Sorted on P Orbitals and first ionisation energy?

[Paseo Del Norte]

Again, I find my self placing a question on the clinician forum because I am not actually preparing for any of the examinations covered on the other forums. (DAT, MCAT, USMLE and so on.)

My current adventures have me looking at some of the periodic trends. I am trying to wrap my head around the "first" ionisation energy concept. I understand that as we move across a period (left to right), you will have a generally experience an increase in electronegativity as the nuclear charge increases. I further understand that in multi-electron atoms you will have electron repulsion resulting in shielding and thus you have to look at effective nuclear charge (Z-effective).

However, I am experiencing a little confusion about some of the exceptions such as Oxygen. ( I am good with the jump you see with say Boron or Aluminium as you have a lone electron in a P orbital and a filled S orbital that is essentially "happy".) When considering Oxygen, I understand Hund's rule and the fact that as you fill orbitals, the electrons will fill with spin +1/2 or up, then fill with spin -1/2 or down. So, with oxygen, we have three spin ups filled with one orbital that is full with a spin up and a spin down.

I am a little confused as to how this results in a decrease in the first ionisation energy however. If it was a shielding issue, I would go along with it, but then why not have an even lower ionisation energy with Fluorine because it has even more electrons? So, clearly I think there is more to this than shielding?

Just to be clear, this is not a homework assignment and simply stuff I think about on my free time.

Thank you for any assistance.

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[ARAI]

Again, I find my self placing a question on the clinician forum because I am not actually preparing for any of the examinations covered on the other forums. (DAT, MCAT, USMLE and so on.)

My current adventures have me looking at some of the periodic trends. I am trying to wrap my head around the "first" ionisation energy concept. I understand that as we move across a period (left to right), you will have a generally experience an increase in electronegativity as the nuclear charge increases. I further understand that in multi-electron atoms you will have electron repulsion resulting in shielding and thus you have to look at effective nuclear charge (Z-effective).

However, I am experiencing a little confusion about some of the exceptions such as Oxygen. ( I am good with the jump you see with say Boron or Aluminium as you have a lone electron in a P orbital and a filled S orbital that is essentially "happy".) When considering Oxygen, I understand Hund's rule and the fact that as you fill orbitals, the electrons will fill with spin +1/2 or up, then fill with spin -1/2 or down. So, with oxygen, we have three spin ups filled with one orbital that is full with a spin up and a spin down.

I am a little confused as to how this results in a decrease in the first ionisation energy however. If it was a shielding issue, I would go along with it, but then why not have an even lower ionisation energy with Fluorine because it has even more electrons? So, clearly I think there is more to this than shielding?

Just to be clear, this is not a homework assignment and simply stuff I think about on my free time.

Thank you for any assistance.

Oxygen is the exception to ionization. I believe it has to do with the size of the atom. The two electrons in the 2px orbital (which are spinning opposite of each other) push away from each other with greater opposing force due to the small volume of an oxygen atom. This results in reduced electron affinity. So, its easier to remove a single electron from Oxygen than it is from Sulfur.

Fluorine is also quite small, so it follows the same pattern: it's easier to remove electrons from fluorine than it is from chlorine.

I think those are the two exceptions. Every other element in Row 2 should have greater electron affinity than it's corresponding element in Row 3.

 

[Paseo Del Norte]

Oxygen is the exception to ionization. I believe it has to do with the size of the atom. The two electrons in the 2px orbital (which are spinning opposite of each other) push away from each other with greater opposing force due to the small volume of an oxygen atom. This results in reduced electron affinity. So, its easier to remove a single electron from Oxygen than it is from Sulfur.

Fluorine is also quite small, so it follows the same pattern: it's easier to remove electrons from fluorine than it is from chlorine.

I think those are the two exceptions. Every other element in Row 2 should have greater electron affinity than it's corresponding element in Row 3.

