Organic Chemistry Acid-Base Reactions

Welcome back to Educator.0000

Next we are going to talk about acid-base reactions or proton-transfer reactions.0002

Let me show you an outline of the topics we are going to be talking about.0009

First we'll discuss the different definitions we have for acid-base reactions.0012

Then we are going to talk about several things that can have an impact on the acidity of a given molecule--look at trends in the periodic table, inductive effects, resonance effects.0017

We will wrap up by looking at some common acids and bases.0027

Most students coming through general chemistry, having a year of general chemistry, will know that there is two main definitions for acid-base reactions.0032

That is having Lewis acids and bases and Bronsted-Lowry acids and bases.0039

A Lewis acid is something that is described as an electron pair acceptor; so this has to do with electrons; we describe this as an aprotic acid.0044

We will find that the other type of acid, the Bronsted-Lowry, is one we call a protic acid; we will see that next.0056

How can something accept a pair of electrons?--well, it is something that will have a vacancy; that has no octets.0062

Some common Lewis acids are things like BF₃ or AlCl₃.0070

These molecules, when you take a look at their detailed Lewis structure, you find that there is a vacancy on this bond; it is missing an octet.0076

Same thing with the aluminum; it has only three bonds; and we know that atoms want to have a filled octet; and they would readily accept a fourth bond.0088

That is what makes them electron pair acceptors and excellent Lewis acids.0097

A Lewis base then is something that can donate a pair of electrons; so something with either a lone pair of electrons or maybe a pi(π) bond--these would make very good Lewis bases.0103

Let's see an example--if we reacted ammonia (NH₃) with AlCl₃.0113

We said AlCl₃ would be a good example of an acid--a Lewis acid; I'm going to put this in quotes because it's a Lewis acid.0118

Ammonia would be a good example of a base because it has a lone pair of electrons.0126

The reaction e would expect to have happen between these two is that the nitrogen would share its lone pair of electrons with the aluminum and form a bond.0131

We are going to form a new bond between nitrogen and aluminum; and this would be our product of this Lewis acid-base reaction.0140

Because of this new bond, we are going to have some formal charges now on the nitrogen and the aluminum.0149

This nitrogen has one, two, three, four electrons around it; we know nitrogen wants five; so this nitrogen now has a positive charge.0156

This aluminum also has one, two, three, four electrons around it; but aluminum only want three; so it has an extra electron; and we would have a negative charge on the aluminum.0165

So this is the product of our Lewis acid-base reaction.0177

Let's just take a look at one more example--a carbocation is another good example of a Lewis acid; because it is missing an octet and can accept a pair of electrons.0184

Water could be used as a Lewis base; and so same idea--we could use that lone pair to form a bond between the oxygen and the carbon.0196

What does our product look like?--we now have a new bond between carbon and oxygen; and what does this oxygen still have on it?--it still has two hydrogens.0206

Any lone pairs?--well, one lone pair is now being used as a sigma (σ) bond; the other lone pair is still there; so this would be the product of this second reaction.0219

Again, let's check our Lewis structure and check for formal charges; this carbon now has four bonds so it is neutral.0231

But this oxygen has one, two, three, four, five electrons; we know oxygen wants six; it has only five so it is missing an electron; so we would have this positive charge here.0237

These are a few examples of some Lewis acid-base reactions; but in organic chemistry, we typically don't refer to such reactions as acid-base reactions.0248

We are going to have a new name for Lewis acids; we are going to call these electrophiles--things that love electrons; and we are going to call Lewis bases nucleophiles.0257

We will see reactions like this down the line further down in our lessons; but when we do so, we are not going to be calling them Lewis acids and bases necessarily.0272

We are going to be describing them as electrophiles reacting with nucleophiles.0284

For organic chemistry, when we are discussing acid-base reactions, we are talking about the other definition--the Bronsted-Lowry type.0289

In a Bronsted-Lowry definition, a Bronsted-Lowry acid is an H⁺ which we call a proton; it is an H⁺ donor; it should be a Bronsted-Lowry acid.0297

That is why we call this kind of acid a protic acid; because it is one that donates protons; that is a source of protons.0314

A Bronsted-Lowry base is a proton acceptor--something that can take a lone pair of electrons.0324

The reaction--an acid-base reaction then can be described as a proton transfer; a transfer from an acid to a base.0337

The mechanism for that transfer is not the hydrogen atom flying over to the base; because our mechanism when we use curved arrows is showing electrons moving.0346

Instead, it is the lone pair of the base or whatever electron source the base has.0360

It is the lone pair attacking the proton and then breaking the H-A bond and having those two electrons go to the remainder of the acid structure.0366

Our product then will now have the base with a new proton on it and the A group with a new pair of electrons on it.0376

Now we have both the base and the acid were neutral as shown.0387

The base, after you add a proton to it--after you add an H⁺, will now be positively charged; and the A, after losing an H⁺, will be negatively charged.0391

You will have these charges if we started out neutral in both cases.0401

After a base has accepted a proton, we call that new structure the conjugate acid of that base; so a base is related to its conjugate acid by protonation.0407

The structure that remains after an acid has lost its proton is known as the conjugate base of that acid; so an acid is related to its conjugate base by losing a proton or donating a proton.0420

This is called a proton-transfer reaction; and what we have here as shown is an equilibrium.0435

