Enter your Sign on user name and password.

Forgot password?
  • Follow us on:
Loading video...
Lecture Comments (16)

0 answers

Post by Jacob Mohar on December 9 at 09:18:52 PM

Can you explain lattice energy?

1 answer

Last reply by: Professor Hovasapian
Sun Apr 27, 2014 3:18 PM

Post by Rafael Mojica on April 27 at 12:28:53 PM

I need help with molecules that have 2 or more elements!

1 answer

Last reply by: Professor Hovasapian
Tue Dec 3, 2013 12:46 AM

Post by John Wadsworth on December 2, 2013

Those darn crazy lines. :)

0 answers

Post by Michael Amin on November 18, 2013

Hello Mr. Hovasapian,

Shouldn't Iodine be able to bond 7, 5, 3, or 1?

since row 3 starts with the d orbitals, and if we draw the electron configuration then see how many p and s orbitals are full with two electrons then we can see how many of them transfer over to the d orbitals to form more bonds. Finally we can deduce the number of bonds that it can take place.

Never mind, ICl4- lol..... oops didnt see that, so there was only 7 electrons around the central atom but we added 1 more because it was an anion :)

Anyways i was looking forward for more Lewis structures :(

1 answer

Last reply by: Antie Chen
Sun Apr 21, 2013 10:31 AM

Post by Professor Hovasapian on April 20, 2013

Hi Antie,

Yes, starting with row 3 (Principal Quantum Number 3), 5 d orbitals exist for every atom. Elements in row 3 DO use their d orbitals ( for example PCl5). The chemistry of Sodium and Magnesium is governed by the electrons in their s orbitals, but the d orbitals are there. Now, transition metals are different -- these do regularly use their d orbitals, and instead of the octet rule, we often use the 18 electron rule. If you go on to study Inorganic Chemistry in college, you'll spend a fair amount of time on this topic.

The elements in row 2 do NOT have d orbitals, so they can only hold 8 electrons.

The reasons for all of this behavior are Quantum Mechanical, and, again, if you go on to take Physical Chem in college, some of this will be explained fully.

Hope all is well.

Raffi

0 answers

Post by Antie Chen on April 19, 2013

Hey Raffi, I enjoy your great teaching, and have a little confusion about the Row 3.
I look up your answer to another student, you said that "Row 3 and all the other rows below it - these elements have d orbitals available for filling"
How about Na or other elements is in row 3 but don't fill any electron in d orbital?
and the F or other elements have an atomic number before 10 cannot contain more than 8 electron?

1 answer

Last reply by: Professor Hovasapian
Fri Mar 1, 2013 2:33 PM

Post by Erica Rapetti on March 1, 2013

I thought you said that group 3 can have more than 8 electrons? How did Iodine get 12 electrons?

1 answer

Last reply by: Professor Hovasapian
Mon Jan 7, 2013 5:09 PM

Post by Andreea Cirstea on January 6, 2013

Why is there no bracket around the ICl4 Lewis structure with a minus sign?

2 answers

Last reply by: Suresh Sundarraj
Mon Dec 17, 2012 7:16 PM

Post by Riley Argue on October 30, 2012

Great lecture. Thank you.

Bonding & Lewis Structure

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

Transcription: Bonding & Lewis Structure

Hello, and welcome back to Educator.com; welcome back to AP Chemistry.0000

Today, we are going to start on bonding; so I'm just going to talk about some general aspects of bonding.0004

Before I actually launch into bonding and Lewis structures today, and a continuation of Lewis structures and molecular geometry for the next lesson, I wanted to talk a little bit about this particular general section of it.0011

A couple of lessons ago, I had mentioned that there are some certain aspects of chemistry (like, for example, periodic trends, like the trend in ionization energy or the trend in atomic size)--these are things that I have not addressed directly in these lessons.0024

There is a reason why I did that: there is a tendency to actually overstate the case.0041

Most of these are very, very straightforward, and it would certainly behoove you to go ahead and look through these things in your respective textbooks.0046

It shouldn't take very long at all.0057

What I wanted to do is sort of give you...I wanted to hopefully get across what is going on chemically.0060

If that is understandable, then it's very, very easy for you to go to your book and just sort of jump to a particular section that I have not covered formally, and understand exactly what is going on.0068

Everything that I have discussed should allow you to go back to the particular chapters (in bonding or whatever) and look through some of the topics that are in there, that I have not covered formally; and they should be readily understandable.0078

In general, the things that I have skipped--that I have not discussed formally--they are very, very straightforward.0093

I am going to leave them to you; there are some things that I am going to mention, that I will not spend an inordinate amount of time talking about, because there are more important concepts (underlying concepts that are more important) that I want you to wrap your mind around.0099

But again, be aware that there are certain things that I have not discussed formally, particularly in the chapters that involve bonding--mostly because I want to get to Lewis structures, which will be with you for a while throughout your career; molecular geometry, which will be with you; and things like molecular orbitals and hybrid orbitals, and things like that that you will continue to discuss.0113

So anyway, having said that, let's go ahead and just jump right on in.0139

OK, so we are going to discuss something called a covalent bond.0145

Oh, yes--one of the things, of course, that does show up in the bonding, that I did not discuss, is ionic bonding, lattice energies--things like that.0149

They are very, very, very straightforward; our concern, as far as chemistry is concerned, is primarily for the covalent bond, so that is what we are going to be spending most of our time with.0158

Covalent bond: what is it?0170

Well, it's very, very simple: it's when two atoms share a pair of electrons; that is it.0174

That is all; so let's say, if you had a particular atom, like (oh, what shall I pick)...let's just pick ammonia--something like this; now, we'll get to the symbolism in just a minute...something like that...these single bonds represent a shared electron pair.0192

This is equivalent to...there is a pair of electrons between that; there is a pair of electrons there; there is a pair of electrons there; and there is pair that is not shared, that belongs to nitrogen.0212

What this means is that these electrons, in between--well, nitrogen can claim them whenever it needs to claim them, and hydrogen can claim them whenever it needs to claim them.0222

That is it; that is what sharing is--it is pooling your resources so that both of you can participate in the overall resource.0231

That is it; that is the definition of sharing, and that is exactly what they do, and that is all a covalent bond is.0239

A covalent bond is one pair of shared electrons in between two atoms; that is all.0244

Now, sometimes we can share more than one pair.0251

When we share one pair (let me actually write it down here), one pair shared--we call that a single bond.0255

It makes sense, right?--a single bond.0268

We represent it with a single line; that is why you see this line up here.0270

That is a symbol for a pair of electrons.0275

Well, two pairs shared--exactly what you think: it's called a double bond, and how many lines do you think that is?--a double line!0279

An example of a double bond would be the bond between carbon and oxygen in carbon dioxide; it's a double bond.0292

They share two pairs of electrons in between them.0298

They can share three pairs; we call that a triple bond.0301

That is it; nothing more--very, very straightforward.0307

And we represent it with a triple line.0310

A good example is the nitrogen molecule: as you know, nitrogen is diatomic; it is not just nitrogen atom--it is nitrogen molecule, N2; it looks something like this.0313

Each nitrogen is actually...there are three pairs of electrons in between the nitrogen atoms themselves, and then a couple of electrons on each that are exclusive, just to each nitrogen.0324

That is it--that is a covalent bond; it is just a shared pair of electron: a single bond, a double bond, a triple bond: nothing more, nothing less.0334

Now, granted, many physical chemists would actually argue with that, but for our purposes, it's really this simple.0345

OK, now I do want to draw one little picture to show you what exactly happens, energetically, so you understand what is happening.0351

So, if I plot the energy on the vertical axis, then I'm going to make this my 0 point energy; and on this axis, this is something called the internuclear distance (a fancy word for the distance between the two atoms; that is it--internuclear distance).0359

We say "internuclear" because we are actually measuring the distance between the centers of the atoms--the actual nuclei--not the atoms themselves, which are sort of clouds of electrons, you remember.0376

Here is what happens: if this is the 0 energy...well, if the atoms are infinitely far apart, the energy is 0; they are infinitely far apart.0386

As I bring them closer together, closer together, closer together, closer together, well, here is what happens.0396

I bring them closer and closer and closer together; they actually start to drop in energy; they reach a minimum; and then, as I push them even closer together, the energy actually shoots up.0403

What that means is the following: 0 energy--as I bring them closer and closer and closer, the attraction of the positive nucleus of one for the electrons of the other--the desire to share--increases.0415

When that increases, the energy actually drops; remember, all systems seek a point of lowest energy; all systems, in all science, across science--this is a deep, deep universal phenomenon--all systems want to minimize their energy.0428

So, as it turns out, when you bring two atoms together, the energy actually drops--it becomes more stable.0443

That is what lower energy means in science: when we say the energy drops, we mean the situation is becoming more stable.0449

