Chapter 3: Speaking with Pictures: Drawing Structures

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Welcome curious minds to another deep dive.

Today we're embarking on a journey into, well, what I consider the secret language of chemistry.

That's a great way to put it.

If you've ever felt that organic chemistry is like learning a foreign tongue, you're not wrong.

It really has its own vocabulary, its own grammar, and crucially, its own unique way of speaking with pictures.

We're going to learn those words today.

The structures that represent the molecules, you know, the ones too tiny to actually see.

And to guide us, we're drawing directly from the excellent organic chemistry for dummies seconded.

And that analogy really hits home.

I mean, understanding these visual representations isn't just about, you know, acing a chemistry exam.

It's about comprehending the unseen, the fundamental building blocks of, well, everything around us.

Molecules behave in surprisingly bizarre ways, ways that sort of defy our everyday intuition.

How so?

Well, think about quantum phenomena, right?

Articles that seem to like tunnel through barriers or exist in multiple states at once.

Our macroscopic brains can't easily grasp this stuff.

Okay, yeah.

So these models, like the Lewis structure we'll talk about, they're less about seeing the molecule in a literal sense and more about giving us a predictive language, a fundamental map.

It's what chemists use to design new pharmaceuticals or create advanced materials or even just unravel how enzymes work inside us.

Okay, let's unpack this then.

Our mission today is to get a quick but thorough understanding of the fundamental principles of drawing these structures.

Yep.

We'll look at the different ways chemists represent them, the grammar of electron movement, which sounds fascinating, and how to correct for, well, structural flaws using something called resonance.

Resonance, exactly.

Essentially, we're building our foundational vocabulary and sentence structure to truly speak organic.

So let's start with the very foundation, the most complete way to articulate these molecular words,

the Lewis structure.

The blueprint.

Exactly.

Think of it as the chemist's detailed blueprint.

It shows you which atoms are connected and, crucially, where all the electrons actually live.

We represent shared electrons, the ones forming bonds with lines.

Right, the sticks.

Yeah, the sticks.

One line for a single bond, that's two shared electrons.

Two lines for a double bond, meaning four shared, and three for a triple bond, six shared.

And for electrons that aren't shared, what we call non -bonding electrons or lone pairs, they're just indicated with simple dots right on the atom.

And what's really fascinating here, maybe crucial is a better word, is how quickly you need to be able to identify formal charges on these Lewis structures.

Formal charges, okay.

Why?

Well, electrons, of course, are negatively charged.

So if an atom in a structure has too few electrons compared to its neutral state, its normal state, it'll end up with a positive charge.

Makes sense.

And if it has too many, it'll be negative.

There's a super quick and dirty equation to figure this out.

It's really just an electronic counting method.

Okay, give it to us.

You take an atom's natural valence electrons from the periodic table, then you subtract every non -bonding dot you see around it in the structure, and finally you subtract every stick or bond connected to it.

So formal charge, valence electrons,

dot sticks, simple as that.

Okay, let's try an example, maybe one from the source material.

The amide ion, NH2.

So nitrogen, it's in the fifth column of the periodic table, so it has five valence electrons.

In NH2, the Lewis structure shows it has four dots, that's two lone pairs and two sticks, meaning two single bonds, one to each hydrogen.

So plug that into the formula.

Five valence electrons minus four dots minus two sticks equals negative one.

So the nitrogen in the amide ion carries a negative one formal charge.

Okay, that calculation is precise, but as you mentioned, there are ways to speed this up.

To recognize patterns almost instantly, what are some of those common tells for a charged atom?

That's a really powerful shortcut you'll develop with practice.

It becomes almost second nature.

You'll soon be able to just look at a structure and almost instantly tell if an atom is charged.

For example, a neutral carbon will always have four bonds, always.

That's its valency, we say it's the traveling.

It typically wants to form four bonds.

So if you see a carbon with only three bonds shown, you immediately know it's either positively charged if it lacks a lone pair or negatively charged if it has a lone pair filling that missing spot.

Ah, I see.

