Organic Chemistry Structures and Properties of Organic Molecules

Welcome back to Educator.0000

Today, we are going to be talking about the structure and properties of organic molecules.0001

The atoms in organic molecules are held together by bonds.0008

So we are going to talk first, about the types of bonding we can have.0010

Then we will look at the shapes of the organic molecules, and the physical properties that result from those shapes.0014

Atomic orbitals are the spaces--the regions of electron density, where you can find electrons about particular atoms.0021

For the carbon atom, we know we have an s orbital available for bonding, and we have three p orbitals that are available for bonding.0031

The s orbital is a spherically shaped orbital.0043

The p orbitals have this dumbbell shape; we have three of them.0049

One can be aligned along the x-axis, we call that the p_x orbital.0053

One can be along the y-axis, and one can be drawn along the z-axis here.0057

They are typically shaded different colors because that is representing the sign of the wave function.0064

The sign changes at the center there, at the origin, and that is why we typically show them in different colors.0069

But that is the general shape of these atomic orbitals.0076

Atomic orbitals are defined as regions with a high probability of finding the electron density.0080

These are defined by mathematical equations that are called wave functions.0085

I already mentioned that the sign changes at the node.0091

But the other thing to mention is that the node where we have a change of sign, the probability of finding any electron density at this point--it has a zero probability.0094

Now, if we want the atoms to come together to form molecules, the electrons can no longer reside in atomic orbitals.0107

They must be shared between atoms, in new orbitals, called molecular orbitals.0112

The way molecular orbitals are formed is we take two atomic orbitals and overlap them and form new orbitals.0117

What we will find is, if we take two atomic orbitals, they will combine to give two new molecular orbitals.0127

The reason we get these two results is because there are two possible combinations of those mathematical equations.0134

Let's take a look at an example--for example, the simplest molecule we have, just H₂.0140

Let's imagine taking a hydrogen atom and another hydrogen atom, each of them with a single electron, and bringing them together to form a bond.0146

We typically represent a covalent bond as a line drawn between the two hydrogens--we call this a sigma(Σ) bond.0154

What this line represents is--it represents two electrons that are shared in a molecular orbital, called a sigma(Σ) molecular orbital; that is what we refer to as a Σ bond.0161

Now, each of these hydrogens is coming in with an s orbital.0174

Here, we have the s orbital coming from one hydrogen, s orbital coming from a second hydrogen.0179

Let's imagine that, mathematically, those have the same sign.0183

What are the two possible ways that they can combine?0186

Well, they can combine in a positive way, if we add these two--we would describe that as being in-phase.0188

The result would be a molecular orbital that has a shape like this.0195

That is what we call a bonding molecular orbital--the Σ bonding molecular orbital.0200

This is what it looks like when you have a favorable overlap between the two atomic orbitals.0206

OK, but another way they can come together--if you can imagine, the s from hydrogen A subtracted from the s orbital of hydrogen B.0210

We would describe those now as being out-of-phase.0220

They would no longer have a favorable overlap; they would be held apart like this.0222

We call this molecular orbital, an antibonding molecular orbital.0229

We label it as Σ*, so a sigma with an asterisk there, we read that as sigma star.0234

What we have here now is a node where the sign has changed from one sign to the other.0240

That is going to be a higher energy orbital.0248

This a phenomenon that is true in general--that as you increase the # of nodes in your molecular orbital or in your atomic orbital, you are going to increase in energy.0252

It is going to be less stable, the more nodes that you have.0265

Another way to represent this combination of atomic orbitals to give molecular orbitals--is with an orbital diagram like this.0269

We can show one hydrogen coming in with its atomic orbital; the s orbital had a single electron in it.0276

On the other side, we have a second atomic orbital with one electron in it as well.0284

When they combine, and they form a bond to bring those two hydrogens together, we will now have two electrons in this system.0290

They will have two new possible places to reside--we have a Σ orbital, and we have a Σ* molecular orbital.0297

Where will those two electrons want to go?0305

Well, as usual, we are going to fill the orbitals starting with the lowest possible energy; we will pair them up, and we will put them in the lowest energy Σ molecular orbital.0307

That is a kind of graphical representation of what a Σ bond looks like.0318

What if we wanted to form a pi(Π) bond?--pi(Π) bonds come about by overlapping p orbitals.0326

Now, again, let's have the same idea--we bring in a p orbital from one atom and a p orbital on a second atom.0332

Because we have two atomic orbitals, we are going to get, as a result, two new molecular orbitals--because of these two possible combinations.0339

If the signs are matching at the top and bottom when we bring them together, that is where we are going to get our favorable constructive overlap.0347

We are going to get a region of electron density up here, a region of electron density below.0355

This is our bonding molecular orbital; we call this a Π bond or a Π MO.0362

That is what the electron density looks like in a Π bond.0369

You have no electrons in the area where you would find the Σ bond between the two nuclei.0373

Instead, you have a cloud of electrons above and below the Σ bond.0379

OK, how can these p orbitals, what is another way that they can combine?0384

If you imagine flipping one of them over, as if you were subtracting the two, you would now have a mismatch in all directions.0388

This is what we call the Π* molecular orbital.0398

You can see here that there is a node in this direction and a node in that direction, so this now is a very high energy orbital.0401

We can look at it analogous to the way we saw the formation of the Σ bond.0410

We can imagine the p orbital coming in with a single electron from one atom, a second p orbital with a single electron from another atom.0415

When we combine these two, we have two new molecular orbitals that have been formed.0423

One of a lower energy--that is the Π molecular orbital.0428

One of a higher energy--we call that the Π* molecular orbital.0432

We are going to locate these two electrons in the lowest energy orbital possible.0435

We are going to pair them up and put them in the Π molecular orbital.0442

Now, overall, when we take a look at the possible orbitals, then, in which electrons can reside in a molecule.0448

If we want to look at them overall, if we list them in the order of increasing energy, we will find that the most stable place for electrons to be is in the Σ bond, in a Σ molecular orbital.0457

These are going to be the most stable electrons, the lowest energy electrons.0471

That makes them pretty stable, pretty unreactive; these are great bonds to form.0474

They are going to be more difficult bonds to break.0478

These are lower in energy than a Π bond; remember a Π bond has that node so it is going to be higher in energy.0482

If you have a lone pair of electrons--we call those non-bonded electrons; those are in an n-type orbital, and that is a higher energy level still from Π bonds.0490

These two types of electrons--non-bonded electrons and Π electrons, because they are a little higher in energy, these are a little less stable than the Σ electrons.0502

They are going to be more reactive; and we are going to see lots and lot of chemical reactions involving the breaking of Π bonds or involving non-bonding electrons in reactions.0510

Just like the Σ bond is the lowest energy--Σ molecular orbital is the lowest energy, that Σ* then is the highest in energy; and the Π and the Π* have this relationship.0521

Of course, the antibonding orbitals are always going to be the highest in energy; and they are much, much less stable, and they are usually empty.0534

In a normal bonding situation, we will not put any electrons in those Π* or Σ* orbitals.0542

You might think, why is it relevant to even know about those?--why should we even consider those?--what relevance do they have?0550

Let me just very briefly give some justification for that; and that is under the topic of electronic transitions.0559

If we imagine a ground state molecule, a stable molecule, and we have a Π bond in that molecule.0567

We know we are going to have two electrons in that Π bond; they are going to be paired up.0573

They are going to be in the Π molecular orbital; the Π* molecular orbital is empty.0580