At a certain point, it reaches a distance (OK, this is 0): there is a distance between them which is a minimum.0455

If you push them any closer, now the repulsive forces of the respective nuclei--the positive-positive charges--now they become stronger than the attractive forces, and now the energy rises again and shoots up; it destabilizes the molecule.0466

This distance--this is what we define as the bond length; whenever you have a nitrogen atom and a nitrogen atom, when you bring them together, bring them together, from an infinite distance apart, they actually drop in energy.0481

They reach a point where they are perfectly happy sharing those three pairs of electrons.0494

If you push them any closer together than that, they actually will repel each other.0499

You will destabilize the molecule.0503

They will not want to be together: there is an energy minimum: that energy minimum--the distance at which that energy minimum takes place--that is the bond length.0506

This is the energetics involved in a covalent bond.0515

That is all that is happening--that is all that is happening.0518

I hope that makes sense.0521

OK, so let's define a few other terms here: we are going to define something called electronegativity.0523

Electronegativity--it seems a little redundant: an electron is negative, so why are we calling it that?--well, that is the name that was chosen for it.0530

OK, it is the measure of an atom's ability to pull electrons in a covalent bond toward itself.0540

That is it: so it is true that they are sharing the electrons, but let's put it this way: if you have two atoms that are of a different type, like carbon and hydrogen, carbon and chlorine, silicon and fluorine, oxygen and nitrogen...they are not going to share equally.0572

They do share, but they don't share equally.0588

Different atoms have different electronegativities; there are numbers that are assigned to these, and you can see those numbers by just looking in your book.0591

There is a sort of a picture of a periodic table, and it gives you specific values, from 0 all the way to 4.0599

Fluorine has the highest electronegativity; that actually is a measure for how badly these atoms want the electrons.0605

They do share, but they actually are not sharing equally; there is a little bit of a tug-of-war going on.0612

The atom with the higher electronegativity pulls the electrons a little bit closer to itself; that is all that means.0618

Let's do a little pictorial here: if I have a hydrogen molecule, there is a pair of electrons shared between them; it is right down the middle--I mean, it is literally right down the middle.0624

It isn't closer to this hydrogen; it isn't closer to this hydrogen; the reason being, these are the same--they have the same electronegativity.0634

They have the same desire to pull these electrons toward them, so they cancel out.0641

However, if I take something like...well, let's just take nitrogen and oxygen: so if I have a nitrogen-oxygen bond (no worries about anything else), oxygen has a higher electronegativity than nitrogen does, which means that these electrons that it is sharing in the single bond are not exactly in the middle.0646

As it turns out, these electrons spend a greater portion of their time closer to oxygen.0666

Oxygen is electronegative; it will share, but it wants the electrons more, so it actually pulls the electrons closer to it.0674

And literally, the electrons spend more time closer to oxygen than they do nitrogen.0683

They aren't entirely with the oxygen or nitrogen--this is not an ionic bond (if that were the case, it would be ionic, where the electrons are literally just stolen from the other element)--but here, it's a shared covalent bond, but electronegativity measures the extent to which it belongs to one atom or the other.0688

That is all the electronegativity is.0706

OK, each atom has a given electronegativity; now, the Δ electronegativity, or I should say, the difference in electronegativity, is a measure of the polarity of the bond.0709

Polarity is a very important concept--polarity of the bond.0733

I'm not going to talk about it much; I'm just going to define it, because the more we talk about it, it actually might confuse you more; I'm just going to say a couple of words about it--just know that it exists.0737

There is nothing mysterious about it: polarity of a bond (or of "the" bond).0745

Now, here is what polarity means: polarity means a charge separation--a charge separation is when you have one end...it is when you take a positive and negative charge, and you separate them.0752

Positive on one side, negative on the other: that is what polarity means.0773

Polarity is like a magnet: when you take a magnet, you note that a magnet has a positive end and a negative end; it is polarized--that is what polarizing means.0776

It means pulling to opposite ends, literally: the two poles--we have a north pole; we have a south pole; there is a magnetic quality to it.0784

It is polarized: there is a positive and a negative end.0794

We can call it anything we want (north, south, positive, negative, hot, cold...); it's just that there is polarity; OK.0798

It is a charge separation: in other words, there is a plus end and a minus end.0805

In a bond like NO, well, guess what; because the electrons spend a greater part of their time towards the atom that is more electronegative, the oxygen, well, there are more electrons toward the oxygen; that means it has more...it is carrying a little bit extra negative charge.0817

