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chemistry question

can someone please explain to me how can Phosphorus and other elements have different number of bonding regions if the elements is still the same and no ionic charged is included this might be the only thing in chemistry I don't seem to understand like mathematically speaking its impossible I understand that in the third shell the maximum number of electrons is 18 but how can phosphorus for example form PCL3 AND PCL5 like how am I supposed to know how many bonding regions will there be if we are told that we should have 8 electrons max in an atom is my hypothesis of it being about the electron shell rule as in the third shell can have 18 so it can have as many bonding regions as it wants
(edited 6 months ago)
Original post by cr7090121
can someone please explain to me how can Phosphorus and other elements have different number of bonding regions if the elements is still the same and no ionic charged is included this might be the only thing in chemistry I don't seem to understand like mathematically speaking its impossible I understand that in the third shell the maximum number of electrons is 18 but how can phosphorus for example form PCL3 AND PCL5 like how am I supposed to know how many bonding regions will there be if we are told that we should have 8 electrons max in an atom is my hypothesis of it being about the electron shell rule as in the third shell can have 18 so it can have as many bonding regions as it wants

You will have been taught the octet rule at GCSE, but you need to bear in mind that it is really only a general rule of thumb that works most of the time for the elements in the s- and p- blocks of the periodic table (and the rule tends to fail completely when looking at the d- and f- block elements).

There is a complication that isn’t explained at A level, which explains why the elements of groups 5 - 7 on periods 3 and below can form compounds with unusually large numbers of bonds surrounding them, whereas the elements on periods 1 and 2 cannot. I won’t go into it in any depth as I would prefer you stuck strictly to the A level syllabus as it would lose you marks if you decided to bring it up in an exam.

For now, using the fact that the third quantum shell can take up to 18 electrons is perhaps the best way of thinking about it (though this is not the full story). At A level, the most number of electron pairs you will see in the outermost shell of an atom of an element in the p- block is 6 (so not quite forming as many bonding regions as it wants).

Generally you can deduce from the name or formula of the compound how many bonding regions there are (e.g sulphur hexafluoride or SF6 ==> 6 bonding regions).
Reply 2
Original post by cr7090121
can someone please explain to me how can Phosphorus and other elements have different number of bonding regions if the elements is still the same and no ionic charged is included this might be the only thing in chemistry I don't seem to understand like mathematically speaking its impossible I understand that in the third shell the maximum number of electrons is 18 but how can phosphorus for example form PCL3 AND PCL5 like how am I supposed to know how many bonding regions will there be if we are told that we should have 8 electrons max in an atom is my hypothesis of it being about the electron shell rule as in the third shell can have 18 so it can have as many bonding regions as it wants

It's beyond A level, but essentially it's because the d subshell (3d in P's case) becomes available, so some of the 3d orbitals can be used in bonding as well as the 3p orbitals which are used as you'd expect.
https://chem.libretexts.org/Bookshelves/General_Chemistry/ChemPRIME_(Moore_et_al.)/07%3A_Further_Aspects_of_Covalent_Bonding/7.02%3A_Exceptions_to_the_Octet_Rule


From a Lewis structure/electron pairing point of view, you could think of it looking a bit like this (it's a bit more complicated but kind of helps with the numbers):
P.png
where the blue dots are the 5 "outer" electrons from phosphorus, and the red ones are from the other atom (like chlorine) either 3 or 5, forming covalent bonds
Reply 3
Original post by TypicalNerd
You will have been taught the octet rule at GCSE, but you need to bear in mind that it is really only a general rule of thumb that works most of the time for the elements in the s- and p- blocks of the periodic table (and the rule tends to fail completely when looking at the d- and f- block elements).

There is a complication that isn’t explained at A level, which explains why the elements of groups 5 - 7 on periods 3 and below can form compounds with unusually large numbers of bonds surrounding them, whereas the elements on periods 1 and 2 cannot. I won’t go into it in any depth as I would prefer you stuck strictly to the A level syllabus as it would lose you marks if you decided to bring it up in an exam.

For now, using the fact that the third quantum shell can take up to 18 electrons is perhaps the best way of thinking about it (though this is not the full story). At A level, the most number of electron pairs you will see in the outermost shell of an atom of an element in the p- block is 6 (so not quite forming as many bonding regions as it wants).

