Energy jump

[steevey]

So say I have an electron from carbon sharing a bond with Hydrogen. The electron from the carbon atom has enough energy to exist in the second energy level, so what happens to the energy when it fills the shell of the first energy level of the hydrogen atom? How does the cycle continue? Shouldn't the electron lose energy when it goes to the hydrogen atom's first energy state?

[mississippichem]

So say I have an electron from carbon sharing a bond with Hydrogen. The electron from the carbon atom has enough energy to exist in the second energy level, so what happens to the energy when it fills the shell of the first energy level of the hydrogen atom? How does the cycle continue? Shouldn't the electron lose energy when it goes to the hydrogen atom's first energy state?

 

Molecular orbitals are the result of a linear combination of atomic orbitals. If the bond is formed homolytically (by combination of H and C radicals), then an electron in a 1s hydrogen orbital spin pairs with an electron in a 2s carbon orbital to form singly degenerate orbital with [math]a_{1}[/math] symmetry. This [math] 1s-2s\sigma [/math] orbital is lower in energy than either of the atomic orbitals it came from.

 

This is not the case for all molecular orbitals though. Methane for example has higher energy orbitals,[math] \sigma^{*}(a_{1}^{*})[/math], and [math] \pi^{*}(t_{2}^{*})[/math] that are higher in energy than the corresponding atomic orbitals. These orbitals are unoccupied in the ground state and are anti-bonding.

[lemur]

Does the energy level of the molecular bond determine the amount of energy needed to break it (i.e. break the molecule into constituent fragments)?

[mississippichem]

Does the energy level of the molecular bond determine the amount of energy needed to break it (i.e. break the molecule into constituent fragments)?

 

Yes. Sometimes there are ionic forces to deal with, but bond enthalpy (the energy of the bond) is by far the determining factor for the energy required to break the bond.

[steevey]

Molecular orbitals are the result of a linear combination of atomic orbitals. If the bond is formed homolytically (by combination of H and C radicals), then an electron in a 1s hydrogen orbital spin pairs with an electron in a 2s carbon orbital to form singly degenerate orbital with [math]a_{1}[/math] symmetry. This [math] 1s-2s\sigma [/math] orbital is lower in energy than either of the atomic orbitals it came from.

 

This is not the case for all molecular orbitals though. Methane for example has higher energy orbitals,[math] \sigma^{*}(a_{1}^{*})[/math], and [math] \pi^{*}(t_{2}^{*})[/math] that are higher in energy than the corresponding atomic orbitals. These orbitals are unoccupied in the ground state and are anti-bonding.

 

 

So the electron more or less averages out to a single orbital that it can share between both atoms?

 

 

What does the electron configuration look like with such a thing?

[mississippichem]

So the electron more or less averages out to a single orbital that it can share between both atoms?

 

Yes, but there are two electrons per single bond.

 

What does the electron configuration look like with such a thing?

 

 

This is an image of the [math]a_{1}[/math] and [math]t_2[/math] states for methane.

[steevey]

Yes, but there are two electrons per single bond.

 

 

 

 

This is an image of the [math]a_{1}[/math] and [math]t_2[/math] states for methane.

 

What about in an ionic bond?

 

Also also, why does the electron occupy a low region for n between C-H, but all the covalent bonds in these occupy high regions?

Edited February 2, 2011 by steevey

[mississippichem]

What about in an ionic bond?

 

Ionic bonds aren't really bonds. The preferred term is ionic interaction. Ionic interactions are just coulombic forces between positively and negatively charged ions. There is no real localized bond or shared electron pair.

 

Though there can be a significant degree of ionic interaction between atoms that are already covalently bonded to each other. This comes from different effective nuclear charges which in turn leads to "unequal sharing" of electrons. Don't take that analogy too far though; unequal sharing really means that one of the atoms just gets more of the wave function localized around it.

 

Also also, why does the electron occupy a low region for n between C-H, but all the covalent bonds in these occupy high regions?

 

I'm not sure I understand this question... Low region? High region?

Edited February 2, 2011 by mississippichem

[steevey]

. This [math] 1s-2s\sigma [/math] orbital is lower in energy than either of the atomic orbitals it came from.

 

But the "lower energy level" doesn't seem to fit the electron configuration of any of the other bonds.

[mississippichem]

But the "lower energy level" doesn't seem to fit the electron configuration of any of the other bonds.

 

That's why all the other orbitals in methane are anti-bonding. Meaning that they are higher in energy than their constituent atomic orbitals. They decrease the stability of said compound.