Can you accept that if beer is confined in a container and is at equilibrium with a headspace mixture that is 3:1::N2:CO2 (keg of Guiness ready to serve) that were you to disturb the keg to the point that a bubble formed in the beer its N2:CO2 ratio would be the same as in the headspace (3:1)? Were it otherwise Henry's law would have to be different for the interface between the headspace and the beer and the bubble interior and the beer. How could this be? There is only one Henry's law. If you can accept that this is the case we can go forward from there. If not read #67 until you understand it and then we can go forward. Perhaps you are not familiar with Henry's law or the concept of chemical potential. Any p-chem text should get you up to speed on those but there should be enough in #67 to explain things adequately. A couple of other posters here have seen the light. You should be able to as well.
The great nitrogen bubble debate
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BitterSipper
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No, I choose not to get into what is not relevant. It does not matter what the contents of a bubble under pressure inside of the keg or under pressure in a syringe are. I even conceded to you that possibly some bubbles might contain mostly nitrogen, but this is not relevant to the context. It matters how much of each gas is dissolved and the amount that they diffuse during the pour. On a cursory exam of your postings, you even contradict your own statement, even with best case equilibrium scenarios that don't even account for the complete circumstances of pouring a beer.
I saw an easy way to show this to you and I have posed to you fully valid logic that is completely contradictory to your statement but you refuse to address it.
Again, I pose to you a simple logic argument;
Beer poured with only nitrogen dissolved in it forms next to no head
A beer with less nitrogen dissolved in it must have less of a head comprised of nitrogen than the beer with more nitrogen.
Beer poured with CO2 and nitrogen has less nitrogen dissolved than the 100%, but pours with a substantial head.
That head must be comprised of mostly Co2.
If you repeat Henrys law again, I will assume you simply have no supporting argument for your statement.
And in so saying declare with even louder voice that you don't know what you are talking about.
It only contradicts what I say because it is wrong.
Hardly
Except that it isn't. You can verify this for yourself fairly simply. Draw a beakerfull of beer with the spout below the surface of the beer. Equip a good sized syringe with a short length of tubing and have a pair of haemostats handy. Draw a goodly amount of beer into the syringe. As you do so it will foam. Draw it gradually so that no (or very little) air leaks in around the plunger. Hold the syringe vertically. Shake it. As you do so more gas will escape the beer displacing it. When you have recovered a good amount of gas depress the plunger until all the liquid is out of the syringe. You now have a syringe with nothing but gas from the beer in it.
Move the end of the tube to a container of a strong lye solution and carefully draw in 10 - 20 cc of it. Clamp the tube before withdrawing from the lye. Shake. The gas volume from a regular beer will change dramatically from say 35 to 5 ml (what I just got with a Pils amounting to an 86% reduction in volume). With a stout (or any beer drawn on mix) the gas volume will change less dramatically. I just got about a 50% reduction with the stout I have on tap. Clearly, as the theory (based on the Henry Law) predicts, the bubbles in my stout are about 50% nitrogen. And we note that work at the Guiness lab says theirs contain about 75% nitrogen.
You don't understand the theory. You have no experimental evidence. I have both. I think we have to give this one to science.
I am happy to discuss, explain, illuminate to the extent I can any of this but there comes a time when one realizes that he is trying to teach a pig to sing. As the old adage goes one shouldn't do that as he isn't going to succeed and it annoys the pig. I think we are at that point. If I see a logical question I will respond in turn. If not, I won't.
I do hope the other readers were interested by the little experiment I did tonight. Pretty neat and based on the ASBC MOA for determining the amount of carbon dioxide in a beer.
timdillon36
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So if we look at the scientific method, in the few years of this post, we have steps one to three covered.
Its time to test, analyze, and report findings.
I am willing to help in the testing but currently lack the funds for a nitro or beer gas setup but i do have a simple way to test the one theory of carbing a beer and then dispensing at higher presure with a bladder.
Anyone in the Indy area willing to help with this please PM me.
My idea for this experiment:
Make a large batch of a clone of a commercially available stout. I have the equipment to makea 15g. I was also thinking a left hand millk stout clone because bloth reg and nitro versions are available for comparison.
Ferment beer in one or several FV.
Transerfer into 5 seperate 2.5g cornies and one bladder vessel which i can provide.
Keg 1 carb as normal and serve with co2 at high pressure.
Keg 2 carb as normal and push with nitro.
Keg 3 beer gas
Keg 4 remove all co2 from beer using vacuum or wine whip? and put on straight nitro.
Edit
Keg 5 carb as normal and serve with short large diameter line.
Assemble a group of 20-30 taste testers.
Collect data and report results.
Its time to test, analyze, and report findings.
I am willing to help in the testing but currently lack the funds for a nitro or beer gas setup but i do have a simple way to test the one theory of carbing a beer and then dispensing at higher presure with a bladder.
Anyone in the Indy area willing to help with this please PM me.
My idea for this experiment:
Make a large batch of a clone of a commercially available stout. I have the equipment to makea 15g. I was also thinking a left hand millk stout clone because bloth reg and nitro versions are available for comparison.
Ferment beer in one or several FV.
Transerfer into 5 seperate 2.5g cornies and one bladder vessel which i can provide.
Keg 1 carb as normal and serve with co2 at high pressure.
Keg 2 carb as normal and push with nitro.
Keg 3 beer gas
Keg 4 remove all co2 from beer using vacuum or wine whip? and put on straight nitro.
Edit
Keg 5 carb as normal and serve with short large diameter line.
Assemble a group of 20-30 taste testers.
Collect data and report results.
Looking back over the thread it appears that there are many long standing misconceptions about pouring beer with mix. I used to believe, as many appear to, that the CO2 is there to make the bubbles and the N2 to push the undercarbonated (by usual standards) beer hard enough through the sparkler plate to produce the tiny bubbles which I, too, thought contained no nitrogen. I used to argue that one could get the same effect as beer gas by carbing to 1 vol, storing at 1 vol and then temporarily doubling or tripling the CO2 pressure while serving.
What I've gotten from this thread is that the CO2 for carbonation and N2 for pushing model is wrong. The goal is to have a bubble with about equal amounts (partial pressures) of CO2 and N2 in order to dilute the CO2 to the level where the carbonic 'prick' on the tongue is attenuated thus making for a smoother, milder, creamier head. I did not understand this until I did the sums in response to this thread. This, is I think, the most important aspect of the discussion here.
What I've gotten from this thread is that the CO2 for carbonation and N2 for pushing model is wrong. The goal is to have a bubble with about equal amounts (partial pressures) of CO2 and N2 in order to dilute the CO2 to the level where the carbonic 'prick' on the tongue is attenuated thus making for a smoother, milder, creamier head. I did not understand this until I did the sums in response to this thread. This, is I think, the most important aspect of the discussion here.
This is the Brew Science forum. If applying science (math, chemistry) to the properties of beer and brewing offends you, you're in the wrong place!
The entire rest of the forum is available for discussion based on anecdotal evidence and "logical arguments". Those things have value of course, as it's mainly how brewing progressed for millennia, and to a great extent it's how it still progresses.
The entire rest of the forum is available for discussion based on anecdotal evidence and "logical arguments". Those things have value of course, as it's mainly how brewing progressed for millennia, and to a great extent it's how it still progresses.
timdillon36
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I dont know if this is directed at me or in general.
I agree that science and formulas based on theories and laws are very useful. But those theories and laws were not accepted as such without testing with results that back up those hypothesis.
I am in no way educated enough to understand most of the math that has been discussed but i do grasp the concept and would like to test the ideas that have been talked about. If anyone else has acces to equipment to measure results better or has a better plan on how to do this test please speak up.
I make my living working on a professional race car. I have seen time and again the guys at the desk have an idea that should work great on paper but when you take it to the track you find that they were so focused on one part rather then the whole the car is slower.
If we want more than anecdotal evidence maybe some people here that have access to lab grade equipment can test the beer and foam that is poured from different methods.
View this as you want but what i am saying is i am willing to be your lab tech in training. If anyone else that has a more ambitious plan or a better idea on how they plan to carry this to fermentation (or fruition) i believe this needs to get past talking shop.
With science and technology along with testing those ideas we continue to advance and build a better mouse trap.
How many times has palmer been quoted but then proven wrong in the past 20 years.
I do not dispute anything that was said earlier but am being redundent again that it should be tested.
If a bar uses beer gas to serve all beer does that mean the local ipa poured from a bottoms up will now be a nitro ipa?
While I cannot speak for passedpawn I am certain his remarks were not directed at you but out our resident Luddite.
I did post in #85 a simple method for estimating the relative nitrogen/CO2 content of your beer foam/bubbles using nothing more sophisticated than a syringe.
Based on my take on this thread the most important taste test you can do is a comparison (double blind triangle of course) of the heads on the same beer poured with straight CO2 against those poured with mix. Is there a statistically discernable difference in the 'creaminess' of those heads?
I did post in #85 a simple method for estimating the relative nitrogen/CO2 content of your beer foam/bubbles using nothing more sophisticated than a syringe.
Based on my take on this thread the most important taste test you can do is a comparison (double blind triangle of course) of the heads on the same beer poured with straight CO2 against those poured with mix. Is there a statistically discernable difference in the 'creaminess' of those heads?
BitterSipper
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Except that I do, you are equating keeping a discussion simple as ignorance, which is a flaw in your own character.
Again, the diffusion of gas under 3 atmospheres of beer gas is irrelevant to beer in a glass under 1 atmosphere of mostly nitrogen and oxygen, and essentially void of co2. That you keep claiming they are is astounding.
Not to mention that the conditions of pouring a beer are not that of undisturbed progression to to equilibrium, which all of your posts have basically used as a premise, and at least one of your posts has contradicted your own claim.
Great, you're actually moving your experiment towards the conditions of an actual pour of beer, that's good. However, you're still a long ways off. You still aren't diffusing in the proper conditions and you seem to have tried to recreate some of the disturbances of a pour, but I would be aghast if you claimed it was equivalent.
BitterSipper
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I don't, I take issue with the misapplication of principles, and the poor control of variables.
It's an easily understood equivalent. If I can explain to a person that their diatribe is contradicted by basic logic, they should immediately review their statements and address that logic directly. I chose a simple "anecdotal" (not technically, but whatever) route to point out the flaws in his reasoning. If you prefer a plain stating of facts;
There is much more CO2 than nitrogen dissolved in beer under beer gas.
Under 1 atmosphere of pressure of earths air, which is comprised mostly of nitrogen and effectively no CO2, effectively all the CO2 will diffuse out. In addition not all of the already much lower amount of dissolved nitrogen will. I already conceded to him that initially, nitrogen may (in his context) diffuse more quickly. However, there is very little of it and more importantly;
This is ignoring the fact that the beer is highly disturbed during the pouring process. The forcing of the liquid through such a small restriction causes immense conditions. Rapid pressurization and depressurization, high velocities, cavitation etc. What occurs during this process is not explained by simple gas principles (at least not the way they are being applied here), not that they agreed in the first place.
Are you guys having fun with this? I hope so. However, you have
to realize that much research has been done with this topic and
you are essentially trying to reinvent the wheel without the
aid of enough experience in chemistry and without a laboratory.
The solubility of nitrogen in water is not the same at 1 atm of
pressure above the water as it is at higher pressure, nor is the
solubility the same at low and high temperatures. Since CO2
forms carbonic acid in solution, its solubility is affected by
pH, and beer is not neutral. The bubbles that
come out of solution are not the same thing as the foam that forms
above the head, the physical behaviour of the foam is not the same
as the physical behaviour of a bubble. The nitrogen that is in the
foam doesn't have to come solely from the small amount of nitrogen
that was in the liquid, it can diffuse into the foam as it forms
from the bubbles. There are many processes going on including the
initial bubble formation, creaming, disproportionation and drainage,
each of which has its own complex equation. The most important
factors for bubble/foam stability are the solubility of the gas,
which is why nitrogen impacts greatly the stability of the head in
stout, and the initial size of the bubbles that come out of solution
(that is where the widget comes in). There are many many factors
involved which were initially successfully realized in beer purely
by fortuitous experiment but which would probably be impossible
to understand analytically without the input of many minds.
There are many free articles available at the Journal of the Institute
of Brewing, Bamforth's paper "The relative significance of physics
and chemistry for beer foam excellence", 2004
(use http://onlinelibrary.wiley.com/advanced/search and
type in institute of brewing and select publication title and then
select for article title and type in stout on the next line) has
an excellent summary and also a nice table where you can see clearly
the huge difference pure nitrogen versus pure co2 makes in the lifetime
of a bubble. He also references a Ph.D thesis by Ronteltap titled
"Beer Foam Physics" that is also freely availabe (search the author
and title at scholar.google.com).
Ray
to realize that much research has been done with this topic and
you are essentially trying to reinvent the wheel without the
aid of enough experience in chemistry and without a laboratory.
The solubility of nitrogen in water is not the same at 1 atm of
pressure above the water as it is at higher pressure, nor is the
solubility the same at low and high temperatures. Since CO2
forms carbonic acid in solution, its solubility is affected by
pH, and beer is not neutral. The bubbles that
come out of solution are not the same thing as the foam that forms
above the head, the physical behaviour of the foam is not the same
as the physical behaviour of a bubble. The nitrogen that is in the
foam doesn't have to come solely from the small amount of nitrogen
that was in the liquid, it can diffuse into the foam as it forms
from the bubbles. There are many processes going on including the
initial bubble formation, creaming, disproportionation and drainage,
each of which has its own complex equation. The most important
factors for bubble/foam stability are the solubility of the gas,
which is why nitrogen impacts greatly the stability of the head in
stout, and the initial size of the bubbles that come out of solution
(that is where the widget comes in). There are many many factors
involved which were initially successfully realized in beer purely
by fortuitous experiment but which would probably be impossible
to understand analytically without the input of many minds.
There are many free articles available at the Journal of the Institute
of Brewing, Bamforth's paper "The relative significance of physics
and chemistry for beer foam excellence", 2004
(use http://onlinelibrary.wiley.com/advanced/search and
type in institute of brewing and select publication title and then
select for article title and type in stout on the next line) has
an excellent summary and also a nice table where you can see clearly
the huge difference pure nitrogen versus pure co2 makes in the lifetime
of a bubble. He also references a Ph.D thesis by Ronteltap titled
"Beer Foam Physics" that is also freely availabe (search the author
and title at scholar.google.com).
Ray
BitterSipper
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That's my point. I don't claim to know all the conditions that contribute to the forming of the head. What I do know and have been arguing is that it is not as simple as a SS solution moving to equilibrium. And even if it was, the reasoning behind his conclusion under that premise is wrong to begin with.
I mean, he's arguing the content of the head of a beer is equivalent to the gas in a bubble formed in a pressurized keg. His "syringe experiment" use sealed headspace that start out empty, stay sealed, and share no relation to the conditions a beer is poured in.
Even accepting his premise. When challenged he posts a graph showing headspace volume ratio to gas ratio to support his conclusion. Seems he expects people to select how much ratio of beer to head they have and then use that to find the gas mixture. But even then it completely fails to take into account the conditions a beer is poured in, being mostly a full atmosphere of nitrogen, shifting the graph. It also fails to account for any gas being lost while the head is formed and as it's replenished, again shifting the graph. Combined with the fact that it's moot to begin with as a beers head is not formed simply by the process of moving undisturbed to equilibrium.
timdillon36
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Is this project ready to move to protype phase yet?
Of course we recognize that tons of research has been done on this. I saw somewhere what Guiness spent on the development of the widget and it was astronomical. I don't think they would have done that had they not reasonable assurance that success with it would allow them to sell more beer.
One can find out many of the details, and it is a complex problem, by reading the research but then one can get a beer by going to the grocery store and we don't do that here. The Brew Science forum is for those of us who want to understand something of the science of brewing. So what we have done here is given some analysis based on the basic science of the phase interface between gasses and liquids and posed a gedenken experiment or two which would let someone with knowledge of basic chemistry (or physical chemistry if you prefer) understand why it is that even though nitrogen is not very soluble in beer a beer bubble (be it in the foam or in the beer) contains a good bit of nitrogen AND verify by simple experiment that he can do in his own kitchen, that this is indeed the case. If people follow the reasoning (and clearly not all can) and if they do the experiment (suck up bubbles from the foam if you want to see whats in the foam bubbles only) perhaps we have helped dispel the popular notion that mix drawn beer foam contains only CO2. Then they can start to think about and appreciate what the benefits of N2 in the bubbles (in the beer and foam) actually are. Have I had some fun? You betcha!
Your statement that CO2 solubility depends on pH is somewhat misleading. The solubility of CO2 depends on the Henry coefficient for CO2 at the temperature in question. The amount of CO2 that dissolves and the pH of the resulting solution depend on the partial pressure of CO2 and what else is in it. If there is an excess of calcium ion, for example, it is sufficient to know the partial pressure. Yes, the pH and amount of dissolved CO2 are related but the control is the partial pressure of CO2.
