General Chemistry Bonding & Molecular Structure

Hi, welcome back to Educator.com.0000

Today's lecture from general chemistry is going to be on bonding and molecular structure.0002

As always, we will start off with a brief introduction.0010

After that, we are going to then go into the different types of chemical bonds.0014

After types of different chemical bonds, we are then going to discuss yet another periodic trend0019

which is called electronegativity and how it relates to something we call bond polarity.0024

The next two bullets deal with what we call Lewis electron dot structures.0030

These are going to become quite prominent in our future discussions in general chemistry.0035

After discussing Lewis structures, we will go into something we call resonance and formal charges.0042

After that, we then jump into a type of bonding theory called VSEPR theory or valence shell electron pair repulsion.0049

Followed by molecular polarity and bond properties.0058

As usual, we will finish it off with a brief summary followed by some sample problems.0064

In the previous lecture on molecules and atomic structure and periodic trends, we introduced the concept of valence electrons.0072

Remember valence electrons are an atom's or ion's outermost electrons.0082

That is they are the farthest away from the nucleus and the least tightly held.0088

It turns out that we have highlighted them for a reason.0094

It is because they are the primary electrons involved in chemical reactions and in chemical bonding.0097

In this lecture, we are going on to focus our attention on valence electrons and their qualitative representations.0107

Before going into that, we need to first discuss what types of bonds occur.0116

Really for our purposes for general chemistry, there are two types of bonds.0124

The first type is an ionic bond.0127

This exists between a metal and a nonmetal or between a cation and an anion.0131

We actually touched upon this when we first learned about polyatomic ions and naming ionic compounds.0138

Remember that in an ionic bond, it is really an electrostatic attraction occurring,0145

resulting in complete transfer of electrons always from the metal to the nonmetal.0150

That is not the type of bond we are going to be focusing on in this lecture.0157

We are going to be focusing on what is called a molecular bond.0161

A molecular bond exists between nonmetals.0166

Remember we learned how to name molecular compounds earlier on.0168

In a molecular bond, electrons are not being transferred, instead they are shared.0174

This sharing of electrons occurs mutually or not.0180

Because of this sharing, a molecular bond is also referred to as a covalent bond0184

which is how I am going to start to refer to it from now on, a covalent bond.0189

Next we are going to examine this sharing of electrons and0194

see the extent of how mutual or not the covalent bond is.0197

In order to address the question whether or not the sharing of electrons in0208

a covalent bond is mutual or not, we need to now introduce another periodic trend.0213

This periodic trend is what we call electronegativity.0218

I am going to abbreviate it as EN for short.0222

Electronegativity, it is literally a quantity, it is a value that describes the ability of0227

an atom to attract electrons to itself from a chemical bond in a gas phase.0232

Again it is the ability of an atom to attract electrons to itself in a chemical bond.0240

This is usually determined in the gas phase.0246

Electronegativity is more or less like your GPA which is on a scale of 0 to 4 with 4 being the highest.0250

High electronegativity values tend to be those elements that have very strong nuclear charges.0260

They attract electrons quite readily to themselves.0266

Those atoms that have very strong nuclear charges also tend to be very small.0272

What we are going to do right now, we are going to reexamine our periodic table.0277

Let's go ahead and repeat the periodic trends that we have already been through.0283

Going this way remember is atomic size increasing.0288

Going this way it was increasing ionization energy.0295

If elements that have high electronegativity values have strong nuclear charges and tend to be small,0305

that means the highest electronegativity values are going to be right here.0311

In other words, electronegativity is going to parallel ionization energy.0315

Again electronegativity is going to be parallel to ionization energy and opposite to atomic size.0323

There is a little acronym we can remember to help us with this.0330

F-o-n-c-l-b-r-i-s-c-h, fonclbrisch is the pronunciation of this little silly acronym.0334

But it is a nice crude way of remembering the order of electronegativity0345

in case you don't readily have a periodic table in front of you.0350

Fluorine is going to be the highest electronegative atom.0354

Carbon and hydrogen are going to be the lowest; fonclbrisch.0358

We are going to classify a bond with specific words depending on the0366

difference in the electronegativities of the two atoms that share the bond.0373

If the difference in electronegativity is great, then the sharing of electrons we say is not mutual.0379

That means the atom with the higher electronegativity gets a greater share of the electron density within the covalent bond.0387