Similarly, neutral nitrogen is trivalent, three bonds and one lone pair, always.

Neutral oxygen is divalent, two bonds, two lone pairs.

And halogens,

you know, fluorine, chlorine, bromine, iodine,

they're monovalent, one bond and three lone pairs when neutral.

So recognizing those standard counts.

Exactly.

Recognizing these typical valencies lets you quickly spot atoms that don't fit the neutral pattern and therefore must carry a charge.

You don't have to do the tedious calculation every single time.

It's kind of like learning the standard look of a word to spot a misspelling faster.

That makes perfect sense.

So, okay, those full Lewis structures give us every detail, which is great, but I can totally imagine drawing them for really large molecules would be incredibly tedious.

Oh, yeah.

Impractical.

So do chemists have a shorthand, like, you know, texting shorthand, but for these molecular words?

Absolutely.

We actually have a few different levels of abbreviation.

First, there are what we call condensed structures.

Condensed.

Here, the bonds between carbon and hydrogen aren't explicitly drawn.

Instead, the carbon and its attached hydrogens are just grouped together, like writing CH2 or CH3.

Right.

The carbon -carbon bonds might still be shown with lines, or sometimes they're just implied by adjacency.

For instance, but dinonone can be shown as CH3CH2CH3.

Okay, I see the groups there.

And if you have identical groups, like in diethyl ether, you can even use parentheses with a subscript, like CH3CH2TO.

That tells you there are two identical ethyl groups attached to the oxygen.

Oh, that's handy.

Or for really long chains, we could abbreviate repeated units, like CH3CH2 5CH3 for heptane.

That CH2 -5 means five CH2 units in a row.

Clever.

These condensed structures are generally best for straight chains, though.

They can get a bit clunky or ambiguous with rings or highly branched molecules.

But the most common method, really the workhorse, is the line bond structure, sometimes called skeletal structure.

Line bond.

Okay, this sounds important.

It is.

These are incredibly efficient.

The core rule is, well, pretty simple.

Each point or vertex and each end of a jagged line is assumed to be a carbon atom.

Just assumed.

Yep.

And here's the other key part.

Hydrogens attached to these carbons are not explicitly shown.

They're just implied.

Implied?

How?

The assumption is that you, the chemist looking at it, can mentally supply the correct number of hydrogens needed to make that carbon neutral, which means giving it a total of four bonds.

Ah, back to the valency.

Exactly.

So if a carbon atom in a line bond structure already shows, say, two bonds connected to it, two lines coming off it, you just know it must also have two implicit hydrogens attached that aren't drawn.

So it's almost like a little mental puzzle every time you look at one, you have to fill in the blanks.

But for something new to this, is it easy to make a mistake, miscount those invisible hydrogens?

It definitely takes practice to build that intuition, no doubt about it.

But it becomes second nature surprisingly quickly.

Yeah.

However, and this is really crucial, hydrogens attached to non -carbon atoms, like nitrogen, oxygen, or sulfur,

those must be explicitly shown.

Always.

Okay, so carbons hide their hydrogens, but other atoms show them.

That's the rule.

So while hexane is just a simple six -carbon zigzag line, a molecule like, say, ethanolamine, which has an oxygen and a nitrogen, will clearly show the H on the OH group and the two Hs on the NH2 group.

Gotcha.

And what about rings or triple bonds?

Good question.

For rings, we just use polygons.

A pentagon means a five -carbon ring, cyclopentane,

a hexagon for a six -carbon ring, cyclohexane, and so on.

And for triple bonds, because they have a linear geometry, they're always drawn as a straight line segment between two carbons, not jagged like single or double bonds.

Okay, so what does this all mean for you, the listener?

Well, it means a really powerful practical benefit.

When you see a carbon atom in one of these line bond structures, if it has, say, two bonds explicitly shown, you automatically know it has two hydrogens you don't see.

If it only shows one bond, it must have three hydrogens.

And while it definitely takes practice to instantly see those missing hydrogens, the payoff is huge.

First, you'll draw structures much, much faster.