And, because there is a little bit of a deficiency here--that means they are further away from the nitrogen, on average--it carries a partial positive charge.0834

This δ symbol--it means "partial"; so you don't get a full +1 charge or a full minus charge; you get a fraction of a charge.0842

This is a polarity; there is a charge separation in this bond, and we express this with a symbol that looks like this.0852

An arrow with a little perpendicular cross at the other end: this means that this is the positive end, and the electrons are being pulled that way.0861

It is polarized; it is a polarized bond; the electrons are pulled in that direction.0871

If it were the other way, it would go the other way; this is the universal symbol to show that a bond is polar.0876

OK, so now, bond polarity: again, it is all we are doing--a bond polarity.0886

If there is...if you have two atoms of a different nature that are bonded to each other, if they are different, there is going to be some polarity difference between them.0896

One of them is going to be more electronegative than the other.0905

The difference in electronegativity might be small: it might be .1, .2, .3...but there is still some polarity.0908

As we start to get into the .5, .6, .7 range, now the bond is starting to become very polar.0915

Now, the charge distribution--there is definitely a positive end, and there is definitely a negative end, to the bonds.0920

It is a polar bond.0926

We have assigned some numbers to this; I wouldn't follow these numbers too strictly; this is sort of a range; it's just--we want to give you an idea of the difference in electronegativities.0928

Bond polarity of 0, let's say 0.5, let's say 1.0; and these are electronegativity differences that I am describing here...and 2.0.0941

Anything in the 0 to maybe...I don't know...0.7 range--we just call that a straight covalent bond.0955

Anything in the .7 to about the 2.0 range (and again, I say .7; most people will say .9 or 1; I say about .7, roughly to about maybe 1.7 or 2)--they call that a polar covalent bond.0965

A polar covalent bond is exactly what you think: it is when, yes, you have a covalent bond between two species; yes, they are sharing an electron (which is what "covalent" means); but the electronegativity difference in them is so large that the electrons spend a greater amount of time on one side, therefore creating a charge separation--a negative end to the bond and a positive end to the bond.0984

It is a polarized bond; they call that a polar covalent bond.1011

Anything above an electronegativity difference of 2...we usually call that an ionic bond.1014

And again, these are sort of rough estimates: it doesn't mean that, if something is 1.9, it's a polar covalent bond; if it's 2.1, it's an ionic bond.1020

No, bonds are...there is a spectrum, there is a continuum, of behavior, and it all depends on a lot of things.1030

These are just rough numbers.1037

Just to give you an idea, let me do 0 again, and 0.5, 1.0, 1.5, 2.0; so, a hydrogen-hydrogen bond, a chlorine-chlorine bond, an oxygen-oxygen bond...it's going to be an electronegativity difference of 0, because they are both the same.1040

NO...yes, you are talking about a roughly .5 or .6.1060

An electronegativity difference of 1...you are talking about maybe a sulfur-oxygen bond; that is about 1.1068

1.5...you are getting into the carbon-fluorine...that is about an electronegativity difference of 1.5.1074

Up here and beyond--here is where you are getting into your ionic compounds: sodium chloride, magnesium chloride, things like that--purely ionic bonds.1082

They are so...at this point, we don't even call them covalent; I mean, the sodium has given up its electron--there is no sharing going on.1091

That is it: it is just...electronegativity is a measure of the extent to which electrons want to be pulled by a particular atom, the charge separation that exists (the negative end; the positive end).1098

Partial negative end/partial positive end--we call that a polar bond: it has a polarity; that is it.1110

A covalent bond, polar covalent bond, ionic...these are just some terms that you are going to hear; and that is all it is--it's nothing particularly strange.1117

Now, when a bond is polar, we say it has a dipole moment (the bond).1126

We say it has a dipole moment--that is it; I just wanted to throw that out there, because you are actually going to hear the word "dipole moment."1147

Dipole moment just means that it is polarized; that is it--it's just a fancy way of saying that there is a permanent...1153

It's usually used in reference to molecules, because when we actually do what we are going to do a little bit later on--when we do Lewis structures, we are going to analyze each bond in the molecule to see if it's polar.1162

You can actually have a bunch of polar bonds, but the polarity is such...arranged in space such...that all of the directions cancel out, so you end up with a molecule that actually is not polar, which is kind of unusual.1174

The bonds can be polar, but the molecule may not be polar; and we will talk a little bit about that; but you will often hear the word "dipole moment" being used to describe mostly molecules, but you can also use it to describe a bond--any polar bond that has a dipole moment (in other words, it is polarized); that is it.1186