Afterthought: Nobody knows more about "boobles" than Charley B!
One can find out many of the details, and it is a complex problem, by reading the research but then one can get a beer by going to the grocery store and we don't do that here. The Brew Science forum is for those of us who want to understand something of the science of brewing. So what we have done here is given some analysis based on the basic science of the phase interface between gasses and liquids and posed a gedenken experiment or two which would let someone with knowledge of basic chemistry (or physical chemistry if you prefer) understand why it is that even though nitrogen is not very soluble in beer a beer bubble (be it in the foam or in the beer) contains a good bit of nitrogen AND verify by simple experiment that he can do in his own kitchen, that this is indeed the case. If people follow the reasoning (and clearly not all can) and if they do the experiment (suck up bubbles from the foam if you want to see whats in the foam bubbles only) perhaps we have helped dispel the popular notion that mix drawn beer foam contains only CO2. Then they can start to think about and appreciate what the benefits of N2 in the bubbles (in the beer and foam) actually are. Have I had some fun? You betcha!
Your statement that CO2 solubility depends on pH is somewhat misleading. The solubility of CO2 depends on the Henry coefficient for CO2 at the temperature in question. The amount of CO2 that dissolves and the pH of the resulting solution depend on the partial pressure of CO2 and what else is in it. If there is an excess of calcium ion, for example, it is sufficient to know the partial pressure. Yes, the pH and amount of dissolved CO2 are related but the control is the partial pressure of CO2.
Afterthought: Nobody knows more about "boobles" than Charley B!
I'm using these double arrows to indicate reversible reactions.
If you increase the amount of H3O+ (in other words, lowering the pH)
you drive the second reaction to the left, increasing the amount of
H2C03, which in turn drives the first reaction to the left. The principle
reason CO2 has the solubility that it does (much greater than say N2 or Ar)
is because it reacts with the water to form a new substance, carbonic
acid (H2CO3). Breaking down the H2CO3 back to CO2 lowers the overall
solubility of the CO2. The lower pH I'm talking about comes from phosphate
in the wort, not the carbonic acid.
You have this fixation on Henry's law. Henry's law is only relevant for
non-reactive gases, not gases like CO2 or NH3 which react in solution.
At most, for a completely neutral pH, you can only have an "effective" Henry's
law constant which is controlled by the equilibrium constant of the
carbonic acid. Otherwise it doesn't apply at all. For example, if you put
one mole of CO2 above water, no matter what pressure, if you have 1 mole
of NaOH in solution (high pH) all of the CO2 will dissolve, because the reaction
of NaOH with carbonic acid is irreversible, forming sodium bicarbonate.
There is in fact no "constant" for Henry's law for CO2 & NH3 because
its value varies with the pH of the water.
Ray
If you increase the amount of H3O+ (in other words, lowering the pH)
you drive the second reaction to the left, increasing the amount of
H2C03, which in turn drives the first reaction to the left. The principle
reason CO2 has the solubility that it does (much greater than say N2 or Ar)
is because it reacts with the water to form a new substance, carbonic
acid (H2CO3). Breaking down the H2CO3 back to CO2 lowers the overall
solubility of the CO2. The lower pH I'm talking about comes from phosphate
in the wort, not the carbonic acid.
You have this fixation on Henry's law. Henry's law is only relevant for
non-reactive gases, not gases like CO2 or NH3 which react in solution.
At most, for a completely neutral pH, you can only have an "effective" Henry's
law constant which is controlled by the equilibrium constant of the
carbonic acid. Otherwise it doesn't apply at all. For example, if you put
one mole of CO2 above water, no matter what pressure, if you have 1 mole
of NaOH in solution (high pH) all of the CO2 will dissolve, because the reaction
of NaOH with carbonic acid is irreversible, forming sodium bicarbonate.
There is in fact no "constant" for Henry's law for CO2 & NH3 because
its value varies with the pH of the water.
Ray
BitterSipper
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nestar:BitterSipper
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Your reasoning is followed fine, it's just full of flaws.
Flawed assumptions you have made;
-All gas released during a pour stays trapped in the head.
-Calculating diffusion rates in a sealed container when those conditions do not match real conditions.
-The core assumption that it's an undisturbed process to equilibrium.
ETC
Where is all God's creation did you ever get that idea????
Henry's Law gives the relationship between partial pressure of CO2 and carbon dioxide in solution. The equation is
CO2 + H20 <----> H2CO3*; pKhy
and I have given a polynomial expansion for pKhy for CO2 in an earlier post. H2CO3* represents the sum of the aqueous CO2 and the carbonic acid. Water chemists commonly work this way for convenience.
Actually only a small amount of the aqueous CO2 converts to carbonic and some of that dissociates
H2CO3* <---> H+ + HCO3-; pK1 = 6.38 (at 20°C)
The two stages of reaction are rolled into the single pK for convenience. Let's drop the star and just write
H2CO3 <---> H+ + HCO3-; pK1 = 6.38 (at 20°C)
understanding H2CO3 to be the 'effective' carbonic acid whose concentration is the Henry coefficient times the partial pressure of CO2.
Some of the bicarbonate further dissociates
HCO3- <---> H+ + CO3-- ; pK2 = 10.38
If calcium is present
Ca++ + CO3-- <---> CaCO3; pKs (a function of temperature)
There is water present
2H2O <---> H3O+ + OH-; pKw ~ 14 (@25° C)
Thus we have 5 chemical equations for this simple system each with accompanying equilibrium equation. A sixth equation is necessary to insure that the sum of all charges in the solution is 0.
All the equations except the first are functions of pH. pH is thus called the Master Variable and the system is solved by finding the pH (by iterative methods such as root bisection or Newton's) which zeroes the charge (6th equation). There is only one independent variable in all this (at a given temperature) and that is the partial pressure of CO2. Once we have the pH we can then solve for the concentrations of all three carbon containing species and the calcium.
All this can be found in a water chemistry text such as Stumm and Morgan and is compactly summarized in a peer reviewed paper which appeared in Cerevesia about 10 years ago.
For those of you that live in alternative universes, I am afraid that this is the way things work. If you think they work in some other way you are simply wrong. I know that is hard to accept but I can't change the facts as much as some might like me too.
Note that the arguments I sometimes get which are generally themed "Well of course that's what you think because the corporations control the universities and force them to teach that garbage so they can stay in control" are relegated to the tin foil hat bin. IOW ignored.
Speaking as a Moderator: Let's keep the emotion and ad hominem attacks out of the discussion. Avoid using words like "Luddite" or "diatribe." Stick to strictly technical arguments, please. Thanks.
doug293cz
doug293cz
How can you have a technical discussion with someone who has no idea what
he's talking about?
Folia Microbiologica
January 1982, Volume 27, Issue 1, pp 5559
Effect of temperature and pH on absorption of carbon dioxide by a free level of mixed solutions of some buffers
Abstract
The rate of absorption of carbon dioxide by solutions of NaHCO3, KH2PO4,
hydrogencarbonate, phosphate and borate buffers at 20, 30 and 40°C was
determined manometrically. The absorption rate increases for all buffers
tested with increasing pH. The CO2 absorption rate by KH2PO4 and by the
phosphate buffer at low pH is lower than that of water. For other buffers
tested it is equal to or higher than that of water, especially at higher
temperatures.
Ray
That's just how life is, isn't it. Whenever you talk to somebody about a subject, one of the parties knows more than the other. It's not uncommon, though, to find that they both think they are the party with the upper hand wrt the knowledge.
I'd think that in scientific areas, it'd be easy to sort things out quickly. But we're human, and egos get in the way.
I have no idea who's right here, and TBH the discussion doesn't even interest me, but I'd like to think that you smart guys could sort it out with numbers and Stoichiometry.
This is simply untrue. Stumm and Morgan, (Aquatic Chemistry, An Introduction Emphasizing Chemical Equilibria in Natural Waters, Wiley, 1981, New York) for example offer Table 4.7 which is a table of values for the CO2 Henry constant as a function of temperature in pure water, sea water and a NaCl solution. pH is mentioned nowhere.
There is slight variation with ionic strength. As in nearly everything we do in brewing we use the chemistry of ideally dilute solutions. Everyone knows that we are not working with ideally dilute solutions but those who know a little chemistry understand that the differences between considering ionic strength (which complicates the math mucho) and not amounts to pH shift of less than 0.1 in most brewing applications.