That's a massive advantage when you're dealing with complex molecules or working through reaction mechanisms.

Absolutely.

And maybe even more importantly, you'll find it much easier to visualize chemical changes.

It's like gaining a kind of molecular x -ray vision.

And that's practically a superpower for predicting how a molecule might behave in a reaction later on.

It really is.

And you'll often encounter structures in textbooks or papers that actually combine all three methods.

Some parts drawn out fully, like Lewis, some condensed, and some as line bond.

That sounds potentially confusing.

It could be a little jarring at first, yeah.

But chemists usually do this intentionally to emphasize specific parts of a molecule that are particularly important for whatever they're discussing, like the active site of an enzyme or the part undergoing a reaction.

It's all about communicating clearly and efficiently.

Okay, now that we've sort of got the words down the structures and learned some clever abbreviations.

The shorthand.

Right.

The shorthand.

Let's move to the grammar and syntax of organic chemistry,

arrow pushing.

This is where I guess the magic happens.

You could say that.

It's how we show mechanism.

Because when we talk about chemical changes, reactions, we're almost always talking about electrons, right?

Electrons, electrons, electrons.

Always the electrons.

These little guys are the keys to chemistry because they're the adventurous ones, aren't they?

Constantly bustling about, forming new bonds, breaking old ones.

Protons and neutrons, they pretty much stay put, fixed in the nucleus.

Yep, spectators mostly.

So organic chemists concern themselves mostly with what happens to the electrons.

How do we show that?

To show this electron movement, this flow, organic chemists use a very precise set of arrows.

Our source, the Dummies book, outlines five major types you'll encounter.

Okay, five types.

Let's hear them.

First, you have the double -headed resonance arrow.

Looks like line.

This specifically shows the relationship between different resonance structures of the same molecule or ion, which we'll dig into next.

It does not mean equilibrium.

Resonance arrow.

Got it.

Then there's the equilibrium arrow.

That's the one with two half arrows pointing opposite directions, like here.

This indicates a reversible reaction, where reactants and products exist together.

Yeah, and I've seen that one.

Then the simple, single, full reaction arrow.

That just shows the overall transformation of reactants into products in a chemical reaction.

Straightforward.

Okay.

And then, for showing the actual movement of electrons within a mechanism, you mainly use the full -headed curved arrow.

Looks like a curve with a normal arrowhead arrow.

This is by far the most common one you'll draw.

It signifies the movement of two electrons, typically representing a lone pair forming a bond, or a bond breaking to form a lone pair, or a bond shifting.

Two electrons.

Full -headed arrow.

Lastly, there's the half -headed curved arrow, sometimes called a fishhook arrow.

It's a curve with only half an arrowhead, like a fishhook.

This is used specifically for showing the movement of just one electron at a time.

You see these mainly in free radical reactions, which involve unpaired electrons.

Fishhook for one electron.

Got it.

So, full -head for two, half -head for one.

Exactly.

And here's the crucial, like, the absolutely unbreakable rule for drawing these curved arrows, the full -headed or half -headed ones, showing electron movement.

This is the one that will save you points on every single exam.

Okay, I'm listening.

Arrows always originate from a source of electrons.

That means either a lone pair, the dots, or a bond, the sticks, specifically the electrons in that bond.

And they always point toward where those electrons are going, usually to an atom to form a new lone pair, or into the space between two atoms to form a new bond.

From electrons to where they go, makes sense.

They never originate from atoms themselves, especially not from positively charged atoms or nuclei.

Never, ever.

So, for example, if you want to show, say, water grabbing a proton, an H plus ion.

Right, protonation.

That arrow starts from one of water's lone pairs, the electrons on the oxygen, and points to the H plus ion.

You don't draw the arrows starting from the H plus ion.

Absolutely not.

Because an H plus has no electrons.

It's just a naked proton.

There's nothing there to move.

The source material really emphasizes this, basically shouting, never, never, never, never, never, never.

Uh -huh.

Okay, message received.

Drawing an arrow originating from a positive charge is a major fundamental conceptual error.

