Chemistry - Ligands/Coordinate Bonding

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I'm learning transition metal Chemistry for the first time, and I'm getting quite confused.

Usually a dative covalent bond involves the sharing of 2 electrons coming from the same atom, and hence a formal charge appears on the molecule, e.g. CO3^2-, SO4^2- etc. However with ligands, there can often be as many as six coordinate bonds to the central metal ion, yet the effect of these six electrons being donated don't affect the overall charge? Even with neutral ligands like water, shouldn't there be a formal charge per ligand, caused just by dative covalent bonding? Only the charges of central metal ions e.g. Cu^2+ or any charges on the ligands e.g. Cl^- are included in the overall charge of the complex ion.

I hope I'm making sense here!
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sayema1
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(Original post by I <3 WORK)
I'm learning transition metal Chemistry for the first time, and I'm getting quite confused.

Usually a dative covalent bond involves the sharing of 2 electrons coming from the same atom, and hence a formal charge appears on the molecule. However with ligands, there can often be as many as six coordinate bonds to the complex ion, yet the effect of electrons being donated don't affect the overall charge? Even with neutral ligands like water, shouldn't there be a formal charge caused by dative covalent bonding? Only the charge of the central complex ion is included in these cases.

I hope I'm making sense here.
have you checked with revision notes in physics and maths tutor?
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(Original post by sayema1)
have you checked with revision notes in physics and maths tutor?
I've tried with all my resources, but can't seem to find an answer!
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charco
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(Original post by I <3 WORK)
I'm learning transition metal Chemistry for the first time, and I'm getting quite confused.

Usually a dative covalent bond involves the sharing of 2 electrons coming from the same atom, and hence a formal charge appears on the molecule, e.g. CO3^2-, SO4^2- etc.
The above statement is not correct. The formal charge on an ion is not due to dative coordinate bonds.

Formal charge is a concept that is applied to individual atoms within a species to give an idea as to which of the possible resonance structures is more likely to be favoured. This is affected by dative covalency.

The actual charge on an ion is due to the atom or atoms having gained (or lost) one or more electrons.



However with ligands, there can often be as many as six coordinate bonds to the central metal ion, yet the effect of these six electrons being donated don't affect the overall charge? Even with neutral ligands like water, shouldn't there be a formal charge per ligand, caused just by dative covalent bonding? Only the charges of central metal ions e.g. Cu^2+ or any charges on the ligands e.g. Cl^- are included in the overall charge of the complex ion.

I hope I'm making sense here!
Ligands that are neutral have no effect on the overall charge, and as there are no other possible resonance forms it makes no sense to discuss the concept of formal charge on each atom as applied to a transition metal complex.

The actual charge of a transition metal complex ion is simply the sum of the metals oxidation state (ionic charge) and the charges of the individual ligands.

At the end of the day, the charge on an ion is the difference between the total number of protons and electrons.

Formal charge just applies this principle to individual atoms within a species.

Example:

Carbon monoxide, CO
------------------------------
There is a sigma bond and a pi bond between the two atoms, and there is also a dative coordinate bond where two electrons from oxygen are donated into a pi bond.

Carbon therefore has a lone pair(2) half share of two bonding pairs(2) and half share of a dative (1), plus two electrons from the 1s shell. This makes a total of 7 electrons. It has six protons, therefore there is a formal charge of -1 on the carbon atom.7
Oxygen has half share of the three bonds (3) it has a lone pair (2) and 2 electrons in the inner shell = total of 7 electrons. It has 8 protons and hence a formal charge of +1
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(Original post by charco)
The above statement is not correct. The formal charge on an ion is not due to dative coordinate bonds.

Formal charge is a concept that is applied to individual atoms within a species to give an idea as to which of the possible resonance structures is more likely to be favoured. This is affected by dative covalency.

The actual charge on an ion is due to the atom or atoms having gained (or lost) one or more electrons.




Ligands that are neutral have no effect on the overall charge, and as there are no other possible resonance forms it makes no sense to discuss the concept of formal charge on each atom as applied to a transition metal complex.

The actual charge of a transition metal complex ion is simply the sum of the metals oxidation state (ionic charge) and the charges of the individual ligands.

At the end of the day, the charge on an ion is the difference between the total number of protons and electrons.

Formal charge just applies this principle to individual atoms within a species.

Example:

Carbon monoxide, CO
------------------------------
There is a sigma bond and a pi bond between the two atoms, and there is also a dative coordinate bond where two electrons from oxygen are donated into a pi bond.

Carbon therefore has a lone pair(2) half share of two bonding pairs(2) and half share of a dative (1), plus two electrons from the 1s shell. This makes a total of 7 electrons. It has six protons, therefore there is a formal charge of -1 on the carbon atom.7
Oxygen has half share of the three bonds (3) it has a lone pair (2) and 2 electrons in the inner shell = total of 7 electrons. It has 8 protons and hence a formal charge of +1
Thanks for your reply. I must be confusing the definition of formal charge with the overall charge of a polyatomic ion. In any case, I think there's a key point I seem to be missing.