Such a bond is what we call a polar covalent bond.0395

We schematically indicate polarity using partial charges and a dipole arrow.0400

Let's go ahead and take a look at an example.0405

We look at the hydrochloric acid and the bond between hydrogen and chlorine.0409

If I look at fonclbrisch, chlorine is going to be more electronegative than hydrogen.0414

That means chlorine is pulling all of the electron density to it.0425

In other words, all of the electron density accumulates around chlorine.0432

It becomes partially negative due to all that electron density.0437

To represent partial, we use the Greek symbol, δ.0442

If chlorine is δ-, that means hydrogen is going to be δ+0447

because all of its electron density is essentially being removed away from it and toward the chlorine.0451

This produces what we call a dipole moment in the molecule.0460

A dipole moment is going to be represented using a dipole arrow.0469

The dipole arrow is always going to point toward the more electronegative atom.0473

Again a dipole arrow points toward electronegative atom; toward the more electronegative atom.0482

Basically what a dipole arrow is, a dipole arrow or dipole moment, it basically shows where the electron density has accumulated.0491

It points toward the center of higher electron density.0503

Shows area / points toward area of accumulated electron density.0510

Let's go ahead and look at the opposite situation.0531

If the difference in electronegativity is negligible, then we have a pretty much equal sharing of electron density.0535

It is more or less mutual.0543

Such a bond is called a nonpolar covalent bond; nonpolar covalent bond.0545

For example, if you just look at typical hydrogen H₂, we have a difference in electronegativity between the two atoms to be zero.0551

They are identical; this is a nonpolar covalent bond.0559

You may be wondering what exactly is considered negligible and what is considered great.0563

I am going to leave that up to your instructor and up to your textbook that you are using because0570

everybody can have a slightly different gauge of the difference in electronegativity and0575

what they consider to be negligible and what they consider to be rather significant.0581

Now that we know the different types of bonds and what we are exactly dealing with,0591

we want to now introduce a schematic way to represent valence electrons.0596

We first are going to talk about atoms.0601

Basically we represent valence electrons schematically as dots.0603

If you recall, you can just go by the periodic table to determine how many valence electrons exist for main group elements.0611

For main group elements, the number of valence electrons equals to the column number that the element is in.0621

Let's go ahead and look at sodium.0642

Sodium is in column one of the periodic table; it has one valence electron.0647

All I am going to do, I am going to represent that as one dot.0651

That is it; that is the electron dot structure for sodium.0654

Let's go ahead and look at magnesium.0659

Magnesium is in column two of the periodic table; it has two valence electrons.0661

I am going to put one of those dots there.0666

I am just going to put another dot maybe on another side of the element if you will.0668

Boron is an element in column number three.0675

It is going to have three valence electrons so I am going to show that to have three dots.0681

Carbon, column number four, it is going to have one dot here, here, here, and then there.0685

On to nitrogen, nitrogen is column five; it has five valence electrons.0694

One dot, two dot, three dot, fourth dot; where do I go put the fifth dot?0701

Really there is not too much of a physical significance here, but I am going to start pairing up.0705

Remember when we talked about atomic orbitals and Hund's rule.0711

Remember we singly fill orbitals before pairing up; I am just going to follow that.0715

Let's go ahead and apply that to oxygen.0721

Oxygen has six valence electrons because it is in column six.0722

One, two, three, four, five, and six.0726

On to fluorine now is in column seven.0732

One, two, three, four, five, six, and seven.0738

Finally we can do maybe krypton.0743

Krypton is going to be a noble gas in column eight.0748

One, two, three, four, five, six, seven, and eight.0751

These are the representative elements from their respective column numbers and how we represent them in terms of Lewis dot structures.0757

That was pretty simple, representing valence electrons of atoms, but now let's go bigger.0774

Let's go ahead and look at electron dot structures for polyatomic species.0780

Basically single bonds are going to be shown as a single line each containing two valence electrons.0785

Something like this; this is what we call a single bond.0792

There is going to be exactly two valence electrons in there.0799

I am going to represent valence electrons as VE for short now.0804

Single bonds are not the only type of covalent bonding that we can have.0809

We can also have double bonds.0812

For example, in oxygen, you can have a double bond like that.0815

In a double bond, there is going to be exactly four of the valence electrons.0822

You can also have triple bonds; for example, nitrogen N₂ contains a triple bond.0828

