why there are no quadruple bonds existing between hydrocarbons

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Alkane is single bond, alkene is double bond, alkyne is triple bond, but why’s there no quadruple bond?

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9 Answers

Anonymous 0 Comments

Because then all the four available bonds would be used up holding the carbons together, leaving no space for the hydrogen to attach. It’d no longer be a hydrocarbon, it’d be diatomic carbon. There’s some evidence that this can happen (I found a paper, but couldn’t access it), but it’s not at all stable.

Anonymous 0 Comments

Two reasons:

First: if you have a quadruple bond, you don’t have a hydrocarbon, because all four bonds on both carbon atoms are to one another. This would just be a hypothetical C2 molecule with no hydrogen whatsoever.

More fundamentally: carbon simply doesn’t form quadruple bonds with itself or, to my knowledge, with any other element. The exact reasons why get into the details of molecular orbitals, but we can TLDR it a bit. (This is oversimplifying, and depending on how you frame things, you can sooooorta-kinda claim C2 with a quadruple bond exists, but the fourth bond is so weak that it’s hard to “count” it.)

There are four electrons available for bonding in each carbon atom. In the ground state, two occupy the spherical 2s orbital surrounding the nucleus, and two occupy (some superposition of) the three 2p orbitals forming dumbbells oriented around axes at 90 degree angles to one another. To create bonding orbitals in a molecule, you need these orbitals to overlap.

In single-bonded carbon, this is easy: overlap some pair of the 2p orbitals on the same axis. For double-bonded carbon, you overlap two on the same axis (imagine this as the x-axis) between the atoms (a sigma bond) and two of the orbitals at a 90 degree angle to that (a pi bond). Triple bond, same deal: you have one sigma bond between the two carbons, and two pi bonds in the two planes at a right angle to the sigma bond.

But the geometry doesn’t work out for a quadruple bond. To get a quadruple bond, you need to get one of the *s* electrons involved. To do that, you need to change the shape of the carbon atom, because the *s* electrons “can’t reach” on their own (the *p* orbital dumbbells reach further than the *s* orbital spherical cloud). In technical terms, you need to sp^3 hybridize the carbon atom’s orbitals, changing the 2s and three 2p orbitals into four symmetrically-distributed sp^3 orbitals around the carbon, forming a sort of pyramid shape. And there’s no way to get all four of them to overlap.

This isn’t unique to carbon, and it’s why quadruple or higher bonds are exceptionally rare (and mostly only occur in metal compounds, which have more orbitals to work with courtesy of the *d* and eventually *f* orbitals).

Anonymous 0 Comments

Because then all the four available bonds would be used up holding the carbons together, leaving no space for the hydrogen to attach. It’d no longer be a hydrocarbon, it’d be diatomic carbon. There’s some evidence that this can happen (I found a paper, but couldn’t access it), but it’s not at all stable.

Anonymous 0 Comments

Two reasons:

First: if you have a quadruple bond, you don’t have a hydrocarbon, because all four bonds on both carbon atoms are to one another. This would just be a hypothetical C2 molecule with no hydrogen whatsoever.

More fundamentally: carbon simply doesn’t form quadruple bonds with itself or, to my knowledge, with any other element. The exact reasons why get into the details of molecular orbitals, but we can TLDR it a bit. (This is oversimplifying, and depending on how you frame things, you can sooooorta-kinda claim C2 with a quadruple bond exists, but the fourth bond is so weak that it’s hard to “count” it.)

There are four electrons available for bonding in each carbon atom. In the ground state, two occupy the spherical 2s orbital surrounding the nucleus, and two occupy (some superposition of) the three 2p orbitals forming dumbbells oriented around axes at 90 degree angles to one another. To create bonding orbitals in a molecule, you need these orbitals to overlap.

In single-bonded carbon, this is easy: overlap some pair of the 2p orbitals on the same axis. For double-bonded carbon, you overlap two on the same axis (imagine this as the x-axis) between the atoms (a sigma bond) and two of the orbitals at a 90 degree angle to that (a pi bond). Triple bond, same deal: you have one sigma bond between the two carbons, and two pi bonds in the two planes at a right angle to the sigma bond.

But the geometry doesn’t work out for a quadruple bond. To get a quadruple bond, you need to get one of the *s* electrons involved. To do that, you need to change the shape of the carbon atom, because the *s* electrons “can’t reach” on their own (the *p* orbital dumbbells reach further than the *s* orbital spherical cloud). In technical terms, you need to sp^3 hybridize the carbon atom’s orbitals, changing the 2s and three 2p orbitals into four symmetrically-distributed sp^3 orbitals around the carbon, forming a sort of pyramid shape. And there’s no way to get all four of them to overlap.

