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Lecture Comments (13)

1 answer

Last reply by: Professor Hovasapian
Wed Nov 11, 2015 4:12 AM

Post by Jason Smith on November 6, 2015

Hi professor. Any idea why the 3d orbital fills up on the 4th energy level? Like does science have any idea why this happens? Seems really bizarre!!!

4 answers

Last reply by: Professor Hovasapian
Wed Dec 11, 2013 3:20 AM

Post by Tim Zhang on December 8, 2013

In the last topic, the electron configuration of Mn end with 3d, but it is in the 4th row, so the primary number n equal 4, which means instead of 3d, I should write a 4d?

1 answer

Last reply by: Professor Hovasapian
Wed Sep 18, 2013 12:42 AM

Post by Kingsley Lunga on September 18, 2013

You make life really easy Mr Raffi

3 answers

Last reply by: Xinyuan Xing
Thu Apr 23, 2015 8:53 AM

Post by Abdirisak Ashkir on March 30, 2012

It would be better if you would explain it in a more simple way

Electron Configurations & Diagrams

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.

  • Intro 0:00
  • Electron Configurations & Diagrams 1:08
    • Electronic Structure of Ground State Atom
    • Order of Electron Filling
    • Electron Configurations & Diagrams: H
    • Electron Configurations & Diagrams: He
    • Electron Configurations & Diagrams: Li
    • Electron Configurations & Diagrams: Be
    • Electron Configurations & Diagrams: B
    • Electron Configurations & Diagrams: C
    • Electron Configurations & Diagrams: N
    • Electron Configurations & Diagrams: O
    • Electron Configurations & Diagrams: F
    • Electron Configurations & Diagrams: Ne
    • Electron Configurations & Diagrams: S
    • Electron Configurations & Diagrams: Fe
    • Introduction to Valence Electrons
    • Valence Electrons of Oxygen
    • Valence Electrons of Iron
    • Valence Electrons of Arsenic
    • Valence Electrons: Exceptions
    • The Periodic Table

Transcription: Electron Configurations & Diagrams

Hello, and welcome back to, and welcome back to AP Chemistry.0000

Today, we are going to continue our discussion of electrons and electron configurations, and we are actually going to specifically talk about electron configurations and electron diagrams.0004

Last lesson, we spent a lot of time--it was pretty much just talking; we really didn't do any examples--talking about the structure of the atom, in terms of the energy levels that the electrons occupy, and the four different quantum numbers.0015

So now, we are going to get into sort of the practical aspect, from a chemical point of view--what a chemist talks about when we talk about an electron configuration.0031

So, if we say, "What is the electron configuration for silicon?", we are asking you to specify where these electrons are--the electrons that happen to be flying around silicon.0041

And then, we are going to do the diagrams, which are a little bit more specific.0052

They actually tell us--not only where they are--specifically in which sub-sublevel they are, and what the spins are of those electrons; so it's a very nice pictorial representation.0055

Let's just get started.0066

OK, so we have this notation; so we have a notation which can describe the electronic structure, the atomic structure (when we say "electronic," we are talking about the electrons), of any ground state atom.0070

"Ground state atom" just means the atom as it is found in nature.0108

Now, we can use this notation to actually represent excited states, also, and ions, as we will see in subsequent lessons.0114

But right now, we are just concerned with ground states.0120

Here we go: here is we said that the primary level has an s sub-orbital, and that s can hold 2 electrons.0124

This thing--it can hold 2 electrons.0136

The second primary level has an s sublevel, and it has 3 p sublevels...I'm sorry: it has an s sublevel, and it has a p sublevel, and those p's have 3 sub-sublevels.0145

So, it can hold a total of 2, 4, 6, 8: 8 electrons.0163

The third primary has an s sublevel; it has a p, which has 3 sublevels; and it has a d sublevel, and each d has 5 sub-sublevels, which are the actual orbitals, for a total of 8, 10--it has a total of 18 electrons that it can hold.0170

