[habbababba]

I would like to know if the idea that polarizing the anion by the cation in a lattice leads to a more exothermic value for lattice enthalpy is correct. If yes, why? how does polarization, which is viewed as an extent of covalency, lead to the release of more heat?

Thanks.

[mjc123]

Briefly, if it didn't lower the energy, it wouldn't happen. It describes a situation where the minimum energy is attained when the bonding electrons are not 100% on the anion and 0% on the cation, as in ideal ionic bonding, nor shared 50:50 as in a nonpolar covalent bond, but somewhere in between, e.g. 80:20. In fact there aren't really two completely distinct types of bonding, rather there is a continuum from covalent through polar covalent to ionic, with some bonds at or close to one extreme or the other, and some in between.
To look at it slightly differently, to say that the cation polarises the anion - i.e. that it exerts a force on the outer electrons of the anion, pulling electron density towards itself - is to say that the electrons can lower their energy by moving towards the cation. Remember F = -dE/dx.
Incidentally, just to be pedantic, lattice enthalpy is not exothermic, it is endothermic. It is defined as the enthalpy for the dissociation of the solid compound into gaseous ions - not the other way round. When you're writing a Born-Haber cycle, the enthalpy for the step ions(g)  solid is -L.

[habbababba]

Thanks for the input.

I learned that the more ionic character a bond possesses, the stronger the bond and hence more energy is released. A degree of covalency takes the bond to the covalent side of the continuum and as a result wouldn't that mean less energy is released? (In comparison to one that has more ionic character).
Take for instance 2 different compounds such as NaCl and MgI₂. One way to describe the bonding in the 2 compounds is to compare the difference between the experimental lattice enthalpy and the theoretical lattice energy. The experimental value for the lattice enthalpy of NaCl is -780 KJ/mol and the theoretical is -770. The experimental value for the lattice enthalpy of MgI₂ is -2327 KJ/mol and the theoretical is -1944. The difference is larger in MgI₂ and the explanation for this is that the bond in MgI₂ has a higher covalent character than the bond in NaCl and so more energy is released than expected, whereas in NaCl the bond has a greater ionic character. But I thought the more pronounced the degree of covalency in a bond is, the weaker it is compared to one that has a higher ionic character.

[Corribus]

Incidentally, just to be pedantic, lattice enthalpy is not exothermic, it is endothermic. It is defined as the enthalpy for the dissociation of the solid compound into gaseous ions - not the other way round.

Not all textbooks define it this way. There are plenty that define it from the standpoint of an exothermic dissociation. There is no standard convention as far as I know. The important thing is to just be clear whether you are referring to a lattice formation or lattice dissociation enthalpy.

http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Crystal_Lattices/Thermodynamics_of_Lattices/Lattice_Enthalpies

[Borek]

But I thought the more pronounced the degree of covalency in a bond is, the weaker it is compared to one that has a higher ionic character.

Common mistake, most likely stemming from the fact ionic solids have much higher melting temperatures than the solids made of covalently bonded molecules. However, melting point is not related to the internal bonding, but to the intermolecular forces (which don't make much sense in the case of NaCl).

[habbababba]

Common mistake, most likely stemming from the fact ionic solids have much higher melting temperatures than the solids made of covalently bonded molecules. However, melting point is not related to the internal bonding, but to the intermolecular forces (which don't make much sense in the case of NaCl).

I wouldn't say that melting point is not related to the internal bonding for all compounds, take SiO₂ for instance, has a melting point of about 1,700 °C and yet the forces in SiO₂ are covalent bonds, with no intermolecular forces that I know of. But I think I get your point. SiO₂ has a very high melting point compared to many other ionic compounds, so when it comes to comparing the strength of covalent bonds to that of ionic bonds (by comparing the melting point of giant covalent substances with that of ionic compounds), there are instances when covalent bonds require more energy to break such as in SiO₂. So maybe I can refer to this example as an evidence for correcting the common misconception I previously had.

Thank you for the input.

[AdiDex]

But I have something , Sometimes Intramolecular forces also matter
Eg. Melting point of following Compound

NaCl>NaBr >NaI
MgCl₂>MgBr₂>MgI₂