But if you were to introduce some light energy, especially ultraviolet light, that molecule can absorb some of that energy, if the energy amount is exactly equal to this energy gap between the Π and the Π*.0587

Then the molecule can absorb a photon of light and promote one of these electrons up to a higher energy level; so the Π* offers a place for the excited electrons go.0602

This is what an excited state might look like--an excited Π bond that has absorbed some light.0617

Now when this gets very interesting is that the energy gap here between the Π and the Π* will shrink, will become smaller if the Π bond in question is conjugated with other Π bonds.0622

To be conjugated means you have an alternatingΠ bond system: Π bond, Σ bond, Π bond, Σ bond, and so on... such as this system.0636

We have a Π bond here; and then we have a Π bond right next door, and another Π bond right next door.0644

What happens is, because there is a p orbital on each of these Π bonds--there is one Π bond and then here is another Π bond.0649

If you think about those p orbitals involved, we get this long system of Π bonds that are all overlapping and related to one another, and it ends up shrinking this gap.0657

This gap can become so small that, instead of requiring ultraviolet light to be absorbed by the molecule, you could move into a lower energy light--such as visible light.0670

As we increase the # of conjugated Π bonds, so as this system gets longer and more extensive, you will increase the amount of resonance stabilization--that is this interaction between adjacent p orbitals.0681

That is going to decrease the energy needed for this transition from the Π to Π *.0693

If visible light is what is being absorbed, then the molecule in question is going to appear to us as being a colored molecule.0701

So when you look at the structure of things like beta carotene or chlorophyll, you are going to see this extended Π network.0707

It's all because of the existence of those Π* molecular orbitals that the absorption of light is made possible.0716

Let's talk now about the case of hybrid orbitals, because it is not always so simple as having s orbitals and p orbitals overlapping to form Σ and Π bonds.0726

If you consider carbon, we know that carbon's atomic orbitals has four valence electrons.0735

The outermost atomic orbitals for carbon are the s and the three p orbitals.0742

If we were to fill in those four electrons, we would pair two of them up in the lowest energy s orbital; then we would move up to the p orbital, and we would spread those out as such.0748

This is what our stable electronic configuration looks like with carbon.0762

But we know that carbon likes to have four bonds.0766

The way it looks right now, you would think that maybe just because it has these two single electrons, it might be available for bonding; it can only have two bonds.0770

Well, it is easy enough to promote some energy--add some energy into the system and promote an electron so that, in fact, we can have a carbon with four unshared electrons all ready for bonding.0778

OK, so that leads us toward a carbon--that carbon can now have four bonds.0795

Experimentally, however, we observe that a molecule such as methane, CH₄, has four identical C-H bonds--now, how can that be?0808

If this is our atomic orbital picture, you would imagine that a bond formed by overlap with a spherical s orbital would be quite different than a bond formed by overlapping any of these p orbitals.0818

So how is it that we come up with four identical bonds?0831

Well, the answer to this mystery is that the atomic orbitals in carbon undergo hybridization.0836

This is the mixing of atomic orbitals to give new orbitals that we call hybrid orbitals.0845

What would it look like if you took a spherical s orbital and mixed it with a dumbbell shaped p orbital?0851

Well, you would get something that is still dumbbell shaped, but it is elongated on one of the sides; so it is more spherical and not symmetrical.0860

We are still going to have a node here; we are still going to have the sign change.0870

So it looks something like this, and that is what our hybrid orbitals are going to look like.0874

Now, it turns out that there is a variety of different hybridizations that carbon can undergo.0878

We will see that the type of hybridization that carbon undergoes depends on the # of groups around the carbon.0884

We are going to call these groups: regions of electron density.0893

We will look at some examples and see how we can determine the hybridization that is required.0896

Right here, we have three different sample carbon molecules, and each of these carbons has quite a different bonding situation.0904

Now, if you look closely, you will see each carbon has four bonds like carbon typically does; a stable carbon is going to have four bonds; but the electrons are arranged in different ways.0913

For each of these, let's look at them one at a time and see if we can... what observations we can make.0925

Now, for this molecule, even though there are two carbons, let's just consider one of the two carbons when we are describing the hybridization.0930

This carbon has one, two, three, four separate Σ bonds, but we are going to describe it as having four separate regions of electron density.0937

Such carbons are said to undergo a hybridization called sp3, or you could say this carbon is sp3 hybridized.0947

It is called sp3 hybridization because you mix the carbon's s orbital with all three p orbitals, and so the result is called sp3.0956

The result then--because we mixed four atomic orbitals, we are now going to get four new molecular orbitals.0968

Four new hybrid orbitals--and we are going to call these sp3 hybrid orbitals.0977

This carbon no longer has ordinary s and p orbitals; instead, it has four orbitals around it that are all sp3 hybrid orbitals.0989

Now, what kind of geometry are we going to get as a result?0999

Here, we can use VSEPR (Valence Shell Electron Pair Repulsion theory).1002

If you have four groups of electrons, and you want to spread them out as far as possible, you are going to have a tetrahedral geometry.1006

Meaning, we have about a 109.5 bond angle in a molecule such as this.1016

How about the next molecule?--again, let's look at just one of these two carbons and ask about its hybridization.1024

Again, this carbon still has four bonds, but if you look now, it is arranged into just three groups.1031

We have a single bond here, and a single bond here, and a double bond in this region; so this carbon has only three regions of electron density.1035

Such a carbon is described as being sp2 hybridized.1045

The reason we give it that name is because this carbon mixed one of the s orbitals with only two of the p orbitals to form its new hybrid orbitals.1051

That means, one of the p orbitals on that carbon is still there as an unhybridized, pure atomic orbital with a dumbbell shape, and we will see that in just a bit.1062

So what is the result of this hybridization?1073

Well, we took three atomic orbitals and mixed them, so now we are going to get three new hybrid orbitals... so we're going to get three new hybrid orbitals.1075

What are they called?--they are called sp2 hybrid orbitals.1086

Plus, this carbon has the unhybridized p orbital, so it has one p orbital.1097

So every carbon that is sp2 hybridized will have around it three sp2 hybrid orbitals and one p orbital.1106

What kind of geometry do you expect when we have three groups of electrons that we are trying to get apart from each other as far as possible?1115

How about if we spread them out in the corners of a triangle?--that would be the best; we call that trigonal planar.1123

Trigonal planar--meaning it's going to be a flat molecule, a flat carbon with bond angles of about 120 degrees like in the corners of a triangle.1130

OK, and our last example here has a carbon with four bonds, but these four bonds now only have two regions of electron density--a single bond in one direction, a triple bond in another direction.1140

If we have only two regions of electron density to accommodate--you might be seeing a trend here, and you might predict that the hybridization is going to be called sp hybridized.1153

It is called that because what we need to mix is the s orbital along with just one p orbital, which leaves the other two p orbitals there unhybridized.1163

As a result, when we've mixed two atomic orbitals, we get two new hybrid orbitals, so we are going to have two sp hybrid orbitals.1177

What else will this carbon have?--it will still have these two p orbitals--plus two p orbitals.1191

What about an atom that has just two groups of electron density that you want to get as far apart from each other as possible?1204

VSEPR shows us that we are going to try and get them just in a linear direction; 180 degrees from each other would be as far apart as you possibly get.1211

OK, so these are the basic geometries that we are going to be looking for: sp3 sp2, and sp.1223

Let's look at some examples on how we can apply this.1229

So here is an example--what if I asked you to assign the hybridization of the following molecule.1230

What is given here is just a backbone; so we can just get a little extra practice here drawing Lewis structures.1237

