Giant Covalent and Simple Molecular watch

QuidditchFan · 06-01-2014 21:04

I'm confused by bonding in general - particularly intermolecular bonding, because it's just not 'clicking' for me, so could anyone please tell me how to recognise which bonding is in which molecule (van der waal, permanent dipole or hydrogen bonds) - rules would help e.g. hydrogen bonding is only between N-H or O-H.

My main question is - how do you recognise if a compound is simple molecular or giant covalent? I understand that alkanes etc are simple covalent and diamond and graphite are giant covalent. However, I thought SiO2 was simple molecular but it's not. And in a recent exam Q I looked at, it said ICl was simple covalent but I thought it was giant covalent. I've already looked on the internet, textbooks, asked teachers and friends and I still don't understand it! Can you please help me!

BarneyMersich · 06-01-2014 23:22

Okay, so my spec doesn't cover van der waals, (or we just haven't done it yet) but any small molecule that's covalently bonded will me simple molecular. Iodine chloride can't be a network as both I and Cl can only make one covalent bond to get a full outer shell. (Grp7) Carbon and silicon can form 4 bonds, so they are able to make giant networks. Silicon oxide is a network as silicon can bond to 4 oxygens, and each oxygen can bond to 2 silicon atoms. Diamond and graphite are made of only carbon, so they can easily form a network, the difference is the way the carbons are arranged. I have attached some photos, I hope that helps.

As for id id, pd-pd and hydrogen bonds:
Pretty much every molecule can form id-is bonds. The electrons are constantly moving around the molecule, so at a specific instant, there will be more electrons on one end of the molecule than the other. This creates a dipole, but only for that instant.

Pd-pd bonds are between molecules with a dipole charge, eg HCl. Chlorine is a lot more electronegative than hydrogen, so it attracts the shared electrons more.

Here's how we were taught hydrogen bonding:
You need 2 things:
1) a Hydrogen atom attached to a FON (flourine/oxygen/nitrogen) on one molecule.
2) a lone pair on the FON in the other molecule.
So for example in water, there is a hydrogen attached to an oxygen (1) and it hydrogen bonds to the lone pair on the FON (in this case oxygen) on another water molecule. So H2O can hydrogen bond.

I hope this has helped you, if you need anymore help just let me know

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Trouty97 · 06-01-2014 23:28

Van der waal's forces occur between non-polar,simple covalent molecules. They are caused by the uneven distribution of electrons, which induces an instantaneous dipole in the molecule. This then induced dipoles in neighbouring molecules. The electrostatic attraction between these dipoles is the van der waal's force but they are extremely weak compared to all other types of bonding, so not much energy is needed to overcome them.

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BarneyMersich · 06-01-2014 23:33

(Original post by Trouty97)
Van der waal's forces occur between non-polar,simple covalent molecules. They are caused by the uneven distribution of electrons, which induces an instantaneous dipole in the molecule. This then induced dipoles in neighbouring molecules. The electrostatic attraction between these dipoles is the van der waal's force but they are extremely weak compared to all other types of bonding, so not much energy is needed to overcome them.

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Yeah, we were taught these as instantaneous dipole - induced dipole bonds. No fancy names

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charco · 07-01-2014 09:30

(Original post by Trouty97)
Van der waal's forces occur between non-polar,simple covalent molecules. They are caused by the uneven distribution of electrons, which induces an instantaneous dipole in the molecule. This then induced dipoles in neighbouring molecules. The electrostatic attraction between these dipoles is the van der waal's force but they are extremely weak compared to all other types of bonding, so not much energy is needed to overcome them.

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They are only extremely weak in very small molecules. In larger molecules, such as polymers they are strong. In sulphur, phosphorus and iodine they are strong enough to make the structures solid at room temperature. This makes them stronger than the hydrogen bonding in water!

BarneyMersich · 07-01-2014 10:14

(Original post by charco)
They are only extremely weak in very small molecules. In larger molecules, such as polymers they are strong. In sulphur, phosphorus and iodine they are strong enough to make the structures solid at room temperature. This makes them stronger than the hydrogen bonding in water!

