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# OCR Chemistry A - Unit 1 watch

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1. Just a quick one before the exam tomorrow...

How do you work out how many 'electron pairs' a molecule has in order to work out the shape of the molecule?

Thanks in advance and Good luck to everyone who is also doing the exam.
Just a quick one before the exam tomorrow...

How do you work out how many 'electron pairs' a molecule has in order to work out the shape of the molecule?

Thanks in advance and Good luck to everyone who is also doing the exam.
You draw a dot and cross diagram of the molecule.
And then count how many electron pairs are bonded to the central atom.
e.g. in water - H2O

The hydrogen atoms are bonded to the central Oxygen atom forming two bonded pairs of electrons. Water also has two LONE pairs of electrons that arent bonded to anything.
lone pairs repel more than bonded pairs.
A molecule with 2 lone pairs of electrons has a non linear shape with a bond angle of 104.5
Just a quick one before the exam tomorrow...

How do you work out how many 'electron pairs' a molecule has in order to work out the shape of the molecule?

Thanks in advance and Good luck to everyone who is also doing the exam.
Check how many electrons each atom has in their outer shell. Draw the structure to check how many electrons are used to form bonds (if it's regular covalent bond, then 1 electron, but if it's dative covalent bond then it can be 2 (if it donates) or 0 (if it doesn't donate).
Number of electron pairs = (no. of electrons in outer shell - no. of electrons used to form bonds)/2

Hope this helps
4. Just another kinda related question, why do the lone pairs go together, why dont the bonded pairs go in-between the lone pairs to cancel out? Is there a reason and if asked should i just choose whatever is the most awkward. ie. lone pairs together?
5. (Original post by jontylol)
Just another kinda related question, why do the lone pairs go together, why dont the bonded pairs go in-between the lone pairs to cancel out? Is there a reason and if asked should i just choose whatever is the most awkward. ie. lone pairs together?
Both lone pairs and bonded pairs are doubly negatively charged and repel one another. They cannot 'cancel out'.

It's just that LP-LP repulsion is stronger than LP-BP which is stronger than BP-BP repulsion.
6. (Original post by jontylol)
Just another kinda related question, why do the lone pairs go together, why dont the bonded pairs go in-between the lone pairs to cancel out? Is there a reason and if asked should i just choose whatever is the most awkward. ie. lone pairs together?
In the lone pair, 2 electrons always go together as a pair, the bond pair can't go between them. If you try to draw the boxes and put arrows (which represent electrons), may be you can see it. However, when the reaction is going to occur, one of the two arrows will move to another box, and those 2 electrons are going to form bonds with other electrons from other atoms. But if there is no reaction then those 2 electrons will stay together as a pair.

7. (Original post by Zuzia Bulu)
In the lone pair, 2 electrons always go together as a pair, the bond pair can't go between them. If you try to draw the boxes and put arrows (which represent electrons), may be you can see it. However, when the reaction is going to occur, one of the two arrows will move to another box, and those 2 electrons are going to form bonds with other electrons from other atoms. But if there is no reaction then those 2 electrons will stay together as a pair.

Ahh, i just realised my question was badly worded. I was meaning when you have 2 lone pairs, such as in H2O. I know that a bonded pair cant just hop between a lone pair :P Sorry for the confusion xD
8. (Original post by jontylol)
Ahh, i just realised my question was badly worded. I was meaning when you have 2 lone pairs, such as in H2O. I know that a bonded pair cant just hop between a lone pair :P Sorry for the confusion xD
Oh, I guess I have to practise English more (it's my second language). Sorry. And I've just notice charco's answer, I agree with him/her
9. http://www.thestudentroom.co.uk/show...3#post25301883

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