We have two acids and two bases in competition on who is going to donate the proton and who is going to accept the proton.0442

Because if you look at this reverse reaction, the reverse reaction would also be described as a proton-transfer reaction, right?0448

This now is a proton donor, and this is a proton acceptor; and these products would be the result of a proton transfer.0454

We have a competition for this equilibrium; and the rule is that the equilibrium lies in the direction of the weaker acid-base pair... the weaker acid base pair.0462

This is an important rule to keep in mind; we will see lots of examples of this as we move along.0486

If this equilibrium lies in the direction of the weaker acid-base pair, we are going to have to be able to determine which pair that is.0496

How do you decide who the stronger acid is?--well, if you are lucky enough to have a pKa table, you can use that.0504

Or perhaps you can predict that yourself; and that is what we are going to be focusing on in this lesson.0510

That is looking for features that affect the acidity and make something a stronger acid or a weaker acid.0515

Here is an example--let's say we have water acting as an acid and ammonia acting as a base; let's see what that proton-transfer reaction would look like, OK?0521

The acid donates one of its protons to the base; so our mechanism is the base grabs the hydrogen and leaves the electrons behind on the oxygen.0530

What does water look like after it has been deprotonated or has donated a proton?0541

We will now have HO with three lone pairs; and that means we have an O^---one, two, three, four, five, six, seven; oxygen only wants six; so we have an extra electron there.0547

That is the conjugate base; and what does the conjugate acid of ammonia look like?--it is, instead of NH₃, it is going to be NH₄.0560

You could just draw it as NH₄ or I can draw it out to show the new hydrogen for emphasis.0568

That is also going to be charged; this nitrogen has one, two, three, four bonds--four electrons; it wants five; so this is going to be N⁺; we are going to get this ammonium NH₄⁺.0573

If we take a look at this equilibrium and ask in which direction does the equilibrium lie?--does it lie in the forward direction or the reverse direction?0587

What we can do is we can look up the pKa's of the two competing acids; and the forward reaction, water, that has a pKa of about 16.0595

In the reverse reaction, the ammonium is the acid and that has a pKa of about 9.0604

What is the relationship between pKa and acidity?--they have an inverse relationship--the lower the pKa, the higher the acidity.0610

Let's remember that--the lower the pKa, the higher the acidity, and the stronger the acid; so a pKa of 9 is a stronger acid, and a pKa of 16 is a weaker acid.0619

Where does the equilibrium lie?--it lies in the direction of the weaker acid-base pair; so it is going to be going from the stronger acid to the weaker acid.0641

The equilibrium lies to the left--this is one way to describe that; you could also say that the reverse reaction is favored; and we can make that conclusion simply by comparing the two acids.0650

We said that it is going to be the weaker acid-base pair; and it turns out that once you identify who the stronger acid is, that is also going to tell you something about the conjugate bases.0674

If ammonium has a pKa of 9, this is the stronger acid; its conjugate base is going to be the weaker conjugate base.0687

Because an acid and its conjugate base also have an inverse relationship--a stronger acid has a weaker conjugate base; the weaker acid, water, has the stronger conjugate base.0701

Once we identify the weaker on one side of the acid, it will always correspond with the weaker base being on that same side.0716

We don't have to compare both the acids and the bases; we just need to find a difference in one or the other and that should be enough to answer our question.0725

We can just kind of mention here: the stronger acid has the weaker conjugate base; and the weaker acid has the stronger conjugate base; so you can see how we ended up labeling those two bases.0736

This is how we can decide the direction of an acid-base equilibrium.0755

Let's take a look at some factors that might affect the acidity, the strength, of an acid.0761

One thing we can do is we can look at the periodic trends; let's take a look across a row of the periodic table.0767

Carbon, nitrogen, hydrogen acids--all of these have hydrogens on them; so they can possibly donate a proton; they could be potential acids.0772

Given are the pKa's: methane CH₄ has a pKa of about 50; ammonia has 38; and water has about 16.0783

Who is our strongest acid here?--the lowest pKa is the strongest acid; and the pKa of 50 is the weakest acid.0790

Let's make a quick note about pKa's; let's take a look at the difference here; 50 to 16--that is a difference of 34 pKa units.0805

Does that mean that water is 34 times more acidic than methane?--that would be pretty significant.0816

But, no--remember that the pKa table is a logarithmic scale; so we are looking at orders of magnitude--factors of 10.0825

There is a one with 34 zeros after it... times more acidic; so that is a like gajillion--we don't even know what that number is; so this is hugely, hugely more acidic.0834

We need to be able to understand that; we need to be able to rationalize that and have an understanding for these pKa's.0851

The way we are going to answer the question of why is there such a huge difference in their pKa's when you go across the row in a periodic table?0857

What we are going to do is we are going to look at the conjugate bases of each of these... look at the conjugate bases.0865

If water acts as an acid, what is the conjugate base going to be?0874

We are going to lose an H⁺; that is going to give hydroxide, HO^-, as a conjugate base.0878

How about ammonia?--ammonia has one lone pair of electrons; so after we lose an H⁺, we will have NH₂ now with two lone pairs; also a minus because we are losing an H⁺.0886

Methane CH₄--we can see a trend here; we will now have a CH₃ with an extra lone pair and a negative charge.0899

These are our three possible conjugate bases; and what we want to do is try and look for a difference in their stability.0908