We don't have trouble with the guys who understand the chemistry. We have trouble with the guys who think they understand it but really don't. Anyone can read an abstract or even a paper and think he understands it but it is clear in some cases that they don't. For example:
Of course it does. The higher pH buffers are more alkaline that is they contain more, in the phosphate example, PO4-3 and HPO4-2 moieties than lower pH buffers which contain relatively more H3PO4 and H2PO4-. If you do the math I have outlined you will find no inconsistency between what this abstract says and that math. But gird thyself. If you do phosphate you have now 5 more equations to go into the system. This system is discussed in the Cerevesia article.
The same is true of the carbonate system i.e. a carbonate buffer built around 10.38 will have approximately equal amounts of bicarb and carbonate and, will have twice the alkalinity of an equimolar buffer at 6.38.
I don't see how one could possibly conclude that the Henry coefficient for CO2 is any different between these two cases. If we are in the amosphere with constant PaCO2 the solutions equilibrium pH's will be driven by that PaCO2 and the alkalinities of the buffers before exposure. The amount of H2CO3* in the final solutions will be the same, PaCO2*Khy in any of the cases mentioned in the abstract (barring ionic strength considerations).
[
This is hardly surprising as weak acids (carbonic; pKa1 = 6.38) don't transfer protons to stronger acids (pH buffer near first pK of phosphoric acid ~ 2.2?). Again I don't see how one can conclude that this is explained by a shift in the Henry coefficient. With the phosphate buffer the carbonic acid content of the equilibrium solution is, again, pACO2*KHy, again, ignoring ionic strength.
I think the problem here is that you have some understanding of acid/base equilibria but not as much as you think you do. Read a textbook.
Yes, it can all be sorted out with the stoichimetry and equilibrium chemistry but those with, shall we say, 'original' concepts of what the equations are and what they say need to get on board with the rest of the scientific community. Equilibrium chemistry is not everyones' favorite. I remember how much I hated it in school and that was relatively speaking just brushing the surface. They certainly never asked us to solve systems of 10 equations.
Stumm and Morgan get my vote.
BitterSipper
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BitterSipper
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Not that guy, look at his graph here.
Which he posted to support his claim when challenged by (paraphrased simply); "ok more nitrogen comes out at first, but there isn't a lot of it dissolved, so it will quickly run out". Which is very valid reasoning (as he proves himself)
The horizontal axis is the volume of the gas that has escaped the beer (head space, which is meant to represent the head of the beer), presented as a ratio to the starting volume of beer.
So in a cylinder 10 inches of beer forming 1 inch of head 10/1 = .1) And this is really the smallest ratio a good pour is going to have.
Then the vertical axis is pressure or ratio depending on which line you're referencing. The red line is the total "pressure dissolved in the beer" which starts at 4 atm, and goes to 0. It's not terribly relevant.
The blue line is the ratio of nitrogen to co2 remaining, and meant to represent the the ratio of what is diffusing at that point, if you follow .1 head ratio, you get to .5 of a gas ratio.
Meaning there is roughly 2 times as much co2 diffusing at this point than nitrogen. That's just at that point however, you have to take into account that the inch of head that already formed contains higher concentration of nitrogen. But already his own "best case" experiment is not fully agreeing with his statement.
There's also just a couple problems here. First, while pouring a beer, not all the gas remains trapped in the head, a lot of it is lost initially and continually. So that .1 ratio for when your beer was first poured is going to drop right away to at least something more like .2, with the corresponding decrease in the ratio of nitrogen to co2. Then gas continues to be lost from the head as you're drinking, compounding this error in reasoning.
Next, this is actually a graph of a beer in a syringe.... where earths atmosphere is not acting on it. Earths atmosphere is about 80% nitrogen and at 1 atmosphere of pressure, meaning that from the start, the diffusion of nitrogen is being decreased by a rather large fraction than this graph would lead you to believe, where as the CO2 is not (it's receiving essentially no resistance from the atmosphere)
Next, all of this was moot to begin with, because all this describes an undisturbed process to equilibrium, which pouring a beer is not.
I've already said, I don't understand every impact on the formation of the head of a beer, as it is FAR more complicated than just an undisturbed process to equilibrium like he wants to pretend. What I do know is that the way he supports his argument is completely flawed.
It's also why I preferred the approach of simply dissolving the same volumes of each gas independently/separately, pouring 2 beers, and comparing the heads. The results were clear with next to no nitrogen head, and a normal sized head with the co2. But, I am the first to admit that this is also not a perfect experiment, due to not controlling for any interactions.
What I'm really trying to do here is to point out to the other posters
that their objections to your scenarios (by that I mean your Rube
Goldberg thought experiments and your random jargon spewing), are valid.
You have a habit of attempting to quash other points of view with a
combination of dogmatic appeal to authority, and more jargon. For example:
"Were it otherwise Henry's law would have to be different for the interface
between the headspace and the beer and the bubble interior and the beer.
How could this be? There is only one Henry's law. If you can accept that
this is the case we can go forward from there. If not read #67 until you
understand it and then we can go forward."
Of course, there certainly could be a difference between the Henry's law
constants for those two cases, and the reason you don't see that is because
you don't think about the physical situation, you just want to crunch
numbers. Don't feel bad: at least you can crunch numbers! But this is the
same problem many physical chemists have, and they either learn to leave
the math behind or they end up repeating the past.
As for the Henry's Law constant itself, while it may look like an equilibrium
constant, and therefore should be constant at constant temperature (and vary
with temperature) it is in fact a ratio of concentrations in different phases
(gas and liquid) and it is therefore an ad hoc experimental observation that
varies with just about everything. I'll leave you with these observations
from a recent review of Henry's Law constants:
Ray
that their objections to your scenarios (by that I mean your Rube
Goldberg thought experiments and your random jargon spewing), are valid.
You have a habit of attempting to quash other points of view with a
combination of dogmatic appeal to authority, and more jargon. For example:
"Were it otherwise Henry's law would have to be different for the interface
between the headspace and the beer and the bubble interior and the beer.
How could this be? There is only one Henry's law. If you can accept that
this is the case we can go forward from there. If not read #67 until you
understand it and then we can go forward."
Of course, there certainly could be a difference between the Henry's law
constants for those two cases, and the reason you don't see that is because
you don't think about the physical situation, you just want to crunch
numbers. Don't feel bad: at least you can crunch numbers! But this is the
same problem many physical chemists have, and they either learn to leave
the math behind or they end up repeating the past.
As for the Henry's Law constant itself, while it may look like an equilibrium
constant, and therefore should be constant at constant temperature (and vary
with temperature) it is in fact a ratio of concentrations in different phases
(gas and liquid) and it is therefore an ad hoc experimental observation that
varies with just about everything. I'll leave you with these observations
from a recent review of Henry's Law constants:
Ray
AS rayg was mislead by the abstract/paper he referenced into thinking this implied that the Henry coefficient shifts with the pH of the solution I thought it might be interesting to look into the paper's finding in a little more detail as it is related to a question that is often asked here which is "How much do my buffers' pH's change after I open the bottle?".
Suppose we are interested in (and the abstract mentions them as one of the types of buffers of interest) carbonate buffers. To answer this question we need to be able to compute the charges on all the carbo related species at an arbitrary pH. These species are carbonic, bicarbonate at carbonate. The first is an undissociated molecule and thus has no charge. The second is monobasic and thus has charge equal to minus the number of moles of bicarbonate. The third is dibasic and thus has charge equal to minus two times the number of moles of carbonate. Let Ct be the total number of moles of carbon per liter in the solution. Then the charge on carbo species per liter is -Ct*(0*f0 + 1*f1 + 2*f2) where f0, f1 and f2 are, respectively, the fractions of the total carbon found in, respectively, arbonic acid, bicarbonate ion and carbonate ion. These fractions depend solely on the pH (in ideally dilute chemistry) and are easily found by computing, in order
r1 = 10^(pHs - 6.38)
r2 = 10^(pHs - 10.38)
f0 = 1/(1 + r1 + r1*r2)
f1 = f0*r1
f2 = f1*r2
6.38 and 10.38 are the pKs of carbonic acid
From the f's one computes the carbo system anion charge as
-Ct*Qc(pH) = -Ct*(1 +2*f2) mEq/L
One also needs to know the charge on water derived ions. This is
Qw(pH) = 1000*(10^-pH - 10^pH-14) with the first term inside the parentheses representing the equivalents of H30+ ions per liter and the second the equivalents of (OH)- ions per liter.
Finally one needs to know the cation charge. Suppose we are going to make our buffer from NaHCO3 and Na2CO3. Suppose further that we know if a buffer contains equal amounts of species with basicity n and n+1 (1 and 2 here) the buffer pH will be equal to the pK associated with proton loss from n to n-1 basicity, in this case 10.38. Hence we take a liter of distilled water and add 1 mmol of NaHCO3 and 1 mmol of Na2CO3 to it expecting pH 10.38. We have added 2 moles of carbo (one each as bicarbonate and carbonate) and 3 moles of sodium. To find the actual pH of this buffer we compute the sum of the charges in the liter of water at trial pH values
Q (pH) = +3 - Ct*Qc(pH) + Qw(pH)
until the value that zeroes Q is found. This is the actual pH of the solution. When one does this for equal amounts of sodium carbonate and bicarbonate one gets:
Zero_Qt(1,1,10.38)
Original: Carbonic 0.000; Bicarbonate 1.000, Carbonate 1.000
Eqilibrium pH sans atmospheric CO2: 10.231
Equilibrium Carbonic 0.00016; Bicarbonate 1.16986, Carbonate 0.82998
Equilibrium Sodium 2.00000; H30+ 0.00000; OH- 0.17019
PaCO2 in infinitessimally small headspace: 0.000004 Atm
So far we have excluded our buffer from exposure to atmospheric CO2. If we allow the solution to come to equilibrium with atmospheric CO2 we proceed in the same way except we now know what the equilibrium carbonic acid concentration is:
1000*Atm*Khy=1000*Atm*10^-1.41
pKhy = 1.41 at 20 °C
This is the carbonic content at any pH. To find the pH at which our buffer will equilibrate we use exactly the same approach as before i.e 0 the sum
Q (pH) = +3 - Ct*Qc(pH) + Qw(pH) = +3 - (1000*Atm*Khy/f0(pH))*Qc(pH) + Qw(pH)
Note that here we use (1000*Atm*Khy/f0(pH)) for the total carbo at each trial pH. Doing this we find, for 0.0003 Atm PaCO2
Zero_Qt_(0.0003,1,1,10.38)
Original Carbonic ; Bicarbonate 1.0000; Carbonate 1.0000
For 0.000300 atmospheres CO2 equilibrium pH is: 8.768
Equilibrium Carbonic: 0.0117; Bicarbonate 2.8545; Carbonic: 1.4679; Carbo Total: 2.9360
Carbo added as bicarb and carbonate salts : 2.00 thus 0.94 additional carbo entered as CO2
And with 0.0004 Atm PaCO2
Zero_Qt_(0.0004,1,1,10.38)
Original Carbonic ; Bicarbonate 1.0000; Carbonate 1.0000
For 0.000400 atmospheres CO2 equilibrium pH is: 8.649
Equilibrium Carbonic: 0.0156; Bicarbonate 2.8883; Carbonic: 1.4787; Carbo Total: 2.9575
Carbo added as bicarb and carbonate salts : 2.00 thus 0.96 additional carbo entered as CO2
rayg confused (by failure to carefully read the material he pasted) the additional inflowing CO2 from the air with the partitioning between air and aqueous CO2 (not CO2 + bicarbonate + carbonate). The Henry coefficient describes the latter, not the former.
The other things of major import here is that the 'obsession' with the Henry law is illustrated nicely. It is the Henry law that determines the final pH of the buffer. A change of only 0.0001 Atm changes our equilibrium buffer pH by 0.11 pH. The implications with respect to change in ocean pH as more CO2 enters the atmosphere are clear and explain why this is of such concern to climatologists.
Suppose we are interested in (and the abstract mentions them as one of the types of buffers of interest) carbonate buffers. To answer this question we need to be able to compute the charges on all the carbo related species at an arbitrary pH. These species are carbonic, bicarbonate at carbonate. The first is an undissociated molecule and thus has no charge. The second is monobasic and thus has charge equal to minus the number of moles of bicarbonate. The third is dibasic and thus has charge equal to minus two times the number of moles of carbonate. Let Ct be the total number of moles of carbon per liter in the solution. Then the charge on carbo species per liter is -Ct*(0*f0 + 1*f1 + 2*f2) where f0, f1 and f2 are, respectively, the fractions of the total carbon found in, respectively, arbonic acid, bicarbonate ion and carbonate ion. These fractions depend solely on the pH (in ideally dilute chemistry) and are easily found by computing, in order
r1 = 10^(pHs - 6.38)
r2 = 10^(pHs - 10.38)
f0 = 1/(1 + r1 + r1*r2)
f1 = f0*r1
f2 = f1*r2
6.38 and 10.38 are the pKs of carbonic acid
From the f's one computes the carbo system anion charge as
-Ct*Qc(pH) = -Ct*(1 +2*f2) mEq/L
One also needs to know the charge on water derived ions. This is
Qw(pH) = 1000*(10^-pH - 10^pH-14) with the first term inside the parentheses representing the equivalents of H30+ ions per liter and the second the equivalents of (OH)- ions per liter.
Finally one needs to know the cation charge. Suppose we are going to make our buffer from NaHCO3 and Na2CO3. Suppose further that we know if a buffer contains equal amounts of species with basicity n and n+1 (1 and 2 here) the buffer pH will be equal to the pK associated with proton loss from n to n-1 basicity, in this case 10.38. Hence we take a liter of distilled water and add 1 mmol of NaHCO3 and 1 mmol of Na2CO3 to it expecting pH 10.38. We have added 2 moles of carbo (one each as bicarbonate and carbonate) and 3 moles of sodium. To find the actual pH of this buffer we compute the sum of the charges in the liter of water at trial pH values
Q (pH) = +3 - Ct*Qc(pH) + Qw(pH)
until the value that zeroes Q is found. This is the actual pH of the solution. When one does this for equal amounts of sodium carbonate and bicarbonate one gets:
Zero_Qt(1,1,10.38)
Original: Carbonic 0.000; Bicarbonate 1.000, Carbonate 1.000
Eqilibrium pH sans atmospheric CO2: 10.231
Equilibrium Carbonic 0.00016; Bicarbonate 1.16986, Carbonate 0.82998
Equilibrium Sodium 2.00000; H30+ 0.00000; OH- 0.17019
PaCO2 in infinitessimally small headspace: 0.000004 Atm
So far we have excluded our buffer from exposure to atmospheric CO2. If we allow the solution to come to equilibrium with atmospheric CO2 we proceed in the same way except we now know what the equilibrium carbonic acid concentration is:
1000*Atm*Khy=1000*Atm*10^-1.41
pKhy = 1.41 at 20 °C
This is the carbonic content at any pH. To find the pH at which our buffer will equilibrate we use exactly the same approach as before i.e 0 the sum
Q (pH) = +3 - Ct*Qc(pH) + Qw(pH) = +3 - (1000*Atm*Khy/f0(pH))*Qc(pH) + Qw(pH)
Note that here we use (1000*Atm*Khy/f0(pH)) for the total carbo at each trial pH. Doing this we find, for 0.0003 Atm PaCO2
Zero_Qt_(0.0003,1,1,10.38)
Original Carbonic ; Bicarbonate 1.0000; Carbonate 1.0000
For 0.000300 atmospheres CO2 equilibrium pH is: 8.768
Equilibrium Carbonic: 0.0117; Bicarbonate 2.8545; Carbonic: 1.4679; Carbo Total: 2.9360
Carbo added as bicarb and carbonate salts : 2.00 thus 0.94 additional carbo entered as CO2
And with 0.0004 Atm PaCO2
Zero_Qt_(0.0004,1,1,10.38)
Original Carbonic ; Bicarbonate 1.0000; Carbonate 1.0000
For 0.000400 atmospheres CO2 equilibrium pH is: 8.649
Equilibrium Carbonic: 0.0156; Bicarbonate 2.8883; Carbonic: 1.4787; Carbo Total: 2.9575
Carbo added as bicarb and carbonate salts : 2.00 thus 0.96 additional carbo entered as CO2
rayg confused (by failure to carefully read the material he pasted) the additional inflowing CO2 from the air with the partitioning between air and aqueous CO2 (not CO2 + bicarbonate + carbonate). The Henry coefficient describes the latter, not the former.
The other things of major import here is that the 'obsession' with the Henry law is illustrated nicely. It is the Henry law that determines the final pH of the buffer. A change of only 0.0001 Atm changes our equilibrium buffer pH by 0.11 pH. The implications with respect to change in ocean pH as more CO2 enters the atmosphere are clear and explain why this is of such concern to climatologists.
