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We’ve all heard the phrase “opposites attract.” It may or may not be true for people, but it’s definitely true in organic chemistry. In this episode of Crash Course Organic Chemistry, we’re learning about electronegativity, polarity, resonance structures, and resonance hybrids. We’ll practice a very important skill for this course that will help us avoid a lot of memorization in the future: electron pushing. It’ll be a lot of trial and error at first, but we all start somewhere!

Episode Sources:
“THINK BIG! Must the molecules of life always be Left-Handed or Right-Handed?” Smithsonian Magazine.

Series Sources:
Brown, W. H., Iverson, B. L., Ansyln, E. V., Foote, C., Organic Chemistry; 8th ed.; Cengage Learning, Boston, 2018.
Bruice, P. Y., Organic Chemistry, 7th ed.; Pearson Education, Inc., United States, 2014.
Clayden, J., Greeves, N., Warren., S., Organic Chemistry, 2nd ed.; Oxford University Press, New York, 2012.
Jones Jr., M.; Fleming, S. A., Organic Chemistry, 5th ed.; W. W. Norton & Company, New York, 2014.
Klein., D., Organic Chemistry; 1st ed.; John Wiley & Sons, United States, 2012.
Louden M., Organic Chemistry; 5th ed.; Roberts and Company Publishers, Colorado, 2009.
McMurry, J., Organic Chemistry, 9th ed.; Cengage Learning, Boston, 2016.
Smith, J. G., Organic chemistry; 6th ed.; McGraw-Hill Education, New York, 2020.
Wade., L. G., Organic Chemistry; 8th ed.; Pearson Education, Inc., United States, 2013.

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Hi! I’m Deboki Chakravarti and welcome to Crash Course Organic Chemistry!

Whether we’re talking about romantic relationships or magnets, there's that cliche: “opposites attract.” For people, well, it’s complicated. Some psychologists say it’s because we seek what we don’t have, while others say it’s not really a thing. So, who knows.

But for magnets, “opposites attract” has to do with polarity and physics. And that goes for other science too, like organic chemistry. In polar molecules, opposite charges do attract, and positive and negative regions of molecules are drawn together.

To understand polar molecules, we have to understand their reactive sites by thinking critically about those negatively-charged particles: electrons. [Theme Music]. When atoms are bonded together, they don't necessarily share electrons equally. One element could be a little greedier than another.

Electronegativity is the atomic property that helps us think about how much one atom will attract electrons in a bond, compared to other atoms. Because electronegativity depends on two atoms in a relationship with each other, the electronegativity of a specific element can vary a little, depending on what chemical compound it's a part of. That gets a little complicated, but thankfully, scientists of the past have done a lot of work for us.

American chemist Linus Pauling developed a relative electronegativity scale, which ranks the elements from most electronegative (which is fluorine) to least. We can use Pauling’s scale to determine a couple things about molecules. First, electronegativity can tell us about atomic bonding, which is exactly what it sounds like: the bonds that hold atoms together.

A big difference in electronegativity between two atoms means an ionic bond, a small difference means a nonpolar covalent bond, and everything in between is a polar covalent bond. For example, the electronegativity difference between a hydrogen atom and an oxygen atom in a water molecule is about 1.4, in the "polar covalent bond"-zone. Second, electronegativity can tell us about regions of charge in molecules.

Sticking with our water molecule example, oxygen is quite electronegative and attracts electrons through its bonds, making it partially negative. The hydrogens have lower electronegativity and with oxygen stealing electrons away, they're partially positive. We can draw these partial charges with the lowercase Greek letter delta and a plus or minus sign.

Now, we also know that a water molecule has a bent molecular shape. So one side of the molecule has a concentration of negative charge, while the other is positive. We call this a dipole, two regions with a lopsided charge distribution.

To show a dipole, we use a little arrow. The arrow head represents the negative side of the dipole, while the cross represents the positive side of the dipole. It looks like a little plus sign, that’s how I remember!