What is the most significant difference as you move across the row on a periodic table?--you are going to be increasing in electronegativity.0915

Oxygen is the most electronegative atom compared to carbon and nitrogen; and what effect is that having for us?0923

Let's start by writing the fact here--oxygen is more electronegative than the others.0929

Is that a good thing for that negative charge to be on a more electronegative atom?--absolutely; so what we can say is that the oxygen better handles the negative charge.0941

That means that hydroxide is the most stable conjugate base; hydroxide is the most stable of these three species--the C^- versus the N^- versus the O^-; this is the most stable.0957

What is the relationship then between stability and reactivity?--well, the more stable something is, the lower in energy it is; the less reactive it is.0976

Because it is more stable, hydroxide is the weakest conjugate base; it is the least reactive.0986

What do we know about something that has a stable unreactive conjugate base?--what does it tell you about the parent acid?--it must be a strong parent acid; so hydroxide has the strongest parent acid.1001

Does that agree with our pKa data?--it sure does; that had the lowest pKa; and yes, water was the strongest acid.1021

OK, so it's going to be the stability of the conjugate base that is going to answer so many questions about the strengths of various acids.1029

Let's go through the opposite argument for over here on why this is maybe not such a great conjugate base.1039

What we can say here is that carbon is the least electronegative of all of these carbon, nitrogen, oxygen atoms; and so this is the least stable conjugate base.1047

If this is the least stable, unstable conjugate base--that means it is the most reactive conjugate base; it is the strongest conjugate base... this is the most reactive, strongest conjugate base.1071

OK, the reactivity... the relationship between stability and reactivity for molecules is very similar to that of people.1094

The more stable you are, the more calm and cool and collected you are; you are pretty unreactive; the same is true for molecules.1104

But if you are high energy and unstable, that makes you very reactive and very volatile.1112

We are going to find that same relationship is true when we are comparing people and when we are comparing molecules.1117

If this one is the strongest conjugate base, what does that tell you about the parent acid?--we know that this has the weakest parent acid.1123

That is what our pKa data tells us: pKa 50--this is a horrible acid because carbon hates having a negative charge; it is so electropositive.1137

Our goal and what this is demonstrating here and what we will find again and again and again is that the stronger acid... the stronger acid has the more stable and therefore weaker conjugate base.1148

One thing that can help stabilize the conjugate base is the electronegativity of the atom on which a negative charge resides.1177

Let's take a look at an energy diagram that might help illustrate this concept.1184

If we are comparing methane and ammonia and water as acids, we know that water is the strongest acid.1190

That is because when we compare the difference in energies of our conjugate bases, we see that the more stable conjugate base is the one with the negative charge on the most electronegative atom.1198

Right here--this is the most stable conjugate base; meaning it is the lowest in energy.1210

How about the energies of our starting materials?--these are all neutral stable molecule; so these have all about the same energy.1221

There is not any significant difference in the energies of the starting materials.1228

The significant difference in energies is the atoms bearing the negative charge.1233

When we compare these processes, which of these would prefer to be an acid and donate a proton?1238

This process looks the most favorable; this is the least endothermic; this is the most favorable reaction; and that is why water is the strongest acid.1246

Once again, we are going to look for differences in the stability of the conjugate bases; the more stable the conjugate base, the stronger the parent acid.1263

What makes an acid a good acid?--it doesn't mind donating a proton if it doesn't mind where it is going to.1270

If it is going to a conjugate base, isn't it a stable happy place?--that makes the parent acid more likely to donate its proton.1277

Let's see another example; another periodic trend is when we're going down a family or down a column on the periodic table; so let's compare HF, HCl, HBr, HI.1287

Here we have our pKa's; so of these pKa's, who is our strongest acid?1300

The lower the pKa, the stronger the acid; in fact, the more negative the pKa, the stronger the acid.1305

HI is the strongest acid; that is stronger than HBr, than HCl, and than HF; HF is the weakest of all the halo acids; let's see if we can explain this again.1311

The way we are going to explain it is if we ask: why do we have this difference in pKa's?1327

Once again, we are going to look at the conjugate bases; so what does the conjugate base of HF look like?1333

That means HF is going to act as an acid; it is going to donate a proton; and that leaves behind F^-.1341

HCl gives Cl^-; HBr gives Br^-; and HI gives I^-.1348

Where do we go from here?--we try and find something that will explain a difference in stability between these different conjugate bases.1357

What is the most significant difference as you move down a column in a periodic table?--yes, the electronegativity does change as you move down a column.1366

But a more important difference is going to be the size of the atom as you move down and you add a shell of electrons on each new row.1375

It turns out that iodide is the biggest ion; and fluoride is the smallest ion; so that is the fact; that is the difference between these--all four of these; how do we relate that to stability?1386

Let's think of this ion as having a negative charge; a negative charge tells us we have an excess of electron density.1403

If you can spread that negative charge out over a larger surface area of the large iodide--then that is going to be more delocalized; and that is going to be a more stable charge.1410

What we say about I^- is that the negative charge is dispersed or delocalized--that is a very good word.1424

That means it is not in one small location; it is delocalized; it is spread out over several locations; and that means this is the most stable conjugate base.1437

Where do we go from here once we decide which is the most stable conjugate base?--most stable always means less reactive; so this is the least reactive and therefore weakest conjugate base.1453

