Hi Chemists,
I have been trying to get an answer to the following question over on the physics side of the forum:
In a chemical bond - one involving the sharing of electrons - at which energy level do the electrons reside if they were in 2 different levels in their unbonded state. For example, in Hydrogen the electron is at the lowest energy level while in Carbon the outer "shell" is at the second level - so when these elements bond to form methane - CH4 - do the shared electrons have the energy level of the inner shell or that of the second shell? - Please excuse use of the term "shell" - it´s just for ease of communicating the question."
Any ideas?
regards
edward
QUOTE (Edward 3+Feb 27 2009, 10:22 AM)
Hi Chemists,
I have been trying to get an answer to the following question over on the physics side of the forum:
In a chemical bond - one involving the sharing of electrons - at which energy level do the electrons reside if they were in 2 different levels in their unbonded state. For example, in Hydrogen the electron is at the lowest energy level while in Carbon the outer "shell" is at the second level - so when these elements bond to form methane - CH4 - do the shared electrons have the energy level of the inner shell or that of the second shell? - Please excuse use of the term "shell" - it´s just for ease of communicating the question."
Any ideas?
regards
edward
First of all, let me say that I hate chemistry. I like physics, where things are simpler.
It's always the outer electron which is shared.
Oxygen, for instance, has two energy levels and eight electrons. Hydrogen has just one energy level, but the shared electron occupys the outermost level in both atoms, the first level for H and the second energy level for O.
That is, if I've remembered my chemical valences correctly.
I hate chemistry.
I have been trying to get an answer to the following question over on the physics side of the forum:
In a chemical bond - one involving the sharing of electrons - at which energy level do the electrons reside if they were in 2 different levels in their unbonded state. For example, in Hydrogen the electron is at the lowest energy level while in Carbon the outer "shell" is at the second level - so when these elements bond to form methane - CH4 - do the shared electrons have the energy level of the inner shell or that of the second shell? - Please excuse use of the term "shell" - it´s just for ease of communicating the question."
Any ideas?
regards
edward
First of all, let me say that I hate chemistry. I like physics, where things are simpler.
It's always the outer electron which is shared.
Oxygen, for instance, has two energy levels and eight electrons. Hydrogen has just one energy level, but the shared electron occupys the outermost level in both atoms, the first level for H and the second energy level for O.
That is, if I've remembered my chemical valences correctly.
I hate chemistry.
Hi Alex,
Many thanks - let me explain where I´m coming from on this - I had assumed that the bonding electron occupied the inner "shell" in the H atom but the second shell in the O atom. I agree that these can be regarded as the outermost shell in each atom - both the H and the O, but is the outermost shell of the H atom not at the lowest permitted energy level while the outermost shell in the O atom is at the second lowest permitted level? This then leads on to my problem as to what energy level/shell is occupied by a shared electron.
By the way, I like your definition of "simplicity" !!!
regards
edward
Many thanks - let me explain where I´m coming from on this - I had assumed that the bonding electron occupied the inner "shell" in the H atom but the second shell in the O atom. I agree that these can be regarded as the outermost shell in each atom - both the H and the O, but is the outermost shell of the H atom not at the lowest permitted energy level while the outermost shell in the O atom is at the second lowest permitted level? This then leads on to my problem as to what energy level/shell is occupied by a shared electron.
By the way, I like your definition of "simplicity" !!!
regards
edward
An energy level is not a physical location. It is one of the allowed energies of a system.
When H and O atoms (separate systems) come together (new combined system) the new system has different energy levels including the one that forms the atomic bond. Because this energy level is preferred, we know that it is lower than the average of the original system's original outermost electrons.
Another example: Detailed quantum chemistry calculations of H₂shows it has two molecular orbitals energy levels. A lower energy level which forms the H-H bond, which (like all orbitals) can hold two electrons, and an upper energy level which is anti-bonding. So the H₂system has different energy levels than atomic hydrogen.