In yesterday's post I gave an example of how far a buffer's pH might be pulled by atmospheric CO2. In that example the buffer was made by adding a millimole each of bicarbonate and carbonate to DI water. The pH drops from 10.231 to 8.768 if PaCO2 is 0.0003 Atm. I want to point out that a stronger buffer will be less influenced, One made with 5 mmol each would have an original pH of 10.342 and would wind up at 9.405. So the amount your 10 buffer's pH is going to decrease with exposure to air depends not only on the CO2 content of the air but on the strength of the buffer.
End of Science content.
Unfortunately I do feel it incumbent on me to respond to the increasing intensity of the ad hominem attacks despite the moderators' attempts (and I think they were admirable) to stifle them. The only only recent comment worthy of response IMO is the one that says I fall back on authority. Of course I do. I didn't invent this stuff (though I did re-invent a lot of it). The reason I insist on being consistent with authority is because I want anyone who uses my stuff (like the stuff in the stickies) to be confident that he is getting the straight skinny (or at least consistent with current industry practice). What I post is based on good authority. In general these are texts accepted by the industry as having good working models. I use those models. I mention Stumm and Morgan frequently as an authority as that is the most robust, but probably not the best from the point of view of accessibility. The most accessible is Palmer's Water but as he got what he has to say on this subject from me you may want to 'consider the source' if you are sympathetic to the anti duo's stand. Another good one is Faust and Aly 'Chemistry of Water Treatment' (Ann Arbor Press) and finally (though there are doubtless others) MWH's 'Water Treatment Principles and Design' (Wiley). Anyone can quote stuff out of context in an attempt to prove a point as has been done in this thread. I will quote something in context from the last of the references I gave above:
"The pH does not affect the Henry's constant directly but it does affect the distribution of species between ionized and unionized forms which influences the overall gas-liquid distribution of the compound because only the unionized species are volatile" (p1175).
This is the point rayg consistently misses. The buffer example post of yesterday explains exactly how this works and if anyone wants to see in more detail there is enough info in that post and the Sticky on carbonate to allow him to reproduce that work. If the detractors were willing to do that they might come on board.
End of Science content.
Unfortunately I do feel it incumbent on me to respond to the increasing intensity of the ad hominem attacks despite the moderators' attempts (and I think they were admirable) to stifle them. The only only recent comment worthy of response IMO is the one that says I fall back on authority. Of course I do. I didn't invent this stuff (though I did re-invent a lot of it). The reason I insist on being consistent with authority is because I want anyone who uses my stuff (like the stuff in the stickies) to be confident that he is getting the straight skinny (or at least consistent with current industry practice). What I post is based on good authority. In general these are texts accepted by the industry as having good working models. I use those models. I mention Stumm and Morgan frequently as an authority as that is the most robust, but probably not the best from the point of view of accessibility. The most accessible is Palmer's Water but as he got what he has to say on this subject from me you may want to 'consider the source' if you are sympathetic to the anti duo's stand. Another good one is Faust and Aly 'Chemistry of Water Treatment' (Ann Arbor Press) and finally (though there are doubtless others) MWH's 'Water Treatment Principles and Design' (Wiley). Anyone can quote stuff out of context in an attempt to prove a point as has been done in this thread. I will quote something in context from the last of the references I gave above:
"The pH does not affect the Henry's constant directly but it does affect the distribution of species between ionized and unionized forms which influences the overall gas-liquid distribution of the compound because only the unionized species are volatile" (p1175).
This is the point rayg consistently misses. The buffer example post of yesterday explains exactly how this works and if anyone wants to see in more detail there is enough info in that post and the Sticky on carbonate to allow him to reproduce that work. If the detractors were willing to do that they might come on board.
Last warning. Discuss the technical merits. Refute technical arguments with better technical arguments. Do not go after motives, character, or intelligence of other members.
I'm putting this thread in "time out" for a few days to let everyone cool off.
doug293cz
I'm putting this thread in "time out" for a few days to let everyone cool off.
doug293cz
I'm going to open this thread up for continued discussion. Everyone play nice, or it will get shut down again.
Brew on
Brew on
A bit of additional insight: If we have a glass of stout of volume V0 topped by a mass of bubbles with total gas volume Vb and internal pressure Pb there exists a simple relationship between that pressure and the head partial pressure,P0, of the gas in question in the keg. This is easily derived from nothing more than the Henry Law, the Ideal Gas Law and conservation of mass. The gas in the bubbles comes from the beer. The Henry Law tells us how much gas was dissolved in the beer in the keg and, for a given bubble volume the pressure within the bubbles given the concentration of gas within the beer which is, of course, reduced as bubbles form. The formula is
(Pb/P0) = 1/ ( 1 +(Vb/V0)*(1/(KHy*R*T)) )
Thus the ratio of the pressure in the bubbles to the pressure in the keg is a function of the ratio of the volume of the bubbles to the volume of the beer (which we assume does not change appreciably as the foam drains). The two Henry coefficients of interest are, for nitrogen 6.3E-4 molL^-1Atm^-1 and, for CO2, 3.4E-2. R*T has value 24.06 LAtmmol^-1 (so that KHy*R*T is dimensionless as it must be).
If we put 1 ATM CO2 partial pressure and 3 ATM N2 partial pressure into the formula and experiment with different volumes for the bubbles we find that 20.86% expansion results in 1 Atm total gas pressure in the bubbles. At that level of expansion starting from 1 Atm CO2 we would have 0.8 Atm CO2 in the bubbles and, starting with 2 Atm N2 we would have 0.2 Atm of that.
But we know the head on a glass of stout isn't 20% of the volume and we know by experiment that the partial pressure of nitrogen in a stout bubble is greater than 20%. What happened? As soon as the bubbles are exposed to air CO2 starts to diffuse out at a pretty good clip because the partial pressure of CO2 inside is 0.8 atm, the partial pressure of CO2 on the outside is 0.0003 - 0.0004 Atm (there is a large chemical potential gradient across the bubble membrane) and CO2 is highly soluble in the membrane liquid. Conversely, N2 starts to diffuse IN because the partial pressure of N2 outside the bubble is 0.8 Atm and only 0.2 Atm within. But this inward diffusion doesn't amount to much as the chemical potential gradient is much smaller and nitrogen is not very soluble in the membrane.
Should any still wish to claim that the Henry Law does not apply to CO2 because it dissociates, we note that at beer pH nearly all the carbo is in the form of Carbonic acid (99% at typical ale pH).
(Pb/P0) = 1/ ( 1 +(Vb/V0)*(1/(KHy*R*T)) )
Thus the ratio of the pressure in the bubbles to the pressure in the keg is a function of the ratio of the volume of the bubbles to the volume of the beer (which we assume does not change appreciably as the foam drains). The two Henry coefficients of interest are, for nitrogen 6.3E-4 molL^-1Atm^-1 and, for CO2, 3.4E-2. R*T has value 24.06 LAtmmol^-1 (so that KHy*R*T is dimensionless as it must be).
If we put 1 ATM CO2 partial pressure and 3 ATM N2 partial pressure into the formula and experiment with different volumes for the bubbles we find that 20.86% expansion results in 1 Atm total gas pressure in the bubbles. At that level of expansion starting from 1 Atm CO2 we would have 0.8 Atm CO2 in the bubbles and, starting with 2 Atm N2 we would have 0.2 Atm of that.
But we know the head on a glass of stout isn't 20% of the volume and we know by experiment that the partial pressure of nitrogen in a stout bubble is greater than 20%. What happened? As soon as the bubbles are exposed to air CO2 starts to diffuse out at a pretty good clip because the partial pressure of CO2 inside is 0.8 atm, the partial pressure of CO2 on the outside is 0.0003 - 0.0004 Atm (there is a large chemical potential gradient across the bubble membrane) and CO2 is highly soluble in the membrane liquid. Conversely, N2 starts to diffuse IN because the partial pressure of N2 outside the bubble is 0.8 Atm and only 0.2 Atm within. But this inward diffusion doesn't amount to much as the chemical potential gradient is much smaller and nitrogen is not very soluble in the membrane.
Should any still wish to claim that the Henry Law does not apply to CO2 because it dissociates, we note that at beer pH nearly all the carbo is in the form of Carbonic acid (99% at typical ale pH).