Now we have something about the strength of the conjugate base; the weakest conjugate base has the strongest parent acid.1473

Does that agree with our pKa?--it does; this had the lowest pKa; so yes, HI is the strongest acid.1487

What would we say about fluoride?--why is fluoride so much less acidic?--so much less likely to donate a proton?1496

Here we have a very, very small surface area with that negative charge; so we could describe this as an intense negative charge--compact, intense; that makes it unstable.1503

It is an unstable anion to be on such a small surface area.1519

If you are unstable, that means you are more reactive; this is the most reactive and therefore strongest conjugate base; and the strongest conjugate has the weakest parent acid.1523

In this case, we can see the size of the atom having an impact on the stability of the negative charge on the conjugate base.1543

Let's take a look at some inductive effects that we might have to explain differences in acidity; if we ask which is the stronger acid here, we have CH₃OH versus CF₃OH.1554

In this first structure, if we expected this to be an acid, there is two different types of protons that can be donated--one of the hydrogens on the carbon or one of the hydrogens on the oxygen.1566

Which of those hydrogens do you think is going to be the most acidic?1577

We just saw the periodic trends; and we know that because oxygen is so electronegative, that would much prefer to have the negative charge.1581

What we are doing is comparing this OH with this OH when we are comparing the acidity.1588

How are we going to distinguish between these two?--well, once again, let's look at the conjugate bases and see if we can find a difference.1593

CH₃OH--the conjugate base would be CH₃O^-; CF₃OH... let's draw out these fluorines because that is obviously the difference between the two molecules.1603

We either have CH₃ or CF₃; and we have an O^- versus an O^-.1623

This is a case where periodic trends are not going to help us because the negative charge is on the exact same electronegative oxygen in each case.1629

Now we look elsewhere in the molecule to see if there is something that can maybe stabilize or maybe destabilize the negative charge.1637

Clearly, we have these fluorines here; and so let's think about what effect that is going to have on the negative charge; we will start by stating the fact that we know about fluorine.1645

That, of course, is that fluorine is more electronegative than hydrogen--that is what we are comparing it to in this case; of course, fluorine is more electronegative than everything.1654

But we know that fluorine is more electronegative, and what does it mean to be electronegative?--it means that fluorine pulls electron density toward itself.1668

Let's also state that fact: we have an inductive withdrawal of electron density by fluorine... we have an inductive withdrawal of electron density.1677

We can show that with an arrow like this; we could say that the fluorines... each of these fluorines is pulling electron density toward itself.1698

These arrows show the movement of electrons through these simga(σ) bonds; that is what we call an inductive withdrawal--an inductive effect.1707

Here is the tricky part: is that a good thing for the negative charge or is that a bad thing for a negative charge?1716

Typically, when I survey my students on this question, I get about a fifty-fifty split; let's think about what it means to have a negative charge.1723

A negative charge tells me that there is an excess of electron density on that oxygen; that is not a good thing--it would rather be neutral.1730

What are these fluorines doing?--they are helping to pull some of that electron density away from the oxygen, bringing it closer to being neutral.1739

In fact, you could imagine that each of these fluorines is sort of taking on some of that negative charge; it looks like we are delocalizing that charge a bit.1748

Is that a good thing?--that sounds like a good thing; and that is, in fact; this stabilizes... this stabilizes the negative charge.1757

We could again kind of think of it as delocalizing... it delocalizes it.1776

If this is the more stable conjugate base, what does that tell you?--this is the more stable and therefore less reactive, meaning weaker... this is the weaker conjugate base.1785

The weaker conjugate base--this has the stronger parent acid.1807

That was our original question; our original question--which is the stronger acid?--and the acid with the fluorines on it is going to be the stronger acid.1818

This is more likely to donate its proton because the resulting anion on the conjugate base will be stabilized by the fluorines--the inductive effect of the fluorines.1828

We could describe fluorine as an electron withdrawing group; we can abbreviate that EWG; we are going to be seeing that abbreviation a lot down the road.1838

Let's take a quick look at some other common EWGs (electron withdrawing groups).1847

For example, if we have a nitro group, an NO₂ group; I've drawn it out here.1851

This is also something that would pull electron density toward itself; just like a fluorine did, right?1856

We said a fluorine pulls electron density; a nitro would do the same thing because of that N⁺.1862

A cyano group (a CN triple bond) has the same effect; and even though it is not charged like the nitro...1872

Because nitrogen is more electronegative than carbon and because we have some resonance here, there is a partial minus (δ-) on this nitrogen and a partial plus (δ+) on this carbon.1879

Therefore, this also pulls electron density toward itself; it would be an EWG.1890

And a carbonyl--a carbonyl also has resonance that puts a δ+ on this carbon and a δ- on this oxygen; and so that δ+ carbon causes an inductive withdrawal.1896

We can see it can also do stabilization by resonance.1907

Any halide is going to be electronegative and can pull electron density toward itself; so not just fluorine, but also the others--chloride, bromide, iodide.1913

What these have in common is these would all stabilize an adjacent negative charge... these would all stabilize an adjacent negative charge.1922

What if I had a positive charge next door?--now, we are not going to see that on a conjugate base; but what if I had a positive charge, an electron deficient site here? 3231 Would it be a good thing to have a fluorine on that?--or a nitro or a cyano pulling even more electron density away?--that would be a bad thing.1943

