Structure of chorphyll A and B

[scott_alexsk]

Hello, I am not sure this thread should be here, but I will ask anyways. The structure of chorphyll A and B differ by 4 atoms attached towards the head of the molocuel, which I believe is 3 hydrogean and a carbon for B( correct me if I am wrong) and just a hydrogean for A. There is a carbon ring at the top of the molocuel followed by a carbon and hydrogean chain. My first question is, how are the chorphyll molocuels oriented within the cells? I mean is the 'head' facing outwards, or is there some other kind of orientation. My second question is how can the frequencies of light that is effectively obsorbed, differ so greatly between the moloceuls? I understand that each element has its own spectrum, but how does this affinity for certain frequencies of light vary within a complex compound. I mean I would think that each atom would obsorb its specific frequencies of light. Does it act this way, or does every single part of the molocuel obsorb the same frequencies overall? Is it just that the composing atoms obsorb their individual frequencies and as a whole that's what the molocuel can obsorb, or is it something else? By the way when an element is burned are the bands of light it produces the same as the frequencies of light it can effectively obsorb?
Thanks
-Scott

 

[Mike H]

For plants and certain bacteria, chlorophyll molecules are bound to proteins that form an antenna complex around the photosynthetic reaction center. In plants, this is localized in the chloroplast organelle.

The side group (a methyl group for chl a, a CHO group for chl b) does slightly tune the absorption properties (absorption maxima for chl a at 430 and 662 nm, chl b at 453 and 642 nm) which don't differ that much in my opinion. Various derivatives of chlorophyll will absorb at different wavelengths when you look at photosynthetic systems across the board, especially when you consider the influence of the protein environment. Given the alternating single/double structure that composes the porphyrin ring, you can think of the absorption maxima as where the entire structure absorbs considering the delocalization possible. This topic is a favorite of chem textbooks which have coordination chemistry as one or more of their chapters.

 

[scott_alexsk]

Of course my chemsitry or former biology book did not have that material. So I assume that the carbon ring is facing outwards towards the outside of the cell, right? Also does the carbon and hydrogen chain following this carbon ring act as the transporter of heat to the electron transport chain? Also are you saying that as a result of the double bond structure within the carbon ring, light frequencies are obsorbed at all points in the ring, or is it only obsorbed at the point of a specific atoms with the coresponding range of obsorbtion(ie spectrum)?
Thanks for your time.
-Scott

 

[Mike H]

Asking about the orientation of the chlorophyll molecules in relation to the cell is one which I can't definitively answer. I'm not sure if there is one, to be honest. In plants, the chlorophyll molecules are bound up in the thylakoid membranes, which are localized in the chloroplast. The chloroplast is an organelle - I should note that there are multiple chloroplasts per plant cell - and they do not have any preferred arrangement within the plant cell that I am aware of at this moment. Undoubtedly some of the bound chlorophyll molecules, when one averages over the entire cell, will be facing towards the outside of the cell, some towards the center, and some laterally.

The light that is harvested is transferred due to (kind of) the electronic overlap between the chlorophyll molecules until it finally funnels in towards the photosynthetic reaction center (really excitation transfer via resonant energy transfers as is explained by Forster theory). It does not explicitly have to do with the long hydrocarbon (phytol) tail except as that tail influences the electronic structure of the chlorophyll.

Due to the alternating single/double bond structure of the porphyrin ring, you have delocalization of the electrons. So yes, when light is absorbed, it is by the structure as a whole and not by a particular atom or functional group.

 

[scott_alexsk]

I had a feeling that it would obsorb light that way. So this has to do with resonance structure right? Could you explain how this applies to this situation since my knowledge on resonance is limited.
Thanks again,
-Scott

 

[Mike H]

The conjugated pi-electron system of the porphyrin ring - due to the alternate single/double bond structure - allows for multiple resonance structures, which really just indicates delocalization of the electrons. Since there is this sharing of electrons, it's not just a single atom or functional group absorbing light, but rather it's the chlorophyll as a whole.

That's pretty much the entire story.

 

[scott_alexsk]

I have a hard time understanding how the obsorbtion spectrum of the atoms can be combined when there is a resonance structure. Correct me if I am wrong but does not the spectrum produced when you heat an element the same as the light it can obsorb? I mention this because if that is true, then the light that the atoms can obsorb correponds to the difference in certain orbital levels. It seems that the electrons not having a defined postion would not necessarly transfer all of these characteristics from the spectrum obsorbtion (which comes from the difference in frequency of the orbitals), everywhere? This is just my flawed understanding and tell me if the question is more appropraite for another forum.
Thanks,
-Scott

 

[Mike H]

The ultraviolet/visible spectrum of a molecule reflects transitions in the electronic state of the molecule. The electronic structure governs what transitions are possible, so if you took a ring system that was only composed of single bonds and derivatized it so it was now conjugated (alternating single/double bonds), you would have different electronic structures for the purely single-bonded ring system and for the conjugated ring system. It is not a straightforward extension of what is known about individual atoms. For instance, individual carbon atoms have most of their strong spectral features between 60 to 200 nm. UV/Vis-active small organic molecules have strong spectral features that are a bit less energetic, bridging from the UV to the visible. Various pigments and biological cofactors (various porphryin complexes and others) will go from the visible range to heading into the infrared range. The way to really understand this is to look into how the electronic structure changes, how the chemical bonds are made and how metals influence the electronic structure with respect to coordination chemistry.

