Chem EK question 22

[HereGoes]

How does Sulfur have 3d orbitals available when it's configuration is: SULFUR 1S2 2S2 2P6 3S2 3P4 (NO D orbitals!!)

Thanks!
Confused

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[RogueUnicorn]

3d fills after 4s, but doesn't mean it doesn't exist..

 

[HereGoes]

I understand that 3d fills post 4s, but my confusion is to the following question:

22)Which of the following best explains why sulfur can make more bonds than Oxygen?
a)Sulfur is more electronegative than oxygen
b)oxygen is more electronegative than sulfur
c)sulfur has 3d orbitals not available to oxygen
d)sulfur has fewer valence electrons

The answer in EK is C, which confuses me..I thought sulfur was
[Ne]3s(^2)3p(^4), thus no 4s orbital, and no 3d! I am confused as to how Sulfur has 3d bonding available to it, considering it's location on the periodic table.

Thanks in advance for your help!

 

[SoftwareKevin]

I understand that 3d fills post 4s, but my confusion is to the following question:

22)Which of the following best explains why sulfur can make more bonds than Oxygen?
a)Sulfur is more electronegative than oxygen
b)oxygen is more electronegative than sulfur
c)sulfur has 3d orbitals not available to oxygen
d)sulfur has fewer valence electrons

The answer in EK is C, which confuses me..I thought sulfur was
[Ne]3s(^2)3p(^4), thus no 4s orbital, and no 3d! I am confused as to how Sulfur has 3d bonding available to it, considering it's location on the periodic table.

Thanks in advance for your help!

Someone correct me if I'm wrong, but the fact that sulfur has anything in the 3's means that a 3d orbital *does* exist, it just hasn't been filled yet because it's at a higher energy and doesn't get filled until after 4s.

Think back to quantum numbers. the principal quantum number n=3 so l/subshell can be 0 (s), 1 (p), or 2 (d). They all exist even if there's nothing in them... for Oxygen n=2 so l can only be 0 (s) or 1 (p)

Electron configuration does not necessarily tell you what orbitals exist, just which orbitals have been filled.

 

[HereGoes]

Thank you Kevin, this is very helpful.

To take the question one step further, is there a number of "extra" bonds that Sulfur would need to make before it can access its d orbital? (As in, would it's 4s orbital need to be filled when it bonds with another element - ie in order for it to make any bonds with its 3d orbital, it would need to have X number of bonds formed?)

Sorry if this is completely stupid...

 

[Danlee07]

D shell is there to hybridize. 4s2 comes before 3d. Yeah, the orbitals have to be filled before onto the next one. It's one of the rules.

It can hybridize after s is full, and since it's p4, it can hybridize.

For example: SO2. Sulfur is bonded to two oxygens; so the sulfur's hybridization would become a s1 p1 and p up to 3. So you get sp2 (which matches since it's double bonds as well).

 

[catzzz88]

Hi all,

I have this exact same question and was happy to find something when I searched. However, I am still not completely clear on this topic after having read the above explanations.

Any additional examples or explanations would be helpful.

Examples of molecules that include the 3d orbital filling of S, why 3d was necessary to fill and how it was accessed by S, and how many up/down electrons there are in the example would probably completely answer my question.

Thanks in advance...

Best,
C

 

[Gedtomd]

Someone correct me if I'm wrong, but the fact that sulfur has anything in the 3's means that a 3d orbital *does* exist, it just hasn't been filled yet because it's at a higher energy and doesn't get filled until after 4s.

Think back to quantum numbers. the principal quantum number n=3 so l/subshell can be 0 (s), 1 (p), or 2 (d). They all exist even if there's nothing in them... for Oxygen n=2 so l can only be 0 (s) or 1 (p)

Electron configuration does not necessarily tell you what orbitals exist, just which orbitals have been filled.

Thank you Software Kevin for answering this question. Six years later your answer is still helpful!

 

[aldol16]

Not to continue reviving this, but the "existence" of the 3d isn't exactly correct here. 3d is a designation that is given to an electron density map that corresponds to that orbital. In other words, you can think of orbitals as the electrons themselves because all "orbitals" are are electron density maps constructed to contain 90% of the electron density for a particular "orbital" or set of quantum numbers. So the 3d does not "exist" unless there are actually 3d electrons.

But the point here that most missed is that you can only use "orbitals" that are within your highest principal quantum number for bonding. That is, oxygen's highest n is 2 and so it cannot access d orbitals because only n = 3 has them. This is because of energy considerations - using a 3d orbital for oxygen would be so energetically unfavorable for the oxygen that it won't want to do it.