We can make a little note here; and it would destabilize a + charge; that might become important down the road if we ever saw that when we are exploring the effects of EWGs.1958

One thing that we will note about inductive effects is that these are something that decrease with distance.1977

The further away we put that EWG, the less effective it is going to have; because simply there is more bonds to travel through.1983

Inductive effects are looking at the electron in σ bonds being pulled toward a more electronegative atom or electronegative group; and so that is why we see a decrease with distance.1989

If we take a look at these acids, these are all carboxylic acids.2002

We see this group here--a carbonyl with an OH; and that is the group that we are going to have; that is going to be our acidic group--it's the OH there.2008

They all have similar looking conjugate bases where we have an O^- next to this carbonyl; so they will all have very similar conjugate bases.2024

The only difference is what is attached to the carbonyl.2035

Here we have just a CH₃ group; here we've put a chlorine somewhere down the chain; here we've put a chlorine again; here we've put a fluorine.2039

What do we see as effect on the pKa's?--we see that as soon as we put a halogen on here, we lower our pKa; in fact, this is going to be the most acidic with the lowest pKa.2049

Why is this the most acidic?--well, that is because the fluorine is more electronegative compared to the chlorine; so this has the most stable and therefore weakest conjugate base.2064

That fluorine is going to help stabilize that conjugate base the best, right?--it is going to pull electron density away from the conjugate base.2088

That is going to be a good thing; so this is the strongest acid.2095

Compare that to having a chlorine; that chlorine is making it not quite as acidic.2101

How do we compare these two?--they both have a chlorine, but this one is situated a little more closely to where the negative charge will be on the conjugate base; this one is a little further down.2107

We see, if we have... the EWG is farther away and less effective.2121

While they all pull electron density away, the fact that he is not as close to the O^- means that he's not going to have as big an impact on the pKa.2135

Why does this have the highest pKa of all of them?--why is this the weakest acid of all of them?--because this has no EWG.2146

There is nothing there to help stabilize the O^- in addition to... they all have the same O^-, so there is nothing additional here to stabilize the negative charge that the others have.2155

This has the least stable and therefore the strongest conjugate base; and of course, this is the weakest parent acid.2168

If we have to compare EWGs, the closer we can get that EWG to the charge that we are trying to stabilize, the better.2185

Let's take a look at using resonance to help stabilize a conjugate base and what effect that might have on the acidity.2196

If we take a look at these two OH bearing compounds, these both have OH groups; so those are reasonably acidic.2205

But this one has a pKa of 16, and this one has a pKa of 5; now again, that is eleven zeros; that is a huge difference in pKa; let's see if we can explain where that difference comes from.2215

If we look at the conjugate bases, we are comparing an O^- to an O^-.2230

Once again, we can't think of any difference in periodic trends because each of those oxygens is equally electronegative.2234

We might say this carbonyl can have some electron withdrawing effect on that oxygen to help stabilize it.2243

But we just that saw inductive effects maybe have an effect of one or two pKa units--certainly not making it trillions times more acidic like this one is; so what is the difference here?2250

If we take a look at this O^-, I recognize that this lone pair is allylic; it's next to a π bond which means we can have resonance with that lone pair and that π bond.2262

What does that resonance do?--it moves the negative charge to a new location; anytime you can delocalize the charge through resonance, that is going to be a really great effect.2276

What we can say here is conjugate base 2 is resonance stabilized... resonance stabilized; you can say it is stabilized by resonance.2293

Resonance is always a good thing and will always have a big impact on stability.2309

That tells us that conjugate base 2 is the more stable, weaker conjugate base; and that is why conjugate base 2 has the stronger parent acid by far--a pKa of only 5.2315

So resonance will have a tremendous impact on the acidity of a compound.2342

One thing I want to consider though is you might look at the starting compounds and recognize that even the starting acid has some resonance stabilization, right?2349

Because this also had an allylic lone pair; and there is another resonance form we can draw for this; so you might ask which one of these resonance stabilizations is more significant?2362

The answer here is when you think about which resonance form contributes more to the overall picture, that will tell you how significant the resonance is.2380

Here we have an O^-, and here we have another O^-; which of these is the better contributor?--which will contribute more to the overall hybrid?2388

The answer is they are equally contributing; and so what we have in the case of this conjugate base is that we have a large amount of resonance stabilization.2400

In fact, when you have two equivalent resonance forms, that is the best stabilization you can have; so this really delocalizes the negative charge.2417

You could put that down here again; this is so important... delocalize negative charge.2425

While this does have some resonance, because now we've created formal charges, the second Lewis structure is avery small contributor to the overall picture.2435

Even though it exists, what we have here is just a small amount of resonance stabilization; and it turns out that this will be not significant.2445

The more significant resonance that clearly does have an impact on the pKa is the stabilization of the conjugate base.2459

Maybe if we take a look at an energy diagram, we might see that pictorially; it might make a little more sense.2470

We said that the stabilization of this parent acid was small, and the stabilization of the conjugate base was large.2478

If we take a look at an energy diagram, when we compare the parent acids 1 and 2, we might that there is in fact a small difference in stabilization because 2 does have a small amount of resonance.2488

But whatever small difference there is, that is not as significant as the large amount of resonance stabilization.2505

That we have comparing the O^- in the conjugate base 1 and the resonance stabilized delocalized O^- in conjugate base 2.2515