When H and O atoms (separate systems) come together (new combined system) the new system has different energy levels including the one that forms the atomic bond. Because this energy level is preferred, we know that it is lower than the average of the original system's original outermost electrons.
Another example: Detailed quantum chemistry calculations of H₂shows it has two molecular orbitals energy levels. A lower energy level which forms the H-H bond, which (like all orbitals) can hold two electrons, and an upper energy level which is anti-bonding. So the H₂system has different energy levels than atomic hydrogen.
Edward 3,
It really depends on what theory you are using to describe the bonding. If you use molecular orbital theory then you use these atomic orbitals, those around the individual atoms, to compute an orbital that encompasses the whole molecule. Where there are high probabilities of finding these delocalized electrons between two nuclei is where you have the bonding. In valence bond, which is what people are most familiar with, the bonding happens because of overlap in the orbitals. The manner in which the overlap occurs gives the characteristics of the bond. In either case I think to think of it as shared electrons makes this confusing. Rather think of it as creating places of higher electron probability. I don't know if this answers your question. Trippy would explain this better than I and correct anything stupid I have said.
Regards,
Caleb
It really depends on what theory you are using to describe the bonding. If you use molecular orbital theory then you use these atomic orbitals, those around the individual atoms, to compute an orbital that encompasses the whole molecule. Where there are high probabilities of finding these delocalized electrons between two nuclei is where you have the bonding. In valence bond, which is what people are most familiar with, the bonding happens because of overlap in the orbitals. The manner in which the overlap occurs gives the characteristics of the bond. In either case I think to think of it as shared electrons makes this confusing. Rather think of it as creating places of higher electron probability. I don't know if this answers your question. Trippy would explain this better than I and correct anything stupid I have said.
Regards,
Caleb
rpenner beat me to it. And his response is better so read that one.
Hi Caleb, RP,
Many thanks for that - from what you say I am taking it that the energy level of the bonded electron - insofar as it can be defined in these terms - is somewhere between the inner and second shell levels and closer to the inner shell level. I had not realised that there were permitted levels between the inner and outer shells.
Both your responses were most helpful.
regards
edward
Many thanks for that - from what you say I am taking it that the energy level of the bonded electron - insofar as it can be defined in these terms - is somewhere between the inner and second shell levels and closer to the inner shell level. I had not realised that there were permitted levels between the inner and outer shells.
Both your responses were most helpful.
regards
edward
QUOTE (Edward 3+Feb 27 2009, 07:57 PM)
is somewhere between the inner and second shell levels
Not what I was saying.
Not what I was saying.
Hi RP,
I read what you said as meaning that the new "system" i.e. the molecule, had different energy levels from the original constituent atoms - you mean the electrons have different energy levels? No? Or is my error in saying that the energy level of the electrons in the new system is between the two levels of the original system?
regards
edward
I read what you said as meaning that the new "system" i.e. the molecule, had different energy levels from the original constituent atoms - you mean the electrons have different energy levels? No? Or is my error in saying that the energy level of the electrons in the new system is between the two levels of the original system?
regards
edward
Edward 3,
You misunderstood what I was saying. As far as MOT goes the molecular orbitals, what I believe rpenner refers to as "(new combined system) the new system", are constructed using a combination of the atomic orbitals. To reiterate what rpenner stated earlier as far as the molecular orbital energy level it is, as a whole, lower than the average of the atomic orbitals by themselves as the bonding has occurred.
Regards,
Caleb
PS I don't want to put words in your mouth so if what I stated is incorrect you are far more competent at his than I. I will defer to you.
You misunderstood what I was saying. As far as MOT goes the molecular orbitals, what I believe rpenner refers to as "(new combined system) the new system", are constructed using a combination of the atomic orbitals. To reiterate what rpenner stated earlier as far as the molecular orbital energy level it is, as a whole, lower than the average of the atomic orbitals by themselves as the bonding has occurred.
Regards,
Caleb
PS I don't want to put words in your mouth so if what I stated is incorrect you are far more competent at his than I. I will defer to you.