Actually, Fluorine has a higher first ionisation energy than Chlorine. In addition, Oxygen has a higher first ionisation energy than Sulfur. As you move down in period; increase the "n" quantum number, the size of the atomic radius is more significant than the nuclear charge, generally resulting in lower binding energies and lower first ionisation energies.

 

[ARAI]

Actually, Fluorine has a higher first ionisation energy than Chlorine. In addition, Oxygen has a higher first ionisation energy than Sulfur. As you move down in period; increase the "n" quantum number, the size of the atomic radius is more significant than the nuclear charge, generally resulting in lower binding energies and lower first ionisation energies.

Right, the first ionization of Fluorine is higher than Chlorine; but the electron affinity of Fluorine is lower than Chlorine. And the Electron affinity of Oxygen is lower than Sulfur. And I now realize that's probably going to be the same for most of Row 2 vs. Row 3 (i.e. due to the small atomic size of elements in row 2, they will have less electron affinity than those in Row 3).

Still, as I recall, the reason I explained earlier regarding the opposing forces of electrons in the 2px orbital are behind the lower first ionization energy of Oxygen (as compared to Nitrogen).

Hope this helps. I took chemistry 10 years ago. I understand why this is happening, but I've lost the ability to explain it since I haven't talked about it in so long. Does that make sense?

 

[group_theory]

Again, I find my self placing a question on the clinician forum because I am not actually preparing for any of the examinations covered on the other forums. (DAT, MCAT, USMLE and so on.)

My current adventures have me looking at some of the periodic trends. I am trying to wrap my head around the "first" ionisation energy concept. I understand that as we move across a period (left to right), you will have a generally experience an increase in electronegativity as the nuclear charge increases. I further understand that in multi-electron atoms you will have electron repulsion resulting in shielding and thus you have to look at effective nuclear charge (Z-effective).

However, I am experiencing a little confusion about some of the exceptions such as Oxygen. ( I am good with the jump you see with say Boron or Aluminium as you have a lone electron in a P orbital and a filled S orbital that is essentially "happy".) When considering Oxygen, I understand Hund's rule and the fact that as you fill orbitals, the electrons will fill with spin +1/2 or up, then fill with spin -1/2 or down. So, with oxygen, we have three spin ups filled with one orbital that is full with a spin up and a spin down.

I am a little confused as to how this results in a decrease in the first ionisation energy however. If it was a shielding issue, I would go along with it, but then why not have an even lower ionisation energy with Fluorine because it has even more electrons? So, clearly I think there is more to this than shielding?

Just to be clear, this is not a homework assignment and simply stuff I think about on my free time.

Thank you for any assistance.

The reason for the decrease in ionization energy of Oxygen when compare to Nitrogen has to do with the paired p suborbital.

When you have paired p-orbital, there is some slight "repulsion" from electrons in the same orbital (although minimized by having different spin, the repulsion is still there). Nitrogen, with its 3 outer electrons in separate p orbitals, does not have to contend with this sort of electron repulsion. As a result of this slight repulsion, it is slightly easier to removed an electron from a paired orbital than from an unpaired orbital.

In addition, the loss of a 2p electron in Oxygen results in a half-filled subshell (like nitrogen), which is an arrangement of low energy, so the energy level of O⁺ + e^- is lower than expected, and the ionization level is low too. (this is also seen between phosphorus and sulfar, but less pronounced because their orbitals are more diffuse)

That's why Oxygen has a lower first ionization energy compare to Nitrogen. The same trend occurs with phosphorus and sulfur.

As you go from Oxygen to Fluorine to Neon, the electrons that you are trying to removed are occupied in paired orbitals. Although there is greater effective nuclear charge (and hence higher ionization energy level), the change in first ionization energy as you go from Oxygen to Fluorine to Neon is less than the change in ionization energy as you go from Boron to Carbon to Nitrogen (due to the repulsion).

 

[Paseo Del Norte]

Oxygen is the exception to ionization. I believe it has to do with the size of the atom. The two electrons in the 2px orbital (which are spinning opposite of each other) push away from each other with greater opposing force due to the small volume of an oxygen atom. This results in reduced electron affinity. So, its easier to remove a single electron from Oxygen than it is from Sulfur.

