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Acidity is a tricky concept, and it’s not always like how we see it in the movies. As organic chemists, we need to know how to predict the strength of weak acids and bases. In this episode of Crash Course Organic Chemistry, we’ll learn four key factors that we can use to predict relative acidity. This important tool will help us us to predict the products of chemical reactions.

Episode Sources:
Chen, J., Zhao, Y., Li, X. C., & Zhao, J. H. (2019). Pyridine alkaloids in the venom of imported fire ants. Journal of agricultural and food chemistry, 67(41), 11388-11395.
Everts, S. (2014). “An Ant’s Acid Antidote” Chemical and Engineering News, 92(9), 44-45.
Trummal, A., Lipping, L., Kaljurand, I., Koppel, I. A., & Leito, I. (2016). Acidity of strong acids in water and dimethyl sulfoxide. The Journal of Physical Chemistry A, 120(20), 3663-3669.

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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You can review content from Crash Course Organic Chemistry with the Crash Course app, available now for Android and iOS devices.

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

A fun little detail in that sci-fi movie Alien -- you know, the one with Sigourney Weaver -- is how the xenomorph has acidic blood. When the humans cut one open, its blood eats right through the metal floor and leaves a hole. It's creepy!

Xenomorph blood is way stronger than most acids produced by Earth creatures. For example, Tawny ants produce formic acid to protect themselves from the basic venom of fire ants. But formic acid is relatively weak, kind of like vinegar or lemon juice -- which definitely won’t eat through a spaceship.

As organic chemists, we need to know how to predict the strength of weak acids and bases. Not just to dream up cool (and scientifically accurate) aliens, but because acids and bases are molecular hot spots where reactions can take place. So let's add acid-base chemistry to our toolkit for predicting chemical reactions. [Theme Music].

Part of what makes acid-base chemistry tricky is that there are different but overlapping definitions of acids and bases. Arrhenius, Brønsted, Lowry, and Lewis all had their opinions on how an acid and base should be defined. For this video, we’ll stick with the Brønsted-Lowry definition proposed in 1923.

And to help understand this definition, we’ll need to push some electrons around. Last episode, we learned how to use arrows to push electrons within a molecule to understand resonance structures. But we can also push electrons between molecules to show how bonds break and form using the same rules: we start on electrons and point to where a new bond is made.

In the Brønsted-Lowry definition, an acid is anything that loses a proton (also known as a plus-one-charged hydrogen ion). And a base is anything that accepts a proton. Carboxylic acids, like acetic acid and propanoic acid, are Brønsted-Lowry acids.

When we dissolve these acids in water, the water acts as a base and removes a proton to form a hydronium ion and the corresponding carboxylate ion. We say that hydronium is the conjugate acid of water, and the carboxylate is the conjugate base of the carboxylic acid. In organic chemistry, we want to know how readily a molecule will gain or lose a proton.

This is a physical property, like boiling point or melting point, that we describe with Ka -- the acid dissociation constant. Hank talks more about Ka in episode 30 of Crash Course Chemistry if you want to start with the basics. But essentially, it describes the relationships between products and reactants when the rate of the forward reaction is equal to the reverse reaction -- in other words, when the reaction has reached equilibrium.

The size of the Ka tells us if we have more products or reactants. If the Ka is large, the equilibrium favors the product side, and the molecule is a strong acid -- one that's very willing to get rid of a proton. If the Ka is small, then the equilibrium favors the reactants, so the molecule is a weak acid -- one that's a little less willing to get rid of a proton.

So that logic is all fine and good, but it can be a little hard to tell whether the Ka is small or large from looking at a number like 1.8x10-5 . We can get rid of the exponents and make these numbers more manageable by taking the negative log of a Ka to get pKa. Just like the concentration of protons becomes pH with a little math, an acid’s equilibrium constant becomes pKa.

We can use pKa to compare two acids. A lower pKa means the equilibrium prefers the product side and a molecule is a stronger acid. And a higher pKa means the equilibrium prefers the reactant side and a molecule is a weaker acid.

Let's look at a couple examples. We can see that hydrochloric acid has a negative pKa (which is REALLY low). We consider it completely dissociated into products, so it's a strong acid.

Propanoic acid and acetic acid have similar pKa values, which makes sense because propanoic acid only has one extra CH2 group. They're both weak acids. But the pKa of ethanol is dramatically different from acetic acid, even though they're both losing a proton from an OH group.

That's the key though: they're similar, but not the same, when it comes to charge distribution. When acetic acid loses a proton, it forms its conjugate base, acetate. And remember last episode, we talked about the resonance structures of an acetate ion, and how the negative charge is spread out over its two oxygen atoms.

This is like going backpacking with a group of friends and splitting up gear. Like, instead of carrying everything yourself, you take the tent, another friend takes the cooking gear, and another takes the food. Distributing the weight means you all carry some of the burden.

The same is true here: distributing the negative charge over the two oxygen atoms makes it easier for the acetate ion to carry the burden of the negative charge. Because of the resonance stabilization in the conjugate base, it’s not too tough for acetic acid to lose a proton. On the other hand, when ethanol loses a proton, it forms the conjugate base ethoxide.

Ethoxide doesn't have any resonance structures, which means it doesn't have friends to share camping gear with. So its oxygen atom is feeling the full burden of the negative charge. This makes ethanol a weaker acid than acetic acid.