This one...we have the fourth sublevel, our fourth primary: it has an s (let me put it...); it has a p, which consists of 3; it has a d...1, 2, 3, 4, 5...which consists of 5; and it has an f...1, 2, 3, 4, 5, 6, 7...this f consists of 7 sublevels.0195

The total here is going to be 32 electrons.0223

That is it: and now, what I am going to write down is the order in which these orbitals, these sublevels, actually fill in.0228

Now, what you are going to notice is: you are going to see numbers (the 1, 2, 3, 4, 5--those are the primaries); you are going to see the s, p, d, and f (those are the sublevels); but you are not going to see a specific notation for the sub-sublevels.0240

OK, so I'm only going to stop at the sublevels; and then, as a superscript, I'm going to write the total electrons in those sublevels.0252

Here is what it looks like; this is the order of filling.0262

It goes in this order: when you put electrons you move down the periodic table (which--we will actually go through a periodic table later on, at the end of this lesson)--when you go through the periodic table and drop in electrons for hydrogen and helium, then lithium, then beryllium, boron, carbon, nitrogen, are going to fill in this order.0265

Order of electron filling--and it is going from lowest to highest energy.0289

It is: 1s2, 2s2, 2p6, 3s2, 3p6...this is just one thing you have to memorize...4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d10, 6p6, 7s2, 5f14, 6d10, 7p6.0298

Let me tell you what is going on here; let me just pick something randomly.0340

As I drop in more and more electrons as I run down the periodic table, I am filling in these orbitals; I am filling in, so the 1--the primary level, the s sublevel--I fill it in with 2 electrons, and I am done.0345

I move to the second primary: the second primary has two sublevels, s and p; in the s, I stick 2 electrons; I am done; in the p, I stick 6 electrons--I am done.0356

I move to the third level--the third: I fill in the s with 2; I fill in the p with 6.0367

Now notice, you know that the third also has a d; but it doesn't fill up yet.0374

As it turns out, it skips the d; it goes to the fourth level, fills up those two, and then returns to the d.0378

3d: it fills up 10 electrons, and then it jumps back up to 4p.0384

It fills it up with 6 electrons.0389

So, I have...some random one...4p4: OK, if I have something like that, I have the primary; I have the sublevel; and I have the total electrons in that sublevel.0391

I am not breaking this down further into sub-sublevels; I will in a minute, when I do an electron diagram--in that one, I do show you everything; but this particular notation--what we call the electron configuration--the standard electron configuration--only lists the primary and the sublevel, and the total electrons in that sublevel.0408

It does not list the sub-sublevels.0431

OK, let's see: OK, so we have all of these orbitals available; so all of these orbitals are available to us--these are just functions, and these orbitals--these energies--are available to us.0434

Electrons can go anywhere, but electrons are always going to go to the lowest energy first, and fill them up that way.0467

They are going to fill them up until they can't fill them up any more.0475

Once the s is full (it can only accommodate 2 electrons), it moves on to this s; s can only take 2--it moves on to the p; the p can take 6--once that is full, it moves on to the next primary.0477

s, then p, then s, then d, then p, then s, then d, then p; it works its way up to higher and higher energies.0488

These, let's start filling them up and give configurations and diagrams.0496

We will start with hydrogen: hydrogen has one electron, right?0523

Well, the configuration for that is 1s1: primary, the s sublevel, that one electron that it has goes into that s sublevel.0529

It looks like this.0539

I'm just going to do it as that: this is the electron configuration; this is the electron diagram.0545

Helium: helium has 2 electrons: 1s--the s sublevel can accommodate 2 electrons, so we put that second electron there.0554

This is that 1s2; OK.0565

Opposite spin: each orbital can accommodate 2 electrons, but the electrons have to have opposite spin.0573

They can't be both pointing up or both pointing down; they have to be opposite spin.0579