The Orbitron might be of some use.
Granouille: Thanks, I was just about to go digging for that link.
Caleb: Thanks for the vote of confidence.
Edward3:
Consider atomic hydrogen.
1 Electron, 1 proton.
The electron occupies the 1s orbital. The 1s orbital is spherically symmetrical, has no nodal surfaces, it is, in essence a 'solid sphere' with the probability of finding an elecron increasing as you approach the nucleus.
The orbital has a sign, or a phase, it can be + or -. The easiest way to visualize this is to visualize a plucked guitar string, it points up, or it points down (or it's flat).
Like wise, the 1s orbital has a phase, but, it must all be the same phase.
Now, think about 2 Hydrogen atoms sitting seperately.
There are two possible combinations - either both hydrogen atoms have the same phase, or both hydrogen atoms have opposing phases.
If we bring the two Hydrogen atoms together, and sit them beside each other at say 75pm (twice the covalent radius of Hydrogen) then their 1s orbitals are going to overlap, but there's two ways we can do this. We can do it so that the atomic orbitals are in phase with each other, and add to each other, or, we can do it so that they have opposing phases and cancel each other out.
As an analogy, consider two pulses travelling towards each other on a string. Either they can be of the same phase (or in the same direction) in which case, for that breif moment when they overlap, they add to each other. Or, they can be out of phase (opposite directions) at which point, during that breif moment when they overlap, the string is flat.
Something similar happens with the two Hydrogen atoms.
When the two orbitals are in phase with each other, they add up, and the shape of the molecular orbital is, in essence, an oval that's rotationally symmetric about the axis that the two hydrogen atoms are in, and the majority of the electron density lays between the two hydrogen atoms, and the molecular orbital is said to be a bonding orbital - because the electrons are concentrated where they hold they will hold the two nuclei together, they sheild the nuclei from each other. This molecular orbital, for this reason, has the property that it is actually of lower energy than a 1s orbital.
However, where the two orbitals are out of phase with each other, they cancel each other out where they over lap. The end result is kind of like two lobes, of opposing phases that surround each of the atoms, and a plane halfway between the two atoms where there is no electron density (a nodal plan). The majority of the electron density is concentrated 'outside' the molecule - if that makes sense, so the nuclei are exposed to each other, and the electron density is pulling them apart, because of this, the orbital is said to be antibonding, because, in essence, it has the effect of detabilizing the molecule, for this reason, this orbital has the property that it is actually of higher energy than the 1s orbital. This antibonding orbital is also rotationally symmetric along the axis that the hydrogen atoms are on.
Interstingly, the energy of the bonding orbital is lower than the energy of the 1s orbital by the same amount that the energy of the antibonding orbital is.
I hope this is helpful.
Caleb: Thanks for the vote of confidence.
Edward3:
Consider atomic hydrogen.
1 Electron, 1 proton.
The electron occupies the 1s orbital. The 1s orbital is spherically symmetrical, has no nodal surfaces, it is, in essence a 'solid sphere' with the probability of finding an elecron increasing as you approach the nucleus.
The orbital has a sign, or a phase, it can be + or -. The easiest way to visualize this is to visualize a plucked guitar string, it points up, or it points down (or it's flat).
Like wise, the 1s orbital has a phase, but, it must all be the same phase.
Now, think about 2 Hydrogen atoms sitting seperately.
There are two possible combinations - either both hydrogen atoms have the same phase, or both hydrogen atoms have opposing phases.
If we bring the two Hydrogen atoms together, and sit them beside each other at say 75pm (twice the covalent radius of Hydrogen) then their 1s orbitals are going to overlap, but there's two ways we can do this. We can do it so that the atomic orbitals are in phase with each other, and add to each other, or, we can do it so that they have opposing phases and cancel each other out.
As an analogy, consider two pulses travelling towards each other on a string. Either they can be of the same phase (or in the same direction) in which case, for that breif moment when they overlap, they add to each other. Or, they can be out of phase (opposite directions) at which point, during that breif moment when they overlap, the string is flat.