That's cool, I never really thought of that! It's because of the huge amount of electrons they have.

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QuidditchFan · 07-01-2014 16:37

(Original post by BarneyMersich)
Okay, so my spec doesn't cover van der waals, (or we just haven't done it yet) but any small molecule that's covalently bonded will me simple molecular. Iodine chloride can't be a network as both I and Cl can only make one covalent bond to get a full outer shell. (Grp7) Carbon and silicon can form 4 bonds, so they are able to make giant networks. Silicon oxide is a network as silicon can bond to 4 oxygens, and each oxygen can bond to 2 silicon atoms. Diamond and graphite are made of only carbon, so they can easily form a network, the difference is the way the carbons are arranged. I have attached some photos, I hope that helps.

As for id id, pd-pd and hydrogen bonds:
Pretty much every molecule can form id-is bonds. The electrons are constantly moving around the molecule, so at a specific instant, there will be more electrons on one end of the molecule than the other. This creates a dipole, but only for that instant.

Pd-pd bonds are between molecules with a dipole charge, eg HCl. Chlorine is a lot more electronegative than hydrogen, so it attracts the shared electrons more.

Here's how we were taught hydrogen bonding:
You need 2 things:
1) a Hydrogen atom attached to a FON (flourine/oxygen/nitrogen) on one molecule.
2) a lone pair on the FON in the other molecule.
So for example in water, there is a hydrogen attached to an oxygen (1) and it hydrogen bonds to the lone pair on the FON (in this case oxygen) on another water molecule. So H2O can hydrogen bond.

I hope this has helped you, if you need anymore help just let me know

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(Original post by Trouty97)
Van der waal's forces occur between non-polar,simple covalent molecules. They are caused by the uneven distribution of electrons, which induces an instantaneous dipole in the molecule. This then induced dipoles in neighbouring molecules. The electrostatic attraction between these dipoles is the van der waal's force but they are extremely weak compared to all other types of bonding, so not much energy is needed to overcome them.

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(Original post by charco)
They are only extremely weak in very small molecules. In larger molecules, such as polymers they are strong. In sulphur, phosphorus and iodine they are strong enough to make the structures solid at room temperature. This makes them stronger than the hydrogen bonding in water!

You've all helped me, particularly BarneyMersich, so thank you! Just to clarify:

1. Giant covalent structures are normally group 4 and are dependent on the number of bonds.
2. Simple covalent structures are normally group 7/6/5 and again, depends on the number of bonds. There are Van der Waal bonds between molecules and in larger molecules, the higher number of electrons makes stronger VDW bonds which can make structures solid at RTP and can be stronger than H-bonding in water. So is this VDW the bonding in Iodine crystals etc?
3. Dipole - dipole bonds are caused when there is a dipole in one molecule. A dipole in one molecule is because of difference in electronegativities of atoms. Another question - how do you know whether the atoms have a big enough difference in electronegativity to make a molecule polar? I can't really learn electronegativities of all the atoms and if you could give me the most common atoms that come up together with their electronegativities, or how you recognise if there is a dipole in a molecule, that would be great! And does this bare any similarity to ionic bonding. How can I stop getting them confused i.e. is Dipole-dipole bonding only between non-metals, including hydrogen?
4. Hydrogen bonding is between two molecules, one with H-O, H-N, H-F only and the other with the FON - H and a lone pair.

Can any of you please answer the above questions? Thank you again for the help everyone. I think I'm beginning to get it.

BarneyMersich · 07-01-2014 18:07

(Original post by QuidditchFan)
You've all helped me, particularly BarneyMersich, so thank you! Just to clarify:

1. Giant covalent structures are normally group 4 and are dependent on the number of bonds.
2. Simple covalent structures are normally group 7/6/5 and again, depends on the number of bonds. There are Van der Waal bonds between molecules and in larger molecules, the higher number of electrons makes stronger VDW bonds which can make structures solid at RTP and can be stronger than H-bonding in water. So is this VDW the bonding in Iodine crystals etc?
3. Dipole - dipole bonds are caused when there is a dipole in one molecule. A dipole in one molecule is because of difference in electronegativities of atoms. Another question - how do you know whether the atoms have a big enough difference in electronegativity to make a molecule polar? I can't really learn electronegativities of all the atoms and if you could give me the most common atoms that come up together with their electronegativities, or how you recognise if there is a dipole in a molecule, that would be great! And does this bare any similarity to ionic bonding. How can I stop getting them confused i.e. is Dipole-dipole bonding only between non-metals, including hydrogen?
4. Hydrogen bonding is between two molecules, one with H-O, H-N, H-F only and the other with the FON - H and a lone pair.

