3 ms·

H4 is impossible under any temperature or pressure. A hydrogen atom has one, count 'em, one electron to bond with. And bonds require, at minimum, two electrons

H4 is impossible under any temperature or pressure. A hydrogen atom has one, count 'em, one electron to bond with. And bonds require, at minimum, two electrons shared between the two atoms. That's why H2 is pretty stable compared to monatomic H1. You'd have to rip up and throw out hundreds of years of chemistry for H4 to be possible. In other words, the computer spit out nonsense. This reminds me of the French guy in Holy Grail who giggles with his buddies "I told him we already got one.."

_ihaque 3y ago
H4 certainly sounds strange to me (and as other commenters have pointed out, this is a toy example, not a real molecule), but there are weird bond structures out there in real molecules that violate "at minimum, two electrons shared between the two atoms": https://en.wikipedia.org/wiki/Three-center_two-electron_bond https://en.wikipedia.org/wiki/Three-center_two-electron_bond. (The 3C-2E bonds in diborane are not linear, so that doesn't seem like what could be happening here.)
deleted 3y ago
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throwawaymaths 3y ago
This is wrong. For example, there are molecular ions H3+ and CH5+ Some of the most common ions in the universe
jcranmer 3y ago
> And bonds require, at minimum, two electrons shared between the two atoms. That's why H2 is pretty stable compared to monatomic H1. You'd have to rip up and throw out hundreds of years of chemistry for H4 to be possible. Actually, all you have to do is catch up with (checks Wikipedia) 100-year old research. Molecular orbital theory is needed to understand three-center two-electron bonds, and that dates back to the 1920s. However, the concept of "a bond is an electron pair" only predates that discovery by about a decade. It actually turns out that H₃⁺, the classic example of the three-center two-electron bond, predates both of those discoveries. It's also worth noting that chemistry itself as a recognizable field only dates back to maybe the 1760s, and things like bond structure only really start to be worked on in the mid-19th century, with most of our modern understanding dating from various parts in the 20th century. So that's hardly "hundreds of years of chemistry" that would need to be ripped up and thrown out. Let me guess, you never took any chemistry classes past gen chem in college? Molecular orbital theory and two-center three-electron bonds is bog standard organic chemistry (might even be first semester orgo, I'm not sure).
- mchannon 3y ago
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- pred_ 3y ago
  And just because it's exotic, that doesn't mean it's not useful to think about. My impression from some rather limited exposure to quantum chemistry is that the molecule hits a sweet spot where it's both not completely trivial to solve, yet within the capabilities of present-day quantum computers, meaning that it becomes a useful benchmark for algorithm development. See e.g. https://arxiv.org/abs/2303.07417 https://arxiv.org/abs/2303.07417 (which covers H₄ and LiH) and the references therein.
bobbylarrybobby 3y ago
Electrons are fluid. They can agglomerate wherever there's positive charge. I don't think it's ridiculous that there is an electron configuration that permits H4 to exist.
echelon 3y ago
H4 can transiently exist under some extreme initial conditions, it's just not probable and it wouldn't exist for long as it evolves into molecular hydrogen.

_ihaque 3y ago

H4 certainly sounds strange to me (and as other commenters have pointed out, this is a toy example, not a real molecule), but there are weird bond structures out there in real molecules that violate "at minimum, two electrons shared between the two atoms": https://en.wikipedia.org/wiki/Three-center_two-electron_bond https://en.wikipedia.org/wiki/Three-center_two-electron_bond. (The 3C-2E bonds in diborane are not linear, so that doesn't seem like what could be happening here.)

deleted 3y ago

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throwawaymaths 3y ago

This is wrong. For example, there are molecular ions H3+ and CH5+ Some of the most common ions in the universe

jcranmer 3y ago

> And bonds require, at minimum, two electrons shared between the two atoms. That's why H2 is pretty stable compared to monatomic H1. You'd have to rip up and throw out hundreds of years of chemistry for H4 to be possible. Actually, all you have to do is catch up with (checks Wikipedia) 100-year old research. Molecular orbital theory is needed to understand three-center two-electron bonds, and that dates back to the 1920s. However, the concept of "a bond is an electron pair" only predates that discovery by about a decade. It actually turns out that H₃⁺, the classic example of the three-center two-electron bond, predates both of those discoveries. It's also worth noting that chemistry itself as a recognizable field only dates back to maybe the 1760s, and things like bond structure only really start to be worked on in the mid-19th century, with most of our modern understanding dating from various parts in the 20th century. So that's hardly "hundreds of years of chemistry" that would need to be ripped up and thrown out. Let me guess, you never took any chemistry classes past gen chem in college? Molecular orbital theory and two-center three-electron bonds is bog standard organic chemistry (might even be first semester orgo, I'm not sure).

mchannon 3y ago
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pred_ 3y ago
And just because it's exotic, that doesn't mean it's not useful to think about. My impression from some rather limited exposure to quantum chemistry is that the molecule hits a sweet spot where it's both not completely trivial to solve, yet within the capabilities of present-day quantum computers, meaning that it becomes a useful benchmark for algorithm development. See e.g. https://arxiv.org/abs/2303.07417 https://arxiv.org/abs/2303.07417 (which covers H₄ and LiH) and the references therein.

bobbylarrybobby 3y ago

Electrons are fluid. They can agglomerate wherever there's positive charge. I don't think it's ridiculous that there is an electron configuration that permits H4 to exist.

echelon 3y ago

H4 can transiently exist under some extreme initial conditions, it's just not probable and it wouldn't exist for long as it evolves into molecular hydrogen.