Baking Soda

[der wo]

Ferrum,
perhaps you are irrtated by my comments to your posts. I just realize that I understood you wrong. I read "I don't think you should rise the pH to 7 before running." You wrote "reduce". Sorry. perhaps now you understand better why I posted...

[Ferrum]

If you dissolve acetic acid in water, it will for the most part NOT split up into ions. That’s why it is a weak acid. A small part will be CH3COO- and H+ and the rest is simply dissolved CH3COOH. The part that splits up depends on pH however. A weak solution will have a nearly neutral pH and will be split up for a bigger % than a strong solution that will have a lower pH, because it becomes increasing harder to donate the H+ ion as pH drops.

If you take a very diluted vinegar (I’m going to talk about vinegar instead of ethyl acetate because we can assume the ester will be fully broken down) solution, such as our low wines and raise the pH to 11 like der wo, a very big % of the acetic acid will be in acetate (CH3COO-). This would be along the lines of 99.999%+. If you then use any acid to bring down the pH to 7 you will slowly start to give H+ back to the acetate ions and create some acetic acid molecules. This is essentially a buffer solution, because you will have both CH3COOH and CH3COO- coexisting. If you perform a titration using either a strong acid or base, you will notice this in that the pH will not change much at first for both the acid and base. If I had an indicator here at home I would try it because it would be easy to find the equilibrium constant between CH3COOH and CH3COO-, but I don’t have one. Maybe beetroot would work?

I think you are right with your experiment Bagasso, in that at pH the acid form is a very small amount and may not be noticeable. However, it could possibly become noticeable when vinegar carries over into your distillate and turns into ethyl acetate. I don’t know because I don’t reduce the pH back to neutral because I never saw a reason to :p I will do this experiment at one point to find out.

Im not sure what you mean with retaining thiols. I’ve hardly seen anything about those. In general people believe copper removes sulphides but I’m not sure if that works with thiols. I do know they are rather acidic however, much more so than their oxygen analog (alcohols). So I would say having a strongly basic still charge would retain them better.

[der wo]

Thanks Ferrum for correcting some of my thoughts and for posting further details.

A bit off-topic, but perhaps you are the right person to ask this:
I triple distill with a packed column and I use caustic soda. After the last distillation I air it a bit, undiluted and then diluted. Then it doesn't have any ethyl acetate smell anymore. But I think it comes back a bit. Not strong, but it's there. A neutral made a few years ago will have a little bit ethyl acetate smell. I think new ethyl acetate is formed in the meantime. Do you think, this is possible or perhaps even normal? And if yes, would it be possible to prevent this by using hard water for dilution instead of RO water? Would a few Ca+ solve the problem?

[Ferrum]

Whether it is normal I can't say, I don't have enough practical experience to do so. Although I haven't tasted it in better commercial vodka's, I have no idea how long they're aged (probably not long).

Something that comes to mind is having read that aging licors get more acidic over time (dont remember where). For me the most logical explanation, altho it is a speculation, is that ethanol (and higher alcohols) can get oxidized into ethanal and then acetic acid, which will turn into ethyl acetate. If this is true you could prevent it by not over-oxidizing your booze. Of course, whiskeys and such get heavily oxidized by aging in oak for 12+ years, but there the ethyl acetate would not be a big problem.

In theory Ca2+ could coordinate acetate ions and 'protect' them from further oxidation or precipetate. I think the concentration would be far too low for precipitation though. And I feel it would taste wrong to dilute with hard water.

[Ferrum]

Ferrum,
perhaps you are irrtated by my comments to your posts. I just realize that I understood you wrong. I read "I don't think you should rise the pH to 7 before running." You wrote "reduce". Sorry. perhaps now you understand better why I posted...

I completely missed this post, sorry. No I was not irratated at all, but I got confused whether or not you agreed with me :p

[Bagasso]

The acetate-ion is never attached to the base-ion. Except you would boil it down under the solubility point. Acids and salts in solution are always splitted up to ions. Acetic acid is CH3COOH. In solution it is split up into CH3COO- and H+. And for example caustic soda (NaOH) is in solution Na+ and OH-. And their salt sodium acetate (NaCH3COO) is in solution Na+ and CH3COO-. Sodium acetate is very soluble, so "in reality" it doesn't exist unless you boil the solution down. But for example the salt of sulphuric acid and lime is very insoluble. So if you mix those substances in water, you get calcium sulfate, it will settle out.