Something similar happens with the two Hydrogen atoms.
When the two orbitals are in phase with each other, they add up, and the shape of the molecular orbital is, in essence, an oval that's rotationally symmetric about the axis that the two hydrogen atoms are in, and the majority of the electron density lays between the two hydrogen atoms, and the molecular orbital is said to be a bonding orbital - because the electrons are concentrated where they hold they will hold the two nuclei together, they sheild the nuclei from each other. This molecular orbital, for this reason, has the property that it is actually of lower energy than a 1s orbital.
However, where the two orbitals are out of phase with each other, they cancel each other out where they over lap. The end result is kind of like two lobes, of opposing phases that surround each of the atoms, and a plane halfway between the two atoms where there is no electron density (a nodal plan). The majority of the electron density is concentrated 'outside' the molecule - if that makes sense, so the nuclei are exposed to each other, and the electron density is pulling them apart, because of this, the orbital is said to be antibonding, because, in essence, it has the effect of detabilizing the molecule, for this reason, this orbital has the property that it is actually of higher energy than the 1s orbital. This antibonding orbital is also rotationally symmetric along the axis that the hydrogen atoms are on.
Interstingly, the energy of the bonding orbital is lower than the energy of the 1s orbital by the same amount that the energy of the antibonding orbital is.
I hope this is helpful.
QUOTE (Trippy+Feb 27 2009, 05:27 PM)
...
However, where the two orbitals are out of phase with each other, they cancel each other out where they overlap. The end result is kind of like two lobes, of opposing phases that surround each of the atoms, and a plane halfway between the two atoms where there is no electron density (a nodal plan). The majority of the electron density is concentrated 'outside' the molecule - if that makes sense, so the nuclei are exposed to each other, and the electron density is pulling them apart, because of this, the orbital is said to be antibonding, because, in essence, it has the effect of destabilizing the molecule, for this reason, this orbital has the property that it is actually of higher energy than the 1s orbital. This antibonding orbital is also rotationally symmetric along the axis that the hydrogen atoms are on.
Interestingly, the energy of the bonding orbital is lower than the energy of the 1s orbital by the same amount that the energy of the antibonding orbital is.
I hope this is helpful.
Absolutely lovely!
I can see it now, and that helps. I thought I understood it visually before, but you done gone and fixed that.
Thanks much!
However, where the two orbitals are out of phase with each other, they cancel each other out where they overlap. The end result is kind of like two lobes, of opposing phases that surround each of the atoms, and a plane halfway between the two atoms where there is no electron density (a nodal plan). The majority of the electron density is concentrated 'outside' the molecule - if that makes sense, so the nuclei are exposed to each other, and the electron density is pulling them apart, because of this, the orbital is said to be antibonding, because, in essence, it has the effect of destabilizing the molecule, for this reason, this orbital has the property that it is actually of higher energy than the 1s orbital. This antibonding orbital is also rotationally symmetric along the axis that the hydrogen atoms are on.
Interestingly, the energy of the bonding orbital is lower than the energy of the 1s orbital by the same amount that the energy of the antibonding orbital is.
I hope this is helpful.
Absolutely lovely!
I can see it now, and that helps. I thought I understood it visually before, but you done gone and fixed that.
Thanks much!
Hi RP, Caleb, Granouille, Trippy,
Many thanks to all for those most helpful and enlightening responses. I clearly need to do a lot of work on my understanding of electron energy levels - it´s nothing like what my old chemistry teacher tried to reach me over 40 years ago.
best regards
edward
Many thanks to all for those most helpful and enlightening responses. I clearly need to do a lot of work on my understanding of electron energy levels - it´s nothing like what my old chemistry teacher tried to reach me over 40 years ago.
best regards
edward
QUOTE (Cusa+Mar 1 2009, 10:31 AM)
Fair question: How does one electron hold together 2 atoms. Should it not be a "shell" bond?
I propose that chemistry is about these "shell bonds."