That is it: we are done with the 1s.0584

Now, we will do lithium: lithium is the third element in the periodic table.0588

Lithium has 3 electrons: its configuration is, according to the filling that I just did--the order of filling: 1s2, and then it goes to the 2s1; that is it.0594

The order was: 1s; once that is filled, you go to 2s.0609

3 electrons: these superscript numbers--they add to the total number of electrons for that atom.0611

It looks like this (now I'll actually label them).0619

How shall I do this?--I'll do 1s, and then, I don't want to...OK, I have...let me do it this way...1s; so this is the 1s; this is the 2s.0626

I have 1s2, 2s1; that is the electron diagram for lithium; this is the electron configuration for lithium.0646

OK, let's do beryllium.0656

You know what, I'm going to do this a little bit differently: this is going to be 1s; this is going to be 2s.0662

OK, it will give me a little bit cleaner, here.0676

Beryllium is next: beryllium--it has 4 electrons.0679

Well, we fill 1s2, 2s2; so we have the 1s orbital; we have the 2s orbital; and oh, by the way...1, 2, 3...we have the 2p's, also (right?--because the second primary has an s, and it also has a p); but notice, they are not filled yet.0684

They are there (actually, they are all there); but they are not filled yet.0704

We fill in that and that, and then we fill in that and that.0709

That is the electron diagram for beryllium: two electrons of opposite spin in the 1s orbital; two electrons of opposite spin in the 2s orbital.0714

OK, now, we get to our first boron: 5 electrons; we fill 1s2, 2s2, 2p1; 2+2+1=5 electrons.0724

We have: these are...I'm going to move these a little further out so I can see...these are 2p...OK, so I have: 1s here; I have 2s here; 2p here.0740

There is 1; there is 2; there is 3; there is 4; there is 5.0768

That is the electron diagram for boron: the electrons are: 2 electrons in the 1s orbital; 2 electrons in the 2s orbital; 1 electron in the 2p orbital.0773

OK, boron...carbon: we have 5 electrons...we have 6 electrons for carbon.0784

It is 1s2, 2s2, 2p2; we are just following the order of filling and adding one electron at a time--that is it.0794

The total number of electrons is just...add all of those superscripts (they are not exponents; they are superscripts).0803

So, the 1s orbital; we have a 2s orbital; the 2p is made up of 3 sub-sublevels, so we have 1 electron, 2 electron, 3 electron, 4 electron, 5 electron...notice where I put the sixth electron: I put it in the next sub-sublevel.0810

This is called Hund's rule: when you fill up a sublevel--a p--that has sub-sublevels, you have to fill...the lowest-energy configuration is when you have the most number of unpaired electrons.0831

I didn't put it here; the reason is--as it turns out, when these electrons fill out, this is a lower-energy configuration than that.0854

The same spin, but in a different sublevel: once I fill up all of my sub-sublevels, then I go back and fill in the others.0868

So again, this is always going to seek the lowest energy configuration it can.0878

And it will always do this, so when we fill in the p's and the d's and the f's, we fill them in one at a time with electrons of the same spin, and then we go back and finish it off.0882

Carbon...what is next after carbon?0893

We have nitrogen (correct?--yes); nitrogen has 7 electrons: it is 1s2, 2s2, 2p3.0897

We have a 1s orbital; we have a 2s orbital; and we have a 2p orbital, which contains 3 sublevels: 1, 2, 3, 4, 5, 6, 7--three unpaired electrons.0909

Oxygen: we have 8 electrons: we have 1s2, 2s2, 2p4; we have a 1s orbital; we have a 2s orbital; and we have 3 2p sublevels, so we have 1, 2, 3, 4, 5, 6, 7.0925

Now I go back; so here, I have 2 electrons in my 1s orbital, 2 of the electrons in the 2s orbital, 2 electrons in one of the 2p sub-sublevels, and then 1 electron each of the same spin in the other sub-sublevels of the p.0948