Then, if we needed to assign the hybridizations, here is what we have to do.1244

First, we have to make sure the Lewis structure is complete, including the lone pairs of electrons.1247

OK, and then hybridization is something that doesn't describe a molecule--it describes an atom, so hybridization is going to be for each atom.1252

In other words, what is the hybridization of this nitrogen, all three carbons, and this oxygen?--those each have a hybridization that we can describe.1260

Hydrogens cannot undergo hybridization because all they have is a single s atomic orbital, so there is nothing to hybridize with, there is nothing to mix.1271

The hydrogens always brings an s, but all the others can undergo hybridization.1281

Then what we are going to do is we are going to count the regions of electron density to see how many need to be accommodated.1283

What is a region?--that is going to be either a lone pair of electrons or a single bond, or a double bond, or a triple bond; each of those counts as one region.1289

OK, so let's do that; first, let's complete the Lewis structure.1297

If we want to, we can probably do this by inspection.1300

This carbon needs four bonds so we can make that a double bond.1304

Oxygen likes to have two bonds and two lone pairs; so we can add two lone pairs there; that satisfies both of those atoms.1308

And if we put a triple bond from this carbon, now this carbon has the four bonds that it likes.1317

Then nitrogen, in order to be neutral, has three bonds and one lone pair; so that looks like we have completed the Lewis structure quite nicely.1322

OK, so what is the hybridization for each atom?1333

Well, we can start with this carbon--what we are going to do is we are going to count the regions of electron density.1335

We have one, two, three, four Σ bonds; so we must have four regions of electron density; our hybridization, in that case, is sp3--is sp3.1339

How about the next carbon over?--we have one, two, three regions of electron density around here.1355

So three regions... three regions is not sp3; three regions is sp2.1361

Remember, the reason it is called sp2 is we mixed the s plus the two p's; so the way we got these three hybrid orbitals is by mixing one s and two p's.1365

That is where you are going to see the regions of electron density matching the # of atomic orbitals that went into this.1378

So just a little hint you can have is thinking about adding the exponents.1384

OK, so how about this carbon?1388

The triple bond and the single bond--that counts as two regions; so that is going to make him sp hybridized.1391

The nitrogen has a lone pair and a triple bond; so this nitrogen also has two regions; so this is also sp hybridized.1401

Finally, we look at this oxygen--the question would be do you think we should count these lone pairs together as one region or separate?1411

It comes down to whether or not you think they want to be very close together or they want to be far apart.1421

Of course, lone pairs are going to repel each other; so we count each lone pairs separately: one, two, three regions of electron density--that is going to be sp2 hybridization.1426

So this is how you are going to assign the hybridization for the atoms in a given structure.1438

Now, let's see what these molecules are going to look like in three-dimensions.1448

If we wanted to sketch particular atoms or molecules, how do we take that hybridization and translate it to a sketch and some Lewis structures?1451

First, let's take a look at the atomic orbitals--sp3, we said, was tetrahedral.1464

The way we usually draw a tetrahedron is we draw two bonds in the plane--so we will draw those as straight lines, about 109.5.1470

Well, none of us really knows how to draw exactly 109.5.1480

But we know what 90 degrees looks like; we know what 90 looks like; we know what 180 looks like; so 109.5 is somewhere between.1483

So just a little wider than 90 is usually what looks good.1490

OK, and to draw a tetrahedron then, the other two bonds--the other two hybrid orbitals are going to be out in this direction.1494

We are going to draw one as a wedge, and we are going to draw one as a dash; and that is showing some three-dimensionality.1503

The wedged bond is drawn as a wedge, because it is projecting out towards you, projecting up from the page.1511

The dash bond is dashed, because it is actually hidden from view; it is behind the page.1519

So anytime you see something dashed, it is going to be pointing away from you, and in the background.1523

OK, sp2 had one, two, three regions of electron density--so this is our trigonal planar.1529

We can draw those in the plane, and our bond angle is 120 here.1536

Again, 120 looks very much like 109.5, just a little wider than 90 degrees.1544

OK, but what else did--any atom that is sp2 hybridized had not only these sp2 hybrid orbitals; actually, let's label these.1551

It has three orbitals, each called an sp2; in fact, we can label these up here--these are our four sp3 hybrid orbitals.1561

In addition to these hybrid orbitals, remember, an sp2 hybridized atom also has a p orbital that was not hybridized;1574

That p orbital is going to be perpendicular to the plane of the sp2 hybrid orbital.1581

Remember, it is a dumbbell shaped orbital.1588

What we can do is maybe we can draw one half of it as a wedge sticking out at us and the other half as a dash sticking below the plane.1592

Because it's really--if this is the plane of the sp2, it is projecting straight out towards you and straight back.1602

So what we do is we tilt it just a little bit so we can see both parts.1608

OK, or what we can do is we can draw the p orbital very nicely.1612

OK, so this is a p orbital; and in fact, this is our other p orbital; both of these parts was from the p orbital.1617

We could draw the p orbital in the plane; but if we do that, that tips the rest of the molecule.1626

So now that my sp2 plane is on its edge, one bond or one orbital can be in the plane; and the other two become a dash and a wedge.1631

Either of these would be acceptable drawings for an sp2 hybridized atom, depends on what your focus is--whether you really want to see the p orbital, or you really want to see the bonds.1643

A lot of times, we like to draw the bonds in the plane, so the p orbitals get shifted as dashes and wedges like that.1653

OK, and finally an sp hybridized atom has two sp hybrid orbitals that are 180 degrees from each other, linear.1661

Let's draw this nice and big so we have so room to work with... so this is 180 degrees.1671

What else does an sp hybridized atom have?--it also has two p orbitals on this, if this is a carbon let's say.1680

We could draw one of those p orbitals in the plane, perpendicular to the sp line.1689

And then the second p orbital is going to be perpendicular to that still, orthogonal.1698

We could draw one as a wedge and one as a dash.1703

So here is one p orbital--top and bottom; remember, a p orbital always has two parts.1706

And the second p orbital--back and front, without getting too messy.1711

Sp hybridized atoms usually get a little messy; there is a lot going on.1716

So make sure you give yourself plenty of room when you are drawing these so you don't have to fit too much, stuck in there.1719

This is, and remember we had the p orbitals were aligned along the x, y, and z axises.1726

Whichever p orbital--in this case, it's the p_x as drawn, that was hybridized.1733

So the other two p orbitals are the ones that are still there--the one along the y axis, the one along the z axis.1740

What is it going to look like when we bring these hybridized atoms together to be a molecule, and how can we draw those 3D sketches?1748

Well, if you ever need to do a 3D sketch, the very first thing you want to do is determine the hybridization of the carbons involved or any atoms involved.1756

Because it is the hybridization that is going to determine the geometry.1764

Let's take a look at this first carbon and count the regions of electron density.1768

There is one, two, three bonds to hydrogen and then a fourth bond to carbon--four regions of electron density, means we have an sp3 hybridized carbon.1773

And this second carbon has the exact same arrangement with three bonds to hydrogen, one bond to carbon.1784

So if we wanted to draw a sketch, we would start with either one of the carbon--let's do the carbon on the left; and draw it tetrahedrally.1792

If we draw it tetrahedrally so that we have about a 109.5 bond angle; we have it attached to a carbon and a hydrogen--and where are the other two bonds?1803

Well, to be tetrahedral, we are pointing off in this direction--one is a wedge, one is a dash; and that is what a tetrahedral carbon is going to look like.1817