OK, let's see: OK, so now, let's go ahead and talk about something called bond energy.1205

Now again, this is all just sort of a bunch of background stuff; Lewis structures is what we are making our way towards.1214

Bond energy is...again, a little extra something: I actually wasn't going to talk about it, but I decided that it is something that should be mentioned, because it has something to do with enthalpy.1220

Enthalpy, remember, is the heat of a reaction.1234

Bond energy...and it's exactly what you think it is: it is the energy holding a bond together.1237

Or, from our perspective, it is the amount of energy that we need to break the bond; that is it.1246

OK, so I'll do it from our perspective: the energy required to break a bond.1251

That is it--a certain energy associated with it: different bonds have different energies--some are strong; some are weak; OK.1264

Now, we usually (well, let me write this down) average (no, let's see)...take average values for specific bonds, because environment...1271

When we want to take the...there are tables of bond values, of bond energies for specific bonds (carbon-carbon, carbon-nitrogen, nitrogen-chlorine...things like that); all of these have been measured.1308

But, the thing is...the values that are listed in these tables...they are average values; and the reason they are average values is because context matters.1318

Not all bonds...not all carbon-carbon or carbon-hydrogen bonds...have exactly the same energy.1329

It depends on what else is sort of bonded to that carbon--what else is going on with a molecule.1335

I'll give you an example of this; but in general, when we sort of average them out over a bunch of different bonds, we get a certain average value.1342

That is what we sort of use, most of the time.1350

Let me see: we usually take average values for specific bonds, because environment affects the specific energy of a specific bond.1353

For example, if I take HCBr3, this is 380 kilojoules per mole; if I take the CH bond in HCBr3, or CHBr3 (this is not an acid), it is 380 kilojoules per mole.1375

If I take the CH bond in CF3H, well, this one is 430 kilojoules per mole; it's still a carbon-hydrogen bond, but it has a totally different energy.1397

Well, what if I do the CH bond...if I do HCCl3, I end up with 410 kilojoules per mole; so you see different values.1410

When we end up taking an average value for this, we end up with something (oh, these crazy lines are showing up again; there we go--all right)...we end up with an average of 413 kilojoules per mole.1434

So, you see: different values, same bond (carbon-hydrogen, carbon-hydrogen, carbon-hydrogen); we need to take an average.1447

We use 413 kilojoules per mole (and I apologize for those lines).1455

OK, now, what is nice about this is that bond energies can be used to estimate enthalpies; that is what we are going to do.1461

Let's see: Example 1: OK, estimate the ΔH for the following reaction (so bond energies can be used to estimate ΔH values--they're actually very, very good).1474

We have: CH4 (which is methane) + 2 Br2 + 2 F2 goes to CF2Br2 + HBr + HF.1501

And, if I am not mistaken, this should be a 2, and this should be a 2.1518

We have this equation here, and we want to calculate the ΔH for this.1521

Well, how do we do it?--well, we can do it just by sort of...you can get a pretty good value by just doing bond energy differences.1525

Well, here is what we are doing when we are doing a reaction: we are breaking bonds of all of the reactants, creating free atoms, and then we are putting them back together in a different arrangement.1533

So, as it turns out, the ΔH is going to be this: the ΔH is going to be the energy of the bonds broken (so think about this for a second)...it's the energy of the bonds broken, minus the energy of the bonds formed.1545

Let me do this pictorially: the reactants are a certain energy: in order to break the bonds of all of the reactants, I need to put energy into it; therefore, my energy is going to rise.1575

It is going to go up here to create a bunch of free atoms.1588

When those atoms recombine, well, they are going to recombine, and they are going to release a certain amount of energy until they reach the energy of the products.1592

This difference--that is the enthalpy; that is ΔH.1605

That is what is happening: I am putting energy into the reactants to break the bonds; when the bonds form, they release energy back.1611

They reach a certain level where the bonds create energy, but if there is any excess, it is given off as heat.1620

Or, if they don't have any excess--if they actually end up using more--then it's a positive ΔH (negative ΔH: endothermic, exothermic).1625

That is all that is going on here; so I just wanted you to see this pictorial version, but analytically and numerically, it's the energy of the bonds broken, minus the energy of the bonds formed.1635

So, let's go ahead and do that while I erase these--it's just going to take up space, so I don't want to end up leaving this picture here.1646