I don't know, you make it sound like the molecules don't even exist and their atoms are just floating around disconnected. I'm sure there is some explanation but in layman's terms if you dissolve 1kg NaCl to water and boil it down you will get 1kg of NaCl. You won't lose salt during the boil because the chlorine-ion boils away because it isn't "attached" to the sodium.

So if you have mixed vinegar with bicarb, the acetate ion is still there, but alone it cannot evaporate. It needs a H+ to form the acid again and then it can evaporate. Where the H+ comes from isn't important. Perhaps citric acid, perhaps sulphuric acid, it doesn't matter.

But why did the second part of your experiment fail? Because a pH slightly under 7 isn't enough to get vinegar in smellable amounts. Normal cooking vinegar has something like pH 2.5. I think you added way too much bicarb. Before adding citric acid you should have discarded the bicarb sediment or the citric acid is sucked up by the excess bicarb. And citric acid is not really strong. But probably when discarding the sediment, it should work to lower the pH to 3-4. I did this experiment with sulphuric instead of citric acid.

It didn't fail. I wanted to make sure that the liquid dropped to pH 7 would not smell of vinegar. It didn't. I'm not interested in taking it below that since HDNB was just asking about running with such a high pH (11) being caustic.

[der wo]

The acetate-ion is never attached to the base-ion. Except you would boil it down under the solubility point. Acids and salts in solution are always splitted up to ions. Acetic acid is CH3COOH. In solution it is split up into CH3COO- and H+. And for example caustic soda (NaOH) is in solution Na+ and OH-. And their salt sodium acetate (NaCH3COO) is in solution Na+ and CH3COO-. Sodium acetate is very soluble, so "in reality" it doesn't exist unless you boil the solution down. But for example the salt of sulphuric acid and lime is very insoluble. So if you mix those substances in water, you get calcium sulfate, it will settle out.

I don't know, you make it sound like the molecules don't even exist and their atoms are just floating around disconnected. I'm sure there is some explanation but in layman's terms if you dissolve 1kg NaCl to water and boil it down you will get 1kg of NaCl. You won't lose salt during the boil because the chlorine-ion boils away because it isn't "attached" to the sodium.
Not the atoms float around, but the ions. When you solve NaCl in water you have Na+ and Cl- ions. Noone can say if you produced this solution simply by solving NaCl in water or by mixing HCl and NaOH (hydrochloric acid and caustic soda). Then when you boil it down the ions will connect again and form NaCl.
Ferrum, feel free to correct me, if something is wrong...

So if you have mixed vinegar with bicarb, the acetate ion is still there, but alone it cannot evaporate. It needs a H+ to form the acid again and then it can evaporate. Where the H+ comes from isn't important. Perhaps citric acid, perhaps sulphuric acid, it doesn't matter.

But why did the second part of your experiment fail? Because a pH slightly under 7 isn't enough to get vinegar in smellable amounts. Normal cooking vinegar has something like pH 2.5. I think you added way too much bicarb. Before adding citric acid you should have discarded the bicarb sediment or the citric acid is sucked up by the excess bicarb. And citric acid is not really strong. But probably when discarding the sediment, it should work to lower the pH to 3-4. I did this experiment with sulphuric instead of citric acid.

It didn't fail. I wanted to make sure that the liquid dropped to pH 7 would not smell of vinegar. It didn't. I'm not interested in taking it below that since HDNB was just asking about running with such a high pH (11) being caustic.
Ok, you didn't fail, I see. I wasn't sure, what exactly you wanted to prove. You wanted to prove that rising to pH 7 again is ok.
When I wrote HDNB, that pH 11 is similar caustic like pH 3, I wanted to say pH 11 is similar caustic like a very dry fermented wash, what means not very caustic. But of course not everything what can stand low pH also can stand high pH (and vice versa).

[Ferrum]

It is true, any ionic compound (salt) you dissolve in water fully splits up into ions (assuming it dissolves!). NaCl becomes Na+(aq) and Cl-(aq). CH3COONa becomes CH3COO- (aq) and Na+(aq). Etc. When you remove the water by vapourizing, you return the compound into the salt form.