Fluorine is also quite small, so it follows the same pattern: it's easier to remove electrons from fluorine than it is from chlorine.

I think those are the two exceptions. Every other element in Row 2 should have greater electron affinity than it's corresponding element in Row 3.

I'm still not tracking. You stated earlier that it is easier to remove the electron from oxygen compared to sulfur; however, oxygen has a higher 1st energy than sulfur. In addition, Fluorine has a higher 1st ionization energy than Chlorine. Therefore, we are in totally opposite courts on the issue of first ionisation energy.

I'm with you on election affinity just not first ionization energy.

 

[ARAI]

I'm still not tracking. You stated earlier that it is easier to remove the electron from oxygen compared to sulfur; however, oxygen has a higher 1st energy than sulfur. In addition, Fluorine has a higher 1st ionization energy than Chlorine. Therefore, we are in totally opposite courts on the issue of first ionisation energy.

I'm with you on election affinity just not first ionization energy.

Oops! I was talking about affinity and ionization together and mixing up the concepts.

Yeah, it's easier to remove an electron from oxygen vs. nitrogen and I was attempting to say the same about sulfur vs. phosphate and somehow mixed in oxygen vs. sulfur & fluorine vs. chlorine. Don't mind me, gen chem was a long time ago.

 

[bradt9881]

These are things I can understand. But wrapping my brain around this antimatter stuff only confuses me, and makes me feel stupid. I guess I should have taken more than intro physics, hehe.

 

[Paseo Del Norte]

These are things I can understand. But wrapping my brain around this antimatter stuff only confuses me, and makes me feel stupid. I guess I should have taken more than intro physics, hehe.

It's been a rough ride for me as well. I look at the antimatter issue and the concept of "negative" energy like this: Antimatter seems to be a natural consequence of marrying quantum mechanics and relativity. The possibility of antimatter seems to fall out of the solutions. Furthermore, we even experience and use antimatter in day to day life. PET scans for example. So, it's true, it exists and I don't really understand it.

My biggest hurdle so far aside from struggling with the math has been the probabilistic nature of this stuff. It wasn't until fairly recently that I was able to wrap my head around it somewhat. When I first learned about the uncertainty principle and measurement, I simply assumed it was all because our measurements created the probability. Small stuff required small wavelengths to observe, but small wavelengths were very energetic, resulting in a major disruption of the system. Unfortunately, when you look at things like tunneling electron microscopes and certain types of flash memory, it becomes obvious. Things like electrons are probabilistic, it's not simply a measurement problem, the probability is as real as any other physical property. What a hard pill to swallow.

If my questions are any clue, I have been struggling to understand electrons ever since.


ARAI: not a problem.


Group_theory: I appreciate the response and it makes sense. However, what happens with Fluorine? You now have two filled orbitals with even more repulsion. If I had to guess, the nuclear charge essentially cancels this out resulting in a higher first ionisation energy?

 

[ARAI]

Group_theory: I appreciate the response and it makes sense. However, what happens with Fluorine? You now have two filled orbitals with even more repulsion. If I had to guess, the nuclear charge essentially cancels this out resulting in a higher first ionisation energy?

I love talking about this stuff even though I'll admit I'm not that great with it.

I could be wrong, but I thought that since Fluorine has a bigger nucleus than oxygen, the nucleus of the atom has a stronger attraction (gravitational pull) on the electrons that are circulating around it, thus making it more difficult to remove the electrons.

Make sense?

 

[group_theory]

It's been a rough ride for me as well. I look at the antimatter issue and the concept of "negative" energy like this: Antimatter seems to be a natural consequence of marrying quantum mechanics and relativity. The possibility of antimatter seems to fall out of the solutions. Furthermore, we even experience and use antimatter in day to day life. PET scans for example. So, it's true, it exists and I don't really understand it.

