We know that electrons are charges that revolve around the nucleus. Then, when in covalent bonds the electron is shared; does the electron obey the rule?

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    $\begingroup$ We don't know that electrons "revolve" around the nucleus. That is a pre-quantum mechanics view of atomic structure. $\endgroup$
    – matt_black
    Commented Feb 3, 2018 at 13:12
  • $\begingroup$ @matt_black Not exactly, it may be not compatible with Copenhagen interpretation, but it's not so bad for Bohmian mechanics. $\endgroup$
    – Mithoron
    Commented Dec 7, 2018 at 22:04

3 Answers 3


I am guessing that you have an understanding of the atomic orbitals of an atom such as 1s and 3p atomic orbitals. Atomic orbitals are the probability distribution of where an electron is going to be $90$% of the time. Technically, electrons don't revolve around the nucleus but are rather quantum particles that are superimposed in several positions at the same time. However the description of electrons revolving randomly around the nucleus is suffice at this level.

When a covalent bond is formed, the electrons do still continue to revolve around the nucleus in orbitals, however a new kind of orbital is formed which are called molecular orbitals (rather than atomic orbitals). Here is a simple explanation of what molecular orbitals are.

Molecular orbitals are formed by constructively and destructively overlapping the atomic orbitals together. The simplest example of this is when $\ce{H2}$ is formed. Here, each H atom has a 1s atomic orbital. When they form, these atomic orbitals form 2 molecular orbitals. Where one molecular orbital is due to constructively overlapping the 2 atomic orbitals and the other molecular orbital is due to destructively overlapping the 2 atomic orbitals. This shown in the diagram below:

enter image description here

As you see that these 2 molecular orbitals that are formed are quite different to each other. The MO that is formed due to constructive overlap has greater electron density between the 2 positive nucleus. Meanwhile the MO that is formed due to destructive overlapping has decreases electron density between the 2 positive nucleus and in fact have a node between them. Therefore the first MO is lower in energy than the second MO.

Like AOs, the same rules apply to MOs. Each MO is able to hold 2 electrons and the electrons fill the lowest energy orbital first. So in $\ce{H2}$ the 2 electrons will fill the first MO and will revolve around the nucleus in that orbital (well technically only for $90$% of the time).

This theory can be applied to any compound, however it gets more complicated as the number of AOs increases and the number of atoms increases. However essentially, when atoms form covalent bonds, the electrons continue to revolve around the nucleus in MOs rather than AOs.


Electrons are essentially quantum particles, and are controlled by the laws of quantum mechanics. There is absolutely no way to reconcile these laws with our physical intuition of macroscopic world. Electrons do not revolve; rather, they occupy certain atomic orbitals. To put it simply, they just stay there, delocalized over some space around the nucleus, everywhere at once. In a covalently bound molecule, electrons would occupy molecular orbitals and thus get delocalized all over the molecule.

The description of electrons as tiny charged solid particles that revolve around the nucleus, much like planets revolve around the sun, is the approximation known as Bohr atom. Significant for historical reasons as it is, one should not mistake it for coherent description of electron's behavior. Also, I doubt if you can extend it consistently to model covalent bonds.


The best answer is that the electron does not "revolve" around or "orbit" the nucleus the same way that the planets orbit the sun.

Pictures like the following (from Wikipedia user Halfdan) are compelling, but refer to a model of the atom that was beginning to be supplanted by quantum mechanics a century ago.

enter image description here

The modern quantum-mechanical model of the atom incorporates things like the Heisenberg Uncertainty Principle and wave-particle duality. We now consider electrons in atoms to behave much more like waves. We can determine their energy very precisely, but as a consequence of the uncertainty principle, we cannot determine their position or momentum very precisely.

The atomic orbital pictures we are used to seeing for $s$, $p$, $d$, etc. orbitals are constructs describing regions of space where the electron is. Typically, we plot the 90% surface. In one interpretation that means 90% of the time the electron can be found in that region of space. In another interpretation, 90% of the electron-wave is confined to that space.

When atoms come together to form bonds, the electron-waves overlap to produce new orbitals shared by the atoms. Nanoputian's answer provides a graphical representation of how this works.

  • $\begingroup$ There was no much advance in quantum mechanics “a century ago”. Planck’s patches to save physics from the UV catastrophe? Einstein’s theory of photoelectric effect? Even first Schrödinger’s successes were as late as in 1920. What do you refer to? $\endgroup$ Commented Sep 7, 2015 at 16:25
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    $\begingroup$ @IncnisMrsi 1920 is almost a century ago - I think this is only indicated as a rough timescale. $\endgroup$
    – bon
    Commented Sep 7, 2015 at 16:56
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    $\begingroup$ The Bohr model of the atom, which included quantized energy levels (but also still included electrons as particles) replaced the model in my picture in 1913, and it was quickly apparent that it only worked for hydrogen. So, a century ago, classical models of the atom were already on their way out.\ $\endgroup$
    – Ben Norris
    Commented Sep 7, 2015 at 23:06

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