In the triple bond, there is just going to be six valence electrons.0835

Not all electrons are participating in the bonding.0841

There are what we call nonbonding electrons.0846

Basically for general chemistry purposes, we are going to represent nonbonding electrons as lone pairs.0848

For example, when I go back to the Lewis structure for oxygen each oxygen contains two lone pairs like that.0855

That is what we call nonbonding electrons.0864

Also for nitrogen just to finish it up each nitrogen also contains one lone pair each.0868

How exactly am I coming up with these Lewis structures?0877

There happens to be a very systematic procedure for us to follow.0881

Let's go ahead and take a look at them right now.0885

Step number one, we are going to determine the central atom.0889

There is a couple of guidelines we can use to help us do this.0892

The first point is the following.0896

It is often the element that appears once in the formula.0897

It is never hydrogen; it is often carbon.0901

If there is more than one element that appears by itself, then you choose the least electronegative element.0907

That is number one is to determine the central atom.0915

Number two, we are going to arrange all other atoms around the center and0917

connect it to the central atom using only single bonds first; using only single bonds first.0921

Let's go ahead and look at water; water is H₂O.0928

Here oxygen is the only element appearing by itself.0933

I am going to put it right to the center; step one is done.0938

Now I move on to step two; I surround the center with every other atom.0942

I connect it to the center using a single bond only.0946

Step two is done; let's go ahead and move on.0950

Step three, we are going to complete what is called the octet.0953

Remember we looked at this earlier on that why do atoms gain or lose electrons?0957

It is because they want to achieve eight electrons total, a noble gas configuration.0965

This magic number of eight is what we call an octet.0971

Every atom needs to have an octet in a Lewis structure in order for it to be quote and quote acceptable.0975

There are a couple of exception though--elements that don't need an octet.0982

Hydrogen, hydrogen is completely happy with just two electrons.0987

That is what we call the duet rule for that purpose.0992

Beryllium, beryllium is completely happy with just the four electrons.1004

Boron and aluminum, these elements are fine with just six electrons.1005

Phosphorus and sulfur, phosphorus and sulfur and really elements in period three or more, these actually can have more than an octet.1006

This is what we call an expanded octet.1016

Don't worry, we are going to look at a lot of examples where this occurs.1021

We are going to complete the octet when we read this using lone pairs of electrons.1026

When I go ahead and look at H₂O or water, there is two electrons here.1031

There is two electrons here which means there is four electrons.1039

Now I want to give oxygen eight electrons.1042

That means I need four more electrons to get to eight which is two lone pair.1046

I am just going to do it like that.1050

What we need to do now in step four is to perform a valence electron count.1053

Mother nature has given us these elements.1058

We have to limit ourselves to the number of valence electrons provided by each element.1061

We can only use the total number of valence electrons, not any less, not any more.1068

We have to also account for the charge if it is a polyatomic ion.1074

For each positive charge, we are going to delete one valence electron.1079

For each negative charge, we are going to add one valence electron.1083

Don't worry, we are going to do a lot more examples.1087

But let's go ahead and check out water.1089

For hydrogen in water, there is two hydrogens.1092

Each hydrogen is in column one of the periodic table which means I am going to get two electrons just from hydrogen.1096

There is one oxygen.1103

Oxygen is in column six giving me six electrons just from oxygen.1105

The grand total of valence electrons in water is therefore eight.1111

Let's see if we have used eight in this molecule here.1117

I have two from the single bond here, two here, giving me a grand total of four, six, and eight.1123

Yes, this is a valid Lewis structure for water because everybody's octet rule is good to go.1131

Hydrogen's duet rule is satisfied.1139

I have used exactly the allotted indicated number of valence electrons.1142

If too many electrons are used, we need to then resort to multiple bonding.1153

We are going to try double bonds first, followed by triple bonds.1158

But here is the warning.1162

For each multiple bond added, we need to delete a lone pair from each atom sharing the bond.1165

Let's go ahead and look at an example; that is carbonate.1171

Carbonate is CO₃^2-; this is a nice polyatomic ion we can look at.1176

Step one is to determine the central atom.1182

That is going to be the element I am pairing by itself.1184

In this case, it is only carbon.1187

Step two, we are going to surround the central atom with all other atoms.1190

We are going to connect it to the center using a single bond just like that.1195

Step three, we are going to complete everybody's octet using a lone pair of electrons.1201

