Question 1

I don't really understand what he means by  shells or the numbers and letters like F:1S2

Accepted Answer

So electrons exists in clouds in a way around the nuclei of atoms. These regions where we find electrons are represented by the quantum numbers, of which there are four. These include the principal quantum number, the angular quantum number, the magnetic quantum number, and spin quantum number. Due to the Pauli exclusion principal no two electrons can have exactly the same quantum numbers.

The principal quantum number, often represented by n, is another term for the electron shell and are positive integer values (n=1,2,3,etc.) These electron shell numbers broadly correspond to the period, or row, an element is in. So elements in the first row of the periodic table have the first electron shell, or n=1, as their valence shell or outermost electron shell.

The angular quantum number tells you the subshell an electron is in within an electron shell and is represented by the letter 'l' often. The allowed values of l in an electron shell are determined by: l = n-1. So electrons in the first electron shell can only have a value of l=0 since n=1. Electrons in the second electron shell can have a value of l=0 AND 1 since n=2, and so on. The subshell tells us broadly the shape the orbitals which hold the electrons take on and these are represented by the letters s, p, d, and f. The values of l correspond to the subshell letters, so: l=0 is s, l=1 is p, l=2 is d, and l=3 is f.

The magnetic quantum number, represented by ml, tells you the specific orbital within the subshell an electron is in. The values of ml in a subshell range from +l to -l. So if l=1, then ml could be -1,0, or +1. This tells us why the s subshells only have 1 orbital while p subshells have 3, and the other subshells have the amount of orbitals that they do (and by extension how many electrons they can hold).

Finally there is spin quantum number, represented by ms, of which an electron in an orbital can only take on two values: +1/2 or -1/2, broadly referring to the electron spinning clockwise or counterclockwise if you want to think of it like that. This is why orbitals can only hold a maximum of 2 electrons because they can only take on one of two ms values.

Putting all these together you can specify exactly which electron you are referring to in an atom.

For the purposes of electron configuration though, we only care about the principal quantum number (the shell number), the angular quantum number (the subshell or the letter), and the number of electrons in those subshells which are basically a result of the magnetic quantum and spin quantum numbers. Shell numbers are written first, then the subshell as a letter, then the number of electrons in that subshell are written as a superscript to the letter.

For for example neutral hydrogen has the electron configuration of: 1s^(1). The first 1 is the shell number, the s is the subshell the electron exists in, the 1 superscript is the number of electrons in that subshell. The order which you fill subshells for more massive elements follows the aufbau principal.

Sorry for the long response, but that's the stuff you need to know to do electron configurations well. Hope that helps.

Question 2

Where did we learn about electron configuration before??

Accepted Answer

On KA, here.

https://www.khanacademy.org/science/ap-chemistry-beta/x2eef969c74e0d802:atomic-structure-and-properties/x2eef969c74e0d802:atomic-structure-and-electron-configuration/v/introduction-to-electron-configurations

Question 3

does this mean you can't differentiate between atoms just by their electron configuration?

Accepted Answer

If we were dealing with neutral atoms which never had electron transfers, then yes we would have unique electron configurations and you could identify an element solely by them. However because electrons are transferred frequently we can have two particles with the same electron configuration, but from different elements. This is a phenomenon called isoelectronicity. So this means we can have a neon atom, a fluoride anion, and a sodium cation and we wouldn't be able to discern any difference if we only inspected their electron configurations since they are all the same.

Hope that helps.

Hope that helps.

Question 4

I think adding a video on an intro to electron configuration and how the orbitals are filled before this video would help a lot. 
I had to search up a separate video on these topics to understand what was going on.

Accepted Answer

I agree, students need more material before they start doing electron configurations. It's odd because here in the high school chemistry section it's excluded (the extra material), but included in the AP chemistry section.

Question 5

since the postive charged Ca ion has the same configuration as Ar does that mean that the Ca ion is now considered Ar??

Accepted Answer

piggybacking on what zelmen said, the thing that determines which element it is, is the number of protons not electrons. It can happen where an element loses a proton because it is unstable, that is what we call radioactive decay.

Question 6

I tried following his advice on trying some electron configurations for myself, but ended up randomly picking the exception to the Aufbau principle: Rh.

Why is the electron configuration for Rh: [Kr] 4d^8 5s^1 instead of [Kr] 4d^7 5s^2?

Accepted Answer

Yep that element, rhodium, is a good example how general rules, like the order of filling electrons, are not always correct and why we have to keep in mind exceptions. Sometimes there is more nuance in chemistry than a simple rule can handle.

Most of the exceptions to normal electron configuration rules for the transition metals occur in rhodium’s period, the fifth period. Elements like molybdenum and ruthenium also have different electron configurations than what we would predict using the usual rules. The difference is usually that one of the electrons in the s subshell is demoted to the lower d subshell. Except for palladium which demotes both of its s electrons into the d subshell. Whatever the change is though between the rule and what is observed, the goal is still the same, to achieve a lower energy state. Essentially the rhodium atom gains more stability by having that s electron in the d subshell than allowing it to remain in the subshell.

Hope that helps.

Question 7

Can fluoride exist by itself? I understand that it is an ion and that it must take an electron to be negatively charged, but can't it just steal that electron from like potassium and those elemental ions exist and be separated as ions?

Accepted Answer

Yes. Fluoride as part of an ionic compound like sodium fluoride (NaF) is a solid at room temperature and is bound to the cation. But many ionic compounds are soluble in water and dissociate into their constituent ions. So when sodium fluoride is in water, the water molecules essentially separate the sodium cations and fluoride anions from each other and they exist as individual ions.

Hope that helps.

Hope that helps.

Question 8

Are there 18 electrons in the third shell or 8 and if there are 18 electrons in the third shell why do we only fill it up to 8

Accepted Answer

So in the third shell there's s, p ,and d orbitals which, if fully filled, do add up to 18 electrons. But if we have an element in the third row (or period) like sulfur or phosphorus, we're only going to fill up the 3s and 3p orbitals in their electron configurations because that's where their valance electrons are. Valence electrons being the electrons that are important to chemical bonding. Once those 3s and 3p orbitals are filled, we'll have used 8 electrons. Only when we start doing the electron configurations of 4 row transition metals like iron and cobalt do we start back filling the 3d orbital. So third row elements like to have an octet to have an electron configuration similar to a noble gas like argon because it makes them more stable. So third row elements are similar to second row elements like carbon and oxygen who also want to have an octet, however third row elements aren't limited to only have 8 electrons because of that 3d orbital and can have greater than 8. It's just that they don't use that 3d orbital as frequently because once they've reached that noble gas electron configuration, they're stable and don't have an overpowering desire to exceed that octet. Hope this helps.

Question 9

I have a question.

Let's say you had to find the electron configuration of Neon

Well, that's pretty easy: [He]2s^2  2p6

However, what if we had to find the electron configuration of Neon with an extra electron? (Negatively charged ion).

Where would that negatively charged ion go? it can't fit into the p sub shell since the p sub shell can only hold 6 electrons.

Would it become [He] 2s^2   2p^6   3^s1 ?

Accepted Answer

Yup, you are correct. To approach it from another angle, you could think of adding an electron as changing the electron configuration to the next element on the periodic table. So adding an electron to Neon would net the same electron configuration as Sodium (Na).

Course: AP®︎/College Chemistry > Unit 1

Electron configurations of ions

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