One thing that might help is to remember that when light is absorbed, it's the electrons which are excited into a higher electronic state. If you put that atom in a different environment, its electronic structure is changed, which means that its absorption properties are going to be different. So if that carbon atom is in a larger molecule, it's not going to behave exactly as a single carbon atom in the gas phase or in a vacuum.

I hope this makes sense.

 

[scott_alexsk]

It somewhat does. But what I am having a hard time grasping what defines the obsorbtion spectrum of the electrons. Are you saying that the difference in orbital frequecies does not define the obsorbtion spectrum though it may define the emission spectrum? I say this since in the resonance structure that is occurring, the electrons are not really with their atoms anymore.
-Scott

 

[Mike H]

The absorption spectrum corresponds to electronic transitions of the molecule. I am stating that the electronic structure of molecules is more complicated than that of atoms. Therefore, electronic transitions in molecules are not the same as that for atoms. When I say not the same, I don't mean that the basic idea suddenly changes for molecules - you're still measuring the transition between two different electronic states - but that the electronic states for atoms are not the same as molecules. Any good physical chemistry text (e.g., MacQuarrie & Simon) will provide you with the quantitative details.

I gave an example in my last post - we see carbon atoms absorbing at very short wavelengths (between 60 to 200 nm, with many between 65 to about 150 nm), while we see absorption for small organic molecules at longer wavelengths (for instance, anthracene and naphthalene will absorb at wavelengths from about 200 to 375 nm), and we see different absorption features for even more complicated organic molecules (such as metal-binding porphryins such as chlorophyll and heme going from about 400 to 700 nm).

...the electrons are not really with their atoms anymore.

This describes covalent bonding, even without pi-orbital overlap and the possibility of being able to draw resonance structures. Even in a single bond, the electrons are being more or less equally shared, they don't "belong" to an atom any longer. If they went back to their atom, the bond would break.

This is what I was trying to get at earlier. You have a network of truly shared electrons from the pi-orbital overlap. They're the ones that are doing most of the absorbing in a typical UV/Vis measurement. You can't say anything about "oh, it's the electron from this particular functional group" - by its very nature, it's delocalized over the entire molecule.

As I will be away for the Thanksgiving holiday in the U.S., I probably won't be posting over the next few days. I am also not sure I can explain this any more simply. I would strongly recommend finding or refinding a good p.chem. book like I mentioned above to explain how electronic states differ in atoms and molecules if you're still confused.

 

[Mike H]

I probably should have bothered to dig this up earlier from my files since it seems you're still interested in the topic from other posts elsewhere on the forums. Sorry.

...chl a monomer...the red and orange regions of the visible absorption spectrum is due primarily to the S0 -> S1 transition...

...the absorption spectrum in the green and yellow regions is due to the S0 -> S2 transition...

...The S0 -> S(n) n>2 transitions are located in the blue and higher energy regions of the absorption spectrum...

From L.L. Shipman, et al. (1976) Journal of the American Chemical Society. 98(25): 8222 - 8230. The paper discusses absorption spectra for chlorophyll monomers, dimers, and oligmers in solution. S0 is the ground electronic state, S1 is the first excited singlet electronic state, and S2 is the second excited singlet electronic state of chlorophyll a. A similar reasoning exists for chlorophyll b as well, as memory serves.

The point is that you are actually seeing multiple electronic transitions in a UV/Vis spectrum of chlorophyll (ground state to first excited singlet state, ground state to second excited singlet state,, ground state to higher order excited singlet states).. However, this topic just gets more and more detailed and labyrinthe the further you get into it and I was hoping to not have to actually go dig in my boxes full of papers on photosynthesis...

 

[scott_alexsk]

I asked Gukul this question also, but he has not responded to it yet. Is it possible to determine the absorbtion spectrum of a compound without just seeing what it is. I mean is there a distinct way how these "compound orbitals" are formed from the orbitals of elements when bonded? Oh, and I appreciate you just taking the time to respond to my posts. Anything beyond that is very generous. And if you have any idea where I can read these types of papers online or in certain journal, I could attempt to decpiher it myself then ask you questions, rather than you trying to give me all of the information from these articles.
Respectfully,
-Scott

 

[Mike H]

The paper I cited might be able to be found at the pubs.acs.org website, but it will require a subscription. Your school or nearest library may have access. You may have better success with Google Scholar (scholar.google.com) in finding freely available papers that researchers have posted on their web sites.

In principle, one can calculate the energy of the ground electronic state, the excited states, and predict what the absorption spectrum may look like from the energy differences. However, realistically, these quantum mechanical calculations suffer from a size restriction (I think you can currently go up to a molecule of about 200 atoms or so) and it tends to be not so good for predicting excited states of larger molecules like chlorophyll. For something like acetylene or other small organic molecule, you can probably get pretty good results. So while it is in principle possible, it's probably just easier to do the experiment for more complicated molecules, especially when one factors in solvent effects, temperature, and other variables. If you have a pure compound, a UV/Vis spectrum should take no more than a few minutes as long as it's properly soluble.

http://www.chemistry.mcmaster.ca/esam/Chapter_8/intro.html" is an explanation of molecular orbital theory, which of course becomes a bit messier when going from simple polyatomics to large biologically relevant molecules, and where computational methods such as the quantum mechanical ones I mentioned above become critical.

 

[scott_alexsk]

Wow and thanks. It will take awhile for me to get through that entire site. I'll attempt to do some of the practice problems to better understand the situation. You only need an understanding of Alegebra 2 to get through those problems, right? I will continue this thread later, but now I am going to Florida.
Thanks again MikeH,
-Scott