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AP®︎/College Chemistry
Course: AP®︎/College Chemistry > Unit 1
Lesson 5: Atomic structure and electron configuration- The periodic table, electron shells, and orbitals
- Shells, subshells, and orbitals
- Introduction to electron configurations
- The Aufbau principle
- Valence electrons
- Electron configurations of ions
- Electron configurations of the 3d transition metals
- Atomic structure and electron configuration
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Introduction to electron configurations
Electron configurations describe where electrons are located around the nucleus of an atom. For example, the electron configuration of lithium, 1s²2s¹, tells us that lithium has two electrons in the 1s subshell and one electron in the 2s subshell. Created by Sal Khan.
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Why does he write that Lithium's third electron will go to the s subshell at2:42? Isn't the second shell the 3 p's? I think I'm not understanding something important here.(14 votes)
Have a play around with https://www.ptable.com/#Orbital
It should help you see what is going on(15 votes)
Why is there no 1p electron configuration?(7 votes)
The P "dumbbell" shape is not needed for the first shell. By using quantum numbers we see that the first shell only has one orbital and that will defer to the sphere shape and thus it will be 1s.(16 votes)
I can understand nothing in all the concepts of chemistry. Can anyone help me! Or any tips or advice to learn, understand, and score in chemistry. 😩😭(4 votes)
Sometimes when learning a new skill you might get overwhelmed by the steep learning that is required. When learning any new skill it will require a boost of motivation to get past the hardest parts.
Sometimes it may be useful to take a step back, & review concepts that you are a little unsure about.
I hope you find this advice helpful, and remember that if you ever need help on a specific topic, don't feel embarrassed about asking what may seem like "dumb" questions. We all ask questions that we later realize are rather dumb, but asking obvious questions is often how we learn. Sometimes once we learned something we take for granted how much effort people in history have put in to learn & discover just that.
Happy learning! :-)(19 votes)
- Why can't we do 1p instead of 2p for the electron configuration for carbon?(2 votes)
Saying 1p implies using the p orbitals of the first shell or energy level. This problem with this is that there are no p orbitals in the first shell, only the 1s orbital. p orbitals only begin forming in the second shell which is where the 2 in 2p comes from. Remember the first number in electron configuration is the principal quantum number, or the shell number. So even though it's the first p orbital you observe in electron configurations, it only first appears in the second shell.
As to why there aren't p orbitals in the first shell, in quantum mechanic talk, there are no angular nodes in the first shell to form a p orbital.
Hope that helps.(10 votes)
Why does it go from 2s to 2p, instead of to 1p?(3 votes)- We build electron configurations by filling the lowest energy orbitals first then filling progressively higher energy orbitals. This is known as the aufbau principal. So we fill subshells in the order 1s 2s 2p because 1s is the lowest energy, 2s is higher energy, and 2p is highest energy.
We don’t fill a 1p subshell because there is no 1p subshell. The first electron shell only contains a 1s subshell.
Hope that helps.(5 votes)
In earlier classes, we were taught the 2,8,8,2 Electronic configuration for Elements having Atomic numbers from 1-20. What does this have to do with s, p, d, and f subshells?
Why isn't this Electronic configuration applicable for Elements having Atomic numbers greater than 20?(3 votes)- So you're 2,8,8,2 rule looks like it shows the maximum amount of valence electrons elements can have in a period up to atomic number 20 (calcium). So elements in period 1 (hydrogen and helium) need 2 electrons to fill their valence shell. Elements in period 2 (building on the previous 2 electrons) need 8 more electrons to fill the second electron shell which is their valence shell. And so if I'm understanding it correctly.
This is a decent starting point, but later you learn in chemistry that the placement of electrons is more complicated than this. Electron shells are divided into subshells, which are further divided into orbitals.
The subshells which compose electron shells describe the general shape of electron's orbit around the nucleus. And we use the letters s, p, d, and f for subshells. These subshells are themselves composed of orbitals which are the specific orbits of the electrons and each subshell has a certain number of orbitals. s subshells have 1 orbital, p has 3, d has 5, and f has 7. An individual orbital can hold a maximum of 2 electrons.
Each electron shell has a certain amount of subshells (and therefore orbitals and therefore electrons it can hold). The first shell only has an s subshell which means it only has 1 orbital which is why period 1 elements valence shell only holds 2 electrons. The second electron shell is composed of an s and p subshell which has 4 orbitals altogether which is why period 2 element's valence shell holds 8 electrons. This pattern continues, but it starts to get more complicated with the inclusion of the d subshell and the transition metal elements beginning after calcium.
Hope that helps.(3 votes)
For any given element, are there always going to be 2 electrons in each respective orbital?(1 vote)
I think more than 2 electrons can be accommodated in any specific orbital, but the max numbers of electrons that can be accommodated in the outermost orbital is 8.(2 votes)
- So basically shells are energy levels that can be presented through orbitals of different shapes, that can only hold a certain number of electrons? Am I understanding this correctly? So basically there can be S-orbitals in the second shell which only hold 2 electrons, but the second shell can hold more (up to 8?) if it was a p-orbital that formed?(2 votes)
- As electrons are added to the space around the atom's nucleus they are arranged in a way as to minimize repulsions. The pattern that we observe results in 3 classifications as follows (there are other considerations that you will learn about later as well).
A Shell / Energy Level is a region or set of regions that have the same energy.