Then we come to the next carbon; he needs to be tetrahedral as well.1827

We've already drawn one of the bonds in the plane.1830

The second bond to be drawn in the plane can either be either up here 109.5 or down here 109.5, and those would both be OK.1832

Let's draw it up in this direction.1840

But if we draw this bond up here, then the dash and the wedge are down in this direction.1843

OK, and if we take a quick look at a model here--this molecule is called ethane.1853

Ethane then has two tetrahedral carbons, these are both CH₃.1860

These lines that lead then to nothing, we just assume, are hydrogens.1865

So what we have is the carbon-carbon bond is drawn in the plane; we have two C-H bonds that are also draw in the plane.1871

And then each of them has a hydrogen that is projecting out toward you--we draw it as a wedge; these two hydrogens are pointing away from you--we draw those as a dash.1878

So this is what our ethane molecule looks like.1886

OK, but notice that there is more than one 3D sketch we could have for this molecule because we can rotate around this bond--there is free rotation around this Σ bond.1888

So we would actually have some additional 3D sketches we can draw on.1898

For example, because we can rotate around here, we can have--let's just say or.1903

We can have this structure where the first carbon stays the same, but the second carbon has the other hydrogen that's in the plane pointing down here.1911

If this is where you draw the two bonds that are in the plane, then the ones that are out in the back are pointing up here.1923

So really, working with models really helps you understand what a tetrahedral model looks like, and it's going to be much easier to draw the 3d sketches to make them more realistic.1929

But the goal of your 3D sketch is that it actually has the same shape as what you are seeing.1939

OK, how about a carbon that has a double bond; this carbon has one, two, three regions of electron density.1949

So this is an example of an sp2 hybridized carbon; and again, this symmetrical molecule has two sp2 hybridized carbons.1958

To sketch that, we draw a trigonal planar arrangement for the three Σ bonds--hydrogen, hydrogen, carbon are the three bonds that we have there.1966

On the second carbon, we also have a trigonal planar; that is going to share the same plane as the first one; so all these are about 120 degree bond angles.1980

OK, and those are all the Σ bonds that are all being planar.1991

Now, we have to show what this Π bond looks like.1994

How do we get a Π bond?--that is by overlapping a set of two p orbitals.1996

Remember, every sp2 hybridized atom has a p orbital; it is right here.2002

We can kind of draw it as a wedge and a dash on the first carbon, and then a wedge and a dash on the second carbon.2006

Then, one way or another, we want to show some kind of interaction between the top half and the bottom half to represent a Π bond.2013

Remember, the Π molecular orbital had a cloud of electron density above the Σ bond and a cloud of electron density below, so this 3D sketch represents that pretty reasonably.2021

This is a model of this molecule--this molecule is called ethene.2033

You can see that if we draw it, one way to draw it is just completely to put all six atoms in the plane of the board.2038

Then the p orbitals are going to be right here, perpendicular on the first carbon and the second carbon.2047

Here is our p orbitals--they are going to be overlapping above and below the plane, so we are going to draw that as sticking out, sticking back.2052

OK, another possibility though is to draw this with the p orbital and the Π bond illustrated very cleanly--so p orbital on one, p orbital on the other, connection of the two above and below.2060

That's a nice looking Π bond, but if our Π bond is what we are showing in the plane of the board, what does that do to our Σ bonds, to our hybrid orbitals?2075

Well, now we've tilted the molecule like this, so you can see now that these two hydrogens are projecting out toward you--we need to draw those as wedges, and two of them are as dashes.2084

OK, now if you are viewing this molecule perfectly straight on, these are going to be hiding each other.2093

So what we do is we tilt the molecule just a little bit so that the back hydrogens are up a little, or we could tilt it this way so they are down a little; those are both OK.2099

But we need to show, let's show the two wedges coming a little bit in a downward direction but really they are projecting straight out toward you.2107

And then the two hydrogens slightly up, and they are back.2116

Both of those are good ways to draw the trigonal planar sp2 hybridized atoms, and this molecule of ethene.2121

Finally, we could take a look at this molecule. We have an example here of an sp hybridized, two sp hybridized carbons that are attached.2137

It is sp because we just have a triple bond and a single bond, so just two regions of electron density.2147

So if it's sp, that means we have 180 degrees linear Σ bonds, going straight out 180 degrees in either direction.2153

Now we need to represent two sets of Π bonds--two Π bonds which are really two sets of p orbitals; so we are going to draw that.2163

One of them can be in the plane, perpendicular to the sp--there is one Π bond.2175

And where is the second Π bond going to be?2181

That is going to arise from the p_z orbital on one carbon, front and back, overlapping with the p_z orbital, front and back, on the second carbon.2184

Again, without making a huge mess, we can maybe show the interaction of this top and back.2195

What we have is actually like a cylindrical type of a structure with this.2201

Here is another, here is a model of this, a settling unit--this molecule is called a settling.2208

This one piece represents the carbon triple bond; here is one carbon, and the other carbon, and we have the two Π bonds here.2214

So really, what we end up with is a cloud of electrons just completely surrounding this linear arrangement.2221

Then we have an sp hybrid orbital sticking out directly in one direction and one in the opposite direction.2229

Let's look at another example of doing a 3D sketch.2242

Like we said, we need to complete our structure, and then assign the hybridization.2246

Then we do a sketch; we might have a molecule with several different rotations and possible drawings.2251

So it is a good idea to try and draw the maximum # of atoms possible in the plane of the page; that is going to be the simplest drawing you can have.2260

Here is an example, and I see some things that are missing from my Lewis structure; I know that this oxygen needs some lone pairs.2270

OK, I know that this must be a triple bond here, between the carbon and the nitrogen; and this nitrogen must have a lone pair.2282

This is actually the molecule we had on the previous slide where we assigned the hybridization.2289

So you'll remember our hybridization for the nitrogen and the carbon were both sp.2295

The carbon and the oxygen were both sp2; and this CH₃ out here is sp3.2303

As we get a more complex molecule, we need to bring the geometries about each of the atoms together in a cohesive way.2312

What we can do is we can start with this middle carbon--right in the middle.2319

This carbon we know is sp2 so it is trigonal planar.2325

It has a carbon, and a carbon, and an oxygen for the three Σ bonds, for the sp2 hybrid orbitals.2329

Finally, it has a Π bond up to oxygen; so we are going to represent that.2337

Since we chose to draw these in the plane, the p orbital is going to be sticking out in back.2342

On oxygen, we have the same thing--a p orbital sticking out in the back; and then some kind of overlap here.2349

Some instructors like to show two electrons here to represent the two electrons being shared as the Π bond; and everyone has their different style.2355

OK, what else does this oxygen have?--well, he has the lone pairs.2365

Where are those lone pairs oriented?--because this is sp2 hybridized, and we have already shown the p orbital coming out and back from that, the lone pairs are in sp2 hybrid orbitals that are in the plane.2368

One way to draw that is to just draw it as a straight line, just showing geometry.2384

It's not a Σ bond, but that is one way you can draw it.2389

Or another way to draw it is just to draw kind of a hybrid lobe that is in the plane and put the lone pair in there.2394

We will handle the nitrogen that way.2401

OK, now let's come to this carbon; this carbon we said is sp3 hybridized, so that means tetrahedral.2404

We have already drawn one bond at this angle, so where does the second bond go?2410

Well, it needs to be 109.5; we can choose to go up here, or we can choose to go down here; either one is fine.2415

But if this is where our two bonds are that are in the plane, the other two bonds are down here--one is a wedge and one is a dash.2423