This is the electron configuration for oxygen--this configuration right here is what makes oxygen the particularly interesting thing that it is.0965

Hopefully, this is starting to come together; let's keep going a little bit further.0980

Let's do fluorine: fluorine has 9 electrons (I know you probably get the pattern by now; I get that, but it's nice to sort of go through it anyway).0985

It is 1s2, 2s2, 2p5; we have a 1s orbital; we have a 2s orbital; we have 1, 2, 3, 4, 5, 6, 7, 8, 9.0998

And then, we have...neon is the first noble gas that we come to.1020

10 electrons (well, helium is the first; neon is the first in the main group): 1s2, 2s2, 2p6.1027

Notice that: 1s2, 2s2, 2p6--after that--see, p can only accommodate 6 electrons; after this, we have to jump to the 3.1037

This 2s2, 2p6--later, when we get to 3s2, 3p6, 4s2, 4p6, all of the noble gas configurations end that way--the complete s and p are full.1045

That is why they are noble: they don't react, because their outer electrons (what we will define in a minute as their valence electrons)--they are complete; they don't need to be filled.1058

OK, so we have 1s; we have 2s; we have 2p; boom, boom, boom, boom--get in the habit of filling these in one at a time, even though you know that you have six of them.1068

1s2, 2s2, 2p6; OK, now let's do a random one.1086

OK, now, randomly, let's do sulfur.1092

Sulfur has 16 electrons; so, let's fill them up.1100

1s2, 2s2, 2p6, 3s2, 3p...1, 2, 3, 4.1104

2, 4, 10, 12, 16...yes, that is the electron configuration.1113

Let's go ahead and do the electron diagram: we have a 1s; we have a 2s; we have a 2p; we have a 3s; we have a 3p; and I'm going to go ahead and put a (you know what, I'm going to do it over here so I have a little more room--I have a) 3s; I have a 3p; I have a 3d; 1, 2, 3, 4, 5.1118

1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16; there you go: notice, 1s2, 2s2, 2p6, 3s2, 3p4: notice this 3p4--it is the same configuration as oxygen 2p4.1150

If you look at the atom underneath sulfur, it's going to be 4p4; if you look below that, it's going to be 5p4.1175

Everything in one group of the periodic table has the same electronic configuration, but at a higher primary level; that is why they behave similarly, because they have the same electronic structure.1184

That is why their chemistry is similar--this is an explanation for why the periodic table is arranged the way that it is.1198

We will see some more of that in just a minute.1205

OK, let's do iron: let's see, iron...OK, iron has 26 electrons.1209

26 electrons: it is 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d6.1218

These are going to start to get kind of long, so I am going to introduce, with iron, a shorthand notation.1235

This is the same as...I can write this as argon 4s2 3d6; and the reason is because this 1s2, 2s2, 2p6, 3s2, 3p6--that is the electron configuration for argon--the noble gas that is in the row right before it, from your perspective all the way to the right.1241

OK, so we have 1s (let's see here...yes); we have a 1s orbital; we have a 2s orbital; we have 3 2p orbitals; we have a 3s orbital; we (oops, let's make some room here--I hope we have enough) have 3 3p's; we have some 3d's (1, 2, 3, 4, 5); we have (yes, I'm definitely going to need a lot more room here) 4s, 4p's (oops, let's make sure these crazy lines don't start showing up again)--1, 2, 3; 4d--1, 2, 3, 4, 5; and we have a 4f--1, 2, 3, 4, 5, 6, 7.1268

So, we fill them up: there is 1; there is 2; 3, 4, 5, 6, 7...and you notice, I am filling in the p's one at a time.1326

Get into that habit.1337

That is 1; that is 2; 3, 4, 5, 6; now, notice, from 3s2, 3p6--the order of filling is now 4s2.1340

So, I'm going to come here to 4s2; I'm going to fill that up; and then, I'm going to go to 3d6.1351