Although we cannot predict the exact location of the electrons at any time, we can map out the regions of space that they occupy. We can organize these electrons into different subshells based upon the shape of the region they occupy. Finally, within each subshell there are individual orbitals referencing a specific region of space around the atom's nucleus. Each orbital then has 2 electrons, which are said to have different and opposite spins.
In the 1st energy level, we have 1 subshell, which basically means we have 1 type of orbital. This specific type of orbital is called the s orbital, and we have 1 s orbital for every s subshell.
In the 2nd energy level, we have a p subshell in addition to the s subshell. In every p subshell there are 3 p orbitals. This means that we have a total of 8 electrons in the 2nd energy level (2 from the s subshell, and 6 from the p subshell).
As you go further, you will get practice with identifying and conceptualizing electron configuration.
Hope this helps.(4 votes)
- I notice you've written the electron configuration for Li as 1s²2s¹ whereas most periodic tables will display it as simply 2s¹ (ignoring the 1s²).
Is this simply because it is accepted that if electrons are filling the 2s orbital they will, by definition, have already filled the 1s orbital with the maximum amount (2 for the 1s orbital) of electrons?(2 votes)- Periodic tables don't usually have electron configurations in them due to space restraints, but is it possible they are using the square bracket notation? Like you can write 1s2 2s1 out as [He] 2s1, which basically means the electron configuration of He (the previous row's noble gas), and then whatever else is outside the square brackets.
It doesn't save much space here, but imagine if you were writing the electron configuration for say iron out, it takes much less space to write out [Ar] 4s2 3d6 than to write 1s2 2s2 2p6 3s2 3p6 4s2 3d6(4 votes)
- in what grade is this concept usually taught? i mean in what grade is it appropriate to start learning this?(3 votes)
- I think that it starts in college as the grade says College Chemistry.(1 vote)
2:42Video transcript
- [Instructor] In a previous video, we've introduced ourselves to the idea of an orbital, that electrons don't just orbit a nucleus the way that a planet might orbit a star, but really, in order to describe where an electron is at
any given point in time, we're really thinking about probabilities, where it's more likely to be found and less likely to be found. And an orbital is a description of that, where is it more or
less likely to be found. And this diagram shows
us the types of orbitals which can be found in
the various subshells which are found in the various shells. So you have the s subshell, the p subshell that has three
different orbitals in it, you have the d subshell
that has one, two, three, four, five different orbitals in it. And then you have the f subshells. Now each orbital can fit two electrons. So if you're thinking about the subshell, the s subshell could fit two electrons, the p subshell can fit six electrons, the d subshell can fit 10 electrons, and the f subshell can fit 14 electrons, two per orbital. Now the goal of this video is to think about electron configurations
for particular atoms. And to help us with that, we will look at a periodic table of elements. And so first, let's just think about the electron configuration
of the simplest element. If we're talking about
a neutral hydrogen atom, a neutral hydrogen atom, it
has an atomic number of one which tells us it has one proton, and if it's neutral, that
means it has one electron. Now where would that one electron be? Well it would be in
the lowest energy level or the first shell, and that first shell has only one subshell in it. It only has one type of orbital. It only has an s subshell,
and so that one electron in that neutral hydrogen
atom would go over there. So we would say its
electron configuration 1s1, in the first shell which is made only of an s subshell, it has one electron. Now what happens if we go to helium? Well, a neutral helium atom is going to have two electrons. So instead of just having one
electron in that first shell, we can fit up to two there. So its electron
configuration would be 1s2. Now what do you think is going to happen when we go to lithium? Well lithium, a neutral lithium will have three electrons in it, so the first two could go
to the first energy level, the first shell, so the
first two will go 1s2, and then the third electron is going to go into the second shell, and the subshell that
it's going to fill first is the s subshell. So then it'll go to the second shell and start filling up the s subshell. So notice, two electrons
in the first shell and one electron in the second shell. Now what about beryllium? Well, that's gonna look a lot like lithium but now it has four electrons. So two of them are going to
go into the first shell, 1s2, and then the next two are
going to fill up the s subshell in the second shell. I know it's a bit of a mouthful, 2s2. Notice, we have four total electrons which would be the case in
a neutral beryllium atom. But what about boron? Boron gets interesting. A neutral boron would have five electrons. So the first two we're going
to fill the first shell, 1s2, now the second two are then going to go to the second shell and
fill up the s subshell 2s2, and then we're going to start
filling up the p subshell. So let's see, we have one
more electron so we go 2p1. So we're going to have one electron in one of these p orbitals. And then what happens
when we go to carbon? Well it's going to look a lot like boron but now we have one more
electron to deal with if we have a neutral carbon atom, it's going to have six electrons. So then an extra electron is once again going to fall into the p subshell in the second shell because
that can fit six electrons. So we're going to fill the
first shell with two electrons then the 2s subshell with two electrons, and then we have two more electrons for the 2p subshell. Now you can imagine as we get
to larger and larger atoms with more and more electrons, this can get quite complex. So one notation folks often use is noble gas configuration where instead of saying,
okay, this is carbon, they could say that, hey look, carbon is going to have
the electron configuration of helium, remember, the noble gasses are these Group 8
elements right over here, so it's going to have the
electron configuration of helium which tells us this right over here, and then from that, we're
going to also have 2s2, 2s2, and then 2p2. You could just take helium's
electron configuration right over here and put it right over here and you would get exactly
what we wrote before.