OK, so now, let's take a look at the bonds that we are breaking.1659

We are breaking four CH bonds (right? CH4, so 4 CH bonds); that is going to equal 4 times 413 kilojoules per mole (and again, these values for bond energies--they are in your book, or they are in the back of the book, or they are in a CRC table; they are available in a list...or they are available on the Web).1664

OK, so that equals 652 kilojoules; we are breaking 2 Br-Br bonds, so that is going to be 2 times 193 kilojoules per mole, equals 386 kilojoules.1690

Oh, in case you are wondering: "This is kilojoules; this is kilojoules per mole; what happened to the mole?"--well, here.1712

This is actually 4 moles: one molecule of CH4...in other words, 1 mole of methane molecules contains 4 moles of CH bonds.1717

4 moles of bonds, times 413 kilojoules per mole in the bond--that is where the mole disappeared to--it's just simple.1729

2: we are breaking 2 F-F bonds, and that is going to be 2 times 154 kilojoules per mole, for a total of 308 kilojoules.1738

When we add all of these up, we get 2,346 kilojoules; that is how much energy we have to put in to completely separate methane, bromine, and fluorine into free atoms.1752

Now, we are going to form some bonds.1763

Well, what kind of bonds do we form?--we form 2 CF bonds (right?--we form 2 CF bonds), so it's 2 times 485 (I'm just going to write them down); we form 2 CBr bonds; that is going to be 2 times 276; we form 2 HBr bonds--that is 2 times 363; and we form 2 HF bonds, which is 2 times (wow!) 565 (that is a bit of a surprise; OK).1766

And, when I add all of this up, I end up with 3,378 kilojoules.1806

When I take 2,346 minus 3,378, I get a ΔH of -1,032 kilojoules.1813

Negative ΔH: this is an exothermic reaction.1825

When methane, bromine, and fluorine come together to form CF2Br2 and 2 molecules of hydrobromic acid, 2 molecules of hydrofluoric acid, energy is released: 1,032 kilojoules per every mole of methane that reacts.1828

It's highly exothermic.1847

That is it; OK.1849

Now, we are going to get to the heart and soul of bonding: Lewis structures.1854

OK, I'm just going to launch into Lewis structures; I'm not going to say a lot--I think the more I say, the more confusing it is going to get.1861

This is one of those things; so we are just going to jump in and do some Lewis structures.1868

I'll just tell you what it is: it is a notation for representing bonding; that is it.1872

The bonding that takes place in a molecule--this is a notation that represents--that tells you what is going on.1886

Now, it's very, very important to realize that it says nothing--Lewis structures say nothing about the arrangement of atoms in space.1892

They don't tell you anything about geometry; we will get to that next lesson.1901

It says nothing (you know what, I really need to write this more clearly; this is not going to work; all right, let's try this again: Lewis structures)...1905

A Lewis structure is a notation for representing bonding in a molecule.1919

It says nothing about molecular geometry...you know what, I don't like that; I don't want to use the word "molecular geometry"; it says nothing about how the atoms are arranged in space.1950

We don't want to be fancy; we want to understand what is happening.1969

OK, it uses only valence electrons--only, only, only valence electrons.1977

You remember what valence electrons were: they are the number of electrons in the outermost shell, the total number of electrons in the highest primary orbital.1986

1s2, 2s2, 2p5: 2s2, 2p5--that is 7 valence electrons.1997

It is also the numbers on top of the periodic table; so, OK.2003

Here is how we do a Lewis structure; I am going to write the rules, and then we'll just do some Lewis structures.2008

Add up all the valence electrons in a molecule.2016

2: Arrange the atoms (in other words, just sort of put them on a piece of paper next to each other--arrange the atoms).2030

3 (and again, when we run through this, you will see what it is--when we actually do the actual examples): Use a pair of electrons for each bond between 2 atoms.2041

4: Add the remaining electrons in pairs (that is what is important--2 at a time), until the shared total plus the lone pairs, on each atom, is 8 electrons: this is called the octet rule.2073

OK, 5: Hydrogen only needs 2 electrons (it doesn't need the 8; that is the exception).2121

And the last rule: Row 3 elements (phosphorus, sulfur, silicon) can, can, accommodate more than 8 electrons, if necessary.2134

So, if you have any electrons left over after all the other atoms are done (in other words, have 8), put the extras on the central atom (and it is usually the central atom that will be a Row 3 element).2168

OK, let's just do some examples; it's the only way this is going to make sense.2207

