Thanks for my incessant drivel this time goes to @Dred_pirate for tagging me in a thread started by @Merakian_Hash! Here: Propane vs n-Butane
Propane vs. n-Butane vs. Alcohol vs. [Insert Solvent]
WHY you should read this!
Propane and n-butane are very similar, but they definitely extract cannabis differently. Ethanol and other alcohols work, but they also extract more or less undesirables, depending on the alcohol. They key differences between these molecules lie in the differences between their van der Waals interactions, so it is important to understand precisely just what those are!
By the way, you don’t have to read my interpretations of these concepts, but definitely surf around, starting here: Wikipedia: van der Waals force
As you might guess from the plural (interactionS or forceS) there are several different types of atomic and molecular interactions that make up the broad category called “van der Waals”, named after the Dutch scientist, Johannes Diderik van der Waals, which is why it is unusually capitalized. It is somewhat misleading (but not wrong) to call these interactions “forceS”, because they are actually just different applications, permutations and combinations of aspects of the fundamental Electromagnetic Force in relatively “weak” intermolecular ways that are highly dependent on distance. All of any van der Waals force is observed only at very short distances between atoms and molecules, and has attractive & repulsive components:
《…》 Closer than 0.4 nanometers, the total force is repulsive.
->… .<- Between 0.4 and 0.6nm is where the attractive behaviors dominate.
||…_…|| Any further apart than 0.6nm, all van der Waals force is too weak to observe.
There are 4 major components of intermolecular interactions:
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Exclusion - Pauli (repulsive fermion behavior)
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Orientation - Keesom (attractive/repulsive, electrostatic permanent/permanent polarity)
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Induction - Debye (attractive, permanent/induced polarization)
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Dispersion - London (attractive, instantaneous random multipoles)
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The Pauli exclusion principle just says that 2 fermions (particles with half-integer spin states) with equal characteristics in the same quantum system (such as 2 electrons in the same atom) cannot exist in the same quantum state (same energy, shell, orbital, and spin) at the same time. Basically, this is what keeps atoms and molecules “hard” (discrete and relatively incompressible). This principle gives rise to an overall repulsive force that keeps atoms and molecules from intermingling inside one other, despite all the vacuum space involved… much the same way a concrete block repels a human skull, and that no, you are not actually dissolving into the couch.
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The electrostatic Keesom force is the one we all know and love, between 2 permanently polar moietes. The electronegative regions of polar molecules get snuggly with the electropositive regions of other polar molecules. That said, MANY molecules have some degree of permanent polarity, including cannabinoids (which are alcohols and often acids), fats (fatty acid), and entire groups of molecules like aldehydes, ketones, etc., along with the usual suspects like water, ions, salts, acids, bases, et al.
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The inductive Debye force is where things get a little unusual, because permanently polar molecules can cause (induce) instantaneous polarity in non-polar molecules. This is the main reason that polar solvents like water (aka: The Universal Solvent) can keep a little bit of totally non-polar material, such as carbon dioxide or butane, completely and truly dissolved. This is also the main key to unlock the mysteries posited in this thread!
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The dispersive London force is probably to what most people are referring when they say “van der Waals force”. It is usually predominant of the three attractive van der Waals forces (orientation, induction, dispersion) between atoms and molecules that are not small and highly polar (like water). This is the “sticky” force that keeps non-polar and barely polar molecules together with each other and amongst themselves. It is due to the natural fluctuations of electron fields around atoms and molecules, which create momentary or instantaneous polarity in just about everything in the universe. Naturally, this is also very important for extractions, in general.
To show you how important this effect really is, I will paste a Wikipedia table here:
Ahem! Now that that’s over with, let’s see what it means for extractions!
Propane (and CO2) vs. Butane(s)
As you might surmise, propane has slightly less dispersive force, because it is smaller. It has only 3 carbons in CH3-CH2-CH3 configuration. Although they are single bonds and can therefore spin, there is not a lot of variation between V and A except possibly an extremely brief — .
Carbon dioxide is H2C=O=CH2, which is a perfectly straight, stiff, short molecule. Double bonds cannot spin or flex (vibrate out of plane),
so CO2 is just stuck as >—<
…with the >< representing the hydrogens that cannot budge.