No electrons there to push.

Okay, clear as day.

So we have Lewis structures, the shorthand, the arrows for grammar, but you hinted earlier that Lewis structures aren't always perfect.

That's right.

They're incredibly useful models, but sometimes a single Lewis structure falls short.

It can't accurately describe the exact location or distribution of certain electrons,

specifically

lone pair electrons and pi electrons, those electrons found in double or triple bonds.

Okay.

They can be more spread out or delocalized than a single drawing suggests, and this is precisely where resonance structures come into play.

They help us capture that more nuanced, more accurate picture of electron distribution.

Resonance, so it's like an upgrade to the Lewis structure idea.

In a way, yes.

It's a way to handle situations where one Lewis structure just isn't enough.

Take the carboxylate anion, RCO2, as a prime example.

If you draw one Lewis structure for this, it'll inevitably show one carbon -oxygen double bond and one carbon -oxygen single bond.

Okay, I can picture that.

And the negative charge would seem to be located entirely on the oxygen with the single bond.

But experimentally, when we measure the actual molecule, we find that both carbon -oxygen bonds are identical in length, somewhere in between a typical single and a typical double bond.

Oh, weird.

And the negative charge isn't stuck on one oxygen, it's equally shared delocalized across both oxygen atoms.

The actual molecule isn't literally flipping back and forth between two different structures.

It exists as a single, unchanging entity, which is a hybrid, a blend of all the valid resonance structures we can draw for it.

A hybrid, not flipping, okay, that's a key distinction.

It really is.

Resonance structures are our way of depicting this electron delocalization on paper, using familiar Lewis structure rules, because a single Lewis diagram just can't capture that blended reality.

So it's not like the molecule is indecisive, it's just that our drawing system needs a full story across.

That's a perfect analogy.

And to draw these resonance structures correctly, there are three fundamental rules you absolutely must follow.

Okay, lay them on me.

Rule one,

atoms are fixed, they cannot move.

You're only changing the distribution of electrons.

The atomic nuclei stay in the same place.

Got it.

No moving atoms.

Rule two,

only lone pair electrons and pi electrons, those in double or triple bonds, can move.

Single bonds, sigma bonds, stay put.

You cannot break single bonds when drawing resonance structures.

If you do, you're drawing a different molecule or breaking it apart and not showing resonance.

Okay, only lone pairs and pi electrons move.

And rule three, you absolutely cannot break the octet rule, especially for second row elements like carbon, nitrogen, and oxygen.

The sum of an atom's bonds plus its lone pair electrons cannot exceed eight, which usually means no more than four bonds lone pairs total.

The octet rule again, don't give carbon five bonds.

Exactly.

No Texas carbons.

Okay, so with those rules,

how do you actually find these alternative resonance structures?

Is it just trial and error?

It can feel like that at first, but it's really about recognizing a few common recurring patterns in the molecule.

And of course, using that arrow pushing grabber we just talked about, there are four main patterns our source highlights.

Four patterns.

The first pattern is when you have a lone pair right next to a double or triple bond.

We say it's allylic to the pi bond.

Lone pair next to a pi bond.

Yeah.

In this case, the lone pair electrons move towards the pi bond to form a new pi bond and simultaneously,

the electrons in the original pi bond get pushed over onto the next atom, becoming a new lone pair.

This usually requires drawing two curved arrows to show the complete electron shift and avoid breaking the octet rule.

Two arrows usually.

Okay, pattern one, what's next?

Pattern two involves a positive charge, occasion, next to a double bond, a triple bond, or even a lone pair.

Again, allylic position or next to a lone pair.

Positive charge next to pi electrons or a lone pair.

Right.

Electrons are attracted to positive charge.

So if you have a occasion next to a double bond, the pi electrons from that bond will shift over towards the positive charge, moving the double bond and the positive charge.

Or if the occasion is next to an atom with a lone pair, that lone pair can move in to form a new pi bond, neutralizing the original planchen, but often creating a charge elsewhere.