It doesn't matter which is which because this is another example where we are tilting the molecule just slightly... when we have a tetrahedral at them.2429

If these are the two bonds we are drawing in the plane, then the wedge and the dash are totally aligned with each other.2438

We tilt it a little bit to the left or right; it doesn't matter which way you draw the wedge.2447

And we have hydrogens in these positions.2452

OK, and then finally, we come to this carbon; he is sp hybridized.2455

OK, and here is where we really need to be careful with our geometry.2460

Because we know sp hybridized means linear, and we have already shown the first bond in this direction, so the second bond has to continue in that same direction.2465

Right?--this is the bond angle that has to be 180 degrees.2474

Be careful from thinking that linear always means horizontal on a page; linear means it is 180 degrees going in either direction.2478

This nitrogen then is also sp hybridized; and so the lone pair is also linear.2490

Let me show you, instead of drawing a straight line like this.2495

We could draw the hybrid orbital pointing off in this direction to indicate that that is where you would find the electron density that is in non-bonded electrons.2499

Finally, we want to show our p orbitals; now, where do the p orbitals go?2509

Because the 180 degrees is right here, they are linear; that means the p orbital needs to be perpendicular to that.2512

See how the molecule gets all tilted to its side?2518

We connect the top and the bottom, and then the second Π bond comes from the p orbital out in back.2523

So it all starts from recognizing what the hybrid orbitals look like, and then bit by bit, we can bring them together as molecules, as simple molecule and more complex molecules.2530

But ultimately, our goal is that the drawing that you have should very closely match the structure that you have.2543

Here, we have our linear C-N triple bond, and see how it is kind of pointed off in this direction.2557

Our C-O double bond is pointing straight up, and our CH₃ is over here; let's tilt this a little bit.2564

We chose to draw it with the hydrogen pointing in the plane, and then here is our dash and our wedge hydrogen.2571

So really, an accurate 3D sketch should very much look like the shape of the molecule in reality.2578

Let's try one more example; this one is a little trickier--this molecule is known as allene, and it is a classic molecule to study 3D sketches for.2584

When you draw an allene, you see that it has two carbon-carbon double bonds that are directly connected to one another.2594

The CH₂s on the end are sp2 hybridized, pretty straight forward, because they have three regions of electron density--two hydrogens and the double bond.2607

This middle carbon has a double bond here and a double bond here, so this is sp hybridized.2618

But what is interesting about this is usually our sp hybridized atoms are triple bonds, and so what is different is we have double bonds here.2624

Let's see what that sketch would look like; we could start with that middle carbon--we know he is sp hybridized.2634

The Σ bonds are going to be linear, 180 degrees from each other.2640

So this will, overall, be a linear molecule with respect to the carbons.2645

Then we can draw one of the Π bonds; let's draw the Π bond to the left--that is a p orbital here, and a p orbital here; and we show the Π bond.2650

Then we can finish off this carbon; we have already shown the Π bond.2661

If the Π bond is up here, what does that do to our two hydrogens?2665

These two hydrogens can no longer be in the plane because they have to be perpendicular or orthogonal to that p orbital, so the hydrogens are out and back.2670

How do we arrange this second Π bond?--the second Π bond has to come from the other p orbital that is on this carbon.2683

Remember every sp hybridized carbon has two p orbitals.2691

So the second p orbital is out and back; there must an aligned p orbital on the last carbon here, and that is how we form our second Π bond.2695

One Π bond is going to be above and below the molecule; the other Π bond electrons are going to be in front and behind the molecule.2707

Now, what does that do to this carbon?--where do the hydrogens go here?2715

Because we have shown the p orbital coming out in back and front and behind the plane, that means the sp2 hybrid orbitals are in the plane.2720

What we get for this molecule of allene is a very interesting structure.2731

Because it has a twisted structure where the hydrogens on the left are pointing out and back--a wedge and a dash.2734

We have a linear carbon in the middle, and the hydrogens on the right are going to be in the plane, up and down.2747

OK, here is one set of p orbitals, and then here is the second set of p orbitals, like this.2756

Let's talk about hybridization, and how it relates to resonance.2768

Because a lot of times, when we draw a Lewis structure, there may be more than one possible Lewis structure that we can draw.2774

We call this resonance--when there is more than one structure we can draw, more than one arrangement of electrons.2782

We know we have this resonance because we have an allylic lone pair--a lone pair that is next to a Π bond.2790

That lone pair can be shared, and the Π bond can move up, and we can get this new resonance form.2794

OK, the question that we have here is--what is the hybridization?2801

What does that mean for the hybridization and geometry of this nitrogen?2809

Well, you might think that--well, it has two hydrogens, a lone pair, and a single bond--that is four regions of electron density.2813

So ordinarily, you might expect that nitrogen to be sp3 hybridized.2823

But when you look at this second resonance form, we see that it has just three regions of electron density--the hydrogen, and the hydrogen, and the double bound.2827

So this nitrogen is clearly sp2.2835

We also know that in resonance, electrons are the only things that move and not the atoms or the Σ bonds.2839

In fact, this nitrogen cannot be sp3 in one resonance form and sp2 in another; in reality, it is sp2 the entire time.2849

This the one exception where the exercise of simply counting the regions of electron density can be a little tricky.2859

Because if you have a lone pair of electrons that can be involved in resonance, specifically an allylic lone pair of electrons.2867

Those electrons must be in a p orbital in order to have resonance delocalization.2875

Let's take a look at what this resonance picture really looks like.2882

If we consider the geometry of this molecule, and the resonance hybrid, we know that the middle carbon is sp2 hybridized.2888

The middle carbon is trigonal planar, and sp2 hybridized.2899

But the nitrogen is also sp2 hybridized, and the oxygen is sp2 hybridized.2903

What we have is a p orbital on the oxygen, a p orbital on the carbon, and a p orbital on the nitrogen.2911

Those three p orbitals are aligned, and that is what's sharing the electrons.2918

That is what a resonance picture looks like--you have a partial bond; sometimes it is drawn as a dashed partial Π bond here.2924

Because the C-N relationship has a partial double bond character, and the C-O bond has a partial double bond character.2932

We can draw it like this; so this is what the actual hybrid looks like.2941

This is what resonance always has--is electrons that are involved in p orbitals; so this lone pair must be in a p orbital.2944

Which means it cannot be sp3 hybridized; that has no p orbitals.2959

It needs sp2 hybridization, so one of p orbitals is free; and that lone pair then can interact with the other p orbitals from the carbon-oxygen double bond.2963

OK, so because we are not able to move our atoms, it cannot be tetrahedral in one case and then trigonal planar in the other, because the atoms are fixed throughout resonance.2975

It is always going to be sp2 hybridized; we cannot have sp3 on any atom that is undergoing resonance.2985

Now let's talk about physical properties that we might expect for organic molecules, and perhaps how they relate to their 3D structure as well.2996

Physical properties are things like water solubility, boiling points, that sort of thing, melting points.3005

All these properties are based on intermolecular forces--forces between molecules.3010

For example, if we take a look at a sample of methane, that is a liquid.3016

If we heat that, and we heat it until it reaches its boiling point, these molecules are going to gain enough energy to be separated from one another.3021

Each of these methane molecules, CH₄, will now no longer have the close contact that they did in the liquid stage.3035

They will be freely flowing and independent in the gas phase; and this is what the phenomenon of boiling looks like.3043

OK, notice, just a quick point--notice that my methane molecule has stayed intact during the boiling process.3055

