- When comparing the relative stabilities of two charged atoms that are in the same row of the periodic table, electronegativity is the most important factor
- When comparing the relative stabilities of two charged atoms that are in the same column of the periodic table, polarizability is the most important factor
- Larger (more polarizable) atoms are better able to stabilize negative charges
In the previous section we saw that molecules are more stable when their charges are less ‘concentrated’. Consider F– and Cl–, represented below. Both bear a full negative charge, but Cl– is larger so the negative charge is spread out over a larger area. Because the charge on Cl– is less concentrated, Cl– is more stable than F–.
This is a general trend: larger atoms are better at stabilizing negative charge than smaller atoms. But what if we consider the relative electronegativities of F– and Cl–? F is more electronegative, so based on electronegativity alone, we would have incorrectly predicted that F– is more stable than Cl–. Both atom size and electronegativity affect stability for all ions, but because atom size changes much more drastically down a column of the periodic table than across a row, we find that:
- When comparing the stability of two charged atoms from the same row of the periodic table, electronegativity is the more important factor
- When comparing the stability of two charged atoms from the same column of the periodic table, atom size is the more important factor
F and Cl are in the same column of the periodic table, so the difference in their atomic sizes overrides their electronegativity difference when predicting their relative stabilities.
This trend is often referred to as polarizability. The polarizability of an atom is a measure of how easily its electrons are distorted (or polarized) by the presence of an external electric field (e.g. an ion). Larger atoms have more electrons and hold onto their outer electrons less tightly, so are more polarizable. The more polarizable (larger) an atom is, the better it is at stabilizing a negative charge.
Avoid this common error:
Only consider polarizability for the atoms that bear a formal charge. A common error is to consider polarizability instead of electronegativity when comparing the induction in two molecules. For example, the two molecules below are identical except for the halogen. The molecule on the left is more stable because F is more electronegative so the electron density is more spread out. We do not consider polarizability differences between F and Cl here because these atoms do not bear the formal negative charge.
Interactive: