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Author Topic: Why is Teflon Hydrophobic?  (Read 17474 times)

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ET

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Why is Teflon Hydrophobic?
« on: December 22, 2013, 03:11:05 AM »

PTFE, while known for it's hydrophobicity, it's rather counter intuitive because basically there should hydrogen bonding between the water molecules to the fluors, and not just hydrogen bonds, the strongest attraction. Is it because of the charge distribution over the carbon chains?
Can anyone explain please?

Thanks!
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Archer

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Re: Why is Teflon Hydrophobic?
« Reply #1 on: December 22, 2013, 04:27:18 AM »

As I understand it there is no dipole in PTFE to allow for hydrogen bonding because the molecule is perfluorinated.

Look up the dipole for carbon tetrafluoride compared to fluoromethane
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Enthalpy

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Re: Why is Teflon Hydrophobic?
« Reply #2 on: December 23, 2013, 10:06:48 AM »

No global dipole moment because of symmetry, but hydogen bonds don't need one, do they?

Now I also wonder about ET's remark - provided I got it properly. From what I read, hydrogens stick to electronegative atoms like oxygen at an other molecule, because these atoms have a bumρ of electron density at the opposite side of the chemical bond, say C-O or H-O. This should have happened with C-F as well, shouldn't it?

To make a comparison: how is CO2 dissolved in water? Always as the acid? Or as well through hydrogen bonds to CO2's symmetric oxygens?

I've searched the water solubility of:
CH4: 30ppm mass
CH2F2: 0.44%
CF4: 3.8ppm in mol
so a dipole moment helps a lot to dissolve, and fluorine doesn't.
« Last Edit: May 08, 2014, 05:16:10 AM by Dan »
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kubgk

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Re: Why is Teflon Hydrophobic?
« Reply #3 on: April 08, 2014, 03:02:40 PM »

Hi,

small doubt relating to this. How fluorinated molecules reduce the surface energy ? 
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Re: Why is Teflon Hydrophobic?
« Reply #4 on: April 09, 2014, 09:59:24 AM »

I apologize in advance for the length of this post, but this is such an interesting topic to mull over. I'll stating up front that these are just my speculations. I haven't delved into the literature to verify whether anything here is true or false, so YMMV.

Here is some initial food for thought:

http://www.chemguide.co.uk/qandc/ptfe.html

I encourage you to read that before you read what I have to add.

The author of that webpage points to an apparent discrepency between what many authors say about the low magnitude of PTFE intermolecular (van der Waals) forces and the high melting point of PTFE (~327 C, compared to, say regular old low density polyethylene at around 105-115 C). Why is this an apparent discrepency? Well, let's look briefly at the two facts:

(1) Most websites and other popular literature sources point to the very low magnitude of van der Waals (London dispersion) forces in PTFE as being responsible for the nonreactivity and low friction coefficient of PTFE surfaces. Many like to point out that PTFE is the only substance to which a Gecko's toes won't adhere.* The stickiness of a Gecko's toes is commonly attributed to van der Waals forces between the high-surface-area toes and the surface that the Gecko is trying to climb. I don't think we have any reason to doubt this latter fact (about Geckos). In turn, the lack of susceptibility of the PTFE surface to instantaneous induced dipole moment changes - that is, the lack of van der Waals forces at the PTFE surface - is usually attributed to the very strong C-F single bond, and the extremely low polarizeability of fluorine. Fluorine is so electronegative that it holds its electrons very close to the nucleus, which lowers the probability that some other nearby dipole can temporarily pull an electron farther away. Hence the "nonstickiness" of perfluorinated polymer materials.

Sounds good, right? Uh-oh:

(2) PTFE has a very high melting point. The temperature at which a phase change for a substance occurs is well known to all chemists to be highly influenced by the amount of intermolecular forces holding it together, and typically a higher phase change temperature means MORE and STRONGER intermolcular forces - all things being equal of course.

So, if the non-stickiness of PTFE is attributed to very low magnitude intermolecular forces and the high melting point of PTFE is attributed to very high magnitude intermolecular forces, something clearly ain't right.

The author of the article I linked to solves this discrepency by essentially rejecting (1) completely. From the horse's mouth - er, fingers - "However . . . several web sites talk about PTFE having very weak van der Waals forces. That has got to be completely untrue. If it had very weak van der Waals forces, it would be a gas - not a fairly high melting point solid!" [Emphasis not mine.]

On the surface this seems reasonable.  After all, the melting point is an actual piece of indisputable data that is hard to rationalize without assuming that the intermolecular forces in PTFE must substantial in magnitude.

However, I maintain that this discrepency is due to an inappropriate comparison of bulk and surface properties, as well as a need to explore in more detail what gives rise to the high melting point of PTFE. Let me address these one at a time.

(A) Phase changes and intermolecular forces.  The temperature at which a phase change occurs for a substance is due to the strength of the intermolecular forces that hold the substance together. For boiling points, applying this rule is usually pretty simple. You look at the structure of the molecules making up the substance, determine whether there are likely to be dipole-dipole interactions, hydrogen bonds, or London forces, and make some arbitrary but straightforward assessment of the overall strength of these interactions based on the magnitude and number of the forces identified as important. It's typically understood that a higher molecular weight corresponds to higher boiling point, for example, because higher molecular weight substances tend to have more overall molecular-scale surface area - more places for molecules to stick together due to transient dipoles. (Many students assume higher weight substances are harder to boil because of gravity, but this couldn't be more wrong.)

Melting points, while still generally dependent on intermolecular forces, are not so straightforward to explain. No longer can we just look at the constituent molecule and make an assessment of the relative strength of intermolecular forces expected, because the way a solid is put together matters. The crystal structure, density, particle size, and so forth all make a huge difference on bulk properties, even among solids which are compositionally identical (and moreso if they aren't). I don't want to reinvent the wheel here, so let me link to an excellent description about how molecular shape and symmetry, in addition to the usual electronic factors, play a role in dictating the strength of intermolecular forces of a solid, and the resulting melting point: http://www.masterorganicchemistry.com/2010/07/09/chemical-tetris/ .

Essentially the point that is being made there is that while two substances may be chemically similar with respect to molecular polarity, the number of relevant intermolecular forces in the solid state can vary substantially simply because of the way the molecules pack together. A branched alkane, for example, will have a higher melting point than a linear alkane of the same molecular weight due to the effeciency of packing in the former: linear alkanes pack much more closely together than branched alkanes, which means the shared surface area is larger (or, the average distance between molecules is smaller). Strength of intermolecular interactions is due not only to relative differences in charges (permanent or transient) but also to the distances between them. If you compare hexane to 2-methylpentane, the reason the latter melts almost sixty degrees colder is because due almost wholly to packing efficiency.

The same kind of logic can be applied to polymers. Low density and high density polyethylene differ primarily in the extent of branching along the polymer backbone. HDPE has less branching and as a result packs together much more efficiency (it is higher density). This translates into better shared surface area between adjacent polymer strands, and in turn a higher melting point (~120 to ~130 C) compared to LDPE (~110 C), which packs less efficiency and has less shared surface area between adjacent strands.  Molecular weight, too, plays a role. Ultra high mol. weight polyethylene (UHMWPE) has an even higher melting point (~135 C) because higher weight strands have a higher degree of tangling and such that increases the relative surface area even more.  All these differences in a series of polymers that are nearly identical electronically (i.e., as far as polarity and polarizability is concerned).

The relevance of all this to PTFE is as follows: the melting point of PTFE is 327 C. Corresponding melting point of LDPE is about 110 C. What has stronger intermolecular forces? The clear answer is PTFE. That's kind of indisputable. But it's pertinent to ask why. Fluorine is heavier than hydrogen and thus PTFE strands might be assumed to be higher molecular weight than LDPE. But that's neither here nor there. The real reason is mass and electron density. The specific gravity of LDPE is ~0.92 and the specific gravity of HDPE is ~0.95, accounting for the slightly higher melting point of the latter. The specific gravity of PTFE is a whopping 2.2. Granted, the molecular weight of fluorine probably has a lot to do with that, but so does the packing state. PTFE has so few free volume holes because of the way the individual polymer strands join together, deriving predominantly from fluorine sterics and fluorine-fluorine repulsion. Let me quote from the original article linked to above: "The [fluorine-fluorine] repulsions lock the [PTFE] molecules into a rod-like shape with the fluorines arranged into very gentle spirals - a helical arrangement of the fluorines around the carbon backbone. The rods will then tend to pack together a bit like long thin pencils in a box." So, what this means is that even if the C-F bond is very strong and fluorine very resistant to polarization, the fact that PTFE is so dense and the polymer strands organized so tightly and regularly ensures a high surface area contact between strands. I.e., even though fluorine holds its electrons tight, the fact that everything is crammed so close together in PTFE greatly increases the frequency of transient dipole interaction, and thus gives rise to a high melting temperature.

This is all basically a round-about way of saying that yes, the intermolecular forces in PTFE are large, but they are large because of frequency of occurance and packing structure of the solid - that is, the intermolecular forces are strong despite the low polarizability of fluorine. This gives us a perfect way to rationalize the apparently contradictory behavior that occurs at the PTFE surface.

(2) Surface properties vs. bulk molecular properties. Several times now I've italicized the word "surface", and I've done this for a reason. What happens at the surface of a solid is a function of the respective free energies of the solid and whatever the solid is interfacing with (even if it's just air). The apparent discrepency that's been discussed above is rooted in the fact that two very different processes are under consideration: cohesion vs. adhesion. Because of tight packging, as we've seen, PTFE has a high energy of cohesion. But, this still doesn't change the fact that the fluorines in PTFE are not very sticky, on a per fluorine basis. At an interface, where packing density perhaps is not so important, the inherent stickiness of the substance is predominantly responsible for determining the magnitude of adhesion when another substance comes into close contact.
 
In fact, and this is just conjecture, it may be argued the packing density that gives rise to high energy of cohesion in PTFE directly contributes to the low energy of adhesion and especially the low coefficient of friction, due to mechanical effects as well as a higher available surface area for adhesive bonding. If there is a high packing density, this would (I would think) give rise to a smoother, more uniform surface. Thus there would be less surface area available for bonding at the interface, certainly in the event that the bonding substance is a liquid or other substance - like denaturing proteins in a cooking chicken breast - that can conform to whatever geometry is available. Rougher surfaces take more energy to move parallel to each other under an applied perpendicular force - there's your friction aspect. (I believe that Teflon pans lose some of their stickiness as they age due to formation of micro-scale pits and other physical surface defects.)

Even beyond that - let's pretend for a moment you had a perfectly flat surface made of PTFE and a perfectly flat surface made of LDPE. We can agree that the PTFE material would make a good pan and the LDPE would not because PTFE has a higher melting point owing to its much greater polymer packing density and corresponding strength of intermolecular forces. LDPE has a correspondingly low force of cohesion between its molecules and falls apart at a pretty low temperature - this despite the fact that the C-H electrons on LDPE are inherently more sticky (polarizeable) than the C-F electrons in PTFE. There is simply a lot more C-F's bumping up against C-F's in PTFE than there are C-H's bumping up against C-H's in LDPE.

On the other hand, at the interface of our ideally flat, structurally identical surfaces, the PTFE material I think would have roughly the same amount of C-F's available for forming London bonds as the LDPE has C-H's available. What we do here is essentially normalize the materials for density of bonding area, and the sticky LDPE wins out now by a fair margin when it comes to adhesion potential. (And in fact, we err on the conservative here, if my conjecture about the roughness of the LDPE vs. PTFE surface is correct - a real LDPE surface would have a higher surface for London bonding than an PTFE surface of the same area). This is somewhat borne out in the numbers: despite its higher melting point, PTFE has a lower surface free energy (19.4 dyne/cm) and higher water contact angle (109.2 deg) than polyethylene (31.6 dyne/cm, 96 deg), polypropylene (~30.5 dyne/cm, 102.1 deg), polystyrene (34 dyne/cm, 87.4 deg), and polybutadiene (29.3 dyne/cm, 96 deg) (source: http://www.accudynetest.com/polytable_03.html). Notably, the four comparison polymers are all exclusively C-H containing polymers and thus might be expected to have superficially similar degrees of stickiness despite having very different melting points and other bulk molecular properties, and this is what the numbers kind of show, at least when compared to PTFE. Granted, this is probably a simplification of what is going and it would be easy to nitpick against this argument, particularly since things like contact angle also depend on the cohesive properties of the liquid being dropped on the substrate.

Anyway, to sum: in the bulk, PTFE forms weak intermolecular forces due to the high polarizeability of fluorine, but the dense packing in the solid state ensures that it forms a large number of them, which translates into a high melting point. The cohesive forces holding PTFE together are very strong. However, at the interface between PTFE and a binding substrate, I think this packing density may not be as important, and the inherent resistance of fluorine to sticking creates a low energy of adhesion compared to other materials. That is, the intermolecular forces between the PTFE surface and another substance at the interface are especially weak. The tight packing of PTFE polymer strands may also give rise to a more uniform, smooth surface which limits the extent of mechanical adhesion and provides less surface area for dispersive adhesion compared to other solids.

* I'm not sure if that's the case; there are a few substances with a lower coefficient of friction than PTFE.
« Last Edit: April 09, 2014, 10:10:15 AM by Corribus »
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Enthalpy

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Re: Why is Teflon Hydrophobic?
« Reply #5 on: April 11, 2014, 01:28:56 PM »

Having invested much time trying to understand the melting point of alkanes, I fully agree with the ease of packing. I believe to see an other factor, seemingly less known: the ease of the molecule to change its conformation, for instance rotate around its bonds. And an other one: the global degree of symmetry of the molecule.

