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### Course: Chemistry archive > Unit 1

Lesson 5: Electron configurations- Shells, subshells, and orbitals
- Introduction to electron configurations
- Noble gas configuration
- Electron configurations for the first period
- Electron configurations for the second period
- Electron configurations for the third and fourth periods
- Electron configurations of the 3d transition metals
- Electron configurations
- Paramagnetism and diamagnetism
- The Aufbau principle
- Valence electrons
- Valence electrons and ionic compounds
- Valence electrons and ionic compounds
- Atomic structure and electron configuration
- Introduction to photoelectron spectroscopy
- Photoelectron spectroscopy
- Photoelectron spectroscopy

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# Electron configurations for the second period

Writes out quantum numbers for all elements of the second period. Introduces Hund's rule, and connects blocks in periodic table with electron configuration. Created by Jay.

## Want to join the conversation?

- What is a "ground state condensed electron configuration" for an element? Do we use the most recent/previous noble gas and simplify the configuration accordingly?(16 votes)
- Ground state - most stable, all electrons are in their natural state

Condensed - use closest noble gas with a smaller atomic number (so you're correct)

For example, calcium's ground state condensed electron configuration would be [Ar]4s^2(21 votes)

- Why do opposite spins of electrons in same orbital make it more stable?(8 votes)
- Electrons have charge, and a spinning charge generates an magnetic field. If you do the right hand rule, you'll see that electrons with opposite spins have opposing magnetic fields. Essentially, these magnetic fields "cancel" each other out, thereby making it more stable than if electrons were spinning in the same direction. The magnetic force would push electrons away from each other when electrons have the same spin.(20 votes)

- How would one state Hund's rule in a phrase?(3 votes)
- If two or more orbitals of equal energy are available, electrons will occupy them singly before filling them in pairs.(22 votes)

- If there's only one electron in an orbital, do we assume the spin is +1/2? Does it matter?(5 votes)
- It doesn't really matter.

It is just a convention to assign +½ to the first electron in an orbital.(8 votes)

- i suddenly feel like the hunds ruls as described in the videos suddenly contradicting what we have learnt before !

why its being said that there wecan only be one electron in px orbital when the previously learnt rule was that we cud accommodate 2 electrons in each orbital.(2 votes)- Hund’s rule states that we fill electrons singularly into the orbitals of the same subshell before we begin pairing them.

For an s subshell this doesn’t affect the filling at all since there is only one orbital in an s subshell. For a p subshell however, there are three orbitals now. So using Hund’s rule we have to fill those three p orbitals with three electrons of the same spin before be begin spin pairing with the fourth electron.

Hope that helps.(3 votes)

- Does Hund's Rule apply only to p orbital or all the upcoming ones too like d ,f , etc?(2 votes)
- Hund's rule applies to all of the orbitals. 1 electron occupies each orbital, and only after all of the orbitals are filled does the orbital get filled with two electrons.(3 votes)

- Not sure if this question is answered, but Be has config 1s2,2s2, though it is not a Noble gs it is very tempting to say that it has filled subshells.Why we don't call it a complete octet or rather filled subshell(2 votes)
- So neutral beryllium does have a completed 2s subshell and gains a certain level of stability from this. We can tell this because beryllium has a noticeably higher first ionization energy compared to both lithium and boron. It is just that this level of stability isn't had pronounced as an actual noble gas.

Hope that helps.(3 votes)

- Can we write Neon(1s^2 2s^2 2p^6) as [Be]2p^6 ?(2 votes)
- Technically you can but it is not practised as a notation as noble gas notations are completely filled configurations. I don't think you will get marks if you write it in this way, but it's not a wrong method.(3 votes)

- What do the L, Ml, and the other letters stand for? Or where can I find a video to explain that concept better?(2 votes)
- What is the hund's rule ?(1 vote)
- If you have multiple orbitals of the same energy (so the three p orbitals or the five d orbitals) you singly occupy each orbital with an electron of the same spin before you pair any of them up.(4 votes)