We have a large amount of resonance; so there is a large difference in energy; and so once again, this transformation from conjugate base 2 is the least endothermic and the most favorable.2525

What we are going to be looking for, something to bear in mind, is that the most significant resonance stabilization that we can look for.2541

Of all the resonance that we are looking for is we want to find a way to delocalize the negative charge of the conjugate base; we want to be able to move that conjugate base around.2552

The most significant resonance stabilization is delocalization of the negative charge.2567

If there is a way you can move that negative charge to a different location because of resonance, that is going to be something that will lead you to the correct answer and the correct conclusion.2577

Let's see another example; how about if we compare these three compounds?--again, we are looking at three compounds that bear OH groups.2589

But we are trying to see which of these OH's is going to be most likely to donate a proton and therefore be the strongest acid.2600

Because we are comparing three, let's also try to decide which would be the least acidic; which would be the least likely to donate a proton.2606

As usual, the answer is going to reside, since these are all neutral stable molecules, the answer is going to reside in the structures of the conjugate bases.2613

Conjugate base 1 has this O^-; conjugate base 2 has this O^-; conjugate base 3 has this O^-.2622

What we are going to do is try and find a difference in their stability; do any of these have resonance?--that might be a good thing to look for.2629

This one has some resonance because we have a carbonyl; this has some resonance; we can draw an O^-C⁺; so we have that resonance.2640

Does this one have resonance?--can we take this negative charge, take that lone pair, and move it in here?--would that be a way to delocalize that negative charge and move it around?2656

No, we can't do this; because this carbon already has four bonds; it has no place to move its electrons; so that would just be five bonds; so this has no resonance.2666

This has some; we will see what relevance that has here; and how about this last one; does this have any resonance?2679

If we take a look at these lone pairs, we see that they are allylic to a π bond or next to a π bond; so yes, this does have resonance.2685

We can draw a new Lewis structure that uses one of those π bonds; this carbon now will have a negative charge.2696

Are there any other resonance forms?--there are actually; because this lone pair is still allylic.2708

We can move that in; any time we have allylic resonance, we can make the lone pair a π bond and make the π bond a lone pair.2715

Any more resonance?--yes, in fact, we can continue moving this around; we will just say et cetera here because we have made our point that we can move that negative charge around.2727

When we look at the two possible resonances for conjugate base 2 and conjugate 3, what I see in conjugate base 3 is these resonance forms actually relocate the negative charge.2735

They delocalize the negative charge; so that is going to be excellent resonance.2748

Let's say that about conjugate base 3; We have very good resonance stabilization; and that is because we have delocalized the negative charge.2751

This is most definitely going to be the most stable conjugate base; the most stable conjugate base is the weakest conjugate base.2772

The weakest conjugate base has the strongest parent acid; so this has the strongest parent acid; and so we would expect #3 to be the most acidic.2786

how about if we compare 1 and 2 now; if we are looking for who is the least acidic, is there any difference between 1 and 2?2806

2 has some resonance, but notice it doesn't help to move that negative charge; but is this something that stabilizes the negative charge or is it something that makes it worse?2816

Having a positive charge there would help take some of the electron density away from that oxygen, wouldn't it?2827

In fact, we saw the carbonyl acting as an EWG; and so this in fact would be a good thing for the negative charge.2835

Let's make the point first that the resonance... because it does have a resonance form, but the resonance does not delocalize the negative charge; so it is not as good as conjugate base 3.2850

But the inductive effects of the carbonyl, if it is acting as an EWG, the inductive effects makes it more stable than conjugate base 1.2868

Even though this resonance doesn't help move the negative charge from the oxygen, it helps pull some of that electron density away inductively.2893

The fact that we have an EWG, much like we saw the fluorine doing that behavior, means that this is going to be the second most acidic compound; and then compound 1 is the weakest.2902

Because this has no resonance, this is the least stable and therefore strongest conjugate base; so this has the weakest parent acid; so compound 1 is the least acidic.2916

In this example, we have a combination of both inductive effects and resonance effects.2944

When we have both of those acting like we do in this case, it is the resonance effects that are generally going to win out.2951

Because the resonance effects are going to be ones that really delocalize charge and add stability.2958

How about if we turn it around in looking at an acid-base reaction and try to answer the question which is the stronger base?--same question is asking which is more basic?2968

We are comparing CH₃OH versus CF₃OH; and how are we going to answer this question?2981

Let's take the same approach we did for the acids; and the acids, we looked at the conjugate bases; so when we are comparing bases, we are going to look at the conjugate acids.2989

In other words, let each of these be a base and see where it takes us; so let's protonate in order to look at the conjugate acids.3001

Just like an acid donates a proton, being a base means that you accept a proton.3017

If we imagine this reacting with some acid, it can take a proton from that acid; so what does the conjugate acid of this base look like?3021

It will now have two hydrogens on that oxygen and just one lone pair; that gives us an oxygen with just one, two, three, four, five electrons; it wants six; so it will give us an O⁺.3032

If you add an H⁺ to a neutral molecule, you will end up with a positively charged molecule; so that's conjugate acid 1.3045

Conjugate acid 2, same thing except... this is an oxygen... except instead of a CH₃, we have a CF₃.3055

Let's draw our conjugate acids; and let's look for a difference in their stabilities; so just like we did for the cases of deciding who is a stronger acid.3067