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Deformations store heat energy in the liquid, but are usually prohibited in the solid, so they favour the liquid. More so if the deformation is easy, that is, a low temperature can do it. You can understand it from the entropy if you wish, or just by the ease of packing: a "liquid" molecule that has many conformations has less chances to fit the already existing solid surface hence being caught.

Interestingly, a single methyl branch on an alkane destabilizes the most favourable conformation, that is, it eases the rotation around the bonds of the main chain, while geminal methyls hamper the rotation. This goes in the same direction as the ease of packing and its influence on the melting point.

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The global symmetry of the molecule permits an impinging "liquid" molecule to find more easily a place in the already existing solid surface. Cubane for instance fits in 6 possible orientations, and has a very high melting point for a C8. As well, branched alkanes with as many carbons on both sides of the beanch freeze more easily. More impressive MP : +101°C for 2,2,3,3-Tetramethyl-Butane but -10°C for -Pentane.

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As for Ptfe, its difficulty to rotate around C-C bonds must explain much why it stays solid while Pe is already liquid. Both because more contact points are at the right place, and because the rotations store little heat.

So I fully agree with Corribus that the melting point relates not only with individual intermolecular forces. Small atom-to-atom intermolecular forces are perfectly compatible with a high melting point.

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Coefficient of friction: this is a property of the surface, that is, where the molecules are broken, disturbed, dirty, oxidized, and the like. Again, I agree with Corribus that it results not only from intermolecular forces of the sound macromolecules.

Imagine two rubbing objects: small chips are torn away, chemical bonds are broken again and again. I would expect that the reactivity of fluorine saturates snappily the new pending bonds and makes them less sticky.
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Corribus

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Re: Why is Teflon Hydrophobic?
« Reply #6 on: April 12, 2014, 07:20:39 AM »

Having invested much time trying to understand the melting point of alkanes, I fully agree with the ease of packing. I believe to see an other factor, seemingly less known: the ease of the molecule to change its conformation, for instance rotate around its bonds. And an other one: the global degree of symmetry of the molecule.
Yes, I think all of these factors are related to some degree. "Ease of packing" isn't a particularly easy thing to define rigorously and probably relates to a lot of different molecular structure attributes like some of those you mention.  I think the important take home message is that solids can't be treated simply as very viscous liquids in most cases, which is why melting points are more complicated than boiling points. The same general rules apply, of course, but solids are just more complex.

Quote
As for Ptfe, its difficulty to rotate around C-C bonds must explain much why it stays solid while Pe is already liquid. Both because more contact points are at the right place, and because the rotations store little heat.
That's an interesting additional point I hadn't considered.
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Enthalpy

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Re: Why is Teflon Hydrophobic?
« Reply #7 on: April 13, 2014, 12:04:53 AM »

I just remember that the coefficient of friction jumps with the temperature for Ptfe. It's low at room temperature, but very soon, like +50°C, it increases to 0.3 or 0.4, an absolutely banal value for a polymer. That's in the temperature range of a crystal transition of Ptfe, something like 19 -CF2- per 360° twist becoming 17.

As well, the coefficient of friction jumps with the mechanical load pressure for Ptfe.

These are some reasons why Ptfe isn't used alone when mechanical designers want low-friction bearings.

So clearly, many other important factors weigh in, in addition to intermolecular forces.

More generally, dry friction is poorly understood. Some theories exist, but they add half a dozen effects with >10 tweeking parameters, so they could explain anything, being it observed or not. Engineers need progress from physicists here, first with bare measures, because these are scarce, and only then with theories.
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Enthalpy

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Re: Why is Teflon Hydrophobic?
« Reply #8 on: May 07, 2014, 11:35:43 AM »

Comparing alkanes and silanes supports the idea that molecules that deform more easily melt more readily. See the attached picture - it displays sharp atoms, but their limits are of course fuzzy, more so for hydrogen.

In both cases, the silane and alkane homologues have similar boiling temperature and heat, so intermolecular forces are as strong, but the silanes melt 100K before the alkanes.

Symmetry can't explain it here, since the molecules have identical shapes. The perfection of packing neither, because the molecules that melt more readily absorb more heat for that: it can't be a matter of attraction energy.