Can any of you please answer the above questions? Thank you again for the help everyone. I think I'm beginning to get it.

So:
1) yeah that sounds about right, as long as you remember diamond, graphite and silicon dioxide, you should be fine, at least that's on out spec (OCR salters B)

2) simple molecular doesn't really depend on number of bonds: as long as it isn't a network but it's covalent it will me simple molecular. As long as you can look at something like water, and realise that all the bonds are used up, so it can't be a network you should be okay.
Also,
Careful with wording:

"in larger molecules, the higher number of electrons makes stronger VDW bonds"

I know what you meant, but the mark schemes (especially on my exam board / exam) are very strict on wording. In larger molecules the higher number of electrons causes the molecule to make stronger VDW bonds with adjacent molecules. the way you wrote it suggested the electrons are doing the bonding. Rememer, these are intermolecular, so thats bonds between molecules, as opposed to intramolecular, which are bonds within a molecule (covalent etc).

3) dipole dipoles are when there's a dipole in both molecules.
As for electronegativity:
The general rule of thumb is the top right side of the periodic table is the most electronegative, the bottom left is the least. Things like H-Cl are polar, as Cl is very electronegative and H is not. (Joke: how are bonds like bears? Some are polar, some are not)
This is nothing like ionic bonding. These are due to bonds in molecules having a dipole. A dipole is only a small charge, usually written as δ- and δ+ (that's a delta) this is caused by the shared electrons being closer to one atom than the other (the more electronegative one)
ionic bonding is where there is a charge due to loss or gain of electrons. In an atom, and there are no shared electrons. In fact, the two atoms/molecules ionically bonded are only 'as one' when in solid crystal form. When dissolved in water they split into the ions.

4) yes that's right, but the second one doesn't necessarily need a Hydrogen I don't think: a hydrogen attached to a FON, can hydrogen bond to the lone pairs on another FON.

That's quite an essay, I may have missed something. Anything you're still unsure about just post!

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Hospes · 07-01-2014 18:16

(Original post by QuidditchFan)
You've all helped me, particularly BarneyMersich, so thank you! Just to clarify:

1. Giant covalent structures are normally group 4 and are dependent on the number of bonds.
2. Simple covalent structures are normally group 7/6/5 and again, depends on the number of bonds. There are Van der Waal bonds between molecules and in larger molecules, the higher number of electrons makes stronger VDW bonds which can make structures solid at RTP and can be stronger than H-bonding in water. So is this VDW the bonding in Iodine crystals etc?
3. Dipole - dipole bonds are caused when there is a dipole in one molecule. A dipole in one molecule is because of difference in electronegativities of atoms. Another question - how do you know whether the atoms have a big enough difference in electronegativity to make a molecule polar? I can't really learn electronegativities of all the atoms and if you could give me the most common atoms that come up together with their electronegativities, or how you recognise if there is a dipole in a molecule, that would be great! And does this bare any similarity to ionic bonding. How can I stop getting them confused i.e. is Dipole-dipole bonding only between non-metals, including hydrogen?
4. Hydrogen bonding is between two molecules, one with H-O, H-N, H-F only and the other with the FON - H and a lone pair.

Can any of you please answer the above questions? Thank you again for the help everyone. I think I'm beginning to get it.