My biggest hurdle so far aside from struggling with the math has been the probabilistic nature of this stuff. It wasn't until fairly recently that I was able to wrap my head around it somewhat. When I first learned about the uncertainty principle and measurement, I simply assumed it was all because our measurements created the probability. Small stuff required small wavelengths to observe, but small wavelengths were very energetic, resulting in a major disruption of the system. Unfortunately, when you look at things like tunneling electron microscopes and certain types of flash memory, it becomes obvious. Things like electrons are probabilistic, it's not simply a measurement problem, the probability is as real as any other physical property. What a hard pill to swallow.

If my questions are any clue, I have been struggling to understand electrons ever since.


ARAI: not a problem.


Group_theory: I appreciate the response and it makes sense. However, what happens with Fluorine? You now have two filled orbitals with even more repulsion. If I had to guess, the nuclear charge essentially cancels this out resulting in a higher first ionisation energy?

Several factors - electrons in different subshells don't "screen" out the nuclear charge so you have stronger nuclear charge (so the orbitals are essentially "closer" to the nucleus, this requiring higher ionization energy).

For fluorine, you have 2 filled p-orbitals so each of those orbitals have their own repulsion (the repulsion is not "magnified" because the two orbitals are orthogonal to each other). So you have a higher nuclear charge without any additional screening, and an additional filled subshell.

Also, oxygen/sulfar - by removing an electron, you are putting the electrons with half-filled subshells - which is a low energy state (O⁺ or S⁺), compare to removing an electron from Fluorine or Chloride, which leaves a paired subshell intact (at least for first ionization energy level).



As for weirdness, you are only scratching the surface of weirdness. As you study more quantum mechanics, the more bizzare it gets. The more you study and learn, the more you are confused and baffled. But there are real world application of this stuff - modern electronics would not be possible without this knowledge. MRI and how atoms interact would be impossible without understanding this. For more fun, look up on the quantum mechanics take on the double-slit experiment with electrons, and what Richard Feynman conclusion was. Also look up quantum tunneling (as a result of the uncertainty principle)..

Paseo Del Norte - I assume you are self-studying out of an intro to physical chemistry book and not an intro general chem book?

 

[Paseo Del Norte]

I love talking about this stuff even though I'll admit I'm not that great with it.

I could be wrong, but I thought that since Fluorine has a bigger nucleus than oxygen, the nucleus of the atom has a stronger attraction (gravitational pull) on the electrons that are circulating around it, thus making it more difficult to remove the electrons.

Make sense?

Gravity is pretty innocuous on the atomic level. In fact, at my level I've never used any concept that took gravity into account. (Considering the chemistry of atomic orbitals.)

I do think that because Fluorine has more protons in it's nucleus, the columbic attraction of electrons to the nukes will be stronger. However, my question is how does this all pan out because we increase the amount of repulsion by adding electrons, thus increasing the amount of shielding. In addition, filling another P orbital should result in additional repulsion. How is it that Fluorine does not suffer from the same issue as oxygen? I suspect, the amount of attraction simply overwhelms the added repulsion. I'm just not sure exactly if having two filled P orbitals causes something funny to happen, for lack of better words.

 

[Paseo Del Norte]

Several factors - electrons in different subshells don't "screen" out the nuclear charge so you have stronger nuclear charge (so the orbitals are essentially "closer" to the nucleus, this requiring higher ionization energy).
Makes sense, thanks.


For fluorine, you have 2 filled p-orbitals so each of those orbitals have their own repulsion (the repulsion is not "magnified" because the two orbitals are orthogonal to each other). So you have a higher nuclear charge without any additional screening, and an additional filled subshell.
Also makes sense, thanks.


Also, oxygen/sulfar - by removing an electron, you are putting the electrons with half-filled subshells - which is a low energy state (O⁺ or S⁺), compare to removing an electron from Fluorine or Chloride, which leaves a paired subshell intact (at least for first ionization energy level).