Note that bases are often an ionic compound, in which case they fully split up into ions. This does not mean theyre necessarily strong bases however! Sodium acetate (CH3COONa) fully splits into ions and is a weak base. NaOH also fully splits but is a strong base.

Acids are mostly non ionic compounds. Non-ionic compounds can also dissolve in water but dont split into ions. They can however react with water.
CH3COOH dissolves very well in water and for a *small part* reacts with water to form CH3COO- and H3O+ (often simply H+). Thus it is a weak acid.
HCl dissolves very well in water but almost fully reacts with water to form Cl- and H3O+. Thus it is a strong acid.

[Bagasso]

I think you are right with your experiment Bagasso, in that at pH the acid form is a very small amount and may not be noticeable. However, it could possibly become noticeable when vinegar carries over into your distillate and turns into ethyl acetate. I don’t know because I don’t reduce the pH back to neutral because I never saw a reason to :p I will do this experiment at one point to find out.

I was just pointing out a possibility to anyone who feels uncomfortable running something through their still at pH 11.

Im not sure what you mean with retaining thiols. I’ve hardly seen anything about those. In general people believe copper removes sulphides but I’m not sure if that works with thiols. I do know they are rather acidic however, much more so than their oxygen analog (alcohols). So I would say having a strongly basic still charge would retain them better.

In my experience, when I raise the pH of a wash (not low wines) and strip I get a very strong burnt rubber smell in the distillate compared to when I strip wash from the same batch at the normal pH.

I was reading or seeing something about quick fixes for wine faults and how adding a little acid (ascorbic acid?) made them less volatile and if you can't smell them then you can't taste them, or something along those lines.

Pretty weak, I admit, but I was just thinking that this could explain the differences that I was seeing.

[Bagasso]

It is true, any ionic compound (salt) you dissolve in water fully splits up into ions (assuming it dissolves!). NaCl becomes Na+(aq) and Cl-(aq). CH3COONa becomes CH3COO- (aq) and Na+(aq). Etc. When you remove the water by vapourizing, you return the compound into the salt form.

I think I understand what is being said but the question is how does that affect the distillation process? We never fully vaporize all the water so what does this mean to us?

In the above example, is Cl-(aq) ever vaporized?

ETA: Maybe a better question is if CH3COONa becomes CH3COO- (aq) and Na+(aq) what is there to keep CH3COO- (aq) from coming over regardless of pH?

[Ferrum]

I've never raised the pH of a wash, since i've heard it can release ammonia. I use quite a lot of DAP most of the time so I don't feel like I should try that. Adding base to biological stuff can make all sorts of funky things happen anyway that I don't feel like it's a great idea. Maybe egg or oyster shells are whatever people use would be ok since it's mild but I've not needed it so far. Ive used sodium carbonate for low wines because it is extremely cheap and easily basic enough to do the job, so I would not use NaOH and run my low wines at pH 11 either. I don't feel like it necessarily wrong though aside from having to be careful (Love NaOH for cleaning ).

For the other point raised, no it is not possible to vapourize Cl- or any ions in our conditions. Therefore, if you add CH3COONa to water (or add a decent base to a vinegar solution which would give the same effect) you will not be able to vapourize it. This however assumes the pH is high enough. Conveniently, if you add any noticeable amount of CH3COONa to water the pH will be sufficiently high so that no noticeable amounts of vinegar will be formed. However, if you add an acid you can drop the pH to any level you like. If you do this, you will indeed form CH3COOH that can be distilled.

Dropping the pH to 7 will not turn much CH3COO- into CH3COOH, but it will happen. It is most likely not noticeable though. But keeping the pH a bit above 7 ensures CH3COO- and maybe sulphides stay in their basic (ionic) form and thus ensures they won't leave the boiler. I doubt if it's necessary to keep a high pH but I see even less reason drop it before distilling.

About the ascorbic acid, I can't really think of a reason why it would reduce volatility. Ascorbic acid is mostly used as antioxidant to remove chlorine from water or protect wine from oxygen exposure afaik.