1. Yes. But polymers are also giant covalent structures made up of simple monomer repeating units. If you still haven't covered polymers then don't worry. It would also be better to go through the specs of your relevant exam board. Mostly they'll list all the giant covalent structures that students would be assumed to know at the exam.
2. Yes and yes. Also remember that vanderwall is a general term that covers both instantaneous dipole - induced dipole intermolecular bonds (also called London forces or dispersion forces) and permanent dipole permanent dipole bonds so in the exam it's better to specify what of the two is present. My teacher advises us to use the VDW term only if both are present. Yes instantaneous dipole induce dipole bonds are present in iodine and not pds. Hope you understand why.
3. Leave aside metals and metallic ions in organic chemistry. With a bit of experience it's easy to identify a significant polar bond. Remember that alkane and alkene chains do not have polar bonds on their own (significant polar bonds ).
A molecule is polar if it fulfils two conditions. a It has at least one polar bond. b it is asymmetrical in the sense that the centre of positive and negative charge don't coincide and cancel each other. So chloromethane is polar but tetrachloromethane is not.
BUT a molecule dies not need to be polar to have pdpd forces between them as only polar bonds are needed. So both the above mentioned molecules may show pdpd forces of attraction between each other.
4. Yes. It's between an highly electropositive hydrogen atom and an highly electronegatuve atom. Those are the only possibilities.

In chemistry it's important to ask yourself why you are studying something.

The study of intermolecular bonds is important to explain differences in melting and boiling temperatures of organic substances. Hydrogen bonds are the strongest generally and the more the number of these bonds the higher the value. Also remember that one with fluorine is stronger than one with oxygen which is stronger than one nitrogen. Hope you understand why. Then comes permanent dipole bonds and it's number and strenght.
Then comes induced dipole instantaneous dipole which depends on the number of electrons and surface area (if same number of electrons like in structural isomers)
It's also important to explain solubilities. Like dissolves in like concept may explain why polar substances dissolve in each other and nonpolar dissolves in each other, but is a simplification. When you learn about the properties of different homologous series it'll be more clearer.

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BarneyMersich · 07-01-2014 19:01

Okay, so, as I said I didn't know what Van Der Waal was, I know realise I was slightly incorrect. Whenever you see write Van Der Waals, I mean instantaneous dipole induced dipole. Sorry about that!

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gingerbreadman85 · 07-01-2014 23:34

(Original post by charco)
They are only extremely weak in very small molecules. In larger molecules, such as polymers they are strong. In sulphur, phosphorus and iodine they are strong enough to make the structures solid at room temperature. This makes them stronger than the hydrogen bonding in water!

It's more accurate to state that each individual VDW force is quite weak (excepting if the atoms involved are massive like iodine where the large orbitals involved can result in very significant temporary dipoles) but in molecules with more atoms there are more VDW forces between any two molecules, in the case of polymers many more, adding up to result in a stronger overall force.

gingerbreadman85 · 07-01-2014 23:35

(Original post by BarneyMersich)
Okay, so, as I said I didn't know what Van Der Waal was, I know realise I was slightly incorrect. Whenever you see write Van Der Waals, I mean instantaneous dipole induced dipole. Sorry about that!

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Sometimes also called "London Dispersion Forces".

3 names for the same thing!

BarneyMersich · 07-01-2014 23:49

(Original post by gingerbreadman85)
Sometimes also called "London Dispersion Forces".

3 names for the same thing!

So is Van Der Waals id-id, or is it a collective for id-id & pd-pd ?

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gingerbreadman85 · 08-01-2014 00:01

(Original post by BarneyMersich)
So is Van Der Waals id-id, or is it a collective for id-id & pd-pd ?

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Just id-id, not pd-pd.

BarneyMersich · 08-01-2014 00:10

(Original post by gingerbreadman85)
Just id-id, not pd-pd.

Ah great so I was correct the first time thanks haha

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BarneyMersich · 08-01-2014 00:26

(Original post by gingerbreadman85)
Hmm, double checking there is some controversy on that one.

Most A-level specifications describe VDW and dp-dp as different, however id-id and pd-pd both come under the category of VDW.

So for the sake of the exam, they are different.

My spec doesn't even cover it so I'm not too worried, I just don't want to post incorrect info on the thread. It's useful to know though, thanks!