Just to dispute your last sentence, here is the relevant portion of
the University of Wisconsin's summary of carbonic acid chemistry, available
online free by searching the exact phrase "Chemistry of Carbonic Acid
Equilibria in Water":
Furthermore, no one has claimed "that the Henry Law does not apply to CO2
because it dissociates". What has been said is that the precise value of
the Henry's Law constant that you keep throwing out as exact is not
exact; it is temperature and pH dependent. The reason that point has been
brought out is because you continue to generate scenarios that include
lots of impressive looking formulas and "exact" values that not only are
not exact but entirely miss the true picture of what is happening when
a can of beer is opened, and your posts give the impression of correctness
in part by quoting these exact values, and they are entirely misleading.
Henry's Law applies to equilibrium conditions *only*. A can of beer is at
equilibrium before it is opened (assuming that the temperature of the
beer itself has been held constant), and it is at equilibrium after it has
been opened only after all the excess gas in solution has left, and the
rate at which the gases in solution are leaving the solution and the rate
at which they are entering solution *do not change*. If you see bubbles
coming out of solution, the solution is still saturated or supersaturated
with gas, and is not at equilibrium. When the solution is at equilibrium,
you will see nothing. The entire process of gas release and foam development
have nothing to do with Henry's Law, as another poster has tried to tell
you. Gases bubbling out of solution are *phase changes*, and by definition
phase changes are not an equilibrium process, which is why chemists have
had (and still have) such a hard time modeling them mathematically.
The referenced piece here says exactly what all the other references I have posted say as of, of course, it would be expected to and thus supports what I have been promulgating. The Henry coefficient does not change with pH. The distribution of the different carbo species does however and this is why, as I mentioned in previous posts, one has to consider a system of 5 equations,
1)Henry law
2) carbonic <--> bicarbonate
3) bicarbonate <--> carbonate
4) dissociation of water
5)electrical neutrality
in order to solve these systems. I have explained this is several posts and, of course, it is explained in my references and yours. If you cannot read and understand any of this material then you aren't going to be able to understand what is at play here and there isn't really much point in me repeating how things really are with new references each time only to get the response that this isn't the way things are with the addition of another reference that illustrates that indeed this is the way they are.
Now with respect to things like equilibrium.
I have been careful to point out in many of my posts that the numbers we can come up with practically use thermodynamics to guide us to insight, that thermodynamics has its limitiations and that, therefore, our models are not perfect. Now perhaps you are not familiar with what a mathematical model is. It is a set of equations or an algorithm which adequately represents a real system to the extent that the numbers produced by it give an adequately accurate prediction of what will be found in the actual system. We are brewers here, not PHD's at Harvard's Baker Lab. The models we use, including the one in the multiple references offered by both of us are adequate for the purpose at hand as evidenced by their wide acceptance in the brewing and other industries (water treatment, environmental studies etc.). At the same time we have to recognize that thermodynamics tells us only that a reaction is feasible but not how long it will take for that reaction to come to equilibrium nor indeed even if it will take place. We all know that 2H2 + O2 ---> 2H2O is thermodynamically feasible and happens very fast but only if a spark is provided. We also know that H2CO3 ---> H2O + CO2 (bubbles forming in beer) is thermodynamically feasible, happens pretty fast but not that fast and does not require a spark (though agitation speeds it up). With some perspective, some knowledge of the chemistry and physics and some experience one learns which thermodynamical models are applicable to a particular situation. I posit models that are applicable to the problems I apply them to based on this experience but usually, as I did here, verify by experiment. Stout bubbles contain about 1/3 nitrogen after a few minutes.
As I said in my last post the new formula depends only on the Henry Law and the Ideal Gas Law. Someone with basic knowledge of chemisty/physics should appreciate that the model would be improved if the non ideal behaviour of the gasses involved was considered, for example, the Henry coefficient adjusted for the actual temperature and the ionic strength of the solution considered (I have mentioned all these previously) but someone with experience in these matters should appreciate that at the level of accuracy of interest here to do so would be folly.
A glass of beer with foam on top is not at equilibrium - nothing ever is - but it is approaching equilibrium assymptotically just as mash pH approaches a value it never completely reaches though the rate of change becomes very slow.
There are no phase changes here. Nothing is boiling or freezing or precipitating. There are four phases in this case: The beer, the bubble wall, the gas in the interior of the bubble and the air above. The N2 and CO2 move between these phases. To say that the Henry law's prediction of partition of the gas contents in those phases is not valid here is beyond absurd as that is precisely what the Henry law does and especially as the results a model based on its use are confirmed by experiment.
At least by early experiments. I'd be posting lots more experimental data but my stout keg kicked and it's going to be another week or so before the new one is equilibrated to beer mix (I keg with 1 vol CO2 and draw on 3:1::N2CO2 mix).
1)Henry law
2) carbonic <--> bicarbonate
3) bicarbonate <--> carbonate
4) dissociation of water
5)electrical neutrality
in order to solve these systems. I have explained this is several posts and, of course, it is explained in my references and yours. If you cannot read and understand any of this material then you aren't going to be able to understand what is at play here and there isn't really much point in me repeating how things really are with new references each time only to get the response that this isn't the way things are with the addition of another reference that illustrates that indeed this is the way they are.
Now with respect to things like equilibrium.
I have been careful to point out in many of my posts that the numbers we can come up with practically use thermodynamics to guide us to insight, that thermodynamics has its limitiations and that, therefore, our models are not perfect. Now perhaps you are not familiar with what a mathematical model is. It is a set of equations or an algorithm which adequately represents a real system to the extent that the numbers produced by it give an adequately accurate prediction of what will be found in the actual system. We are brewers here, not PHD's at Harvard's Baker Lab. The models we use, including the one in the multiple references offered by both of us are adequate for the purpose at hand as evidenced by their wide acceptance in the brewing and other industries (water treatment, environmental studies etc.). At the same time we have to recognize that thermodynamics tells us only that a reaction is feasible but not how long it will take for that reaction to come to equilibrium nor indeed even if it will take place. We all know that 2H2 + O2 ---> 2H2O is thermodynamically feasible and happens very fast but only if a spark is provided. We also know that H2CO3 ---> H2O + CO2 (bubbles forming in beer) is thermodynamically feasible, happens pretty fast but not that fast and does not require a spark (though agitation speeds it up). With some perspective, some knowledge of the chemistry and physics and some experience one learns which thermodynamical models are applicable to a particular situation. I posit models that are applicable to the problems I apply them to based on this experience but usually, as I did here, verify by experiment. Stout bubbles contain about 1/3 nitrogen after a few minutes.
As I said in my last post the new formula depends only on the Henry Law and the Ideal Gas Law. Someone with basic knowledge of chemisty/physics should appreciate that the model would be improved if the non ideal behaviour of the gasses involved was considered, for example, the Henry coefficient adjusted for the actual temperature and the ionic strength of the solution considered (I have mentioned all these previously) but someone with experience in these matters should appreciate that at the level of accuracy of interest here to do so would be folly.
A glass of beer with foam on top is not at equilibrium - nothing ever is - but it is approaching equilibrium assymptotically just as mash pH approaches a value it never completely reaches though the rate of change becomes very slow.
There are no phase changes here. Nothing is boiling or freezing or precipitating. There are four phases in this case: The beer, the bubble wall, the gas in the interior of the bubble and the air above. The N2 and CO2 move between these phases. To say that the Henry law's prediction of partition of the gas contents in those phases is not valid here is beyond absurd as that is precisely what the Henry law does and especially as the results a model based on its use are confirmed by experiment.
At least by early experiments. I'd be posting lots more experimental data but my stout keg kicked and it's going to be another week or so before the new one is equilibrated to beer mix (I keg with 1 vol CO2 and draw on 3:1::N2CO2 mix).
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Correction: I don't draw stout with 4 Atm (absolute) pressure as I think I may have stated in some earlier posts. It's actually 3 (30 psig, 45 psia). This means equilibrium partial pressures in the keg headspace are, for a 3::1 mix, 3/4 Atm for the CO2 and 9/4 Atm for the N2. Sticking those in the new model we find the ratio of bubble volume to beer volume is 9.0%. Putting that into the individual Henry law relationships for N2 and CO2 we find the N2 partial pressure inside the bubble to be 32% of an atmosphere and the CO2 to be 68%. This is in better agreement both with my measurements and those reported by others. Also note that this is before the CO2 starts to diffuse out enriching the N2 even more.