I'm going to go to 1, 2, 3, 4, 5, 6; these orbitals are there, but they are not filled in--there are no electrons occupying them.1357

The order of filling must be maintained when you do this; the 3d orbital--yes, it is at the 3 level, but energetically, the 4s fills up before the 3d does.1369

That is why it looks the way that it does.1380

OK, so now, let's define what we mean by "valence electrons."1383

Valence electrons--it is the total number of electrons in the highest primary energy level.1389

You remember--the "highest primary"--those are the numbers (1, 2, 3, 4, 5) that are the beginning of the electron configurations (1s2, 2s2, 2p6, 3p6) those integers...highest primary energy level.1406

OK, so we'll just do some quick examples.1424

For oxygen, we had 1s2, 2s2, 2p4; the highest primary here is 2, and the total number of electrons is 6, so there are 6 valence electrons.1428

For iron, we had an electron configuration of argon 4s2 3d6; well, now notice: the highest primary--that is the definition (right?--"in the highest primary"); so the highest primary here is 4.1442

Even though the 3d comes afterward, and in a ground state the 4s2 is of lower energy, these are actually the valence electrons.1457

This is just 2: 2 valence electrons on iron.1465

Let's try something a little bit more interesting--how about arsenic?1470

The shorthand notation for arsenic is 4s2 3d10 4p3.1475

The highest would be the 4 primary; so we take 3+2, so we have 5 valence electrons in arsenic.1481

Notice: the 4s2 4p3 happens to be the same as...the sp configuration is the same as nitrogen, which happens to be the 2s2 2p3, which is also why arsenic is below nitrogen, in the same group as, and it has a similar chemistry as, nitrogen.1490

It is the valence electrons that determine the chemistry, and those are the electrons that we are going to be interested in.1508

Next couple of lessons, when we start talking about bonding, Lewis structures and things like that, it is the valence electrons that we are going to count.1514

So now, what I'm going to do is: I'm just going to list a couple of exceptions to the electron configuration and how we fill one at a time, because there are a couple--there are 4 elements that you should be aware of, that do things slightly differently.1522

Here are the exceptions: I mean, for the most part, they are not altogether that important; but again, they might come up, and it's good to know.1537

Chromium and molybdenum, copper and silver; this is how they are arranged in the periodic table--chromium on top of molybdenum, copper on top of silver.1549

The expected electron configuration for chromium would be (based on everything that we have done, it would be) argon, 4s2, 3d4.1561

The true electron configuration, as it turns out, is argon, 4s1, 3d5 (OK, this one--I definitely don't want these random lines here, because I want to make sure we see all of the numbers).1576

So, as it turns out, one of these electrons in the 2 actually jumps up to the 3d, so that the 3d--so that the d sublevel, which has 5 sub-sublevels--so that there is 1 electron in each sublevel; and it leaves only 1 electron in the s.1589

As it turns out, this is a lower energy configuration than that one.1607

That is why that electron that is here jumps up to this one.1612

The best way to handle it for anything chromium in that group, and copper group--basically, do the expected, and then just take one of these electrons and put it over there.1615

That is the best way to do it.1625

Copper: you are going to get an expected of argon, it should be (oh, here we go with these lines again) 4s2, 3d9; the true is actually argon, 4s1, 3d10.1626

Again, the d orbital wants so badly to be filled, when it gets that close to being filled, that it literally pulls an electron away from the 4s, and it leaves that there.1646

OK, so now that we have sort of discussed this and thrown a bunch of symbolism around, let's take a look at the periodic table and see what sort of information it can actually give us.1656

We can look at a periodic table and literally read off the electron configuration.1666

That is what is nice about this; so let's take a look at our periodic table here.1670

All right, so notice how this is arranged: we said...these numbers here (1, 2, 3, 4, 5, 6, 7)--those are the primary energy levels; those are the quantum number n.1677

