Stability of matter

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TL;DR Summary
Water molecule not breaking up upon measurement
This is going to be a rather simple question about the understanding of covalent bonds.

H2-Molecule.png


Let's take the simplest molecule - H2 which is gas.
According to quantum mechanics, the two atoms in the molecule each share 1 electron with the other atom and those 2 electrons exist in the form of electron cloud(orbitals).
This sharing of electrons is what keeps the atoms of the molecule together and makes the hydrogen gas stable.

According to QT, if an electron's range of positions is measured, the electron takes a single, definite discreet value, i.e. it is found at a specific location where the amplitude of its wavefunction is greatest.

So, the relevant question is if the theory is right, why isn't this covalent bond that keeps the 2 H atoms together(by sharing 2 electrons) breaking up when I observe, measure or interact with such gas? Or with water, which is two atoms of H and 1 atom O?
Upon measurement, the 2 shared electrons will eventually be found within the atom that they originally came from, thereby separating and breaking the covalent bond.
But this never happens in practice and water and hydrogen gas continue to be water and hydrogen gas(and not separate, independent atoms).

Even if we continuously measured the shared electrons of molecules of water and H2 gas(with photons or other electrons or other particles), the covalent bond between the molecule would never cease. Is there any other process that keeps atoms and molecules of matter together?
 
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PeroK
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Let's take the simplest molecule - H2 which is gas.
According to quantum mechanics, the two atoms in the molecule each share 1 electron with the other atom and those 2 electrons exist in the form of electron cloud(orbitals).
This sharing of electrons is what keeps the atoms of the molecule together and makes the hydrogen gas stable.

According to QT, if an electron's range of positions is measured, the electron takes a single, definite discreet value, i.e. it is found at a specific location where the amplitude of its wavefunction is greatest.

So, the relevant question is if the theory is right, why isn't this covalent bond that keeps the 2 H atoms together(by sharing 2 electrons) breaking up when I observe, measure or interact with such gas? Or with water, which is two atoms of H and 1 atom O?
Upon measurement, the 2 shared electrons will eventually be found within the atom that they originally came from, thereby separating and breaking the covalent bond.
The Hydrogen molecule is a four-particle system. A measurement of electron position would disrupt the molecule from its ground state into a superposition of energy states. This is other side of the measurement process: that it generally disrupts the microscopic system you are measuring.

A measurement of electron position, however, does not automatically break the covalent bond unless you supply enough energy to break the bond. Simply gaining some information about where the electrons are does not in itself reduce the four-particle system to two two-particle systems.

Moreover, the electrons are indistinguishable, so there is no sense in which you can identify the electron that belongs to each proton.
 
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A long digression started by a misunderstanding of chemical bonding has been removed from this thread.
 
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I found what I was looking for. With hydrogen gas and water, the force that keeps the atoms together is not a covalent bond with shared electrons, but molecular dipoles. The force of attraction is generated by electric dipoles. This gives rise to the peculiar setup of molecules with atoms being arranged at very specific angles to each other. This arrangement provides the different chemical properties of the elements.
 
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According to quantum mechanics, the two atoms in the molecule each share 1 electron with the other atom and those 2 electrons exist in the form of electron cloud(orbitals).
This sharing of electrons is what keeps the atoms of the molecule together and makes the hydrogen gas stable...why isn't this covalent bond that keeps the 2 H atoms together(by sharing 2 electrons) breaking up when I observe, measure or interact with such gas?
That is about as good a description of chemical bonding as can be done without the math, but like most math-free descriptions it is too simplified to build on.

We have a quantum system made up of two electrons and two protons. Because the potential energy of the system is lower when they are bound together as a hydrogen molecule than when the protons are more widely separated, the bound configuration is stable. A measurement of the position of any of these particles necessarily adds energy to the system, but if the amount added is less than the binding energy the molecule stays together.
 
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According to QT, if an electron's range of positions is measured, the electron takes a single, definite discreet value, i.e. it is found at a specific location where the amplitude of its wavefunction is greatest.

This is wrong in two ways.

First, it is impossible to measure the position of a quantum object with infinite precision. So you don't get "a single, definite discrete value". You get a finite range of values.

Second, the actual measurement result does not have to be where the amplitude of the wave function is greatest. The wave function only gives probabilities. If you measure a large number of identically prepared systems, you will expect most of the results to be where the wave function is greatest, but some of them will not be.
 
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why isn't this covalent bond that keeps the 2 H atoms together(by sharing 2 electrons) breaking up when I observe, measure or interact with such gas?

Because you're not measuring the position of the electrons. More precisely, you're not having an interaction that results in localizing the electrons with enough precision to disrupt the structure of the molecule.

The size of the finite range in which an interaction can localize the position of the electrons is inversely proportional to the energy of interaction: more precise localization requires higher energy. If you compute the energy that would be required to localize an electron within a distance range significantly smaller than the size of the molecule, you will find that it is more than enough energy to ionize the molecule--to kick the electron out of the molecule altogether, which does indeed count as a "disruption" of the molecule. But ordinary observation of hydrogen gas does not do anything even close to that; the ordinary energy of interaction is much lower.
 
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I found what I was looking for.

Where? Please give a reference. Particularly since whatever source it is appears to be misleading you; see below.

With hydrogen gas and water, the force that keeps the atoms together is not a covalent bond with shared electrons, but molecular dipoles. The force of attraction is generated by electric dipoles.

This is just wrong as a description of how individual molecules are held together in either water or hydrogen gas. It is a (sort of) feasible description of how ionic bonds work, in compounds like salts; but water and hydrogen gas are not salts and the bonds in their molecules are not ionic bonds.

In a polar molecule such as water, the individual molecules, held together by covalent bonds, are indeed tiny electric dipoles, because the different types of atoms (hydrogen and oxygen in the case of water) have different electronegativities, so the electrons have a higher probability of being at one end of the molecule than the other (the oxygen end, in the case of water). So there is an attraction between different molecules in a large body of water, due to the dipoles, that affects its chemical properties (for example, water's high melting and boiling points and high heat capacity). But this does not change the fact that the two H and one O atoms in a single H2O molecule are held together by covalent bonds.

And hydrogen gas, H2, is not a polar molecule at all--both atoms obviously have the same electronegativity, so electrons have an equal probability of being at either end of the molecule and there is no dipole. So any talk of "electric dipoles" in hydrogen gas is simply wrong.

This gives rise to the peculiar setup of molecules with atoms being arranged at very specific angles to each other. This arrangement provides the different chemical properties of the elements.

This is nonsense. The bond angles in particular compounds have nothing to do with electric dipoles. They have to do with the geometries of molecular orbitals.
 

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