# Electron Configuration in the Periodic Table

by Jared Rovny
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00:01 Now let's apply this to something much interesting or may be practical.

00:04 Something that you've seen which is the periodic table.

00:06 The great thing about just these 4 quantum numbers which include the 4 different possible shapes of the orbitals the l quantum number, is that just these 4 numbers describe the entire shape of the periodic table and why it looks the way it does.

00:23 So take a look at the periodic table we have written here.

00:26 Don't forget that in our an electron configuration notation, which we've written to the left here, you always write the energy level and the shape of the energy level l.

00:36 And then you put the number of electrons that are in that shape with a little super script above the l, the spd or f.

00:43 Looking at the periodic table, you can see exactly these numbers that we just described played out.

00:48 So for example in the far left column we see these two blue columns.

00:53 And there's two of those because this is our 's' orbital, that perfectly spherical orbital.

00:59 So that far left column is only going to be fit two electrons.

01:04 Because again that's spherical orbital can only fit a spin up and a spin down electron.

01:08 Going to the next section which is all the way to the right of the periodic table that's that pink area.

01:14 We have the 'p' orbitals.

01:16 The lobe structure.

01:18 And now you can see why they are fixed columns in that section of the periodic table.

01:22 We have 6 columns as we have said, we can fit 6 electrons in the p shape.

01:28 And so have a 6 rows or 6 columns rather in the pink section.

01:32 We have 2 columns in the blue section.

01:34 And we also have now the 'd' section, the d orbital shape.

01:38 And that's in the yellow, right in the middle.

01:40 And now we've already said we will have 10 columns because these atoms can fit or this particular orbital rather, can fit 10 electrons.

01:49 Again spin up and spin down for each different orientation.

01:52 Finally, written in the blow here sort of in a subscript is this green row.

01:57 We the f orbitals.

01:59 The highest orbital shape that we experienced in the periodic table.

02:03 And as we already saw the number of electrons you can fit there is going to be 14.

02:08 And you can calculate this for yourself again in the way that we just described.

02:12 So looking at this periodic table, we can describe the electronic structure for any atom that appears in the periodic table.

02:18 So first of all I should point something out which is that there seems to be a bit of incongruity here which is that the helium on the far right seems like it might be placed on the far left with the blue column.

02:33 And the reason is put on the far right is convention having to do with noble gases which I'll discuss in a minute.

02:38 But for this structure of the table that we just described you can think of the helium is being in the far left column in terms of the electronic structure.

02:46 But let's count the electrons in the helium atom.

02:48 The way we describe the number of electrons in the helium atom is 1 that 1 being for the energy level which tells us which row of periodic table we are in.

02:57 So we are in the first row, in the first energy level for our electrons.

03:01 In the s-orbital shape.

03:03 Since again the helium belongs sort of in the first column, it's still in that circular orbital shape.

03:09 And then we have two electrons, because in that blue column of two rather than being the hydrogen, the first would have 1 electron.

03:17 Helium is in the next column so it has two electrons.

03:21 And so the electron structure for helium is 1s2.

03:25 We can keep going and use the same sort of analysis for higher element numbers.

03:29 So for example for fluorine here, we have a structure that is 1s2, so we're counting.

03:35 We're just sort of reading it as you would a book.

03:37 You start in the top left.

03:38 You read to the right and then you go to the next row.

03:40 You start in the top right and then go to the next row and so on.

03:45 So doing this for fluorine, we start in the top left at hydrogen, we read along and we have the 1s2.

03:49 Just like we had for helium.

03:52 And then we go to the next row.

03:53 And we keep going until we get to fluorine.

03:55 So then in the next row, we have to count the next row of the blue column.

03:59 That's a 2s2.

04:00 2 because we're now in the second energy level.

04:03 S because we are still in that blue shape, the s or circular orbits.

04:07 And 2 because the size of those two columns fits two electrons.

04:11 And then reading just through left or right, we skip over the entire d section because we haven't gone low enough in our energy levels to get there yet.

04:19 So we're going to read all the way across and jump to the p orbitals, that pink section.

04:23 And start just counting up how many electrons we have until we get to fluorine.

04:27 P could fit in principle 60 electrons but we're not going all the way to the right of the periodic table just yet.

04:34 We count only up to fluorine.

04:36 So we only count 5 different boxes or 5 electrons.

04:40 And so using the exact method, just reading through your periodic table, you can find the electronic structure of fluorine exactly as written.

04:47 1s2s2 2p5.

04:51 Moving along we can see another one.

04:53 So for example phosphorous, we can go down.

04:55 And we'll just use the exact same counting method that we described.

04:58 We start at the top, move left, move down and we keep building up our orbitals.

05:01 We have 2s2 and then 2p6.

05:04 And then we go down another row and we have the three level the 3s2.

05:07 And then we get to the 3 p.

