On this page
Whenever you study the p-block elements for JEE or NEET, a familiar rule appears: compounds in higher oxidation states are more covalent than those in lower oxidation states. The classic textbook example is phosphorus pentachloride (PCl₅) compared to phosphorus trichloride (PCl₃).
But why does simply adding two more halogen atoms change the nature of the bonds? The answer lies in how a central atom's charge affects its neighbours, a principle governed by Fajans' rules.
How does oxidation state affect covalency?
To understand the shift from ionic to covalent character, we need to look at the central atom as a positively charged ion (a cation).
In a pentahalide like PCl₅, the central phosphorus atom is in a +5 oxidation state. In the corresponding trihalide, PCl₃, it is in a +3 oxidation state.
According to Fajans' rules, two things happen when an atom reaches a higher positive oxidation state:
Higher charge: The +5 state has a significantly greater positive charge than the +3 state.
Smaller size: As electrons are lost or drawn away, the remaining electron cloud is pulled tighter to the nucleus, making the +5 ion smaller than the +3 ion.
Together, a high charge packed into a small volume results in a very high polarizing power. The central atom becomes extremely effective at pulling negatively charged electrons towards itself.
How does polarizing power create a covalent bond?
When a central atom with high polarizing power sits next to a halide ion, it exerts a strong pull on the halide's outer electrons.
Instead of the electrons staying neatly around the halide, the electron cloud is dragged towards the central atom. This distortion is called polarization.
Because the +5 atom in a pentahalide polarizes the halide's electron cloud much more strongly than the +3 atom does, the electrons are pulled into the space between the two nuclei. When electron density is concentrated between two atoms, it becomes a shared electron pair. This shared state is the defining feature of a covalent bond.
If you are struggling to picture how these electron clouds distort, the Concept Library covers every concept in chemistry with its own lesson. Fermi's AI tutor works through the ideas alongside you; if you are unsure why a +5 charge causes more distortion, it holds a line of reasoning and decides whether to explain the electrostatic pull or ask a guiding question, checking your understanding as it goes.
What is the real-world proof of this rule?
This increase in covalent character is not just a theoretical rule; it radically changes how the substance behaves in the real world. Ionic compounds have strong electrostatic forces holding them in a rigid lattice, giving them high melting points. Covalent molecules are held together by much weaker intermolecular forces, so they melt at lower temperatures.
Consider antimony from Group 15. It forms both a trihalide and a pentahalide with chlorine:
Antimony trichloride (SbCl₃) is a crystalline solid at room temperature and melts at 73.4 °C (Wikipedia).
Antimony pentachloride (SbCl₅) is an oily liquid at room temperature and melts at just 2.8 °C (Wikipedia).
The shift from a +3 to a +5 oxidation state increases the covalent character so much that it turns a solid crystal into a liquid that melts barely above freezing.
How do you use Fajans' rules in an exam?
When a question asks you to compare the covalent character of two compounds, use this simple framework:
Find the oxidation states: Identify the charge on the central atom. The higher the positive charge, the higher the polarizing power.
Compare the cation sizes: If the charge is the same, the smaller central atom is more polarizing.
Compare the anion sizes: For the halides themselves, larger anions (like iodine) hold their outer electrons more loosely than small ones (like fluorine). A larger anion is easier to polarize, meaning iodides are generally more covalent than fluorides.
Whether you are predicting the Lewis acidity of boron trihalides or mapping out oxidation numbers from structure, Fajans' rules provide the bridge between the theoretical oxidation state and the physical reality of the bond.