Mitch Raemsch
You don't know what you're talking about. Stop spamming this thread and go away.
I propose that chemistry is about these "shell bonds."
Mitch Raemsch
You don't know what you're talking about. Stop spamming this thread and go away.
Hello everybody! I was just passing by...
A new chemical bond is just a new shell or orbital, different from what existed before. You may view it as the result of an energy optimization, which helps understand why electric fields, extreme pressure etc can distort it.
However, quantum theory also tells that the electron's wave function can be expressed as a linear combination of all the eigenfunctions, or shells if you prefer - but they are in infinite number.
Now, chemists need to have a description and a vocabulary to think and try to make forecasts about bonds. For that, chemical bonds which are in essence infinite combinations at best and whose form is known from a computer optimization are of no use. So chemists twist the reality by pretending bond to be simple combinations of few orbitals, using perturbation theory where it better shouldn't be since the perturbation is way too big.
This is where the track gets highly slippery. Because bonds are complex but chemists (or any usable theory) need manageable explanations, dozens of cases and exceptions and rules have been invented, like "covalent" or "sp3 hybridization" - which don't work very well and make chemistry complicated.
It's just an attempt to put simple ideas on a complicated world - a bit like rules for bissextile years. Admit that these models are essentially overstretched and make peace with them.
He who hates chemistry should have a look at quark theory. They would be happy to have shells there.
A new chemical bond is just a new shell or orbital, different from what existed before. You may view it as the result of an energy optimization, which helps understand why electric fields, extreme pressure etc can distort it.
However, quantum theory also tells that the electron's wave function can be expressed as a linear combination of all the eigenfunctions, or shells if you prefer - but they are in infinite number.
Now, chemists need to have a description and a vocabulary to think and try to make forecasts about bonds. For that, chemical bonds which are in essence infinite combinations at best and whose form is known from a computer optimization are of no use. So chemists twist the reality by pretending bond to be simple combinations of few orbitals, using perturbation theory where it better shouldn't be since the perturbation is way too big.
This is where the track gets highly slippery. Because bonds are complex but chemists (or any usable theory) need manageable explanations, dozens of cases and exceptions and rules have been invented, like "covalent" or "sp3 hybridization" - which don't work very well and make chemistry complicated.
It's just an attempt to put simple ideas on a complicated world - a bit like rules for bissextile years. Admit that these models are essentially overstretched and make peace with them.
He who hates chemistry should have a look at quark theory. They would be happy to have shells there.
I have a question: what creates the atomic shells?
Do protons create electron shells? Can a proton be shellless? like in an accelerator?
Mitch Raemsch
Do protons create electron shells? Can a proton be shellless? like in an accelerator?
Mitch Raemsch
Enthalpy,
My advisor put it best, he may not have been the originator, when he said. "Some theories are too good to be true while others are too true to be good."
Caleb
My advisor put it best, he may not have been the originator, when he said. "Some theories are too good to be true while others are too true to be good."
Caleb
Hi Friend,
Herre i discuss with you some chemical bond concept with nano tube. the Chemical bond referrs to the forces holding atoms together to form molecules and solids. This force is of electric nature, and the attraction between electrons of one atom to the nucleus of another atom contribute to what is known as chemical bonds. Although electrons of one atom repell electrons of another, but the repulsion is relatively small. So is the repulsion between atomic nuclei.
Herre i discuss with you some chemical bond concept with nano tube. the Chemical bond referrs to the forces holding atoms together to form molecules and solids. This force is of electric nature, and the attraction between electrons of one atom to the nucleus of another atom contribute to what is known as chemical bonds. Although electrons of one atom repell electrons of another, but the repulsion is relatively small. So is the repulsion between atomic nuclei.
Hi Friend,
Here i discuss with you some chemical bond concept with nano tube. the Chemical bond referrs to the forces holding atoms together to form molecules and solids. This force is of electric nature, and the attraction between electrons of one atom to the nucleus of another atom contribute to what is known as chemical bonds. Although electrons of one atom repell electrons of another, but the repulsion is relatively small. So is the repulsion between atomic nuclei.