I decided to go check the stout and it is far from eqilibrium at this point but I'd thought I'd check it anyway and the N2 content in the bubbles is, at this point, 5% if that. Bled the headspace (duh!) and drew a couple of pints so now it is sitting under 25:75::N2:CO2 and will come along a bit faster. The reason I am posting now is that I've refined the test procedure a bit and, as I think that the test procedure is probably the most significant thing to come out of this thread I want to make people aware of how they can perform it if they want to.
You will need a horse size syringe (100 mL) and a short piece of tubing plus a smidge of a higher alcohol such as hexanol or octanol. If you can't get those try a wee bit of liquid detergent.
1)Pull the plunger out and moisten the tip (just a small drop or 2) with hexanol or octanol. This will collapse the foam.
2)Connect a short piece of tubing to the tip of the syringe
3)Draw a glass of stout.
4)Push the plunger all the way into the syringe
5)Lower the tip of the tubing far enough into the glass that the tip is in the liquid.
6)Draw the plunger back just enough that a bit of liquid enters the syringe. Expel it. This is to purge the tube of any air (and the N2 in it)
7)Carefully pull the tip of the tube back until it is in the foam
8)Draw the plunger back very slowly being sure to suck up only foam, not beer (unless you want to check the beer gasses).
9)Bring the plunger back to 50 - 70 mL. The foam will collapse so you will have a fair amount of liquid in there as well as gas.
10)Holding the syringe vertically withdraw the tip from the foam.
11)Carefully expel all the liquid and not the plunger reading. That is the volume of gas you have collected.
12)Quicky move the tip of the tube into a strong solution of lye. Draw in a few mL but leave the tip submerged.
13)Shake the syringe without letting the tubing end leave the lye. More lye solution will be drawn in.
14)When things are stable remove the tip from the lye solution and again expel all the liquid
15)The CO2 has been removed. Read the volume in the syringe again. This is the N2 volume in the beer (plus anything else of which there should be precious little).
I hope a few guys will try this and have fun with it. Note that it is designed to keep air out of the equation and so be very careful that no air sneaks in. In particular be careful when pulling back the plunger that none sneaks in around the plunger seal.
[EDIT] It occurred to me last night that if one, whilst holding the syringe vertically, poured some water into it from the top that this would insure that anything sneaking around the plunger would be liquid and thus largely solve the problem of air ingress by that route. Experiment confirmed this but I noted that there can still be some air trapped between the driving bands of the plunger. This can be worked out too.
You will need a horse size syringe (100 mL) and a short piece of tubing plus a smidge of a higher alcohol such as hexanol or octanol. If you can't get those try a wee bit of liquid detergent.
1)Pull the plunger out and moisten the tip (just a small drop or 2) with hexanol or octanol. This will collapse the foam.
2)Connect a short piece of tubing to the tip of the syringe
3)Draw a glass of stout.
4)Push the plunger all the way into the syringe
5)Lower the tip of the tubing far enough into the glass that the tip is in the liquid.
6)Draw the plunger back just enough that a bit of liquid enters the syringe. Expel it. This is to purge the tube of any air (and the N2 in it)
7)Carefully pull the tip of the tube back until it is in the foam
8)Draw the plunger back very slowly being sure to suck up only foam, not beer (unless you want to check the beer gasses).
9)Bring the plunger back to 50 - 70 mL. The foam will collapse so you will have a fair amount of liquid in there as well as gas.
10)Holding the syringe vertically withdraw the tip from the foam.
11)Carefully expel all the liquid and not the plunger reading. That is the volume of gas you have collected.
12)Quicky move the tip of the tube into a strong solution of lye. Draw in a few mL but leave the tip submerged.
13)Shake the syringe without letting the tubing end leave the lye. More lye solution will be drawn in.
14)When things are stable remove the tip from the lye solution and again expel all the liquid
15)The CO2 has been removed. Read the volume in the syringe again. This is the N2 volume in the beer (plus anything else of which there should be precious little).
I hope a few guys will try this and have fun with it. Note that it is designed to keep air out of the equation and so be very careful that no air sneaks in. In particular be careful when pulling back the plunger that none sneaks in around the plunger seal.
[EDIT] It occurred to me last night that if one, whilst holding the syringe vertically, poured some water into it from the top that this would insure that anything sneaking around the plunger would be liquid and thus largely solve the problem of air ingress by that route. Experiment confirmed this but I noted that there can still be some air trapped between the driving bands of the plunger. This can be worked out too.
I'm here for some troubleshooting assistance. I've scoured the web, and this forum, and still havent been able to nail down a solution to my problem.
I can appreciate all the science you guys are discussing, but I'm not sure it's given me an answer for what I'm dealing with.
Here's my Guinness setup.....
Old, converted fridge into kegerator.
The Guinness kit from KegWorks, with coupler, faucet, regulator, and a 5lb bottle. Around 5' of gas and beer line.
My 5lb bottle was filled locally, and I requested 75% nitro, and 25% CO2. I have no way to verify this was done correctly.
Based on what I've seen on the web, I set my regulator at 30psi.
My Guinness half barrel sat in the cold fridge for a week before I tapped it.
The moment I tapped the keg, I poured a pint. The head had large bubbles in it. After letting it sit for a few minutes, the head resembled pancakes as the first side cooks and the bubbles rise to the surface. Large bubbles. I get a decent cascade after the pour, but the bubbles and head texture are all off.
Now, a week later, I tried dropped the regulator pressure, I bled the keg off, and it poured really flat, so..... back up to about 20-25psi. I'm still getting big bubbles.
Most of the research I've searched for was regarding big bubbles in a nitro beer, and the cause, but I'm unable to find an answer.
The beer tastes fine, but I miss the creamy head texture and would like to resolve this.
Is it possibly my gas mix?
Thanks in advance!
I can appreciate all the science you guys are discussing, but I'm not sure it's given me an answer for what I'm dealing with.
Here's my Guinness setup.....
Old, converted fridge into kegerator.
The Guinness kit from KegWorks, with coupler, faucet, regulator, and a 5lb bottle. Around 5' of gas and beer line.
My 5lb bottle was filled locally, and I requested 75% nitro, and 25% CO2. I have no way to verify this was done correctly.
Based on what I've seen on the web, I set my regulator at 30psi.
My Guinness half barrel sat in the cold fridge for a week before I tapped it.
The moment I tapped the keg, I poured a pint. The head had large bubbles in it. After letting it sit for a few minutes, the head resembled pancakes as the first side cooks and the bubbles rise to the surface. Large bubbles. I get a decent cascade after the pour, but the bubbles and head texture are all off.
Now, a week later, I tried dropped the regulator pressure, I bled the keg off, and it poured really flat, so..... back up to about 20-25psi. I'm still getting big bubbles.
Most of the research I've searched for was regarding big bubbles in a nitro beer, and the cause, but I'm unable to find an answer.
The beer tastes fine, but I miss the creamy head texture and would like to resolve this.
Is it possibly my gas mix?
Thanks in advance!
Obviously you need to verify that they gave you beer gas and not just plain CO2. That would be an obvious explanation. The other thing that comes to mind is to make sure that the sparkler plate is properly installed in the faucet which must be a special faucet for stout which I assume to be the case as you evidently bought a kit. 20 - 25 psig is the right pressure range for sure so that shouldn't be a problem. Perhaps just a little patience. Set things up properly and draw a few pints to see if things settle in.
First of all I'd like to thank AJ for all the information provided here. I've always suspected that the "nitro is there only for pushing the beer at higher pressure" mantra was misleading and you proved that.
I have one question though: do you think that the explanation you gave about the N2 ratio in beer foam also relates to nitro coffee foam?
Nitro coffee uses 100% nitro and also presents the creamy head and cascade effect similar to a Guinness. In order to achieve that effect in less time the coffee guys are using a higher pressure (45 psi) and a 0.5 micron diffusion stone. This probably results in more (and faster) N2 dissolved and more N2 bubbles I guess.
Also, have you considered using diffusion stones for achieving faster "nitrogenation" of your stouts?
As everything I know about nitro coffee I learned today I can hardly offer up myself as an expert on the subject but from what I have read today it would seem that the basic principles that make stout foam what it is would apply to coffee treated similarly.
I note from a couple of the web articles that I read this morning that CO2 is sometimes used with coffee. It is always used, of course, with Guiness - perhaps because that carbonic prick is desired with beer.
No because when the beer goes into the keg it is already at 1 volume CO2 from conditioning in the unitank. This, with a sparkler plate, produces a good, fine head. Not as good and fine as when the N2 equilibrates a couple of weeks later but I am willing to wait that long rather than fiddle with plumbing stout mix to the unitanks.
