that curiosity going.

Thank you so much for joining us for this Deep Dive.

We really appreciate you exploring these fundamental ideas with us.

It's been a pleasure.

Thanks for tuning in.

Until next time, keep diving deep.