Say for instance we have the complex ion: [Cu(H2O)6]^2+ where each H2O ligand is datively-covalently bonded to Cu^2+. Correct me if I'm wrong, but any single covalent bond whether dative or not will involve the sharing of two electrons between two atoms. Therefore the lone pair of electrons on oxygen from H2O will be shared with Cu^2+. Now because both electrons are coming from one atom (oxygen), yet these two electrons are still being shared between oxygen and Cu^2+, the H2O should lose 1 electron per coordinate bond i.e. having a +1 charge.

I can see the above being very similiar to any other dative covalent compound forming, for e.g. ammonia having one lone pair can share it's electrons with H+, and a +1 charge will be created, due to the transfer of positive charge (NH4)^+.

So I really don't understand how complex ions can be any different?
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charco
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(Original post by I <3 WORK)
Thanks for your reply. I must be confusing the definition of formal charge with the overall charge of a polyatomic ion. In any case, I think there's a key point I seem to be missing.

Say for instance we have the complex ion: [Cu(H2O)6]^2+ where each H2O ligand is datively-covalently bonded to Cu^2+. Correct me if I'm wrong, but any single covalent bond whether dative or not will involve the sharing of two electrons between two atoms. Therefore the lone pair of electrons on oxygen from H2O will be shared with Cu^2+. Now because both electrons are coming from one atom (oxygen), yet these two electrons are still being shared between oxygen and Cu^2+, the H2O should lose 1 electron per coordinate bond i.e. having a +1 charge.

I can see the above being very similiar to any other dative covalent compound forming, for e.g. ammonia having one lone pair can share it's electrons with H+, and a +1 charge will be created, due to the transfer of positive charge (NH4)^+.

So I really don't understand how complex ions can be any different?
You are correct in your analysis, but as I explained earlier, it does not make much sense applying formal charge "rules" to complex ions as there is no ambiguity of structure.

The formal charge "rules" are mere accountancy tools that are not reflected in any reality. They are similar to oxidation states in that they just a man made construct to help guide us along.

Complex ions share electrons between the ligand and the metal ion. The origin of the shared pair is immaterial. The whole structure is a complex and as such may carry positive, negative or no charge.
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(Original post by charco)
You are correct in your analysis, but as I explained earlier, it does not make much sense applying formal charge "rules" to complex ions as there is no ambiguity of structure.

The formal charge "rules" are mere accountancy tools that are not reflected in any reality. They are similar to oxidation states in that they just a man made construct to help guide us along.

Complex ions share electrons between the ligand and the metal ion. The origin of the shared pair is immaterial. The whole structure is a complex and as such may carry positive, negative or no charge.
So I guess I'll just have to accept it the way it is! I understand now that in reality the structure and reasoning is far more complex than needed for A-Level standard, but I watched a video recently and he briefly mentioned that dative covalent bonds are formed within empty higher level orbitals of transition metals. Is this true and does it have any relation to the charges in coordinate bonding?
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(Original post by I <3 WORK)
So I guess I'll just have to accept it the way it is! I understand now that in reality the structure and reasoning is far more complex than needed for A-Level standard, but I watched a video recently and he briefly mentioned that dative covalent bonds are formed within empty higher level orbitals of transition metals. Is this true and does it have any relation to the charges in coordinate bonding?
I'm sorry, but once again you are mixing things up.

The bonding in 1st row transition metal complexes does not involve the 3d orbitals, rather molecular orbitals formed from the 4th shell.

There are simple models, such as hybridisation of 4s 4p and 4d orbitals forming octahedral 4sp3d2 orbitals and more complex ('scuse the pun) molecular orbital, rather more mathematical, models.

There are also ligands that donate lone pairs along sigma orbitals and accept them back into molecular pì orbitals (back bonding) giving much stronger bonds.

The very simple model of a ligand donating a pair of electrons into a coordinate bond is useful, but limited. Hence, it is not useful to start trying to calculate formal charges. Any results you get are meaningless. With this model the transition metal ion will invariably have a formal charge of -4 for 2+ ions in an octahedral field.
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(Original post by charco)
I'm sorry, but once again you are mixing things up.

The bonding in 1st row transition metal complexes does not involve the 3d orbitals, rather molecular orbitals formed from the 4th shell.

There are simple models, such as hybridisation of 4s 4p and 4d orbitals forming octahedral 4sp3d2 orbitals and more complex ('scuse the pun) molecular orbital, rather more mathematical, models.

There are also ligands that donate lone pairs along sigma orbitals and accept them back into molecular pì orbitals (back bonding) giving much stronger bonds.

The very simple model of a ligand donating a pair of electrons into a coordinate bond is useful, but limited. Hence, it is not useful to start trying to calculate formal charges. Any results you get are meaningless. With this model the transition metal ion will invariably have a formal charge of -4 for 2+ ions in an octahedral field.
Thank you, that actually makes sense now.

I'd just like to know, what is special about transition metals in that they are able to form ligands using much higher energy levels (e.g. 4p, 4d etc.)?
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(Original post by I <3 WORK)
Thank you, that actually makes sense now.

I'd just like to know, what is special about transition metals in that they are able to form ligands using much higher energy levels (e.g. 4p, 4d etc.)?
Nothing special. Many other elements do this as well.
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