Why does d subshell start at 3d and not 4d?

Hey all,

I’m going onto do A-Level chemistry next year and have started to look over some content.

One thing I can’t seem to understand- Why does the d subshell start at 3 although it is placed in the 4th period? To demonstrate what I mean, I’ve attached an image of the periodic table labelled with subshells.

Also, would the electrons in the 3d subshell of an element such as krypton be in the 3rd shell as the subshell is labelled 3d or the 4th shell as the subshell is in the 4th period and why?




Thanks a lot!

I seem to misunderstand the question. When you say why the d subshell start at 3, do you mean why it is in the 3rd orbital?

For the second question, it depends which subshell they are in. If they are in the 3d subshell, then they will be in the 3rd orbital and if the electrons are in the 4d subshell, then they are in the 4th orbital. Sorry I am not at home but if you explain your question more, I will try and explain to you how subshells work once I get hold of my laptop.

If you’re asking why 4d fills before 3s it’s because it has a lower energy level and therefore in ions 4s would be removed or added before 3d if that makes sense, however when you write the configuration you write it in ascending shell order

How far have you got with content?

If you’ve self taught that’s amazing

Hey, thanks for your reply.

Basically I don’t get why the 3d orbital supposedly comes after the 4s orbital when writing electronic configuration. For example, the electromic configuration of krypton would be:

1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 .

Why isn’t the 3d^10 orbital written before 4s^2. As it’s written after it, does this mean that the 10 electrons in the 3d^10 orbital or infact in the fourth shell from the nucleus?

Thanks a lot!

Okay that makes a lot more sense now. What you need to understand is that when writing, you can write the 4s before 3d or after 3d. It doesn't make a different (in the OCR spec it doesn't at least). However, it is good to write them in sequence so usually 4s is written after 3d. However, what you need to realise first is that GCSE basically lied to you so forget whatever you learnt. You need to start realizing that when the subshells fill, they fill in terms of their energy level. The lower energy levels fill first then the higher ones in order. This is something which is covered under the "aufbau principle". However, when electrons move energy levels or are shared, the electrons from 4s are removed before 3d electrons since 4s is a lower energy level and thus less energy is needed. This is probably the main reason why 4s is written before 3d. Therefore, this has nothing to do with where they are in shells or periods.

Also I would recommend you to not try and learn this on your own. Developing the wrong understanding isn't good for you and it can confuse you more. Just wait for your teacher to go through this and let them explain. However, if you want some reading to do and think you understood what I said, have a look at the following links. They do a great job of explaining what happens and why it does.

https://www.chemguide.co.uk/atoms/properties/3d4sproblem.html
https://chem.libretexts.org/Textbook_Maps/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Electronic_Structure_of_Atoms_and_Molecules/Electronic_Configurations/The_Order_of_Filling_3d_and_4s_Orbitals

If you need something else to be explained, feel free to ask but I think the above links cover more than what is required for A-levels so you should be good

I got a 9 too but am pretty stressed for the a level in case I do really bad or the increase in difficulty is astronomical

Wow! Thanks that’s really useful. You explained that really well. Thanks for taking the time to go through that with me in such detail! I’ll definitely take your advice on board - I have one more question if that is okay as what you just said was so useful. At GCSE we were taught that 2 non metals bond covalently whilst a metal and a non metal bond ionically. I understand from what I’ve been learning that atoms bond on a spectrum from ionically to covalently, depending if they have similar electronegativities or not. My book said that atoms with medium filled shells bond covalently whilst atoms needing to only gain one or two electrons bond together ionically. Is this correct? So two atoms with 6 and 2 or 1 and 7 in their outer shells will be considered to be bonded ionically whilst all other atoms bonding will be considered to bond covalently. Thanks a lot!

EDIT - Just read through those links and they’re great - I tried a google search for it but it said it was due to complex mechanisms of quantum physics whereas those sites both explain it really well so thanks!

Yeah don't google stuff that much. Stick to your book or at max A-level websites.

In terms of bonding, the general rule is like you mentioned it depends on electronegativity of each element. However, that generally what you learnt at GCSE about this is true. Mostly, non-metals and metals form ionic bonds while two non-metals form covalent bonds. You can figure this out when you look at the "pauling electronegativity values". However, these values can be used to estimate but aren't definite so it might be that even if the difference is large there is covalent bonding.

Also, just read you thinking biology seems alarming. Just remember to keep revising the stuff throughout the year. If you don't you will forget stuff and learning that stuff again can be pain especially in year 13 when you are under immense pressure from university applications, interviews, mocks etc.