But, since it’s pretty tough to take the proton off of ethanol in the first place, if we do form the conjugate bases, ethoxide is a much stronger base than acetate. Another great way to see how resonance stabilization can affect acidity is to consider two kinda similar ring compounds: phenol and cyclohexanol. Phenol’s pKa is about 10 and cyclohexanol’s pKa is 16.

In cyclohexanol's conjugate base, that negative charge is stuck on oxygen, because there are no double bonds or other oxygens to share the burden. So, like ethoxide, the conjugate base is less stable and really wants a proton back. But in phenol's conjugate base, phenoxide, the negative charge can be pushed around the benzene ring to make four different resonance structures.

This stabilizes the conjugate base, so phenol is more acidic. Resonance stabilization is one of four major factors that help us understand the role of pKa in our reactions. Another key is the atom that loses the proton.

Within a row on the periodic table, more electronegative elements stabilize negative charge better, and within a group, larger elements form more stable conjugate bases. Bigger atoms have more electrons, which end up in orbitals that are pretty diffuse. So with bigger atoms, the electron cloud is easily smeared out and distorted, a property we call polarizability.

Imagine a cup of water as the electron density on a small atom. If we pour the water into a Frisbee, which is like a large atom, it has more space to move around and spread out. Atom polarizability affects acidity and pKa, because this smeared-out-ness stabilizes negative charges.

It’s always useful to look at an example, so let's go back to phenol and compare it with its closely-related cousin thiophenol, which has a sulfur atom instead of an oxygen. The only difference in structure is the size of the sulfur atom compared to the size of the oxygen atom. To be precise, sulfur has 8 more electrons than oxygen.

So in thiophenol's conjugate base, electrons are more smeared out on the polarizable sulfur atom, stabilizing the conjugate base. And the more stable the conjugate base, the more acidic the acid, making thiophenol more acidic than phenol. The third key to pKa is hidden in covalent bonds.

The inductive effect has to do with electronegativity throughout a molecule, with more electronegative atoms pulling the negative charges toward them through bonds. As an example, let’s compare two other very similar compounds: acetic acid and trifluoroacetic acid. Both of their conjugate bases have resonance stabilization across the two oxygen atoms.

But the resonance stabilization is all that acetate has going for it, giving acetic acid its pKa of 4.76. Now, trifluoroacetate also contains fluorine, an electronegative atom. Those three fluorines pull on the negative charge and stabilize the conjugate base even more.

Sort of like having even more backpacking buddies who carry the poles and other pieces of the tent for you! So the products side of the equilibrium is favored, and the pKa of trifluoroacetic acid is 0.23! The last, and least powerful, key to understanding pKa also has to do with orbital shapes…..specifically the s character of hybrid orbitals.

Remember in episode 4, we talked about orbitals, places where we’re most likely to find electrons around atoms. And we talked about orbitals combining when atoms bond to form hybrid orbitals. The three most common hybrid orbitals for organic compounds are sp3, sp2, and sp, and this is approximately what they look like. sp hybrid orbitals can be thought of as being 50% s-orbital and 50% p-orbital.

So we can say that they have more s character than sp2 and sp3 hybrid orbitals, which combine additional p orbitals into the mix. So, sp orbitals are more similar to a plain old s orbital. These shapes means their electrons are closer to the nucleus, and the atom can stabilize a negative charge better.

By contrast, sp2 and sp3 hybrid orbitals hold the electrons a little farther away from the nucleus and the atom doesn’t stabilize a negative charge as well. To look at this in action, let's compare the pKa of ethane, ethene, and ethyne. The carbons in ethyne have sp hybridized orbitals, which means the conjugate base has more s character and can stabilize the negative charge better.

So ethyne is the most acidic of this trio -- but, to be totally realistic, a pKa of 25 is not very acidic at all in the grand scheme of things and will definitely not be eating through spaceships or anything like that. We can see the importance of the s character of hybrid orbitals by looking at the acidity of different protons within a molecule. For example, this molecule has an alkene at one end (with the protons highlighted in blue) and an alkyne at the other (with a proton highlighted in red).

So which of these two protons is more acidic? By pulling off a proton from each end of the molecule, we can create two different conjugate bases with a negative charge in different places. Using what we know about hybrid orbitals, the alkene end has an sp2 hybridized carbon, and the alkyne end has an sp hybridized carbon.

That means the alkyne end has more s character and can stabilize the negative charge better, so that conjugate base is more stable, and that proton is more acidic. So throughout this episode, we've learned four key things that we can use to predict relative acidity. Each of these factors stabilizes negative charge in a conjugate base, which makes the corresponding acid more acidic.

Number 1: Atom identity -. More electronegative and larger elements stabilize charge better. Number 2: Resonance stabilization -.

If we can draw multiple Lewis structures for conjugate bases, they are more stable. Number 3: The inductive effect -. Electronegative atoms can pull negative charge toward themselves through covalent bonds.

And Number 4: The s character of the orbital. More s character stabilizes negative charge better. Acidity is definitely tricky, but it's an important part of organic compounds and will help us predict the products of chemical reactions.

In our next episode we’ll start to use all of the tools we’ve learned so far, and start forming covalent bonds at our molecular hotspots! Thanks for watching this episode of Crash Course Organic Chemistry. If you want to help keep all Crash Course free for everybody, forever, you can join our community on Patreon.