We're not breaking up the carbons and the hydrogens; we're not having hydrogens and carbons breaking up.3062

So we are never breaking our Σ bonds, our Π bonds, our intermolecular bonds.3069

But what we do need to do is we need to disrupt the intermolecular attraction between the methane molecules.3074

OK, if those CH₄ molecules are strongly attracted to each other, if there is a very strong attraction to one another, what will that do to the boiling point?3082

It's going to be more difficult to pull those molecules apart from one another.3092

So we are going to have to put more energy, and to do that, meaning, increase the temperature even more.3096

That's going to increase the boiling point if we have a strong attraction between molecules.3100

What are the types of interactions--we call these nonbonding interactions between molecules--what are the types that we can have?3107

Well, there's three that we will be discussing; one is called dipole-dipole.3114

This is the kind of interaction that we get when we have polar molecules.3118

Another type of nonbonding interaction, that's actually quite strong, is called hydrogen bonding.3125

That's when we have either an O-H or an N-H as part of our molecular structure.3130

The last one is called van der Waals forces or London forces, and this is what we look for especially when we are dealing with nonpolar molecules.3138

OK, so let's look at each of the three of these and see what effect they can have on the physical properties of molecules.3149

A dipole-dipole attraction is what we have when we have two polar molecules near each other.3159

A polar molecule can be represented like this, where we have two poles.3166

We have a part of the molecule that is electron-deficient or partially positive and a part that is electron-rich that is partially negative.3171

You can imagine that if we had two such molecules, they are going to be aligned in such a way that the electron-deficient region is going to be attracted to the electron-rich region, and so on.3181

So when we have a polar molecule, there is quite a strong attraction between polar molecules.3195

Just like you would imagine the attraction that two magnets have to each other, because they each have poles.3201

Let's look at a series of compounds here.3209

OK, how about the boiling points of sodium chloride; this has a boiling point of 1413 degrees C.3214

This next molecule has a boiling point of 76; and this last one has a boiling point of 36.3224

Let's see if we can use the concept of dipole-dipole interaction to explain this.3230

OK, first of all, we have sodium chloride.3237

This is actually not even a polar molecule; we know this is an ionic molecule.3240

That can kind of be described as the most extreme dipole you have when something is so polar that it is actually ionic.3248

So we call this attraction, not a dipole-dipole, but we call this an ion-ion interaction.3255

The sodium cations are so strongly attracted to the chloride anions that you have to put in a huge amount of energy in order to boil this.3264

So even to melt it is very difficult; and if you want to put it in the gaseous phase as separate ions, it is going to take incredible heat.3278

If you ever tried to salt something that you still have in the frying pan, you might notice that when the salt crystals land on the hot frying pan, they stay totally solid.3284

It would be very difficult to heat those to melting.3296

How about these two?--we have a 76 degree boiling point here, and a 36 degree boiling point--how can we explain the difference there?3301

Well, because of the presence of this carbon-oxygen bond, oxygen is much more electronegative than carbon, so it pulls electron density toward itself.3308

That makes this a polar bond; and this becomes then a polar molecule.3317

This last molecule--this line drawing just represents a carbon chain.3325

We have carbon-hydrogen bonds and carbon-carbon bonds--all of those bonds are nonpolar so, as a whole, this is a nonpolar molecule.3329

That is why it has such a low boiling point.3338

So the overall trend here is that as you increase polarity, you increase your boiling point.3341

One last thing to note is--when we talk about the idea of having a polar molecule.3352

Remember you need to not only look for polar bonds, but you also need to consider the geometry of the molecule.3358

For example, if we have carbon tetrachloride--is a tetrahedral molecule.3368

Well, how do I know that it's a tetrahedral?3376

Because this carbon has four chlorines around it; that makes it sp3 hybridized.3379

This is where our hybridization is going to pay off--we will be able to understand the geometry of the molecule.3386

Even though carbon-chlorine bonds are polar, we have these polar bonds pointing in equal and opposite directions so we have no net dipole; so as a whole, this is a nonpolar molecule.3392

In making the decision on molecular polarity, we need to consider not only the polar bonds, but also the geometry of the molecule.3416

Of course, this, since it has no polar bonds, the geometry of the molecule is irrelevant, because there could be no net dipole moment here.3424

OK, how about hydrogen bonding?--this is really another example of a dipole-dipole; it's just an extremely strong one.3435

That is when you have a hydrogen on any of the three most electronegative atoms on the periodic table.3442

An N-H bond is going to be an extremely polar bond in the direction of the nitrogen, which puts a lot of ¥ä+ character on the hydrogen, ¥ä- character on the nitrogen.3449

Oxygen is even more electronegative so that is an even stronger dipole.3462

Of course, HF has this as well, but HF is its own molecule; it cannot be part of a larger molecule.3467

So that is not as relevant when we are looking at the boiling point and the physical properties of organic molecules.3477

For example, let's take a look at this series of compounds.3485

We have water; we have a carbon chain with an ?OH; we have a carbon chain with an oxygen in the middle; and then we have a carbon chain with no oxygens.3488

OK, so right off the bat, we can see that we have another example of something that is completely nonpolar down here.3499

That's why we have such a low boiling point, -42 degrees C.3507

If we compare these> two molecules, they are isomers of each other--they both have two carbons and an oxygen.3514

So we expect them to be polar molecules; but because this one has an O-H bond, that means it can hydrogen bond--can form hydrogen bonds.3522

In other words, one molecule of ethanol is very strongly attracted to another molecule of ethanol, and so it's going to be more difficult to separate them.3536

OK, this molecule has an oxygen, but there is no hydrogen directly attached to that oxygen so it cannot hydrogen bond.3545

Look at this huge difference in boiling points--about 100 degree difference because of the lack of that hydrogen bonding--so really, really strong effect here.3553

You might also think--well, is this also nonpolar?--do these dipole moments cancel?3561

This is a polar bond, and this is a polar bond; but do they cancel each other?3569

Would we describe this also as a nonpolar molecule?3573

Well, what do we expect for the hybridization of that oxygen in the middle?3577

We know that oxygen has two lone pairs and two bonds, so this is actually going to be sp3 hybridized.3581

Rather than calling it tetrahedral, now we call it a bent molecule because really all we look in geometry is the location of the atoms.3590

In this bent molecule, this polar bond is pointing up and to the right; this polar bond is pointing up and to the left.3599

So the left and right components do cancel, but they both have this upward component with nothing to cancel so this is actually a polar molecule.3606

That explains why this molecule with the oxygen has a higher boiling point than the molecule with all just carbons and hydrogens--there are increase in polarity.3615

OK, and look at water; we know that water has a boiling point of 100 degrees.3627

Why is it such a huge boiling point?--a very high boiling point for such a small molecule.3634

Because it has two O-H bonds, so this has lots of hydrogen bonding going on with water.3639

So it has an extremely high boiling point there.3640

By having two O-H bonds, one way we can describe water is being a hydrogen bond acceptor.3655

Because it has an oxygen that can participate in hydrogen bonding, and it is a hydrogen bond donor.3670

You'll hear those words when a hydrogen bond is being described; and water participates in both directions.3678

What you can show is that we have this water molecule because these bonds are so polar, the very large ¥ä- charge on this oxygen, very significant partial positives on these hydrogens.3684

So what is the interaction going to be between water molecules?3696

We are going to have a very strong attraction between the ¥ä+ hydrogen of one and the ¥ä- oxygen of the other.3700

OK, and it's so strong that we actually draw this dotted line here to represent a hydrogen bond.3711