What is the difference between these two?--well, once again, we see that we have fluorine versus hydrogen.3078

Let's make the note; let's start by stating the facts--fluorine is more electronegative than hydrogen.3083

That means that it withdraws electron density... and it withdraws electron density inductively, right?3097

The effect we have going on is something like this: each of those carbon-fluorine bonds are polar in the direction of the fluorine; it pulls electron density3105

That makes this bond polar as well; and it pulls electron density; it is an EWG.3118

Here is the question: is that a good thing or a bad thing?--in this case, we are looking at a positive charge.3126

That tells us that this oxygen is electron deficient; it is missing electrons.3132

What are those fluorines doing? they are pulling even more electron density away, making it even more positively charged; that doesn't sound like a good thing; that sounds like a bad thing.3137

What we can note is that this destabilizes; this destabilizes the + charge and makes conjugate acid 2... remember we are looking at conjugate acid 2... the less stable.3152

Less stable means more reactive; let's put that in there--more reactive and therefore stronger conjugate acid; because this is unstable, it is now going to be the stronger acid.3175

The stronger conjugate acid has the weaker parent base; conjugate acid 2 has the weaker parent base.3192

Who is the stronger base?--was our original question; who is more basic? 1 is the stronger base; because there is nothing to destabilize the conjugate; this looks much better.3211

Therefore this structure 1 doesn't mind getting protonated as much because it is going to a more stable conjugate.3228

Let's try another example of looking at who is the strongest base; amines are good bases.3239

Each of these nitrogens has a lone pair; and so we are asking which of them is most willing to be protonated--most likely to be protonated?--that would make it a stronger base.3249

It is possible we can look at the conjugate acids, but in this case, the answer is not going to be found in looking at those conjugate acids.3262

Because there is a different right away in looking at the stability of these three parent bases.3270

Let's take a look at resonance for example; this first amine has no stabilization; no resonance stabilization; nothing special about this structure.3278

But when we have a nitrogen attached to a benzene ring, this lone pair is now allylic or we call it benzylic when it is next to a benzene ring; and so it can have resonance.3294

That makes an N⁺ and a C^-; and are there any other resonance forms?--there sure are.3311

This resonance form is still allylic, and we can go on and move that negative charge around the benzene ring carbon; so there is a difference right away between these two amines.3320

Let's take a look at this third structure; this amine now has a substituent, a group, attached to this carbon; let's see what effect, if any, that has on the resonance.3332

Again, because it is benzylic, we know we can draw a resonance form here--NH₂⁺ and C^-.3343

This can continue down; we have our carbonyl down here--this is an aldehyde; and this can continue down to put a negative charge at this bottom carbon.3355

When we put this lone pair at this carbon right next to the carbonyl, look what can happen--it is now allylic to that carbonyl π bond.3376

We can have a new resonance form that moves the negative charge into the carbonyl.3385

Being able to draw a new resonance form is usually a significant thing, but let's take a look at this--does this look like it's significant resonance?3399

It actually is because now we've managed to put the negative charge on an oxygen--the more electronegative oxygen; that is a better place to be than putting it on the carbon.3408

There is an additional resonance form; we can actually move the negative charge up here as well; but I think again this is enough to demonstrate a difference by having this carbonyl group attached.3418

What we can conclude now is that this third structure is the most stable because of resonance; and if it is the most stable, that means it is the least reactive; and that makes it the weakest base.3428

Right off the bat, if we can find a significant difference in the stability of our starting compounds, then that can lead right away to a difference in their basicity.3453

One way we can describe this is we could say that if we were to protonate this, if we were to protonate this nitrogen, what does that do to all this resonance?3467

It takes away that resonance, doesn't it?--because that would take away this lone pair.3479

Protonation of this compound would be a bad thing; it would be unfavorable because we would lose all this resonance stabilization.3484

What we could say is that the... just a comment we can make about this is that the lone pair is unavailable.3490

This is not a very good base because the lone pair is not around to be protonated and to react with acid; we could say that it is tied up in resonance.3500

What is another good way to describe it?--we could say that it is delocalized.3516

Remember the actual structure is a hybrid of all these different resonance forms.3520

In all these other resonance forms, there isn't even a lone pair on that nitrogen; the lone pair is spread out over all these carbon atoms and this oxygen atom as well.3526

The lone pair is unavailable, and protonation... if we were to protonate it and add an H⁺, protonation would lose that resonance.3537

We would cause all that resonance stabilization to disappear; that is unfavorable.3553

This kind of resonance makes this particular amine very stable and very unreactive as a base.3559

Now when we are comparing these two, we would see that again this does have some resonance so this lone pair is also tied up.3568

Compared to this amine, this has no resonance stabilization; it has no delocalization of that lone pair.3575

Because these had no stabilization, this is the most reactive strongest base.3583

What we could say here, the way we would describe this and compare this to the other ones, is we could say that the lone pair is very available.3596

This lone pair has nothing better to do than to sit around and wait for an acid to react with; so that makes him a much stronger base.3608

This is described as an alkyl amine--when we have a nitrogen on an ordinary alkyl group, just a carbon chain like this.3614

These guys are called aryl amines--when you have a nitrogen on a benzene ring.3624

We will find that this trend holds true--that aryl amines are going to be much weaker amines than akyl bases; because of this resonance effect.3634