These two columns (right here--1, 2)--they occupy the s suborbital.1693

1, 2, 3, 4, 5, 6: these 6 columns--they represent the p suborbital.1702

1, 2, 3, 4, 5, 6, 7, 8, 9, 10: what you know as the transition metals--they occupy the d suborbital.1714

1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14: these are the f suborbitals.1724

The only reason they are down here--these actually belong right in here; so, if we were to take these and stick them right in here, it would actually end up making the periodic table two pages wide.1731

That is the only reason they put them down here: because they wanted to fit the periodic table on one page.1743

Now, when we say, "What is the electronic configuration for phosphorus?", we literally just count: 1s1, 1s2, 2s1, 2s2, 2p1, 2, 3, 4, 5, 6; 3s2, 3p1, 3p2, 3p3; so the electron configuration for phosphorus is 1s2, 2s2, 2p6, 3s2, 3p...1, 2, 3.1749

Or, the shorthand notation would be--go up and to the right--neon 3s2, 3p3.1780

Let's do the electronic configuration for manganese: OK, manganese is right over here; so we get 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d 1, 2, 3, 4, 5.1789

Our electron configuration, in shorthand notation, is argon, 4s2, 3d5.1811

We can literally just count electrons; that is what we are doing.1819

These atomic numbers right here--these are (oops, wow, that was interesting; all right, so we have 25 here)--these numbers represent the number of electrons in a neutral atom.1822

When we do the electron configuration, that is what we are doing; we are counting the number of electrons.1836

We are telling you where these electrons go; that is all that we are doing.1840

Notice: carbon: 2s2, 2p2; silicon--3s2, 3p2; germanium--4s2, 3d10, 4p2--4s2, 4p2.1844

Everything in a column has the same electron configuration, except at the next higher primary energy level.1856

Beryllium is 2s2; magnesium is 3s2; calcium is 4s2.1865

Oxygen is 2p4; sulfur is 3p4; selenium is 4p4; that is why they are arranged the way that they are arranged.1872

These fill first--the 1, the 2, the 3; 4--after the 4 fills up, and then the 3d fills up--that is why they are here; that is the whole idea.1883

This periodic table gives you a lot of information.1894

These numbers right up here: 1, 2 (notice), 3, 4, 5, 6, 7, 8--they tell you the valence electrons in that particular group.1898

Don't worry about these--these actually don't tell you the valence electrons in that group; this is an old numbering system.1908

Nowadays, the new numbering system is the Arabic numbers that you see on top: 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18; that just tells you the total number of electrons that you can fill up with the s, the d, and the p.1916

This right here--these are the s electrons: the 1s, the 2s, the 3s, 4s, 5s, 6s, 7s.1932

These represent the d block--the d sublevel has 5 sub-sublevels; each one of those sub-sublevels is an orbital; each one of those orbitals contains 2 electrons.1944

5 sub-sublevels, times 2--10 electrons; that is why they are 10 long.1959

These are the p's: when you fill up the p's, the p sublevel has 3 sub-sublevels; 3 orbitals per p sublevel--that means 6 electrons in total, because each orbital can contain 2 electrons of opposite spin.1965

That is what the periodic table is doing: you can use the periodic table to do your electron structure.1983

You can use the periodic table to read off valence electrons--at least in the main group elements.1989

That is all that is going on here; actually, you can do it for this, too, because in this particular case, let's say we had iron.1999

Well, it's here; well, you know that the 4s2 is the one that is occupied; this is the 3d6.2007

Well, you know that the valence electrons are the ones in the highest primary, so it's just 2 valence electrons.2013

These numbers--again, it's an old numbering scheme; we are not going to be using that.2018

So, with that, we will go ahead and stop it here.2022

Next time, we will continue on with a discussion of...we'll get a little bit greater into depth with...electron configuration, and we'll also start talking about general notions of bonding.2025

We will spend several lessons on that.2037

So, with that, I thank you for joining us here at

We'll see you next time; goodbye.2043