As for weirdness, you are only scratching the surface of weirdness. As you study more quantum mechanics, the more bizzare it gets. The more you study and learn, the more you are confused and baffled. But there are real world application of this stuff - modern electronics would not be possible without this knowledge. MRI and how atoms interact would be impossible without understanding this. For more fun, look up on the quantum mechanics take on the double-slit experiment with electrons, and what Richard Feynman conclusion was. Also look up quantum tunneling (as a result of the uncertainty principle)..

Thanks, I have experienced some of this because I have been working my way through some of the open courses that can be downloaded from Stanford, MIT and so on...


Paseo Del Norte - I assume you are self-studying out of an intro to physical chemistry book and not an intro general chem book?

Pretty accurate. I have taken some general chemistry and have "Chemistry and Chemical Reactivity" second edition by Kotz and Purcell. I also have a couple of Gen Chem books and a basic P Chem review book on my iPad.

More recently, I had to take an additional chemistry course as most of my credits were quite old. I think it was more of a survey type course for allied health professionals. Do not get me wrong, I thought is was a great course and the instructor was outstanding and well qualified with a Ph. D. in orgainc chemistry. We covered many subjects and there was a strong emphasis on basic math, bonding, reactions, equilibrium, and so on. We also had a weekly lab and were required to do our own lab reports.

However, the textbooks and the instruction pretty much glossed over many of what I consider fundamental concepts. Of course, we had to do electron configurations, Afbau principle, orbitals, using quantum numbers and so on. However, I want to make a little more sense on my own if possible.

I would continue on with additional course work in chemistry, math and physics; however, I am currently a respiratory student and simply do not have enough time to take these courses. After I graduate next year, I plan to take additional course work however.

Thanks again for the dialogue. I would probably not plug up this forum with my questions; however, I have had limited to no success asking questions on nursing forums and such. I met a premed who was a Chem major on another forum that has been helpful, so I drop questions over here because I know many of the current physicians and pre-med majors take a fair amount of chemistry or are chemistry majors.

 

[group_theory]

Gravity is pretty innocuous on the atomic level. In fact, at my level I've never used any concept that took gravity into account. (Considering the chemistry of atomic orbitals.)
.

Gravity has no role at the quantum level (that we know of yet - if I knew the answer to the role of gravity at a quantum level - hello Nobel Prize for physics 😛) ... the electromagnetic force is 10³⁸ times stronger than gravity.

 

[ARAI]

Gravity has no role at the quantum level (that we know of yet - if I knew the answer to the role of gravity at a quantum level - hello Nobel Prize for physics 😛) ... the electromagnetic force is 1038 times stronger than gravity.

Thanks for the explanation.

I always imagined that the more dense a nucleus became, the more it attracted mass in the form of electrons. And that was supposed to explain why the more dense nucleus of Fluorine was able to overcome the additional repelling force of another electron.

Ok, so I'm definitely wrong, but I thought it was a good theory. lol. If I wasn't in the middle of a PA program, you'd have me signing up for some a physical chemistry courses next semester.

 

[Paseo Del Norte]

Thanks for the explanation.

I always imagined that the more dense a nucleus became, the more it attracted mass in the form of electrons. And that was supposed to explain why the more dense nucleus of Fluorine was able to overcome the additional repelling force of another electron.

Ok, so I'm definitely wrong, but I thought it was a good theory. lol. If I wasn't in the middle of a PA program, you'd have me signing up for some a physical chemistry courses next semester.

If I had to guess, there is gravity at the nuclear level. Electrons and Protons have mass, so there is probably some sort of interaction. However, as Group_theory stated, there is no role for gravity at the atomic level that we know of. This is probably the holy grail of modern science. (Unifying all the forces.)

I though perhaps you confused gravity for the Coulomb force? (The force of charged objects.) It is similar to gravity in that it's an inverse square law regarding the distance between objects. It's also very strong as Group_theory stated. However, even it is weak compared to the weak and strong forces. Pretty crazy to imagine, but it would have to be, otherwise I imagine it would be near impossible to place protons so close to one and other.