Generally you can deduce from the name or formula of the compound how many bonding regions there are (e.g sulphur hexafluoride or SF6 ==> 6 bonding regions).


thank you for your reply I honestly revise stuff above my a level syllabus a lot like for example I know how to proof hat methane has a bond angle of 109.5 using yxz graph and more stuff ,I don't think they would mark me down but yeah .from your last explanation do you mean as in 6 pairs of electrons should be the maximum output so SF6 with sulphur providing 6 electrons and F providing one each no lone pairs ( which is quite strange as S has 5 electrons in its outer shell meaning the extra one its either taken from the second shell which is means there's going to be an odd number of electrons in the second shell so that theory is wrong my second one would be it takes an electron from F in someway maybe a dative covalent bond?)
Reply 4
Original post by bl0bf1sh
It's beyond A level, but essentially it's because the d subshell (3d in P's case) becomes available, so some of the 3d orbitals can be used in bonding as well as the 3p orbitals which are used as you'd expect.
https://chem.libretexts.org/Bookshelves/General_Chemistry/ChemPRIME_(Moore_et_al.)/07%3A_Further_Aspects_of_Covalent_Bonding/7.02%3A_Exceptions_to_the_Octet_Rule


From a Lewis structure/electron pairing point of view, you could think of it looking a bit like this (it's a bit more complicated but kind of helps with the numbers):
P.png
where the blue dots are the 5 "outer" electrons from phosphorus, and the red ones are from the other atom (like chlorine) either 3 or 5, forming covalent bonds


this was literally my hypothesis I tried to prove it wrong couple times but yeah it sounds and looks correct
thank you
Original post by cr7090121
thank you for your reply I honestly revise stuff above my a level syllabus a lot like for example I know how to proof hat methane has a bond angle of 109.5 using yxz graph and more stuff ,I don't think they would mark me down but yeah .from your last explanation do you mean as in 6 pairs of electrons should be the maximum output so SF6 with sulphur providing 6 electrons and F providing one each no lone pairs ( which is quite strange as S has 5 electrons in its outer shell meaning the extra one its either taken from the second shell which is means there's going to be an odd number of electrons in the second shell so that theory is wrong my second one would be it takes an electron from F in someway maybe a dative covalent bond?)

Fair enough. The concept in question I had in mind was orbital hybridisation (and even then, that isn’t considered to be the actual cause of “hypervalency” - hybridisation is just a useful model that works well for the prediction of the shape of the molecule), just don’t mention it in an exam as you will lose marks for doing so.

At A level, the absolute most number of electron pairs (including both lone and bonding pairs) in an atom’s outermost shell will be 6. Just bear in mind that not all elements are capable of forming compounds in which they have this number of electron pairs in their outer shell - sulphur is one of few examples.

Sulphur actually has 6 outer electrons rather than 5 as you have said (since it is in group 6/16), so it would be forming 6 bonding pairs in SF6 (in which each bonding pair has 1 electron coming from the sulphur and one coming from the fluorine) and there would be no lone pairs in the outermost shell in the sulphur atom
(edited 6 months ago)
Reply 6
Original post by TypicalNerd
Fair enough. The concept in question I had in mind was orbital hybridisation (and even then, that isn’t considered to be the actual cause of “hypervalency” - hybridisation is just a useful model that works well for the prediction of the shape of the molecule), just don’t mention it in an exam as you will lose marks for doing so.

At A level, the absolute most number of electron pairs (including both lone and bonding pairs) in an atom’s outermost shell will be 6. Just bear in mind that not all elements are capable of forming compounds in which they have this number of electron pairs in their outer shell - sulphur is one of few examples.

Sulphur actually has 6 outer electrons rather than 5 as you have said (since it is in group 6/16), so it would be forming 6 bonding pairs in SF6 (in which each bonding pair has 1 electron coming from the sulphur and one coming from the fluorine) and there would be no lone pairs in the outermost shell in the sulphur atom

yeah I totally understand it and my bad I meant to say phosphorus instead of S as P was my example , thank you so much for your explanation if you don't mind me asking what course are you doing ?
and also you see for giant covalent structures am I good to go by just assuming that only Si and SiO2 can have a giant covalent lattice ( as if run through couple of questions 3 pdfs full of questions and that was my conclusion after doing all the questions )and also that there are no other intermolecular forces I should be talking about when speaking with giant covalent lattice
thanks
Original post by cr7090121
yeah I totally understand it and my bad I meant to say phosphorus instead of S as P was my example , thank you so much for your explanation if you don't mind me asking what course are you doing ?
and also you see for giant covalent structures am I good to go by just assuming that only Si and SiO2 can have a giant covalent lattice ( as if run through couple of questions 3 pdfs full of questions and that was my conclusion after doing all the questions )and also that there are no other intermolecular forces I should be talking about when speaking with giant covalent lattice
thanks

At present, I’m not actually doing any course yet as I still have roughly a week until I start at uni (where I will be doing an accelerated master’s in chemistry), though if you mean which A level course did I do, I did Edexcel A level chemistry.

There are more substances with giant covalent structures: carbon (specifically diamond) and germanium are examples. When discussing their melting points, the strength of the covalent bonds is the main thing you want to use to justify the high melting points, as a lot of energy will be needed to break all of them upon melting. You could also mention that there are many covalent bonds, too, but that detail was more important on the old A level spec. There is no need to discuss any other intermolecular forces as they are irrelevant.
Reply 8
Original post by TypicalNerd
At present, I’m not actually doing any course yet as I still have roughly a week until I start at uni (where I will be doing an accelerated master’s in chemistry), though if you mean which A level course did I do, I did Edexcel A level chemistry.

