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Polarising effect A level Chem

I’m confused about this concept and hoping someone can explain what I’m doing wrong

So when we have a small positive ion and a small negative ion they form a strong ionic bond as the outer electron is closer to the nuclei and all that so this increases the lattice enthalpy as we need more energy to break this bond

However, when we have a small positive ion and a LARGE negative ion, there is a larger polarising effect and this polarising effect means that the compound has some covalent character and thus the lattice is stronger and more energy is needed to break the bond

I feel like it’s a contradiction because how are both small and big negative ions resulting in stronger lattice?

I hope the question makes sense

Reply 1

TypicalNerd

Original post by Aeshakhan

(Referring to the part I've put in bold) it's not so much to do with the outer electron being closer to the nuclei - it's more a matter of the bond length being shorter and so by Coulomb's law (which states that electrostatic force is directly proportional to the product of the two charges and inversely proportional to the square of the distance between them), the electrostatic force is expected to be stronger if you have smaller ions.

Let's use the entries in the old Edexcel chemistry data book to try illustrating the point (note that all entries in brackets are the theoretical value - everything not written in brackets is an experimental value and also there is a typo in the table and it should say Cu^+ for the entries for copper compounds, rather than Cu^2+).

Lattice enthalpies.png

Let's look at the halides of magnesium (MgF2, MgCl2, MgBr2 and MgI2) as they best illustrate what is going on. You can see as the size of the anion increases, the experimental lattice enthalpy becomes less endothermic. This is a result of Colomb's law, as explained briefly above.

However, when you compare the experimental and theoretical values of the lattice enthalpies, the disparity between the two gets larger as the anion gets larger and you will notice that the theoretical lattice enthalpies are always underestimates of the experimental ones. This is due to the polarising effect, which is what makes the bonding stronger than expected in the halides with larger anions. But the polarising effect has less of an effect on the lattice enthalpy than the separation of the ions, so the lattice enthalpies (both theoretical and experimental) of the halide salts of any element would be expected to decrease, descending group 7.

So in light of the above, let us rephrase the statements you gave as I rather suspect the following are what I imagine that your teacher or source was trying to articulate:

So when we have a small positive ion and a small negative ion the ionic bonding is strong, because the separation between the ions is small and so the electrostatic force of attraction between them is large. But there is little covalent character and so the experimental and theoretical values tend to be very close.

However, when we have a small positive ion and a LARGE negative ion, there is a polarising effect and this polarising effect means that the compound has some covalent character and thus the lattice is stronger than you would expect and so more energy is needed to break the bond than you would expect. But this does not mean that compounds with larger anions would be expected to have larger lattice enthalpies than those with similarly sized cations and smaller anions.