These factors actually make propane and CO2 more non-polar than butane, because the momentary polar regions are literally smaller (and therefore disappear/change more rapidly) than those momentary dipoles possible in n-butane, which is a big wiggly
CH3-CH2-CH2-CH3, giving rise to:
N, U, and their inverses,
with several very brief L variations between.
Simply put, propane and CO2 are also, for these same reasons, harder and much harder, respectively, to induce into polarity.
Consequently, propane and CO2 can only dispersively “adhere” to small parts of large molecules like cannabinoids for very brief times. For example, it takes up to 4 molecules of propane just to surround the top and bottom of the C5 “tail” on THC and CBD, and that’s not including the front and back in the 3rd dimension!
Also, although propane (and CO2 even more rarely) can be very briefly induced, they are mostly repelled by slightly electronegative regions like the alcohol -OH group(s), and the O= in the carboxylic O=C-OH, and by electropositive regions like the acidic H in the O=C-OH region.
Therefore, propane and, to an even greater degree, CO2 are not as well-suited to dissolve cannabinoids as longer alkanes like n-butane. Luckily there are lots of terpenes and terpenols that have already dissolved the cannabinoids with their larger dipole moments, giving propane and CO2 at least a shot at moving the bulky cannabinoids.
As an aside, the small-but-wiggly nature of n-butane is one proposed reason why it has a slightly higher affinity for water than other alkanes. The area between the two middle HCH groups is a sort of easily induced region in the electron cloud that may mimic the natural +_+ shape of water molecules! A similar hypothesis goes for isobutane, as I will describe.
Just to mention isobutane, its shape is that of a triangular star composed of
HC-(CH3)3
…where all three of the CH3 groups radiate symmetrically outward from the central CH group, but bent slightly toward the side opposite the lone H on the central C. This gives the whole molecule a mild tetrahedron 3D shape. This shape, strained by “steric hindrance,” to have most of the hydrogen (9 of them!) all crammed together on 1 side, gives rise to a semi-permanent and more easily induced dipole! In other words, isobutane is actually more willing to dissolve slightly polar materials from the herb than n-butane or any of the super-non-polars like CO2 and propane.
Some folks truly believe in isobutane’s ability to give them better extractions than n-butane alone, and this explanation may lend some credit as to why, considering the mild polarity in cannabinoids. However, it also could explain why isobutane extracts performed above -40° (and even above -60°C) have a higher tendency to autobudder!
Personally, I think n-butane by itself is the perfect mix of [non-polar, flexible, easily (temperature) controlled rates of dipole induction] to effectively extract almost all of the cannabinoids and other desired compounds from cannabis.
Alcohols
I’m sure most of you are like “Oh shit, alcohol! Here he goes again! ,” but I’m actually getting tired, so I’ll keep this as brief as I can. These are more about hydrogen bonding than van der Waals interactions, anyway, so no worries! Promise!
You guys know that alcohols are amphiphilic (amphipathic), right? That just means they are both polar (hydrophilic, lipophobic) and non-polar (lipophylic, hydrophobic) at the same time. The OH groups are capable of hydrogen bonding, which is the strongest of the weak bonds, and which is orders of magnitude stronger than any van der Waals force!
Well, that’s the blessing and the curse. It is one of the reasons why alcohols are so awesome with very short contact times and low temperatures, but also why they grab a lot of grunk when contact time and temps are too high. They require cold to become selective, but they are very selective… much more so than CO2 and the alkanes.
The non-polar-ness (lipophylic character) of alcohols varies according to their size; larger alcohols (e.g. isopropanol, THCannabinOL) are more non-polar in character than smaller alcohols (e.g. methanol and ethanol), which have more polar character.
Amphiphylic character of alcohols also varies with the number of alcohol groups (-OH, denoted by the OL, usually at the end of name) on the molecule, relative to its size. CannabiDIOL, for example, has 2 OH groups, so even though it is the exact same size and composition as THC, CBD is more polar than THC! Considering the relative power of hydrogen bonding, it is plain to see why a selective alcohol is the best solvent for extracting CBD… especially with the lower abundance of CBD in plants than THC.
Done! See?
Edit: Apparently backslash doesn’t show up here, @sidco @Future ? Is that a bug, or is it a prompt for some script?
Edit: Wow! The autocorrupt built into the publishing is strong with this one! No more ASCII chemistry explanations…