So, water contains two polar covalent bonds. And because of its bent shape it also has a dipole and is considered a polar molecule, one with two clear regions of different charge. To really anchor these ideas about electronegativity, let’s look at another molecule: carbon dioxide, you know, the gas that plants use during photosynthesis.

If we draw the structure of CO2, we can see that its electron pair geometry and molecular shape are both linear. If we look at Pauling’s electronegativity scale, we can see that carbon and oxygen have a difference in electronegativity of about 1. Both bonds are polar covalent and the oxygen has a partial negative charge because of its greater electronegativity, while the carbons are partially positive.

So, is carbon dioxide polar? Because the electrons are being pulled away from each other equally and symmetrically, this particular tug-of-war isn't lopsided, and carbon dioxide doesn’t have a dipole. Therefore it isn’t polar!

It’s a nonpolar molecule. To be clear, in a carbon dioxide molecule, the carbon-oxygen bonds are polar covalent, but the whole molecule is nonpolar. Electronegativity is another key piece of our organic chemistry toolbox.

It reinforces the core idea we’ve been learning: that the structure of molecules informs what they can do and how they react with other chemicals. Organic chemists need to have a working knowledge of the types of bonds and electronegativity, even though we don't calculate it every single time. So let’s get comfier with dipoles and continue building our chemical intuition by looking at 1-chloropropane.

From Pauling's electronegativity scale, we can see that the electronegativity difference between the carbons and hydrogens on most of the chain is really small, so those are all nonpolar covalent bonds. But there's a bigger electronegativity difference between carbon and chlorine, so that’s a polar covalent bond. We have a partial positive charge on the carbon and a partial negative charge on the chlorine, which means the electrons of the carbon-chlorine bond are being pulled unequally in that one spot.

And that means… we have a dipole, folks. This molecule has a region that is polar and other regions that are non-polar. But since it has a polar covalent bond, it's a slightly polar molecule.

It's definitely not as polar as water based on the difference in electronegativity and that three carbon chain, but it's more polar than carbon dioxide. And the polar end creates a molecular hotspot, which is raring to go for chemical reactions. In fact, whether we’re talking about reactions caused by cleaning products or carried out in plastic factories… electronegativity and molecular dipoles are often involved.

Positive regions on a molecule can attract negative regions on another, and chemical bonds can be made and broken. As we get into reactions, a big leap we have to take from general chemistry to organic chemistry is keeping track of electrons and moving them really precisely, which we call electron pushing or arrow pushing. If we can learn and remember some guiding principles about how electrons like to move, we can look at any chemical reaction and make reasonable guesses about its products.

Basically, practicing pushing electrons will help us avoid memorizing every single reaction in organic chemistry. You might want to grab a pencil and a really good eraser, because mistakes will happen and electron pushing is a lot of trial and error at first. All chemists have to start somewhere!

We'll start by looking at acetate, the anion of acetic acid. I'll draw it out with all lone pairs of electrons, because that'll help us keep track of them this whole time. One of the oxygens has one bond and three lone pairs of electrons, giving it a formal negative charge.

Remember, formal charges are the difference between a neutral atom’s valence electrons and the total electrons surrounding the atom in the molecule. So in this case, a neutral oxygen atom has 6 valence electrons. The single-bonded oxygen we're looking at has 6 electrons in lone pairs, which counts as 6 charges.

It also has 1 bond, which counts as 1 charge because the electrons are shared between two atoms. 6 - 7 is -1, which is the formal negative charge! But we can also draw another valid Lewis structure for acetate, where the other oxygen atom has the single bond and formal charge. These drawings are called resonance structures, which are representations of the compound that differ, in this case, in the placement of lone pairs and pi bonds.

We indicate multiple resonance structures with a resonance arrow, which is an arrow with arrowheads on both sides. And we're not flipping the acetate ion around to get its two resonance structures — the atoms stay in place. We’re actually changing where we show the electrons between the two oxygen atoms.