There are more substances with giant covalent structures: carbon (specifically diamond) and germanium are examples. When discussing their melting points, the strength of the covalent bonds is the main thing you want to use to justify the high melting points, as a lot of energy will be needed to break all of them upon melting. You could also mention that there are many covalent bonds, too, but that detail was more important on the old A level spec. There is no need to discuss any other intermolecular forces as they are irrelevant.


so how do you know about hybridisation and other stuff unless you teach yourself which is what I am doing but your level of chemistry understanding is fascinating you need to teach me your ways
also I Understand that but my main question was is giant covalent lattice only formed in SiO2 and Is and diamond and graphite or other compounds too because for example SiCl4 is a tetra , non polar same thing as SiO2 non polar , so why can SiO2 form it and SiCl4 not
(edited 6 months ago)
Original post by cr7090121
so how do you know about hybridisation and other stuff unless you teach yourself which is what I am doing but your level of chemistry understanding is fascinating you need to teach me your ways
also I Understand that but my main question was is giant covalent lattice only formed in SiO2 and Is and diamond and graphite or other compounds too because for example SiCl4 is a tetra , non polar same thing as SiO2 non polar , so why can SiO2 form it and SiCl4 not

I absolutely taught myself stuff beyond the A level spec back in the day, but in my experience, others who did the same had more difficulty than I did when distinguishing between the A level material and the non-A level material, thus lost marks (hence why I am being very cautious how in-depth my explanations are).

Ok, for A level purposes, assume that only diamond, graphite, silicon and silicon dioxide are giant covalent.

SiCl4 cannot form a giant covalent structure, because the chlorine atoms don’t have any more vacant bonding sites, so it is only able to exist as simple molecules. With SiO2, each oxygen is able to form two covalent bonds, so each oxygen atom can ‘bridge’ between two silicon atoms and so you get a network of silicon and oxygen atoms that assumes a giant covalent structure.
Reply 10
Original post by TypicalNerd
I absolutely taught myself stuff beyond the A level spec back in the day, but in my experience, others who did the same had more difficulty than I did when distinguishing between the A level material and the non-A level material, thus lost marks (hence why I am being very cautious how in-depth my explanations are).

Ok, for A level purposes, assume that only diamond, graphite, silicon and silicon dioxide are giant covalent.

SiCl4 cannot form a giant covalent structure, because the chlorine atoms don’t have any more vacant bonding sites, so it is only able to exist as simple molecules. With SiO2, each oxygen is able to form two covalent bonds, so each oxygen atom can ‘bridge’ between two silicon atoms and so you get a network of silicon and oxygen atoms that assumes a giant covalent structure

yeah your right I might just do that with biology and with chemistry I will try my best to distinguish between a level an non a level so maybe I will try to chill with the non a level content
by the way did u have to attend an interview for UNI
Original post by cr7090121
yeah your right I might just do that with biology and with chemistry I will try my best to distinguish between a level an non a level so maybe I will try to chill with the non a level content
by the way did u have to attend an interview for UNI

For Oxford and Imperial, yes. They interviewed me. On my first attempt, I had 4 interviews with Oxford and on my second, I had 3 interviews with them. Imperial only required 1 interview of me in my one and only attempt.

For Durham, UCL and St Andrew’s, no. They just gave me offers without interviews.
Reply 12
Original post by TypicalNerd
For Oxford and Imperial, yes. They interviewed me. On my first attempt, I had 4 interviews with Oxford and on my second, I had 3 interviews with them. Imperial only required 1 interview of me in my one and only attempt.

For Durham, UCL and St Andrew’s, no. They just gave me offers without interviews.


idk why I cant get past the interview stage with UCL and the top unis they ask me questions like "tell me about a time when I demonstrated leadership.. and one time when you had to put ur work aside and help someone else by affecting your own work flow " these types of questions kinda useless because the only way to prove such abilities is by actually witnessing those abilities like you cant really talk about a time when you helped someone out or talk about how much of a good leader you are that's kinda arrogant and just sounds bad , what would you advise me to do ?
Original post by cr7090121
idk why I cant get past the interview stage with UCL and the top unis they ask me questions like "tell me about a time when I demonstrated leadership.. and one time when you had to put ur work aside and help someone else by affecting your own work flow " these types of questions kinda useless because the only way to prove such abilities is by actually witnessing those abilities like you cant really talk about a time when you helped someone out or talk about how much of a good leader you are that's kinda arrogant and just sounds bad , what would you advise me to do ?

Some unis honestly don’t bother to interview for some courses - chemistry seems to be one that only Imperial and Oxford bother to interview you for.

In my experience, I was expecting to be asked those questions in my interviews but never was. Nonetheless, beforehand I had guessed possible questions they could ask about times I had shown leadership etc and between then and doing my interviews, I found opportunities to demonstrate such qualities and wrote down a record of exactly what I did so I could remember the important things to bring up in an interview if it came up.

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