Electrons always move towards the positive center to try and stabilize it.

Okay, electrons move towards positive charges, makes sense.

Pattern three.

Pattern three involves a double or triple bond that includes an electronegative atom, most commonly oxygen or nitrogen.

Electronegative, meaning it likes electrons.

Exactly.

Atoms like oxygen and nitrogen are electron hogs.

They'd actually pull electron density towards themselves.

So in a CO double bond, for instance, like an acetone, the pi electrons are already pulled towards the oxygen.

A valid resonance structure can be drawn by moving those pi bond electrons completely onto the oxygen atom as a lone pair.

So the pi bond breaks and the electrons jump onto the oxygen.

Yep, forming a lone pair.

This leaves the carbon with a positive charge and the oxygen with a negative charge.

You see this pattern a lot with carbonyl groups.

See you.

Okay.

And the last one, pattern four.

The final common pattern is having alternating double and single bonds within a ring.

Think conjugated systems in a cycle.

Like benzene.

Benzene is the absolute classic example.

In these cases, you can simply push the pi electrons around the ring.

Each double bond shifts over one position,

effectively moving all the double bonds around the cycle.

It shows how those pi electrons aren't fixed between specific carbons, but are delocalized around the entire ring.

Alternating double bonds in a ring, just push them around, got it.

Those four patterns seem like powerful tools for spotting resonance.

They really are.

If you can spot one of these patterns, you know resonance is possible and you know which electrons need to move and where to push the arrows.

What's also fascinating, and the source points this out, is that sometimes you can have more than just two resonance structures for a molecule.

Definitely.

It walks through an example, two hexadino, which actually has four significant resonance structures.

It seems like applying one pattern can sometimes create a situation where another pattern applies.

That's a very common scenario.

You might move a lone pair next to a double bond, pattern 1, which creates a negative charge on one atom and shifts the double bond.

But maybe that new double bond is now next to an electronegative atom, pattern 3, allowing another shift.

Or maybe moving a double bond towards a caucasian pattern 2 puts the positive charge next to another double bond, allowing a further shift.

So it can be like a chain reaction of electron movement.

Exactly.

You follow the patterns, push the arrows according to the rules, and see where it leads you.

Sometimes it stops after one step, sometimes you can draw several valid structures.

Okay, so if we can draw multiple resonance structures, are they all equally important?

You mentioned the hybrid idea.

Do some structures contribute more to that final blend than others?

Yes, absolutely.

Some resonance structures are more stable than others.

And the more stable ones contribute more significantly to the actual structure of the molecule, the resonance hybrid.

We generally assess the relative stability,

and therefore importance, using three main factors.

Three factors for stability.

First, structures with the fewest formal charges are generally the most stable and most important.

Neutral structures are usually better than charged ones.

For example, the primary structure of acetone, the neutral one with just the CO double bond, is far more stable and contributes much more to the hybrid than its resonance form where you have a C plus and an O.

Minimize charges.

Got it.

Second, if you do have charges,

consider which atoms carry the charges.

Negative charges are more stable on more electronegative atoms, that's oxygen, nitrogen, halogens.

They're better equipped to handle extra electron density.

Conversely, positive charges are generally more stable on less electronegative, more electropositive atoms, like carbon.

So a structure where a negative charge is on an oxygen is usually better than one where it's on a carbon, all those being equal.

Put charges on the atoms that handle them best.

Makes sense.

And finally, the third factor, which is often considered the most important and can actually trump rule number two, is filled octets.

The octet rule again?

Yes.

A resonance structure where every atom, especially second row elements, has a complete octet of electrons.

Eight valence electrons, through bonding and lone pairs, is often the most significant contributor to the hybrid.

This holds true even if it places a charge on a seemingly less preferred atom, according to rule two.

Wait, so a full octet is more important than putting the charge on the right atom?

Often yes.

For instance, you might have one structure where carbon has only six electrons, a positive charge, and another where nitrogen has a positive charge, but everyone has a full octet.

The structure with all filled octets, even with the positive charge on the more electronegative nitrogen, is frequently the more stable and significant contributor.

