Zenarthra
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I understand that experimental lattice enthalpies are more exothermic than theoretical values, in magnesium halides the ionic bond is polarized and therefore there is covalent character. If this makes them more exothermic, then does it mean covalent character makes the ionic bond stonger?


Thanks!
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Borek
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(Original post by Zenarthra)
does it mean covalent character makes the ionic bond stonger?
Generally speaking there are no purely ionic nor purely covalent bond, (almost) every bond has some mixture of both characters (with possible exclusions being bonds in molecules like N, or between carbons in ethane/ethene/ethyne).

That means you can't say what you just said. You can try to ask question like "does it mean covalent character makes the bond stronger?" - see the difference?
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Zenarthra
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(Original post by Borek)
Generally speaking there are no purely ionic nor purely covalent bond, (almost) every bond has some mixture of both characters (with possible exclusions being bonds in molecules like N, or between carbons in ethane/ethene/ethyne).

That means you can't say what you just said. You can try to ask question like "does it mean covalent character makes the bond stronger?" - see the difference?
Ahh ok, so does it then?
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charco
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(Original post by Zenarthra)
Ahh ok, so does it then?
There are two forms of covalent bonding, simple and giant.

If the ionic structure tends to covalent character moving towards simple covalent then it must weaken the bonding. Example - aluminium chloride
If it produces an enhanced network then it would strengthen the bonding. Example - AgCl

The 'prediction' is horse before cart, using experimental and theoretical values.. i.e. use the experimental data and make the theory to fit.
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Zenarthra
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(Original post by charco)
There are two forms of covalent bonding, simple and giant.

If the ionic structure tends to covalent character moving towards simple covalent then it must weaken the bonding. Example - aluminium chloride
If it produces an enhanced network then it would strengthen the bonding. Example - AgCl

The 'prediction' is horse before cart, using experimental and theoretical values.. i.e. use the experimental data and make the theory to fit.
Sorry i dont understand, what is tends to covalent character and enhanced network?
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ChemistryBud
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http://www.thestudentroom.co.uk/show....php?t=2708652

Check this and let me know if you still need help
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Zenarthra
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(Original post by ChemistryBud)
http://www.thestudentroom.co.uk/show....php?t=2708652

Check this and let me know if you still need help
Thanks, but still dont understand why more covalent character makes ionic bonds stronger.
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ChemistryBud
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(Original post by Zenarthra)
Thanks, but still dont understand why more covalent character makes ionic bonds stronger.
Ok so Charco & Borek have already given the answer but I'll try to simplify things.

When the enthalpy is greater than expected (more exothermic) it is because the covalent character strengthens the lattice interactions (in contrast to if there was ionic bonding only).

However, because 'pure' ionic bonds are stronger than 'pure' covalent bonds; if ionic interactions are weak and the bonds are predominately covalent, there is a a decrease in the lattice enthalpy.

For some lattices it is hard to predict if the enthalpy will be higher or lower than expected. As a result, we collect the experimental data before saying if the covalent character increases or decreases the lattice strength.

In reality there is no such thing as a 'purely' ionic or 'purely' covalent bond', a bond will always have some ionic character (even if it barely contributes to the bond strength) and vice-versa.

Does this make sense yet? Please say if not.
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Zenarthra
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(Original post by ChemistryBud)
Ok so Borek has already given the answer but I'll try to simplify things.

When the enthalpy is greater than expected (more exothermic) it is because the covalent character strengthens the lattice interactions (in contrast to if there was ionic bonding only).

However, because 'pure' ionic bonds are stronger than 'pure' covalent bonds; if ionic interactions are weak and the bonds are predominately covalent, there is a a decrease in the lattice enthalpy.

For some lattices it is hard to predict if the enthalpy will be higher or lower than expected. As a result, we collect the experimental data before saying if the additional covalent character increases or decreases the lattice strength.

Does this make sense yet? Please say if not.
For the part in bold how?
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ChemistryBud
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In addition to the electrostatic interactions between ions you also have a sharing of electrons due to polarization of the halide anion. That is, both ionic bonding and covalent character respectively.

I think this relates to your question about charge density and again I would recommend:
http://www.chem.uky.edu/courses/che514/Fajan.pdf
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Zenarthra
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(Original post by ChemistryBud)
In addition to the electrostatic interactions between ions you also have a sharing of electrons due to polarization of the halide anion. That is, both ionic bonding and covalent character respectively.

I think this relates to your question about charge density and again I would recommend:
http://www.chem.uky.edu/courses/che514/Fajan.pdf


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I'm on phone so can't new thread, but just wanted to ask at zero kelvins all molecules are stationary, but why does this mean that the entropy is zero at this temp?