OK, examples; let's start with H2, hydrogen gas molecule.2212

OK, if you look in a periodic table, hydrogen has 1 electron in its valence shell; there are 2 hydrogens; therefore, the total number of electrons is 2.2219

Well, we arrange the atoms; we write them down on a piece of paper; we take an electron pair (2), and we put them in between to form a single bond.2230

Now, this hydrogen has 2; this hydrogen has 2; hydrogen only needs 2; all 2 electrons are used up; we are done--that is the Lewis structure for hydrogen (also represented with a single line).2238

Let's do CH4, methane.2255

Carbon: it brings 4 valence electrons; hydrogen brings 1 valence electron; there are 4 hydrogens, for a total of 4.2257

We have a total of 8 electrons to distribute.2265

We draw C; we draw H; there are 4 of them, so they are all going to be bonded to the C (that's usually how it is).2270

We put a pair for each--we always start like that, by making a single bond--a single bond to the central atom.2279

And then, we fill in the rest: 2, 4, 6, 8; we have used up all 8 electrons; that is 2; H has 2; H has 2; H has 2; carbon has 8.2288

Everything is taken care of; that is the Lewis structure; it is represented this way.2303

Again, this says nothing about how these are arranged in space; this molecule is not flat.2310

It isn't just 1 carbon with 4 H's around it, arranged in a plane like this; it is not; it is actually tetrahedral.2315

We'll get to geometry later.2322

Let's try NO+: this is a perfectly good species--it just happens to be an ion.2325

Well, when you see a + charge on an ion, that means you have lost an electron.2333

If you see a negative charge on an ion, that means you have gained an electron, so you have to add that to the count.2337

Nitrogen has 5 valence electrons; oxygen brings 6 valence electrons.2343

This + charge means you lost an electron, so you have a total of 10 electrons to distribute in the molecule when you bond.2349

Let's go...we write the N; we write the O; we put 2 right there.2359

Now, going down the list of rules and how to deal with Lewis structures, once we have come this far, now we fill in the other electrons around each atom until each atom has 8.2366

Let's go: that is 2; that is 4; 6; 8; 10; well, now I have used up all of my 10 electrons; I don't have any more.2379

Well, the oxygen has 8, but the nitrogen only has 4.2389

This is not going to work; I need another arrangement.2393

I do N; I do O; I start again.2398

This time, I'm going to go ahead and put 2 pairs in between each.2402

That is 2, 4, 6, 8, 10; well, this has 8--good; oxygen is good; now, this has 6--no, still not enough--I need 8.2406

N...let's try 3 pairs in between...O, 2, 4, 6, 8, 10; all of the electrons are used up.2422

8 around nitrogen; 8 around oxygen; remember, we are sharing, so sharing means you can participate in both.2434

There you go; that is our Lewis electron structure, and you can write it like this.2441

Triple bond: you need to put the lone pairs on--see, these pairs right here--they are with this atom; they are not shared; they need to be there.2446

This whole thing is an ion, so you put a bracket around it, and you put the charge of that ion; that is the Lewis structure.2456

You can leave it in this form, as long as you do this; that is fine; the dots don't matter.2462

We prefer the lines, but the dots are not incorrect.2468

PCl5 (phosphorus pentachloride, right?--remember mono-, di-, tri-): phosphorus brings 5 electrons; chlorine brings 7 electrons (7 valence electrons); there are 5 chlorines here, so 7 times 5 is 35.2474

It's 5 times 7...35 plus 5; we have a total of 40 electrons to distribute in this particular structure--free electrons that we can use.2494

I take P, Cl, Cl, Cl, Cl, Cl; I make my single bonds first: 2, 4, 6, 8, 10; now I fill in the rest: 12, 14, 16, 18, 20, 22, 24, 26, 28, 30, 32, 34, 36, 38, 40.2504

Chlorine has 8; chlorine has 8; chlorine has 8; chlorine has 8; chlorine has 8; phosphorus has 10--it's more than 8.2533

It's OK--phosphorus is a third-row element; it can accommodate more than 8 electrons.2542

This is the Lewis structure for phosphorus pentachloride.2547

When we do the line structure: Cl, Cl, Cl, Cl, Cl--I'm going to leave it up to you and/or your teacher to decide whether you want the lone pairs.2550

I just said that it would be nice to put the lone pairs on; it's a little tedious to sit here and do 2, 4, 6, 6 times 5...30 lone pairs of electrons; you can leave them off if you want, if your teacher says it's OK.2561