Here i discuss with you some chemical bond concept with nano tube. the Chemical bond referrs to the forces holding atoms together to form molecules and solids. This force is of electric nature, and the attraction between electrons of one atom to the nucleus of another atom contribute to what is known as chemical bonds. Although electrons of one atom repell electrons of another, but the repulsion is relatively small. So is the repulsion between atomic nuclei.
QUOTE (Cusa+Mar 2 2009, 06:40 AM)
I have a question: what creates the atomic shells?
Do protons create electron shells? Can a proton be shellless? like in an accelerator?
Mitch Raemsch
Electron shells are simply descriptions of the most probably arrangement of electrons around a nucleus. The "pictures" of orbitals you see are arbitrarily defined, usually to include 95% or 99% of the electron density.
I believe under extreme conditions such as in a particle accelerator a hydrogen nucleus can be stripped of its electrons to form a naked proton but I'm not sure.
Do protons create electron shells? Can a proton be shellless? like in an accelerator?
Mitch Raemsch
Electron shells are simply descriptions of the most probably arrangement of electrons around a nucleus. The "pictures" of orbitals you see are arbitrarily defined, usually to include 95% or 99% of the electron density.
I believe under extreme conditions such as in a particle accelerator a hydrogen nucleus can be stripped of its electrons to form a naked proton but I'm not sure.
Not so extreme situations, also. At 6000K, a fraction of hydrogen atoms are completely ionized -- the Sun's visible surface is opaque because of this effect. The ionization energy of hydrogen is about 13.6 electron-Volts and Boltzmann's constant tells us this is the average energy of a gas particle at 0.158 million K. But the thin gasses of the Sun's corona at about 10 million K and the Sun's core is even hotter.
But state-of-the-art particle accelerators are much more extreme, effectively heating tiny bits of matter to a million million K or more.
But state-of-the-art particle accelerators are much more extreme, effectively heating tiny bits of matter to a million million K or more.
Heh... I'd consider the sun's surface to be 'extreme', but I guess that's because I'm used to dealing with nice cold room-temperature protons in a bottle. 
Thanks for the clarification, I didn't know that in any case.
Thanks for the clarification, I didn't know that in any case.
QUOTE (Edward 3+Feb 27 2009, 03:22 PM)
Hi Chemists,
I have been trying to get an answer to the following question over on the physics side of the forum:
In a chemical bond - one involving the sharing of electrons - at which energy level do the electrons reside if they were in 2 different levels in their unbonded state. For example, in Hydrogen the electron is at the lowest energy level while in Carbon the outer "shell" is at the second level - so when these elements bond to form methane - CH4 - do the shared electrons have the energy level of the inner shell or that of the second shell? - Please excuse use of the term "shell" - it´s just for ease of communicating the question."
Any ideas?
regards
edward
Hi Friend,
The purpose of this thread is to explain some of the Physics and Chemistry behind matter, which may prove potentially useful in cooling applications. I have seen several threads regarding various topics touching on Physics and Chemistry such as a thread where it was suggested some miriacle reaction could absorb all the heat from a CPU for months; threads on why water is used as a coolant and not something else; threads on what else we could make waterblocks and heatsinks from.
Thanks
Parkar
I have been trying to get an answer to the following question over on the physics side of the forum:
In a chemical bond - one involving the sharing of electrons - at which energy level do the electrons reside if they were in 2 different levels in their unbonded state. For example, in Hydrogen the electron is at the lowest energy level while in Carbon the outer "shell" is at the second level - so when these elements bond to form methane - CH4 - do the shared electrons have the energy level of the inner shell or that of the second shell? - Please excuse use of the term "shell" - it´s just for ease of communicating the question."