Oxygen here needs six more electrons because it already has two nearby.1209

Same thing for this oxygen; same thing for this oxygen.1215

The central carbon has six nearby electrons already.1221

It only need two more electrons or one lone pair.1224

Let's go ahead and do a valence electron count now.1231

Don't forget this is a polyatomic ion.1234

I am going to put that in brackets and show the charge.1236

I have three oxygens each contributing six electrons giving me a grand total of eighteen valence electrons just from oxygen alone.1241

I have one carbon contributing four electron.1253

That is going to be four valence electrons just from the carbon.1256

Remember we must account for the charge; a 2- charge means we add two electrons.1259

Remember for negative charges, we add one electron per negative charge.1267

For cations, we take away one electron per positive charge.1273

My grand total of valence electrons is going to be twenty-four valence electrons that I must use in this compound.1278

When we go ahead and look at carbonate now over here, do we have exactly twenty-four?1287

The answer is no; it turns out that we have twenty-six electrons total.1292

We are over our limit.1298

What we are going to do right now is we are going to resort to multiple bonding.1300

We are going to try double bonds first followed by triple bonds.1303

What I am going to do, let's do this in red.1307

I am going to try one double bond first just like that; what else?1310

We have to delete a lone pair from each atom that shares the bond.1316

Let's go ahead and do that right now.1322

There is one lone pair gone from oxygen.1325

The one lone pair gone from the carbon.1327

Let's go ahead and recount.1332

Yes, this Lewis structure, the way we have it here, has exactly twenty-four valence electrons.1334

Every oxygen has an octet; so does the carbon.1340

This is a valid Lewis structure for carbonate.1344

If you follow these steps every time, you will be fine.1349

But let's go ahead and look at one more example.1355

Let's go ahead and look at nitrogen, N₂.1359

Nitrogen, let's go ahead and connect it with a single bond just like that.1363

Now I am going to complete everybody's octet using lone pair of electrons.1368

Each nitrogen here needs six more electrons to get eight.1373

Remember they share those two electrons in the bond.1377

When we go ahead and do the count, we have two nitrogen each contributing1380

five valence electrons which means ten valence electrons are allotted to us.1384

When I look at this molecule and when I count the number of valence electrons, we are too many.1390

We have used twelve; we are going to resort to multiple bonding right now.1395

Sorry, we used fourteen; we used fourteen electrons; we only can use ten.1402

I am going to try a double bond.1408

When I do that, I have to delete a lone pair from every atom involved.1410

Let's go ahead and recount; when we recount, this is twelve electrons.1415

I am still way over my limit; double bonds have failed us.1420

Now we resort to triple bond.1424

I have to remove another lone pair of electrons from each atom sharing the bond.1428

Here we go; each nitrogen has an octet.1433

I have used exactly ten valence electrons.1437

Again follow this basic five-step procedure; you will get it right every time.1441

You notice that for carbonate we put the carbon-oxygen, the double bond on the twelve o'clock position if you will.1451

I could have put the double bond along any other one with oxygen.1458

That is what we call resonance structures.1466

Resonance structures are basically different Lewis structure of a same molecule.1468

For carbonate, this is what we drew originally.1472

Again I could have easily put it down here.1480

I am just putting my lone pairs in right now.1489

Or I could have easily put it right here.1492

These are what are called resonance structures.1504

These are Lewis structures of the same molecule.1506

In chemistry, we like to indicate the presence of resonance structures with a double ended arrow like that.1509

We like to keep track of electrons around an atom relative to its nuclear charge.1521

This is what we call the formal charge.1526

The formal charge is equal to the following.1530

It is equal to the number of valence electrons minus the number of lines and dots attached.1532

Let's go ahead and look at this structure for the carbonate ion.1538

Let's go ahead and look at this oxygen here.1547

This oxygen right there, let's go ahead and use this equation.1552

Formal charge here is equal to number of valence electrons.1556

Six minus the lines and dots attached, the number of lines and dots attached.1560

I have six dots and one line attached giving me a grand total of seven.1565

That formal charge is -1.1571

We usually like to indicate the formal charge just as a tiny superscript around the atom.1573