The drive to achieve a full octet is incredibly strong energetically.

Wow.

Okay.

Filled octets is the top priority.

That's a really crucial hierarchy to remember.

Definitely.

Fewest charges, then charges on the best atoms, but above all, maximize filled octets.

And as people are learning this, practicing drawing these structures, what are some common mistakes they make?

The ones that cost points on exams?

Ah yes, the common pitfalls.

There are a few classic ones the source highlights.

First, simply forgetting charges.

When you move electrons, you often create or move formal charges.

You must correctly calculate and draw these formal charges on each resonance structure.

And crucially, the net charge must be conserved across all structures.

If you start with a neutral molecule, all its resonance forms must also have a net charge overall,

even if they contain separated positive and negative charges within the structure.

If you start with an ion with an anxious one charge, all resonance forms must also have a net next one charge.

Right.

Charge conservation.

Don't lose or gain charge overall.

Second, and we've hit this hard already, but it bears repeating, breaking the octet rule.

Especially for second row elements, carbon, nitrogen, oxygen, fluorine, you can never have more than eight electrons around them.

Soma bonds plus lone pair electrons, eight, usually meaning four things attached.

The infamous Texas carbon, a carbon with five bonds is, as the source puts it, a major organic chemistry nodo, a fundamental red flag.

Instructors see that immediately.

Okay.

Definitely avoid the Texas carbon.

What else?

Third, moving single bonds, sigma bonds.

Remember rule two of resonance.

Only lone pair and pi electrons move.

If you draw an arrow that implies moving or breaking a single bond, you're not trying resonance anymore.

You're drawing a different molecule altogether or indicating a reaction step, not resonance structures of the same species.

Single bonds are the skeleton.

They stay put.

Exactly.

And finally, a more subtle one, perhaps, but important.

Not following electron flow correctly with arrows.

Arrows must consistently show electrons moving from a source,

lone pair pi bond, to a sink,

an atom or the space between atoms to form a bond.

They need to flow logically.

You can't have arrows crashing into each other or pointing nonsensically.

Follow the patterns we discussed.

So four big pitfalls.

Forgetting charges, breaking octets, moving single bonds and messing up the arrow flow.

Good things to watch out for.

Definitely.

Avoiding those common errors goes a long way.

So we've covered quite a journey today, haven't we?

Feels like we've built a solid foundation.

I think so.

From learning the basic words of organic chemistry, those Lewis structures and their various clever abbreviations, like condensed and line bond, to understanding the crucial grammar of electron movement through arrow pushing.

Mechanics.

And finally, diving into the really nuanced world of resonance structures, which kind of correct for the imperfections of our simpler models and show electron delocalization.

Yeah, it ties a lot together.

This deep dive, drawing from organic chemistry I for dummies, should really equip you, the listener, with the foundational tools to start confidently speaking organic chemistry, or at least reading it.

And remember, like any language, mastery truly comes with practice.

Drawing structures, pushing arrows, identifying resonance patterns, doing it over and over is key.

These skills are absolutely essential, not just for, you know, drawing pretty pictures on paper, but for deeply understanding why chemical reactions happen the way they do.

Right.

And maybe more importantly, understanding how you can predict what might happen in a reaction you haven't seen before.

And this really raises an important question, maybe something for you to ponder after this.

Ooh, I like a final thought.

How am I truly mastering this visual language of chemistry?

Thinking in terms of these structures and electron movements, how might that open up new avenues for you to think about designing novel molecules?

Yeah.

Maybe new life -saving drugs or new materials, or even just understanding the incredibly intricate molecular dance that's happening inside your own body right now in your biological processes.

Wow.

That's a fantastic thought to leave everyone with.

Thinking about structure not just as representation, but as a tool for design and understanding.

Precisely.

Well, thank you so much for guiding us through that.

And thank you, our listeners, for joining us on this deep dive into the visual language of organic chemistry.

Keep exploring, keep questioning, and keep being incredibly curious.

We'll catch you on the next deep dive.