This isn’t unique to carbon, and it’s why quadruple or higher bonds are exceptionally rare (and mostly only occur in metal compounds, which have more orbitals to work with courtesy of the *d* and eventually *f* orbitals).

Anonymous 0 Comments

Because then all the four available bonds would be used up holding the carbons together, leaving no space for the hydrogen to attach. It’d no longer be a hydrocarbon, it’d be diatomic carbon. There’s some evidence that this can happen (I found a paper, but couldn’t access it), but it’s not at all stable.

Anonymous 0 Comments

Two reasons:

First: if you have a quadruple bond, you don’t have a hydrocarbon, because all four bonds on both carbon atoms are to one another. This would just be a hypothetical C2 molecule with no hydrogen whatsoever.

More fundamentally: carbon simply doesn’t form quadruple bonds with itself or, to my knowledge, with any other element. The exact reasons why get into the details of molecular orbitals, but we can TLDR it a bit. (This is oversimplifying, and depending on how you frame things, you can sooooorta-kinda claim C2 with a quadruple bond exists, but the fourth bond is so weak that it’s hard to “count” it.)

There are four electrons available for bonding in each carbon atom. In the ground state, two occupy the spherical 2s orbital surrounding the nucleus, and two occupy (some superposition of) the three 2p orbitals forming dumbbells oriented around axes at 90 degree angles to one another. To create bonding orbitals in a molecule, you need these orbitals to overlap.

In single-bonded carbon, this is easy: overlap some pair of the 2p orbitals on the same axis. For double-bonded carbon, you overlap two on the same axis (imagine this as the x-axis) between the atoms (a sigma bond) and two of the orbitals at a 90 degree angle to that (a pi bond). Triple bond, same deal: you have one sigma bond between the two carbons, and two pi bonds in the two planes at a right angle to the sigma bond.

But the geometry doesn’t work out for a quadruple bond. To get a quadruple bond, you need to get one of the *s* electrons involved. To do that, you need to change the shape of the carbon atom, because the *s* electrons “can’t reach” on their own (the *p* orbital dumbbells reach further than the *s* orbital spherical cloud). In technical terms, you need to sp^3 hybridize the carbon atom’s orbitals, changing the 2s and three 2p orbitals into four symmetrically-distributed sp^3 orbitals around the carbon, forming a sort of pyramid shape. And there’s no way to get all four of them to overlap.

This isn’t unique to carbon, and it’s why quadruple or higher bonds are exceptionally rare (and mostly only occur in metal compounds, which have more orbitals to work with courtesy of the *d* and eventually *f* orbitals).

Anonymous 0 Comments

The short answer is that it’s the wrong shape. If you were to imagine the orbitals of carbon forming a 3D shape, it would be a triangular pyramid (aka a d4 if you’re into dice). Bonding wise, one electron can be considered to live at each point of the pyramid. It’s easy to see how you could align two of these so that one, two, or three of the points are able to touch at the same time, but there is no way to make it work so that all 4 points on both pyramids are touching.

Another comment goes into the more involved description about orbitals and hybridization and so on, but it really just boils down to carbon isn’t the right shape.

Anonymous 0 Comments

The short answer is that it’s the wrong shape. If you were to imagine the orbitals of carbon forming a 3D shape, it would be a triangular pyramid (aka a d4 if you’re into dice). Bonding wise, one electron can be considered to live at each point of the pyramid. It’s easy to see how you could align two of these so that one, two, or three of the points are able to touch at the same time, but there is no way to make it work so that all 4 points on both pyramids are touching.

Another comment goes into the more involved description about orbitals and hybridization and so on, but it really just boils down to carbon isn’t the right shape.

Anonymous 0 Comments

The short answer is that it’s the wrong shape. If you were to imagine the orbitals of carbon forming a 3D shape, it would be a triangular pyramid (aka a d4 if you’re into dice). Bonding wise, one electron can be considered to live at each point of the pyramid. It’s easy to see how you could align two of these so that one, two, or three of the points are able to touch at the same time, but there is no way to make it work so that all 4 points on both pyramids are touching.

Another comment goes into the more involved description about orbitals and hybridization and so on, but it really just boils down to carbon isn’t the right shape.