Drawing resonance structures can help us understand where the electrons could be, but in real life, things are a little mushier. All molecules are a blend of their resonance forms, which we call a resonance hybrid. The resonance hybrid is more stable than any individual resonance form, but it’s really hard to draw partial bonds, so we usually pick one form to draw at a time.

And we know resonance hybrids are a thing because we've measured bond lengths in these molecules. For example, in an acetate ion, both carbon-oxygen bonds are shorter than average single bonds and longer than average double bonds. The charge and pi bond are being equally shared across all three atoms.

In fact, a common theme with resonance structures is electrons shared across the p-orbitals of three atoms. So even though identifying resonance forms can be really tricky, we can approach these puzzles three atoms at a time. To start, we need to find a pi bond, and an adjacent p orbital with 0, 1, or 2 electrons in it.

These three atoms are involved in resonance, so let’s number and box them. Now, whatever is at atom 1 moves to atom 3, and the stuff at 3 goes to 1. In this case, our double bond moves to atom 3 and our extra lone pair and negative charge moves to atom 1.

Now that we have two different resonance structures as our beginning and end, we're ready to fill in the middle of the puzzle by pushing electrons. In this episode, we’ll only be pushing pairs of electrons, which we represent using a curved arrow. This arrow always starts on a pair of electrons, and points to where they end up.

To get started, let’s show how a lone pair of electrons on oxygen becomes a double bond. We’ll start our arrow on that pair of electrons, and point it to where the double bond forms. When we move a lone pair off an atom, it has to form a neighboring pi bond.

But we can’t stop here! This would give a carbon with 5 bonds, which is not organically acceptable because it exceeds carbon’s maximum valence electrons. We have to keep pushing electrons until we have a molecule that follows rules we've already learned about the maximum number of electrons and bonds. We're doing science here, not science fiction.

So, next, let's draw an arrow to move a pair of electrons away from that double bond to a neighboring atom, which means our lucky recipient is the other oxygen atom. The middle of the puzzle is filled in, and we've shown how our two resonance structures are connected!

An important thing to notice is that this whole time we only moved electrons in double bonds or lone pairs. Electrons in single bonds don’t go anywhere. We're going to have to understand resonance structures of many different molecules, so let’s try another example.

Here we have a carbon with a positive charge, called a carbocation. We’ll start by finding our pi bond, and an adjacent p orbital with 0, 1, or 2 electrons in it. The positive charge on carbon is a p orbital with zero electrons—let’s number our three atom unit.

Now, let’s swap what we have at atoms 1 and 3 to find our other resonance structure:. Nice. Let’s see if we can move between these structures with arrow pushing.

We always begin our arrows on electrons, so I’ll start my arrow on the double bond. This moves the electrons away from that carbon and gives it a formal positive one charge. And we’ve done it!

In both of these examples, the two resonance structures are very similar and contribute the same amount to the resonance hybrid. But that's not always the case for molecules with resonance structures, like these two:. Resonance hybrids are usually more like a weighted average, sort of like grades.

Big tests might count more than weekly quizzes toward your final letter grade in a class, and some resonance structures matter more than others. Here are a few key guidelines to figure out which resonance structures contribute more to the resonance hybrid. Number 1: Neutral resonance forms are preferred (when possible).

So, for example, this: is better than this:. Number 2: Keep an octet of electrons on oxygen and nitrogen (but carbon can be short). And number 3: Negative charges are preferred on more electronegative elements, and positive charges are preferred on less electronegative elements.

For example, both of these resonance forms have an octet, but one has the negative charge on the more electronegative element. So structure A is our winner, and contributes more to the resonance hybrid. If pushing electrons is still sort of brain-bendy, don’t worry.

We’ll be doing plenty more of it for practice. In this episode, we learned that:. Electronegativity describes how atoms attract electrons differently, which can lead to dipoles in some molecules.

Resonance structures have different arrangements of electrons. And resonance hybrids are a mixture of resonance structures. Next time we’ll combine electron pushing with acid-base chemistry, and learn what resonance structures can tell us about the strengths of organic acids and bases.

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