I personally like to circle them like that.1579

This oxygen here is identical; it is three lone pairs and one bond.1583

This is also -1.1588

Let's look at the carbon.1593

The carbon is going to be four valence electrons minus...1594

I have how many lines attached total?--four.1601

This carbon has a 0 formal charge.1606

This oxygen here, that is going to be six minus the number of lines and dots attached.1609

I have two lines attached and four dots giving me a grand total of six.1615

That one is also a 0 formal charge.1620

Notice what do the sum of my formal charges equal to?1626

The sum of my formal charges is equal to -2 which is the overall charge of the ion.1630

The nice thing about formal charges is that when I add up all formal charges,1637

it must equal to the overall charge of the molecule or ion.1641

This is yet another way to determine if the Lewis structure that you have drawn is valid or not.1650

Remember it must equal to the overall charge of the molecule or ion, the sum of the formal charges that is.1658

We have already assigned the formal charge to each element in carbonate.1670

Again the sum of the formal charges equals to the overall charge of the molecule or ion.1677

It turns out that not all resonance structures for a molecule or ion are made equal.1684

In terms of carbonate, all three were identical.1691

We always had one carbon-oxygen double bond and two carbon-oxygen single bonds.1694

All three were identical.1699

But now let's go ahead and look at sulfuric acid.1701

Sulfuric acid is an example of a molecule that can have resonance.1705

But we are going to find that the resonance structures are not identical.1709

For example, when you look at sulfuric acid, one way to draw its Lewis structure is right here with all single bonds.1713

I am just going to go ahead and put in the lone pairs to complete everybody's octet just like that.1721

Yet another way to draw sulfuric acid is with two double bonds like this.1729

Remember that sulfur is one of those atoms that can have more than eight electrons.1739

In this case, it has ten electrons.1744

Everybody else has an octet and hydrogen's duet rule is good.1747

Is it the resonance structure on the left or the resonance structure on the right?1752

Which one best portrays sulfuric acid?1756

We have to go through what makes up a good resonance structure.1761

It is all about formal charges.1765

Good resonance structures are going to maximize the number of covalent bonds.1767

They are going to minimize separation of charge.1773

That means formal charges of zero are ideal.1776

They are going to obey the rule, the octet rule as much as possible obviously.1780

Finally if there is a negative formal charge, we are going to try as1786

best as we can to place it on the most electronegative atoms.1789

If we go ahead and look at it, here we have -1 formal charge, -1 formal charge.1798

Sulfur is actually having a +2 formal charge right here.1804

That is equal to six minus four.1809

But when we look at the Lewis structure on the right, this is 0.1813

This is 0, 0, 0 and now sulfur actually has a formal charge of 0.1818

You cannot beat that; you cannot get better than all 0 formal charges.1826

The Lewis structure on the right, the resonance structure on the right, is what we call the major contributor.1831

The resonance structure on the left is what we call the minor contributor.1842

We have seen how we can use formal charge to determine what is a quote and quote good Lewis structure.1849

Now we are going to use Lewis structures to help us determine how a molecule physically looks like.1856

That is what is its shape?--what is its geometry?1863

The theory that helps us to determine this, that is going to give us some guidelines is what we call VSEPR theory.1870

VSEPR theory stands for valence shell electron pair repulsion.1878

Basically the premise behind this theory is the following.1885

That a molecule is going to adopt a configuration, basically a shape or geometry, which minimizes electron-electron repulsion.1888

Remember electrons are the same charges; they don't want to be close together.1898

In other words, the conformation is going to maximize the distance between thereby electrons.1903

There is different types of electron-electron interactions.1910

Lone pair-lone pair interactions are going to be the worst because it turns out that lone pairs occupy the greatest volume.1913

We want to avoid lone pairs being too close together.1927

After lone pair-lone pair interactions is going to be lone pair-bonding pair interactions and finally bonding pair-bonding pair interactions.1933

Right now in the next slide, we are going to introduce molecular geometries, their names, and their respective bond angles.1946

What I mean by bond angle is if you have a central atom A surrounded by two atoms like this,1955

basically the bond angle is going to be the angle of that.1962

Bond angle just like in a typical geometry class.1967

Let's go ahead; now this looks like a lot of material.1973

This looks very sometimes intimidating; but again it is not too bad.1978

Let's go ahead and just take each molecule one by one.1984

Number one, the steric number.1987

The steric number is just a fancy name for the number of electron groups around the central atom.1990