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gingerbreadman85 · 08-01-2014 01:24

(Original post by BarneyMersich)
Ah great so I was correct the first time thanks haha

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Hmm, double checking there is some controversy on that one.

Most A-level specifications describe VDW and dp-dp as different, however id-id and pd-pd both come under the category of VDW.

So for the sake of the exam, they are different.

QuidditchFan · 08-01-2014 17:58

(Original post by BarneyMersich)
So:
1) yeah that sounds about right, as long as you remember diamond, graphite and silicon dioxide, you should be fine, at least that's on out spec (OCR salters B)

2) simple molecular doesn't really depend on number of bonds: as long as it isn't a network but it's covalent it will me simple molecular. As long as you can look at something like water, and realise that all the bonds are used up, so it can't be a network you should be okay.
Also,
Careful with wording:

"in larger molecules, the higher number of electrons makes stronger VDW bonds"

I know what you meant, but the mark schemes (especially on my exam board / exam) are very strict on wording. In larger molecules the higher number of electrons causes the molecule to make stronger VDW bonds with adjacent molecules. the way you wrote it suggested the electrons are doing the bonding. Rememer, these are intermolecular, so thats bonds between molecules, as opposed to intramolecular, which are bonds within a molecule (covalent etc).

3) dipole dipoles are when there's a dipole in both molecules.
As for electronegativity:
The general rule of thumb is the top right side of the periodic table is the most electronegative, the bottom left is the least. Things like H-Cl are polar, as Cl is very electronegative and H is not. (Joke: how are bonds like bears? Some are polar, some are not)
This is nothing like ionic bonding. These are due to bonds in molecules having a dipole. A dipole is only a small charge, usually written as δ- and δ+ (that's a delta) this is caused by the shared electrons being closer to one atom than the other (the more electronegative one)
ionic bonding is where there is a charge due to loss or gain of electrons. In an atom, and there are no shared electrons. In fact, the two atoms/molecules ionically bonded are only 'as one' when in solid crystal form. When dissolved in water they split into the ions.

4) yes that's right, but the second one doesn't necessarily need a Hydrogen I don't think: a hydrogen attached to a FON, can hydrogen bond to the lone pairs on another FON.

That's quite an essay, I may have missed something. Anything you're still unsure about just post!

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(Original post by Hospes)
1. Yes. But polymers are also giant covalent structures made up of simple monomer repeating units. If you still haven't covered polymers then don't worry. It would also be better to go through the specs of your relevant exam board. Mostly they'll list all the giant covalent structures that students would be assumed to know at the exam.
2. Yes and yes. Also remember that vanderwall is a general term that covers both instantaneous dipole - induced dipole intermolecular bonds (also called London forces or dispersion forces) and permanent dipole permanent dipole bonds so in the exam it's better to specify what of the two is present. My teacher advises us to use the VDW term only if both are present. Yes instantaneous dipole induce dipole bonds are present in iodine and not pds. Hope you understand why.
3. Leave aside metals and metallic ions in organic chemistry. With a bit of experience it's easy to identify a significant polar bond. Remember that alkane and alkene chains do not have polar bonds on their own (significant polar bonds ).
A molecule is polar if it fulfils two conditions. a It has at least one polar bond. b it is asymmetrical in the sense that the centre of positive and negative charge don't coincide and cancel each other. So chloromethane is polar but tetrachloromethane is not.
BUT a molecule dies not need to be polar to have pdpd forces between them as only polar bonds are needed. So both the above mentioned molecules may show pdpd forces of attraction between each other.
4. Yes. It's between an highly electropositive hydrogen atom and an highly electronegatuve atom. Those are the only possibilities.

In chemistry it's important to ask yourself why you are studying something.

The study of intermolecular bonds is important to explain differences in melting and boiling temperatures of organic substances. Hydrogen bonds are the strongest generally and the more the number of these bonds the higher the value. Also remember that one with fluorine is stronger than one with oxygen which is stronger than one nitrogen. Hope you understand why. Then comes permanent dipole bonds and it's number and strenght.
Then comes induced dipole instantaneous dipole which depends on the number of electrons and surface area (if same number of electrons like in structural isomers)
It's also important to explain solubilities. Like dissolves in like concept may explain why polar substances dissolve in each other and nonpolar dissolves in each other, but is a simplification. When you learn about the properties of different homologous series it'll be more clearer.