Any ideas?
regards
edward
Hi Friend,
The purpose of this thread is to explain some of the Physics and Chemistry behind matter, which may prove potentially useful in cooling applications. I have seen several threads regarding various topics touching on Physics and Chemistry such as a thread where it was suggested some miriacle reaction could absorb all the heat from a CPU for months; threads on why water is used as a coolant and not something else; threads on what else we could make waterblocks and heatsinks from.
Thanks
Parkar
As the person who started this thread, I´m not sure I would agree with what you define as its purpose - but carry on - the original question has been fairly well disposed of !!
Hello all.
I suppose I am still thinking of the whole thing in a classical point of view when I say that when it came to the sharing of electrons that one of the elements will attract the electron more. Can't remember the term for that off hand right now.
Example would be water where the oxygen, in a manner of speaking, steals the electrons from the two hydrogen. So the electrons would be in the 2nd energy level.
I suppose I am still thinking of the whole thing in a classical point of view when I say that when it came to the sharing of electrons that one of the elements will attract the electron more. Can't remember the term for that off hand right now.
Example would be water where the oxygen, in a manner of speaking, steals the electrons from the two hydrogen. So the electrons would be in the 2nd energy level.
Electronegativity, I think.
Electronegativity, giving rise to polar bonds.
One point to recognise is that 'covalent' and 'ionic' bonding are extremes on a spectrum. Very few bonds are either one or the other; the rest fall on this continuum and can be viewed as 'covalent with ionic character' or vice versa. Even the humble C-H bond, which is 'covalent', has some polarity and ionic character, though it's not pronounced enough to influence its reactivity much. But take the C=O bond - it is strongly polar and this dominates its reactivity under many conditions.
Kaeroll
One point to recognise is that 'covalent' and 'ionic' bonding are extremes on a spectrum. Very few bonds are either one or the other; the rest fall on this continuum and can be viewed as 'covalent with ionic character' or vice versa. Even the humble C-H bond, which is 'covalent', has some polarity and ionic character, though it's not pronounced enough to influence its reactivity much. But take the C=O bond - it is strongly polar and this dominates its reactivity under many conditions.
Kaeroll
I just like to add some interesting facts I found out this week.
Ionization of chemical species can be easier in condensed phase (e.g. aqueous solutions or mixtures) compared to gas phase.
Apparently the energy levels changed with the environment not being vacuum. And it made it easier to lose electrons.
In the molecule, the outer shells of the atoms are no longer what they were originally but become molecular orbitals.
Near other molecular orbitals, the rising of bands begins to become a hint.
The electrons once ionized in mixture, become "solvated". They can attach to form dissociating anions as well as detach easily.
Ionization of chemical species can be easier in condensed phase (e.g. aqueous solutions or mixtures) compared to gas phase.
Apparently the energy levels changed with the environment not being vacuum. And it made it easier to lose electrons.
In the molecule, the outer shells of the atoms are no longer what they were originally but become molecular orbitals.
Near other molecular orbitals, the rising of bands begins to become a hint.
The electrons once ionized in mixture, become "solvated". They can attach to form dissociating anions as well as detach easily.
QUOTE
Ionization of chemical species can be easier in condensed phase (e.g. aqueous solutions or mixtures) compared to gas phase.
This is because aqeous solutions are also Hydrogen ion acceptors/donators. This is why they're are able to transfer electrons. Also the spactial seperation of molecules is much closer.
QUOTE (->
| QUOTE |
| Ionization of chemical species can be easier in condensed phase (e.g. aqueous solutions or mixtures) compared to gas phase. |
This is because aqeous solutions are also Hydrogen ion acceptors/donators. This is why they're are able to transfer electrons. Also the spactial seperation of molecules is much closer.
In the molecule, the outer shells of the atoms are no longer what they were originally but become molecular orbitals.
Molecular orbitals are always molecular orbitals.
QUOTE
The electrons once ionized in mixture, become "solvated". They can attach to form dissociating anions as well as detach easily.
Electrons do not ionize, but molecules get ionized by accepting or loosing electrons.
Ionization happens all the time even in current photolithography, precisely in such environments: http://en.wikipedia.org/wiki/Photoresist
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