konichiwa2x said:
What determines the ionic nature of a bond? I am a bit confused regarding this.
And likely for good reason - this is one of those questions you can stump most high school chemistry teachers with - it is very poorly understood and taught.
Is it the polarising ability of the cation or the greater electronegativity of the anion?
First, there is more than one property that determines the polarity of a bond. Second, the properties of both the species involved in the bond are important. And third, 'electronegativity' is not defined on an ion, it is defined for an atom (typically).
What are these properties and how do they affect the polarity of the bond? There are two approaches to answering that. The more popular approach is through Fajan's Rules (as opposed to Pauling's approach). The difference between these approaches is the Fajan treats the polar covalent bond as a perturbation (or small deviation) from the purely ionic bond, while Pauling treats it as a perturbation from a perfectly non-polar covalent bond.
In Fajan's approach, you start with a (hypothetical) "purely ionic compound" and then redistribute charge according to certain rules to make the final compound more covalent than your starting condition. This starting point requires that you are apply the rules to ions (rather than neutral atoms), so you must be familiar with the (periodic trends in) ionic radii before you can use Fajan's Rules successfully.
Regarding fajan's rules, smaller the size of the cation (and greater the charge) and smaller the size of the anion, greater is the covalent nature of the bond.
No, this is not right. Go over Fajan's Rules again. It takes a smaller cation and/or a larger anion to make a more covalent bond (effective charge being constant). NaI is much more covalent (or much less polar) than NaCl (since I
- is bigger than Cl
-).
So in this case, the bonding pair of electrons is pulled closer to the core of the cation, resulting in accumulation of a slight negative charge for the cation?
No, there is never a net negative charge on the cation. And since you start from the case of a perfectly ionic molecules, there is no such thing as a bonding pair.
If you compare LiI with KI, because of the smaller size of the Li
+ ion than the K
+, the former is better able to "suck" some partial negative charge out of the anion, I
- (according to Fajan's Rule#1). As a result, the positive charge (+1) on the Li
+ ion is reduced more than the positive charge on the K
+ ion would be. Also, Li
+, by sucking out more negative charge from I
-, makes the anion not-so-negatively charged anymore. As the result, the net partial charges on Li(\delta +) and I(\delta -) in LiI are smaller than the net partial charges on K(bigger \delta +) and I(bigger \delta -) in KI. In other words, LiI is more covalent (or non-polar), and KI is more ionic (or polar).
But in some other cases, the shared pair(s) of electrons are pulled closer to the anion due to the greater electronegativity of the anion.
Any shared pair is always pulled closer to the more electronegative species, typically the non-metal that forms the anion. But you can talk of shared pairs and electronegativities only if your starting point is a covalent molecule formed from neutral atoms (ie: if you are using the Pauling formalism), not an ionic molecule formed between ions. There is nothing to share in an ionic molecule - the two ionic species have their octets and are happy.
Note that the Fajan formalism does not create a net negative charge on the cation, it only reduces the amount of positive charge. You always end up with a net negative charge on the ion formed from the more electronegative species (the anion) and a net positive charge on the cation.
So.. does this mean that the bonding pair of electrons can be pulled closer to either to the cation or the anion??
No - as explained above.