## Video transcript

- Let's do electron configurations for the second period. So we find the second period on the periodic table and we go across and the first element we see is Lithium with three electrons. So three electrons to
worry about for Lithium. Let's think about the first two electrons for Lithium. The first two electrons for Lithium are going to go into the first shell. So we talked about this in the last video. The first shell when n is equal to one, the only possible value for l is zero. So we're talking about an s orbital. And there's only one s orbital in the first shell here so I can draw in, let me go ahead and draw that orbital in. So here's the one s
orbital in the first shell. So Lithium has three electrons. The first two electrons for Lithium are going to go into this one s orbital. So we pair up our spins like that. So writing the electron
configuration for Lithium, let me go ahead and we'll start writing it right here. So we have one s two so far. Well Lithium has three electrons but the first shell is full, it's closed. So we have to move on to the second shell to add in Lithium's third electron. So in the second shell, n is equal to two. What are the allowed values for l? L could be equal to zero or l could be equal to one. So we talked about that again in the videos on quantum numbers. So when l is equal to zero, we're talking about an s orbital. So in the second shell, in the second energy level, we also have an s orbital and we also have one of them and we also have to think about l is equal to one, that's
talking about a p orbital. The allowed values for ml would be negative one, zero and positive one. So three possible values means three p orbitals here. So we have three p orbitals in the second energy level as well. So let's draw those in on our
orbital diagram over here. So we already drew in this s orbital in the first shell. Next let's draw in this s orbital in the second shell,
the second energy level. It's of higher energy. So we draw it in here. This is the two s orbital. Then we also have p orbitals in the second energy level, we have three of them. So we draw in our p orbitals in the second energy level. They're of higher energy so here are the two p orbitals and
there are three of them. So one of them is, it would be two px, one of them would be two py and one of them would be two pz. Doesn't really matter which one is which. We'll just draw them in
there like that so far. All right, so Lithium. We've taken care of the two of it's three electrons. It's third electron has to go into this next highest orbital
in terms of energy, so that would be the two s orbitals. We have energy going up this way. So as you get higher and
higher, energy increases. So Lithium's next
electron, as you build up the Lithium atom must go
into this orbital here. The two s orbital. So therefore, Lithium's
electron configuration is one s two, two s one and remember what these numbers mean. So this means that there is one electron and an s orbital in the
second energy level. So we have one s two, two s one for Lithium's electron configuration. Let's do the next element. So that's Beryllium. Beryllium has four
electrons to worry about. So for Beryllium, if
you look at the diagram, Let's see if we can just
make a different color here for Beryllium, so let's
make Beryllium red here. So one more electron. So we can put Beryllium's fourth electron into this orbital and pair up our spins. So let's write the electron
configuration for Beryllium. So it would be one s two and then we have two electrons and then two s orbitals, so we would write two
s two here like that. Now, we've filled the two s orbitals. Remember, each orbital can hold a maximum of two electrons. We filled the two s orbitals so when we move on to the next element, which is Boron over here. So Boron has five electrons. So let's write the electron
configuration for Boron. Well so far we have one s two, two s two but that only takes care of four electrons, we need five. So where does the fifth one go? The fifth one goes into the next available orbital here so we're going to put the electron in, the fifth electron for Boron goes into a two p orbital. So we would write two p one indicating that the fifth electron for Boron went into a p orbital in
the second energy level. So one s two, two s two, two p one, is the full electron
configuration for Boron. All right, so let's do Carbon. So next we have Carbon. Let's use blue for Carbon here. So Carbon has a total of six electrons. We have one more to think about. So we know it's going
to go into a p orbital, a p orbital in the second energy level. The question is which
one of these p orbitals do we put this next electron for Carbon? So we have to think about something called Hund's rule. I'm never going to pronounce the German properly so Hund's rule tells us that our goal is to minimize
electron repulsion here. So let's think about... Let me just go ahead and draw the p orbitals down here. So we already have one
electron right here. Well it doesn't make any sense to put an electron into the same orbital here because that puts the
electrons really close together in space. So if you're thinking about a p orbital, remember a p orbital is shaped like a dumbbell so I'm just saying we have a p orbital on
this axis let's say. So we all ready have, let me use, I'll just use blue here. So we already have one electron in there, it doesn't make any sense to add an electron to that exact same p orbital. That puts them really close together in space and electrons repel. So that doesn't make any sense so we need to take that