This water molecule can donate a hydrogen bond to another water molecule; it can also use its oxygen to form a hydrogen bond with another water molecule.3721

As you can imagine, we just get this huge network of water interacting with other water molecules.3732

Huge interaction between the water molecules that really accounts for a lot of the physical properties that we have for water.3741

The fact that it has such a high surface tension, has such a high boiling point, it's a very ordered structure.3748

Water really likes interacting with other molecules of water because of this hydrogen bonding.3755

So hydrogen bonding is one of the most important things to look for when you are considering such a huge property as boiling point.3760

Because it's such a huge--you can see from these numbers.3766

There is a really significant rise in boiling points when you can have hydrogen bonding.3768

OK, and finally, let's consider nonpolar molecules.3773

If what we are trying to do is disrupt intermolecular attractions when we are, let's say, boiling a molecule sample--whatever could possibly keep nonpolar molecules together?3776

Why aren't they all just gases all the time?3788

Well, we have something known as van der Waals or London forces; these are temporary dipoles that can occur in nonpolar molecules.3792

Let's take a look at a nonpolar molecule; just represent it here.3802

Let's imagine that, at some point, there might be an uneven distribution at the surface--so an uneven distribution of electrons at the surface.3806

They are not perfectly, evenly distributed--maybe we have a slight excess of electron density here or a slight reduction there.3827

Well, when that happens, then if it sees another molecule, it is going to induce the opposite temporary dipole in the neighboring molecule, and we are going to get an attraction there.3836

OK, so what we get, as a result of this, is we get a temporary attraction... gives a temporary attraction.3852

The key is this is happening at the surface.3869

What is observed is that the greater the surface area a molecule has, the greater the van der Waals attractions it is going to have.3873

One observation then, is, all other things being equal, as you increase your molecular weight, you are going to increase the boiling point.3882

Increase your molecular weight, you are going to increase the boiling point; because you will simply have more surface area that can interact.3891

Again, let's take a look at a series of compounds here.3899

These are all hydrocarbons, meaning they contain nothing other than carbon or hydrogen; and so these are all nonpolar molecules.3903

We see that this one has four carbons--one, two, three, four; and the one on this end has five--one, two, three, four five.3913

Simply having more carbons, a higher molecular weight, means it is going to have a higher boiling point.3920

OK, but let's take a look at this middle structure--this is interesting because this is also a five carbon structure, and also nonpolar.3927

But we see a pretty significant difference in boiling point here--how can that be explained?3935

What we have in this case is we have--this first molecule is an example of a straight chain structure, more linear.3941

This second molecule is an example of a more highly branched molecule, and the more branched a molecule becomes--it almost becomes more spherical.3951

This is actually quite a spherically shaped molecule.3959

If you imagine how straight chain molecules will interact with one another, we see that they can have a lot of interaction because they have larger surface area.3963

So this has more surface area.3975

Therefore, we're going to have more van der Waals.3982

What is that going to do to the boiling point?--it is going to increase the boiling point.3986

OK, but a branched molecule, if it is a spherical molecule--imagine throwing those together; they have very little surface area, less contact.3991

That is why it has a lower boiling point; it's easier to take these molecules apart.4002

A lot of times, for van der Waals forces, I like to think of them like Velcro strips--the hook-loop fasteners.4010

If you have a long Velcro strip that is tied together, it is going to be really hard to pull those because there is so much contact surface area.4018

But if you have something like the little dart game, where it's a ping pong ball that has some Velcro on it.4027

You know it's very hard to make those stick and very easy to take them apart.4034

That's because it has such a smaller amount of surface area.4037

If you imagine nonpolar molecules as being kind of sticky, then the more surface area you have, then the higher boiling point you are going to have.4041

Now one thing to mention--I'm not going to talk much about melting points.4049

Less contact is something that is going to affect the boiling point as shown.4057

But we can make a note here that, if a molecule is highly spherical, it will have better packing of the crystals when it is a solid.4061

In the solid state, if you have something that is spherical, you can pack them very tightly.4078

What happens as a result--it's going to be difficult to disrupt that crystalline structure, and you will see an increase in melting points.4084

So melting point is a little more complex than boiling.4093

Because, in addition to just the attraction between molecules, you also have the issue of how well packed, how tightly packed the molecules can fit to one another.4098

Besides boiling point, another physical property that we can take a look at is water solubility.4114

The key to looking at solubility is that like dissolves like.4121

In order to be soluble in a solvent, a solute should have the same characteristics as that solvent.4128

When we consider water, we know that water is polar.4135

Water can hydrogen bond so it's really going to like molecules that can do the same, and they are going to have good water solubility.4139

Again, let's look at a series of compounds here.4146

We've listed these in, kind of, decreasing order of water solubility.4149

This first one is the most soluble in water; in fact, we describe this as being miscible with water.4153

Miscible means that it is soluble in all proportions.4161

This is ethanol--you can never add to much ethanol to water and have it separate out as a separate layer; it will continuously dissolve in water.4168

OK, but looking at this next structure, when you have three carbons--is it now less soluble?4177

In fact, you can only dissolve 8 grams of this structure in 100 mL of water.4182

As soon as you try and add more than 8 grams, it is going to separate out.4187

In fact, because this is a liquid, it is going to separate out as a second layer.4193

It has some water solubility, but a limited amount.4196

In this compound now, when we have one, two, three, four, five, six carbons, we say that this is the least soluble of all them.4200

Less than a gram of this will dissolve; very, very little will dissolve in a 100 mL sample.4207

What we have here is, as we are increasing the # of carbons, increasing the length of this carbon chain--what sort of characteristics do we associate with a carbon chain?4212

Well, this is something that is nonpolar.4224

As you increase your nonpolar component, nonpolar part of the molecule, you lower your water solubility.4228

It's going to be nonpolar; you could even describe this as being hydrophobic.4235

Because water is so polar, that something that is nonpolar will not like water at all.4242

You could even call it greasy; things that have very long carbon chains are things that are going to seem greasy to us--something that won't want to dissolve in water.4247

Here you can see that as you increase your polarity, you increase your solubility.4259

Of course, the examples I've shown all have ?OH groups, which means they can also hydrogen bond with water.4269

Even though this has a very long carbon chain, it is as least slightly soluble.4276

Because this ?OH group is extremely hydrophilic, really likes water because it is not only so polar, but it can interact with water and hydrogen bond with it.4280

OK, and finally, let's take a look at one more example--this molecule is called acetone; this is also miscible with water.4290

You may have worked with acetone before in the lab; it is a commonly used solvent, especially for cleaning glassware.4299

Nail polish remover is usually acetone based, and that is an aqueous solution so it is a fairly common organic solvent.4305

It too is miscible with water so it is completely soluble in water.4315

OK, why is it so well dissolved with water?--well, one thing is that it's polar.4319

It's not only polar because we have a dipole moment with a carbon-oxygen bond, and oxygen being more electronegative than carbon.4326

But also, because this carbonyl, this C=O double bond, can have resonance;--that means there is two ways to draw acetone.4336

The second resonance form is also contributing to the actual structure of acetone, and this resonance form actually has a + or - charge.4347

It has some ionic character there; this carbonyl carbon here has a very highly ¥ä+ character; the oxygen has a highly ¥ä- character.4357

Again, this is an extremely polar molecule; that is what makes it so much like water even though it is an organic molecule.4369

OK, and in addition, it can also be a hydrogen bond acceptor.4376

Acetone itself cannot hydrogen bond with other molecules of acetone because it has no ?OH group, so acetone is a fairly low boiling solvent.4382