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Borek
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(Original post by Zenarthra)
Thanks, but still dont understand why more covalent character makes ionic bonds stronger.
You did the same mistake for the second time.
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Zenarthra
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(Original post by Borek)
You did the same mistake for the second time.
Look, i was just going off a markscheme to a question.
Just forget it.
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Zenarthra
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(Original post by Borek)
You did the same mistake for the second time.
(Original post by charco)
There are two forms of covalent bonding, simple and giant.

If the ionic structure tends to covalent character moving towards simple covalent then it must weaken the bonding. Example - aluminium chloride
If it produces an enhanced network then it would strengthen the bonding. Example - AgCl

The 'prediction' is horse before cart, using experimental and theoretical values.. i.e. use the experimental data and make the theory to fit.
(Original post by ChemistryBud)
http://www.thestudentroom.co.uk/show....php?t=2708652

Check this and let me know if you still need help
Why does the Al3+ ion distort the e cloud but MgO does not?
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ChemistryBud
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:damnmate:

An Al3+ ion on its own will not distort an electron cloud, are we meant to guess the compound you are referring to? In the case of MgO, the electron cloud of the oxygen atom will be distorted but probably not to the same extent as this mystery Al3+ compound...

This relates to charge density, did you take a look at the link I sent?
http://www.chem.uky.edu/courses/che514/Fajan.pdf

Al3+ has a higher charge density than Mg2+ (the oxidation state of Mg in MgO).
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charco
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(Original post by ChemistryBud)
:damnmate:
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Zenarthra
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(Original post by charco)
(Original post by ChemistryBud)
:damnmate:

An Al3+ ion on its own will not distort an electron cloud, are we meant to guess the compound you are referring to? In the case of MgO, the electron cloud of the oxygen atom will be distorted but probably not to the same extent as this mystery Al3+ compound...

This relates to charge density, did you take a look at the link I sent?
http://www.chem.uky.edu/courses/che514/Fajan.pdf

Al3+ has a higher charge density than Mg2+ (the oxidation state of Mg in MgO).
Yes i did read it, and apologies i meant Al2O3 .
So MgO has higher mpt than Na2O because Mg2+ attracts O-2 more strongly than in Na2O.
But in Al2O3 Al3+ attracts O2- even more strongly?
Why then does Al2O3 have a lower mpt than expected?

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ChemistryBud
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(Original post by Zenarthra)
Yes i did read it, and apologies i meant Al2O3 .
So MgO has higher mpt than Na2O because Mg2+ attracts O-2 more strongly than in Na2O.
But in Al2O3 Al3+ attracts O2- even more strongly?
Why then does Al2O3 have a lower mpt than expected?

Thanks!
Haha thank you for the identity of the compound. This is actually a common question and a nuisance because it disagrees with our predictions- that the bonding in Al2O3 will be stronger since the charge density is greater for Al3+ than Mg2+ (Al3+ has both a larger charge and smaller radius).

The accepted answer is that because the electronegativity difference between Al and O is lower than that between Mg and O; bonding in Al2O3 is less ionic and more covalent (it has more covalent character). Now this may sound hypocritical; earlier we said that ionic bonding with covalent character increases the strength of lattice interactions. This is an exception to that trend whereby the additional covalent character actually causes the melting point to be lower than expected ('the bonding tends towards simple covalent rather than enchanced lattice interactions'). As charco pointed out, it is also a great example of why we collect experimental data to confirm/reject our initial predictions and hypothesis (just because it sounds right doesn't mean it is).
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Zenarthra
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(Original post by ChemistryBud)
Haha thank you for the identity of the compound. This is actually a common question and a nuisance because it disagrees with our predictions- that the bonding in Al2O3 will be stronger since the charge density is greater for Al3+ than Mg2+ (Al3+ has both a larger charge and smaller radius).

The accepted answer is that because the electronegativity difference between Al and O is lower than that between Mg and O; bonding in Al2O3 is less ionic and more covalent (it has more covalent character). Now this may sound hypocritical; earlier we said that ionic bonding with covalent character increases the strength of lattice interactions. This is an exception to that trend whereby the additional covalent character actually causes the melting point to be lower than expected ('the bonding tends towards simple covalent rather than enchanced lattice interactions'). As charco pointed out, it is also a great example of why we collect experimental data to confirm/reject our initial predictions and hypothesis (just because it sounds right doesn't mean it is).
Wow great explanation!
please may you also check this thread:
http://www.thestudentroom.co.uk/show...9#post48177979

Thanks dude!
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