electron out of there. That doesn't make any sense. We need to add an electron
to another p orbital. So we'll take this electron
out of there like that. So remember, there are other p orbitals on these other axis here. So here's another p orbital and then here's another p orbital. So we need to add an electron to another one of these. Whichever one, it doesn't really matter. Let's just say we're adding one here. So we're adding an electron to a different p orbital,
whichever one it is, px, py, or pz. And it turns out that keeping the spins parallel helps to minimize
the electron repulsion for pretty complicated reasons and I think they might still be doing research on this and
so we put the electron in a different orbital and
we keep the spins parallel which helps to lower
energy for the atom here. And so that's where we're going to put Carbon's sixth electron. We're going to put it into a different p orbital and we're going to keep the spins parallel like that. So we can go ahead and write Carbon's electron configuration. Just read off everything we have on our orbital notation here. So we have one s two, we have two s two and we have two p two here. So two electrons in the p orbitals in the second energy level for Carbon. Next we have Nitrogen so let's use green here for Nitrogen. So Nitrogen has seven electrons, so one more electron to think about. Let's put Nitrogen right here. So we have so far, one s two, two s two. Now let's think about Nitrogen. So we need to add one more
electron to our diagram. Once again, we're going to
follow Hund's rule here. We're not going to add the electron to one of the already occupied orbitals, we're going to add this
electron for Nitrogen to an unoccupied orbital and we're going to keep the spins parallel to keep everything lower in energy. And so we have three
electrons for Nitrogen, in the two p orbitals. So we write two p three. So we have one, two, three. So we have one s two,
two s two, two p three, would be the full electronic configuration for Nitrogen. Let's move on to Oxygen. So let's pick, let's see here, what color should we pick for that? Let's use orange here for Oxygen. So we have eight, eight total electrons. So for Oxygen, let's see,
let's put Oxygen right here. So so far we have, one s two, two s two. So how many more electrons
do we need for Oxygen? Oxygen has a total of eight electrons, we all ready represented four, so we need to represent four more. Oxygen's eighth electron, now that all of our orbitals are occupied, we can start to pair our spins. So we put Oxygen's
eighth electron in there. So we can start to pair up our spins. We have four electrons
in the two p orbitals for Oxygen, so we write two p four. So one s two, two s two, two p four. Notice if you add these together, two and two and four then you get eight which is the total number of electrons that we had to represent for the electron configuration for Oxygen. Let's move on to Fluorine. So let's use a different
green here for Fluorine. So let's say Fluorine right here. Nine total electrons. So once again, we're pairing up our spins. So we add Fluorine's
ninth electron to there and we can go ahead and
write it right here. So for Fluorine we would write one s two, two s two, and notice we have five electrons now. So two p five would be for Fluorine. And then finally, let's
go ahead and do Neon. So Neon has ten electrons. So we have one more
electron to account for. We have one more space right? The last electron for Neon would go into a two p orbital here. So for Neon we would write, one s two, two s two, two p six. Notice we have no more
places to put electrons, in the first or the second energy levels. We are completely full. So the second shell is now full and if you wanted to add another electron, you would have to open up a new shell. You would have to go to
the third energy level. And so you notice a pattern here emerging on the periodic table. So we said that Hydrogen's electronic configuration over here was one s one. Then we went over here
to Helium was one s two. And then we moved on to
the second energy level. So this was Lithium
here ended in two s one. And Beryllium ended in two s two. And then we filled the s orbital and moved on to the p orbitals. And notice we have over here, Boron's last electron was two p one, Carbon's two p two,
Nitrogen's two p three, Oxygen's two p four, Flourine's two p five and Neon is two p six. So notice we have these
six boxes over here on the periodic table. Those represent our p orbitals. And then over here on the left we have these two boxes representing our s orbital. And so that's the idea. The s orbital, we have one of them, holds a maximum of two electrons. We have these two boxes
on the periodic table. Over here on the right,
we had these six boxes which is the maximum number of electrons we can put into the p orbitals because we have three p orbitals, each one can hold two. So noticing these patterns on the periodic table helps you when you are writing
electron configurations. You can just sit down and look at the periodic table and write them out after you've had enough practice. So make sure to do all of these again and think about electron configurations. Where you're putting your electrons and think about how it relates to the structure of the periodic table.