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(Original post by gingerbreadman85)
Hmm, double checking there is some controversy on that one.

Most A-level specifications describe VDW and dp-dp as different, however id-id and pd-pd both come under the category of VDW.

So for the sake of the exam, they are different.

Thank you for all your help! I'm doing OCR Course A so for me, the induced dipole-dipole interaction is the same as VDW. The three intermolecular forces that we have in our specification are: hydrogen bonds, permanent dipole-dipole forces and van der Waals' forces (induced dipole-dipole). Referring to that, it was still quite easy to understand your explanations so thank you for the clarity. Just about the previous points I raised (and clarification)

1. Group 4 are giant covalent structures. In the specification, it only looks at diamond and graphite and I think I'm okay with those (thanks to all of you!). I haven't learnt about polymers but I should also remember that polymers are also giant covalent structures (if I get this).

2. Simple covalent structures are normally group 7/6/5 and again, depends on the number of bonds. There are Van der Waal bonds between molecules and in larger molecules, the higher number of electrons makes stronger VDW bonds which can make structures solid at RTP and can be stronger than H-bonding in water. I understand this is the vdW bonding in iodine crystals which are broken when the crystal is heated.

3. Dipole - dipole bonds are caused when there is a dipole in one molecule. A dipole in one molecule is because of difference in electronegativities of atoms.Rule of thumb is that top right is most electronegative and bottom left is least electronegative and it's mostly practice which will tell me whether a molecule is polar or note (as long as it's covalently bonded) . Also, molecule must be asymmetrical in the sense that the centre of positive and negative charge don't coincide and cancel each other.

I got confused with ionic because I thought they were similar but now I realise polar covalent bonds are INTRAmolecular forces (which are on the scale between pure ionic and pure covalent.

I've been told that permanent dipole-dipole interactions only occur between polar molecules but it was interesting to know that they can happen between non-polar molecules. Hospes, I think I understand what you mean by this - that a molecule only needs to have polar bonds and even if it is symmetrical and polar bonds cancel out so the molecule overall is not polar, permanent dipole-dipole interactions still occur (because of those polar bonds).

4. Hydrogen bonding is between two molecules, one with H-O, H-N, H-F only and the other molecule with the FON - H and a lone pair. Hydrogen bond occurs between the δ- atom of one molecule (draw it to lone pairs if there) and δ+ atom of the other molecule.

Is that all right?

I cannot thank you enough for all the help you've given me! Though it doesn't come as naturally as the other concepts (I suspect that's down to practice and familiarising myself with them), it has made sense. I think, if I am given a molecule and told to explain the forces and bonds in them, it should be fine.

TheGrinningSkull · 08-01-2014 18:06

(Original post by QuidditchFan)
x

I'm just going to add, thing of hydrogen bonding as just being a stronger form of permanent dipole bonding. And like you've said permanent being stronger than vdW when there's non-symmetry etc.

gingerbreadman85 · 08-01-2014 18:11

(Original post by QuidditchFan)
I got confused with ionic because I thought they were similar but now I realise polar covalent bonds are INTRAmolecular forces (which are on the scale between pure ionic and pure covalent.

Bingo. Polar bonds are in a sense partly ionic in nature, this means they also have higher bond energies than would be expected from a non-polar bond. In fact, the Pauling electronegativity values come from comparing bond energies with expected values for a purely covalent system. In the same way, some ionic substances can have covalent character (something you might meet in A2)

(Original post by QuidditchFan)
Hydrogen bonding is between two molecules, one with H-O, H-N, H-F only and the other molecule with the FON - H and a lone pair. Hydrogen bond occurs between the δ- atom of one molecule (draw it to lone pairs if there) and δ+ atom of the other molecule.

Also, you need to keep the hydrogen bond in line with the H-X bond that it links to, like:
H^δ+-X^δ-:-----H^δ+-X^δ-

Otherwise, looks like you've got a pretty good understanding of things.

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