[ComradeLenin]

I'm a bit confused! Our natural sciences teacher has said that π bonds can only be obtained by lateral overlap of pure p orbitals. So if that's true, I can't have double bonds in molecules in which I find sp3 ibridization, as I don't have any pure p orbital to use for the bond. However, I've found a plenty of molecules in which the central atom has sp3 ibridization but also double bonds, like for example ClO3. Where do I make mistakes?

[Corribus]

First, this isn't a very accurate description of what a pi-bond is, so that's part of your problem. It's good enough for general chemistry I guess but then you shouldn't apply it unusual molecules/ions like chlorates that have more complicated electronic structures or "hypervalent" nuclei. Also in these cases the hybridization model don't work very well, so I don't see the sense in general chemistry textbooks trying to apply this model to them. The hybridization model is fine for molecules that have typical valency - like organic molecules featuring carbon - but don't be surprised if you find cases where you get contradictions. Molecular orbital theory is far superior for describing bonding in these situations.. and more advanced treatments like DFT are even better.

But the fact that you are finding contradictions is a good sign that you're thinking critically about these things. The basic structural models you learn in introductory chemistry classes aren't perfect... or even very good in many cases. They are instructive and can teach you a lot about quantum mechanics, and they can be used to understand qualitative trends, but they have significant inaccuracies and shouldn't be applied broadly. You'll run into problems quickly. Modern physical chemists use more complex treatments for real quantitative calculations of molecular structure and reactivity.

[ComradeLenin]

Thank you very much, I really appreciated your help and you got me interested! Could you give me a more detailed (but still easy ) definition of pi bond? I'm not in desperate need of it, I'd just like to understand better how these things work

[Corribus]

A practical definition of the bond classification may be the number of atomic orbital lobes that interact to form the bond. For example sigma bonds are defined by the interaction of one lobe on one atomic orbital with one (and only one) lobe on another atomic orbital. By contrast, pi bonds result from interaction of two atomic orbital lobes on one atomic orbital interacting with two atomic orbital lobes on another atomic orbital. (Just so, delta-bonds, which you've probably never heard of, are formed by four atomic orbital lobes interact with four atomic orbital lobes, and Phi bonds from 6 to 6). True, pi orbitals are most likely to form between adjacent sets of p-orbitals in a side on fashion, but they may also form between a p orbital and a d orbital, or two d-orbitals. Just like sigma orbitals don't only form between two s-orbitals. This isn't a perfect definition because bonds aren't, strictly speaking, interactions only between pairs of atomic orbitals, but usually a single pair will dominate so using this definition for practical purposes won't lead you astray.

A more formal definition of the bond classification is via electronic symmetry, which is where the sigma/pi/delta designations come from - they have the symmetry analogous to that of their origin atomic orbitals (s, p, d, f). Pi bonds are best defined by the symmetry of electron density distribution with respect to the bond axis, where electron density lies above and below the bond axis, but with a node being on the bond axis.

[AWK]

like for example ClO3. Where do I make mistakes?

First of all, ClO₃ does not exist. Only the presence of such a compound is postulated temporarily in the reaction mechanisms. In the solid and liquid state, it exists as a red ionic compound ClO₂⁺ClO₄^- and in the gaseous state as Cl₂O₆ (O₂ClOClO₃). For such compounds, it is safer to talk about single and double bonds instead of the terms σ and π used in organic chemistry.

[ComradeLenin]

like for example ClO3. Where do I make mistakes?

First of all, ClO₃ does not exist. Only the presence of such a compound is postulated temporarily in the reaction mechanisms. In the solid and liquid state, it exists as a red ionic compound ClO₂⁺ClO₄^- and in the gaseous state as Cl₂O₆ (O₂ClOClO₃). For such compounds, it is safer to talk about single and double bonds instead of the terms σ and π used in organic chemistry.

Sorry, I meant ClO₃^-, not ClO₃

[ComradeLenin]

A practical definition of the bond classification may be the number of atomic orbital lobes that interact to form the bond. For example sigma bonds are defined by the interaction of one lobe on one atomic orbital with one (and only one) lobe on another atomic orbital. By contrast, pi bonds result from interaction of two atomic orbital lobes on one atomic orbital interacting with two atomic orbital lobes on another atomic orbital. (Just so, delta-bonds, which you've probably never heard of, are formed by four atomic orbital lobes interact with four atomic orbital lobes, and Phi bonds from 6 to 6). True, pi orbitals are most likely to form between adjacent sets of p-orbitals in a side on fashion, but they may also form between a p orbital and a d orbital, or two d-orbitals. Just like sigma orbitals don't only form between two s-orbitals. This isn't a perfect definition because bonds aren't, strictly speaking, interactions only between pairs of atomic orbitals, but usually a single pair will dominate so using this definition for practical purposes won't lead you astray.

A more formal definition of the bond classification is via electronic symmetry, which is where the sigma/pi/delta designations come from - they have the symmetry analogous to that of their origin atomic orbitals (s, p, d, f). Pi bonds are best defined by the symmetry of electron density distribution with respect to the bond axis, where electron density lies above and below the bond axis, but with a node being on the bond axis.

What a concise but clear reply, thank you very much