The explanation I'm pleased with is that the bigger silicon atom makes subgroups more mobile in the liquid, while the solid holds them in place. Then:
  • A liquid silane sticks less to a solid surface, because its shape misfits the crystal most often;
  • The more deformable liquid silane stores more heat, its entropy is bigger than the solid, this favours the liquid.
Both explanations must be equivalent, one being understandable and the other computable. Anyway, more mobile subgroups both absorb more heat when melting and favour the liquid, while other explanations wouldn't do both.

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n-Pentane (melts at -130°C) versus neopentane (-17°C) tell the same story.
http://webbook.nist.gov/cgi/cbook.cgi?ID=C109660
http://webbook.nist.gov/cgi/cbook.cgi?ID=C463821
Neopentane's symmetry would promote the solid phase, but this explanation alone fails to tell why neopentane takes less heat (3.1kJ/mol) to melt than n-pentane (8.4kJ/mol). The difference, 2.5*R at -17°C, fits well a few vibrations and rotations freed by melting.

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http://en.wikipedia.org/wiki/Neopentane#Boiling_and_melting_points
noticed the argument about the heat of fusion, and links with little explanations to:
James Wei (1999), Molecular Symmetry, Rotational Entropy, and Elevated Melting Points.
Ind. Eng. Chem. Res., volume 38 issue 12, pp. 5019–5027 doi:10.1021/ie990588m
which may well give the same explanation as here - without the silanes?

Marc Schaefer, aka Enthalpy
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Enthalpy

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Re: Why is Teflon Hydrophobic?
« Reply #9 on: June 30, 2014, 07:20:31 AM »

A heat-resistant polymer without chlorine nor fluorine would be nice; stiffness to vibrations and rotations could be the key.

Marc Schaefer, aka Enthalpy
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Re: Why is Teflon Hydrophobic?
« Reply #10 on: July 29, 2014, 04:47:39 AM »

I just want to add to this old discussion that I have maintained an ongoing email conversation with the ChemGuide site author. He has revised his Teflon article since my original lengthy post above, and so some of the starting points I extracted from his original article may no longer be there.
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Re: Why is Teflon Hydrophobic?
« Reply #11 on: July 29, 2014, 07:59:02 AM »

I have maintained an ongoing email conversation with the ChemGuide site author. He has revised his Teflon article

Great thing to know that our members are trying to make the internet better ;)
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Enthalpy

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Re: Why is Teflon Hydrophobic?
« Reply #12 on: September 29, 2018, 11:32:46 PM »

Pentane isomers illustrate nicely that the packing density doesn't essentially determine the melting point. From
 https://en.wikipedia.org/wiki/Pentanes

mp °C   bp °C   kg/m3   Isomer
----------------------------------
-130     +36     621    n-pentane
-160     +28     616    isopentane
 -17     +10     586    neopentane
----------------------------------


The density is for the liquids at 0°C. Solids at 77K would be more relevant: nice, easy and useful experiment to conduct for undergraduate or high school students.

On the pentane isomers example, we see again that symmetry or deformability (both I suppose) influence the melting point very strongly. That would make it more difficult for software to estimate melting points by the search of optimum solid packing: it should also evaluate the effect of molecules orientation and deformation in the liquid. Maybe the estimation of thermodynamic variables suffices for that.

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Re: Why is Teflon Hydrophobic?
« Reply #13 on: October 01, 2018, 05:49:54 AM »

My thought is that for small molecules, the molecular symmetry plays a much more important role than packing density on the melting point - symmetry impacts the entropic contribution to the phase change because there are more ways for a symmetric molecule to line up and join an ordered solid than there are for a non-symmetric molecule. For large floppy molecules, and polymers, the symmetry contributions are more or less identical, so packing density (and subsequent enthalpic considerations) can play a starring role.

Melting points of small molecules can only be predicted with any reasonable accuracy when molecular symmetry/entropy is taken into consideration.
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Re: Why is Teflon Hydrophobic?
« Reply #14 on: October 01, 2018, 12:34:03 PM »

Can the symmetry be included in a state variable, especially the entropy?

That would be very convenient, and it's the goal of thermodynamics to replace seemingly horrible sets of impact parameters, orientations and so on, by a handful of state variables.

But I don't see any symmetry factor in dQ/T. Any ideas for an additional term?

The question is not only academic... If some day software is to predict melting points, trying all possible orientations, conformations, positions and speeds to check if a liquid molecule sticks to the solid would be a horrible load for a PC, while some sort of symmetry number would compute quickly.

Predicting melting points would be important, but to my knowledge, only two teams work (or did work) on it worldwide, and apparently they only let software search for the best stacking of molecules in a solid, which isn't the full picture.
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