But, in general, Lewis structure shows lone pairs, and it shows shared pairs.2574

Lone pairs; shared pairs...2580

Let's try ICl4-: well, iodine brings 7 electrons; chlorine--there are four of them; each one brings 7 electrons; there is a negative charge on here, so that is an additional electron; 28, 35, so we have a total of 36 electrons that we can distribute among this.2585

ICl4: I is going to be central; let's put Cl here, Cl here, Cl here, Cl...if you're wondering why I didn't put them this way, this way--just for a little change of pace.2608

Arrange the atoms; now, use a pair of electrons to form a single bond between each atom and the central.2623

There is 2; there is 4; there is 6; there is 8; now, let me fill in the rest.2630

10, 12, 14, 16, 18, 20, 22, 24, 26, 28, 30, 32; I have used 32 electrons; I have 4 left over.2635

I'm going to put them here, and I'm going to put them here, on the central atom.2652

Iodine--each one of them has 8, 8, 8, 8--iodine has 2, 4, 6, 8, 10, 12; iodine can accommodate more than 8, because it's not only a row 3 element--I think it's a row 4 or 5 element.2656

So, not a problem; this Lewis structure looks like this.2669

Cl, Cl, Cl, Cl, dot, dot; if you want to avoid the lone pairs on the other ones, that is fine; do not leave off the lone pairs on the central atom.2675

Do not leave off the lone pairs on the central atom.2686

Let me say it again: Do not leave off the lone pairs on the central atom.2690

They are very, very important; if you want to leave these off, that is one thing, but not the central atom.2693

Those lone pairs have to be there; they actually have a lot to do with the chemistry of the compound--and you will understand that when you take organic chemistry.2698

OK, now, let's do BeCl2.2707

Beryllium brings 2 electrons, and chlorine brings 7; there are 2 of them, for a total of 14, 15, 16; we have 16 electrons to distribute.2714

We write Be; we write Cl, Cl around it; we use 2 electrons there, 2 electrons there; that is 4; 6, 8, 10, 12, 14, 16.2725

Well, chlorine has 8; chlorine has 8; but beryllium only has 4.2738

You might think to yourself, "Well, wait a minute; why don't you just do this one?"2743

1, 2, 3, 4, 5, 6, 7, 8 chlorine; this will work out just fine: chlorine has 8; beryllium has 8; chlorine has 8.2749

Yes, that is perfect; as it turns out, this is not the Lewis structure, and here is why.2760

This is the Lewis structure, right here: beryllium, boron, aluminum...these things can accommodate fewer than 8 electrons.2766

Remember, we said that there are some that actually accommodate fewer than 8 electrons: beryllium is one of those.2775

There aren't too many: beryllium, boron, aluminum...I think that is about it; they can accommodate...2782

Well, you are thinking to yourself, "Well, why not this?"2789

Valid question, "Why not?"--great question.2791

Here is why not: when you take a look at electronegativities for chlorine versus beryllium, chlorine is very highly electronegative; beryllium is not very electronegative.2796

The idea that chlorine would actually allow two of its lone pairs to be shared by beryllium--it actually...it won't happen.2807

As it turns out, when there is a huge electronegativity difference, sharing will not take place, especially because the electronegative atom--in order for it to share electrons, it has to release them a little bit.2818

Chlorine is too electronegative; it will not release its lone pair of electrons.2831

It will share the one electron that it does have, but that is it; it will not...2835

So, as it turns out, this Lewis structure for beryllium chloride, the correct one, is this.2839

OK, here we go with the crazy lines again, so let's see if we can write this slowly: dot, dot, dot, dot; I think that is the real key--I think I just have to move more slowly here; there we go.2851

That is the correct Lewis structure for beryllium; beryllium is electron-deficient.2866

It is a highly reactive species precisely because it is electron-deficient.2871

Beryllium, boron, and aluminum can have less than that; these highly electronegative atoms--the ones at the top right of the periodic table (the fluorine, chlorine, oxygen, sulfur, things like that)--they generally will not share if they don't have to.2874

They won't release their lone pairs readily.2892

Certainly, the halogens will not do so.2895

That is why this is the correct structure, and not this.2897

It seems to contradict what we said, but it is in line with the electronegativity difference.2900

OK, thank you for joining us here at Educator.com to discuss Lewis structures.2906

Next time, we will continue with Lewis structures; we will talk about the polarity of molecules, and we will talk about molecular geometry; and we will begin our discussion of hybrid orbitals.2910