Finally, let's take a look at some common acids and bases now that we have seen the sorts of things that can cause a compound to be a stronger acid or a stronger base.3647

Let's summarize some of the things we should know moving forward.3655

Strong acids are defined as those with negative pKa's; these are things that when in water will completely dissociate; these will completely protonate water.3660

These are ones you should already be familiar with; things like sulfuric acid, nitric acid, the halo acids--HCl, HBr, HI (we will see those a lot throughout organic chemistry), C₃O⁺.3670

Any time you see any of these structures as part of your reaction conditions or as a reagent, you should immediately recognize them as strong acids.3684

Most often, it is going to tell you that the very first thing you are going to do in that reaction is protonate something.3693

These are fabulous proton sources, and they are very reactive, very unstable, and they are going to find a place to protonate.3699

Most definitely, if you don't already recognize these as strong acids, you can make up some flashcards so that they will be more familiar to you.3708

Weak acids are things that generally have a low pKa--somewhere below 16 or so.3716

What falls in this category is this arrangement when we have a carbonyl with an OH; these are known as carboxylic acids.3724

Carboxylic acid has a pKa somewhere around 5; it is quite easy to deprotonate a carboxylic acid.3732

Ammonium, NH,₄⁺, is also a decent acid, quite easy to donate a proton; as is water and just an ordinary alcohol.3738

An alcohol has a pKa closer to this 16; a carboxylic acid has a pKa down here closer to 0; but all these compounds are quite resonable to deprotonate; it is quite easy to deprotonate these.3748

Things that we might describe as very weak acids are those with even higher pKa's, above 16.3762

Something like a ketone; these protons next to a carbonyl, we will find are acidic; we can remove them, but it takes quite an effort; it takes a very strong base in order to do that.3768

Amines as well; an amine is not easy to deprotonate, but it can happen; of course, a nitrogen is less likely to lose a proton than an oxygen; so that is why we have a higher pKa here.3783

If we have a triple bond, a carbon-carbon triple bond with a hydrogen on it, this functional group is known as an alkyne--having a carbon-carbon triple bond.3797

Hydrogens on alkynes are also possible to protonate; if you have a very, very strong base, you can protonate in these positions.3807

The other functional group listed we can describe as extremely weak acids, these numbers are so high--above 40, that we could pretty much say that these are not acids.3816

It is going to be so difficult to remove these protons, that for all intents and purposes, we could say that no reaction is going to happen; there is no base that is strong enough to deprotonate these.3828

Of course, there are exceptions to that, but in introductory level organic chemistry, we are never going to see any reactions like this.3836

Here we have an alkene--is what we call it when we have a carbon-carbon double bond, and a hydrogen on an alkene carbon will never deprotonate; we will never be able to remove.3844

The same goes for an ordinary alkane carbon; putting negative charges on these carbons would be so totally unstable that we are never going to see that.3856

Having a triple bond makes it OK; we will see that down the road when we study alkynes; but alkenes and alkanes will never deprotonate.3869

While typically we are not going to be memorizing our pKa table, it i i437 Id when we study alkynes. s very good to have a familiarity with these various functional groups.3877

And know whether or not it's going to be possible or maybe even easy to deprotonate any of these hydrogens.3884

One last example--let's see if we were to ask about the direction of an equilibrium.3892

Let's say we were given a pKa table, and we looked at the following equilibrium.3897

When we looked in the pKa table, it turns out we were able to find a pKa value for every one of these species.3902

We know that for the direction of the equilibrium, we know that the equilibrium lies the direction of the weaker acid-base pair; but how do we use this information to decide who the weakest is?3912

There is something very important to remember for a pKa; remember what that A stands for?--that is the measurement for an acid; pKa's are for acids; pKa tells us something about the strength of an acid.3935

In the forward reaction, which species is behaving as an acid?--the acid is the one who donates a proton; and it looks like the ammonium is acting as an acid.3951

In this reaction, water is acting as a base; it is accepting a proton; so in this reaction, the pKa of water is irrelevant because water is not donating a proton in this reaction.3964

So 9 is the pKa for the forward reaction; and how about the reverse reaction?--who is the acid, and who is the base?--let's take a look at the direction of that proton transfer.3978

H₃O⁺ is going to H₂O; so that means H₃O⁺ is the acid in the reverse reaction; ammonia is reacting as a base in this reaction.3988

Which is the pKa number that is relevant to us?--just the H₃O⁺ pKa because that is the acid for the reverse reaction.3997

Now we've identified... we've narrowed it down to the two competing acids; we've looked up their pKa's.4005

Of course, we learned in this lesson how to predict the relative acidities.4011

So we should actually be able to, even without a pKa table, we should be able to decide which is the stronger acid, NH₄⁺ or H₃O⁺.4016

You think you'd be able to do that?--that's a very good exercise; I think it has to do with the difference in electronegativities of nitrogen and oxygen.4027

Which of these two species is more stable?--I want you to think about that.4034

But we see here that pKa of -2 means that this is the stronger acid; pKa of 9 means this is the weaker acid; and so where does our equilibrium lie... our equilibrium direction?4038

It lies to the left; you could say that the reverse reaction is favored or that the reaction lies to the left.4053

This concludes our lesson on acid-base reactions.4060

We will see you soon back at Educator; thank you.4063