If you ever got some acetone on your skin with nail polish remover--it feels cold, it evaporates very quickly, it's quite low boiling.4393

But, because it has an oxygen with lone pairs, it can participate in hydrogen bonding with one of these+ hydrogens in water.4401

We still can form hydrogen bonds because it is a hydrogen bond acceptor.4415

Water solubility is another physical property that we cannot explain in terms of polarity and perhaps molecular geometry sometimes too.4422

One other thing we can look at with organic molecules, is the fact that we can have isomers of a given structure.4434

Two isomers are compounds that are related because they have the same molecular formula.4444

In other words, they have the same # of carbons, hydrogens, nitrogens, oxygens, and so on; but they are different compounds.4450

Traditionally, the definition of isomer, the initial one that most of us learn in general chemistry are known as constitutional isomers or structural isomers.4457

These are isomers that have different connectivity; so you have the same atoms, but they are put together in a different order.4468

For example, here, we have a four carbon chain and a four carbon chain; but here, the double bond is in the first two carbons, and here, it's between the middle two carbons.4476

These are two different compounds, but they are isomers of one another.4486

They can have quite different connectivity.4493

Here, we have four carbons and an oxygen.4496

We can maybe have it with an ?OH group, or we can have an oxygen sandwiched between carbons, or that ?OH group can be on a different part of the carbon chain and so on.4500

There are several other isomers that we can draw.4508

But what relates them is they all have the four C₄H₁₀O which relates them as isomers.4511

One last example here--both of these structures look quite different, but they both have five carbons and six hydrogens and one oxygen.4518

Isomers can be very, very different compounds from one another as long as they have the same formula.4527

Now, there is another kind of isomerism called stereoisomers, and that is something we are going to be getting into quite a bit in organic chemistry.4534

This is where you have molecules with the same formula and the same connectivity, OK, but they have different spatial arrangement.4543

Only by looking at the three-dimensional shape can you tell that they are different molecules.4554

For example, both of these molecules have one, two, three, four, five carbons; and they have a double bond between the second and the third carbon.4561

One, two, three, four, five carbons, double bond between the second and third carbons.4571

These molecules will actually have the same connectivity, the same order of atoms, and even the same name when we learn about naming.4575

But these are not interconvertible; these are two different structures.4583

Remember, you can't rotate around a carbon-carbon double bond.4586

So these are unique compounds, but the only way you know that they are unique is when you look at their shape in space.4591

Here, the two carbon groups are on the same side of the double bond, we describe that as being the cis isomer.4597

Here, the carbon groups are on opposite sides of the double bond, we call those trans isomers.4607

This is just one brief example into the idea of having stereoisomers; again, more to come later.4613

Sometimes, this is called cis-trans isomerism; it is one type of isomerism that you can have.4619

Finally, as an overview and preparation for the material to come, let's take a look at the idea of having functional groups.4628

This is where we have an arrangement of atoms, put together in a common way.4637

Much throughout organic chemistry, we usually study it by going from one functional group to the next.4646

For each functional group, we will learn that molecules that have this specific, particular arrangement of atoms, very often have the same reactivity, the same behavior.4651

And so, we study those classes of compounds, one at a time.4662

The simplest type of carbon compound we can have is just called an alkane.4666

An alkane, like methane, is when we have nothing more than carbons and hydrogens and nothing else--no double bonds, no other atoms.4671

We could abbreviate just as an R group with a hydrogen; we use the letter R to represent any kind of carbon chain.4680

So you'll see the letter R a lot, to represent just a generic carbon chain.4691

For example, we have this that is called methane.4696

For each of these functional groups, not only will we be studying their behavior and their reactivity, but we will also be studying their nomenclature--how do we name these compounds?4698

Another group that we will be studying early on is called the alkyl halides.4708

That is when you have an alkane with a halogen now attached.4711

This, for example, is R with a chlorine on it because the halogens are often denoted just with the letter X.4716

We could use the abbreviation ?RX to represent any sort of, a wide variety of alkyl halides, and we'll learn the nomenclature for that as well.4723

This would be called the chloromethane because it is a methane with a chlorine attached to it.4734

Aromatic compounds are interesting compounds that have this characteristic six-membered ring with three alternating Π bonds.4740

Those compounds are called aromatic; this specific compound is called benzene.4751

Again, at some point, throughout the course of organic chemistry, you are going to learn about what it means to be an aromatic compound, and what compounds other than benzene itself could be described as aromatic.4758

The abbreviation, -AR, represents some aromatic group, like this benzene six-membered ring.4771

If you were to have a double bond, a carbon-carbon double bond in this structure, we call those compounds, alkenes.4779

If you have a triple bond, we call those alkynes.4786

When we study the reactions of carbon-carbon Π bonds, we usually kind of look at alkenes and alkynes closely together because they have very similar reactivity.4790

An alcohol is what we call it when we have an ?OH group attached to a carbon chain, so ROH represents a variety of alcohols such as methanol.4802

This is when you have methane, and you put an ?OH, we now call it methanol.4812

This is kind of an example of many of the suffixes we will see in our nomenclature, depending on how the name of the compound ends, it tells you what sort of functional group.4818

Every time that we see an ?ol, it tell us that somewhere in the structure, we have an ?OH group.4828

An ether is what we call it when there is an oxygen, but it's not attached to a hydrogen; it just has a carbon on either side; so this called an ether when you have R-O-R.4836

There is a wide variety of functional groups that contain the C=O double bond; the C=O double bond is called the carbonyl.4849

This is a hugely important functional group in organic chemistry.4858

As you can see, it is ubiquitous all throughout organic chemistry in organic molecules and natural products and naturally-occuring compounds4863

Depending on what two groups are on each side of the carbonyl, we can break these compounds up into different classes of functional groups.4872

The simplest ones are when we have just carbons on either side, or maybe a carbon and a hydrogen.4884

When there is a hydrogen, we call it an aldehyde; and when there are two carbons on either side, we call it a ketone.4890

Again, ketones and aldehydes--we will look at their reactions and nomenclatures and all that hand in hand.4896

If you have an ?OH and a carbonyl on the same carbon, this is no longer an alcohol.4902

This entire functional group, the entire set of atoms is taken together as a single functional group--these are called carboxylic acids; they would be abbreviated RO₂CH.4908

If you expand RO₂CH, one of these oxygens has to be a carbonyl double bond and the other is an ?OH.4919

Acetic acid is an example of a carboxylic acid; that is this particular structure right here.4927

These last groups, this last set of compounds are called carboxylic acid derivatives.4935

These are compounds that are derived from carboxylic acid so these are typically studied hand in hand, generally around the same time.4944

If you have a halogen, we call that an acyl halide or maybe an acid chloride, for example, if it's a chlorine.4956

If you have an -OR group here, it's no longer an ether; when there is a carbonyl and an -OR, we call that an ester; notice how similar these names are.4963

A lot of this nomenclature really comes to being very familiar with subtle differences in names.4974

When there is a nitrogen, we call that an amide.4981

When we have an oxygen surrounded by a carbonyls, we call that anhydride.4984

When we have a -CN triple bond in a structure, we call that a nitrile.4990

All those are carboxylic derivatives, and like I said, that this is just a preview of what is to come.4993

Because bit by bit, we will be learning about the nomenclature of each of these functional groups, and then the reactivity and their general behavior and physical properties.4999

So thank you for joining us as at Educator.com, and we will see you again soon.5011