05:09 We go all the way across to the pink.

05:11 But we only have a few of those electrons in the p section.

05:14 The problem here is that imagine what happens if we keep going.

05:17 This just get's bigger and bigger and bigger.

05:19 We kept going lower and lower and lower in the periodic table, the length of our electron configuration notation will become completely impractical.

05:28 So we have a trick.

05:29 What we do now is we look at the previous noble gas.

05:32 So this is the noble gas as I was talking about.

05:35 The noble gas which you can see in the far right of the periodic table, is defined as those columns in the far right, which means that the electron orbitals are filled.

05:44 So what we mean by that is if you've moved all the way to the right, you've counted up all the electrons that can possibly fit in a particular orbital.

05:53 We saw that recently with fluorine, how we were counting along the p-orbital and adding more and more electrons.

05:58 But we stopped before we got to the far right.

06:00 Because fluorine wasn't far enough over.

06:03 The noble gases are the opposite.

06:05 The nobles gases are always after you've counted all of your electrons.

06:08 And so we say that the noble gases have their outer orbital, their outer shell of electrons filled.

06:13 And this is that convention that I mentioned with helium.

06:16 The reason we write helium all the way to the right, even though it's a blue box and blue in this periodic table, is because helium is also a noble gas.

06:24 It's outer shell, that s-shell is completely filled.

06:27 And so helium also behaves like a noble gas.

06:30 So using these noble gases, we know that for example if I were counting my way up to phosphorous, I would read again left to right, top to bottom, keep going.

06:40 And right before I go to the row of phosphorous when I was still at the second row, I know that since I have to get to the phosphorous, I read across that entire row.

06:48 So what I can do instead of counting the entire electronic structure for phosphorous all over again.

06:53 We are writing this big long expression.

06:55 I could instead just start at the prior reference point if you will.

07:00 We write it like this.

07:01 So you see that neon rather is the noble gas that's right before phosphorous.

07:06 We also know what the electronic structure of neon would be because we keep counting the periodic table up to that point.

07:12 So what we do for phosphorous, is we simply write this neon in brackets which represents the entire electronic structure for neon.

07:19 So it's sort of like an abbreviated jumping point, an abbreviated starting point.

07:24 Because again we know very well what the electronic structure for neon is we keep counted out.

07:28 And we can write out all of our orbitals.

07:30 So starting from neon all we do is start there and then continue reading.

07:34 We go to the next row rather.

07:36 Start reading those 3s orbitals, jump all the way across and read the 3p orbitals up to phosphorous.

07:43 So one more example just so we understand this neon structure, rather this noble gas structure.

07:47 Give this one a try.

07:48 Suppose we're talking about titanium.

07:50 So titanium, you can see here is now in the d-level.

07:54 You could reference if you like the prior noble gas the one right before it, and try to write out the electronic structure for titanium.

08:01 I'm going to show it in a moment here so I would recommend pausing.

08:04 Seeing if you can give this a try using exactly the conventions that we just introduced.

08:08 If you did this hopefully it look something like this.

08:11 You start with the argon, the prior noble gas, the one right before the titanium.

08:16 And then you continue reading.

08:17 You go the next row.

08:19 You read across that row to the three -- the two of the 4th level energy electrons in that configuration, those s's.

08:28 And then you move right into the d-structure.

08:29 But just those two first electrons in the d-structure giving you 4s2 3d2 to get the titanium electron structure.

### About the Lecture

The lecture Electron Configuration in the Periodic Table by Jared Rovny is from the course Electronic Structure.

### Included Quiz Questions

1. 6
2. 5
3. 4
4. 3
5. 2
1. The periodic table is arranged in the order of added electrons, and electrons fill the “4s” section before the “3d” section.
2. The “3d” section is put with its corresponding energy level because the electrons will fill lower values of “n” first always.
3. The Pauli Exclusion Principle disallows electrons from filling the “3d” section before the “4s” section.
4. This is just for historical reasons, and the “3d” section would be placed higher if the periodic table were recreated now.
5. Electron configuration notation dictates that the energy level n come before the angular momentum quantum number.
1. By referring each electron configuration from the prior noble gas
2. By referring each electron configuration from the prior angular momentum quantum number
3. By omitting the angular momentum quantum number when it can be easily deduced
4. By omitting the energy level quantum number when it can be easily deduced
5. By listing the energy levels independently of the other quantum numbers
1. [Ar]4s1
2. [Ar]3s4
3. [Ar]4p1
4. [Ar]3s24p4
5. [Ar]5d2
1. To keep with the other noble gases, which have their outer electron shells filled.
2. Each row in the periodic table must span the full width.
3. The Pauli Exclusion Principle disallows it from remaining in the s section.
4. The n=0 energy level can only hold one electron.
5. The 1s section of the periodic table cannot have its outer shell filled.

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