What I mean by an electron group is basically any type of bond.2002

A single bond counts as one group; a double bond counts as one group.2011

A triple bond counts as one group; or a lone pair.2015

Basically one lone pair counts as one group.2024

The basic geometry, the basic geometry is also known as the electron geometry.2031

Basically the basic geometry or electron geometry is everything in this first column.2043

What they have in common is that the central atom E contains zero lone pairs.2051

When we have a situation where there is going to be only two electron groups around a central atom with no lone pairs,2059

that results in a linear geometry or a bond angle of 180 degrees.2068

When we have three bonding groups and no lone pairs around a central atom,2073

we get what is called a trigonal planar geometry and a bond angle of 120.2077

When we come up to four electron groups around a center and no lone pairs, this is what we call tetrahedral.2083

You are going to see that we have what is called a broken wedge and a solid wedge notation.2092

That implies three-dimensional nature of the molecule.2098

Broken wedge means it is pointing away from you into the paper.2101

Solid wedge means it is pointing toward out of the paper.2107

If you look at this molecular here and imagine its three-dimensional nature, we do get a pyramid.2112

A tetrahedron basically is what it is called; let's go on.2118

Five bonding groups around the center atom and no lone pairs, that is what we call trigonal bipyramid.2124

We have two bond angles going on, 120 and 90 degrees.2130

Finally six bonding groups around the central atom, no lone pairs,2135

we get an octahedron or an octahedral geometry with a 90 degree bond angle.2138

What happens as soon as we place lone pairs on a central atom?2145

As soon as we place lone pairs on the central atom, we get everything here.2150

We are going to use a new name to describe their geometries.2157

That is what we call the molecular geometry or the molecular shape.2161

Please make note, there is a difference between electron geometry and molecular geometry.2166

Sometimes they are identical.2171

When they have zero lone pairs like we just covered, electron geometry is the same as molecular geometry.2174

For steric number three, we saw that the electron geometry was trigonal planar.2182

But as soon as we put a lone pair from one of the Xs,2189

that lone pair is a great area of electron density.2195

It is going to push the Xs away from it.2199

We don't have trigonal planar anymore.2203

Instead the geometry name is called bent or angular.2204

We have a bond angle of less than 120.2208

Steric number four, with one lone pair, we get what is called a trigonal pyramid with a bond angle less than 109.2212

If we have two lone pairs, we have a bond angle of much less than 109.2222

Steric number five, one lone pair is what we call a seesaw geometry.2228

If we have two lone pairs, it is called a t-shape geometry.2236

If we have three lone pairs, it is called a linear geometry with a bond angle of 180.2241

Finally steric number six, one lone pair is called square pyramid.2246

Two lone pairs is going to be called square planar.2252

Three lone pairs is t-shape; four lone pairs is going to be linear.2255

You should definitely consult with your instructor and determine which ones he or she exactly would like you to know by heart.2261

Now that we have gone over the names of the geometries and how molecules can look like,2274

we are going to use a molecule shape to determine if it itself is polar.2280

We have already talked about bond polarity, but what about an entire molecule?2286

A step by step procedure is the following.2290

Step one, we are going to draw the Lewis structure.2292

Step two, from the Lewis structure, we are going to determine the molecular shape.2296

Then we are going to assign all bond dipoles.2300

Remember that is the arrow that points toward the more electronegative atom.2302

Then we are going to do the following analysis.2307

If the dipoles cancelled each other, that means there is no net direction of electron density and the molecule is nonpolar.2310

If the dipoles seem to point in a general direction, that is they don't cancel,2319

then we say the molecule is expected to be polar.2325

Let's go ahead and look at a few general examples.2327

For example, if we have EXX and we have one dipole going this way and one dipole going this way,2331

this is a 180 degree bond angle, those dipoles completely cancel each other.2340

This would be an example of a nonpolar scenario.2346

We can do a couple of more examples.2351

EXXX and we can have one dipole, for example, let's go here, one going here, one going here.2353

In this case, all three dipoles, they point in completely opposite directions at a 120 degree bond angle.2366

This would be also an example of a nonpolar molecule.2373

Let's go ahead and look at another one.2378

Now EXX and X; this would be trigonal pyramid geometry--one dipole here, here, and here.2381

In this case, because this is a three-dimensional pyramid, these dipoles don't cancel.2392

They all point away from the central atom.2397

This would be an example of a polar molecule.2405