[Bagasso]

For the other point raised, no it is not possible to vapourize Cl- or any ions in our conditions. Therefore, if you add CH3COONa to water (or add a decent base to a vinegar solution which would give the same effect) you will not be able to vapourize it. This however assumes the pH is high enough.

The question was why, if it isn't "connected" to the base?

About the ascorbic acid, I can't really think of a reason why it would reduce volatility. Ascorbic acid is mostly used as antioxidant to remove chlorine from water or protect wine from oxygen exposure afaik.

There was a question mark there because I can't recall if it was ascorbic or another acid. Maybe it doesn't matter. It is the same question, Why would pH have an effect on volatility of molecules, or the parts that they split into, if they are no longer connected?

That is what I took away from the last few posts. Molecules split into individual ions when dissolved, in a solution they are not connected. They don't connect (reconnect) until you drive the water out so, vinegar in a solution should behave the same no matter what bases are also in the solution.

Obviously that is not the case so, is that explanation purely academic?

[der wo]

I use NaOH because of its way better solubility in low wines and feints. For solving baking or washing soda into them, I would have to heat it or wait a long time.

Good questions, Bagasso... I prefer that Ferrum answers them, you proabably too...


BTW: When I distill low wines at pH 11, there is still some ethyl acetate in the foreshots. Rising pH only reduces it to a large degree.

[Ferrum]

If you dissolve CH3COONa it turns into CH3COO-. In high pH, most of it stays that way. In low pH, most of it turns into CH3COOH. CH3CO- will always stay in the boiler, CH3COOH is easily distilled over. It is not the reconnecting with sodium that’s important for us but the reconnecting with an H+ from an acid. In our low wines we obviously don’t dissolve CH3COONa but it created made by adding soda.

Why does this happen? You used a base first to get the pH up, and then you used an acid to bring it down to 7 again. This basicly means you create a buffer solution. The pH is no longer controlled by the congeners in the low wines but purely by the large excess of acid and base you added. Therefore, the state the congeners are in, CH3COO- or CH3COOH, is purely determined by the pH of the buffer, not by whether it was added in acid or base form, because it is only a small amount compared to the buffer.

Der Wo, are you sure you used the right soda? if they don’t dissolve well it seems like you used calcium salts or something. Both baking soda and washing soda are extremely soluble like most sodium salts and should only require a small stir. For example in cafés (at least here) the beer glasses are washed using a strong soda solution. I paid 75 cents for a bag of 1 kg Na2CO3 while I have to pay like 3 euros for 0.5 kg solid NaOH.

[der wo]

Der Wo, are you sure you used the right soda? if they don’t dissolve well it seems like you used calcium salts or something. Both baking soda and washing soda are extremely soluble like most sodium salts and should only require a small stir. For example in cafés (at least here) the beer glasses are washed using a strong soda solution. I paid 75 cents for a bag of 1 kg Na2CO3 while I have to pay like 3 euros for 0.5 kg solid NaOH.

In water it dissolves well. But in alcohol it doesn't.
Two pages earlier we already debated the solubility:
5g washing soda into 1l 40% low wines is around pH 11.
With baking soda probably something like 15g is necessary.
Dissolving 15g baking soda is more difficult than 5g washing soda.
1g caustic soda is pH 12 (sorry, yesterday I wrote pH 11...). Much easier to dissolve. I guess 0.5g is pH 11.

And somewhere I posted a youtube video. Solving 5g washing soda into 1l feints. There you can see how bad it dissolves:
https://www.youtube.com/watch?v=Y9zNvOseK8s" onclick="window.open(this.href);return false;" rel="nofollow
I am sure that it is washing soda, not lime or chalk.

[Bagasso]

If you dissolve CH3COONa it turns into CH3COO-. In high pH, most of it stays that way. In low pH, most of it turns into CH3COOH. CH3CO- will always stay in the boiler, CH3COOH is easily distilled over. It is not the reconnecting with sodium that’s important for us but the reconnecting with an H+ from an acid.

I think you are mistaken about CH3CO- (assuming you meant CH3COO-). Ethyl acetate does not stay in the boiler. It is the major congener in heads.

ETA: Like I said, the explanation seems like a purely academic exercise. Also, generic terms like "high pH" and "low pH" don't really tell us much. According to wikipedia:

The acetate anion, [CH3COO]−,(or [C2H3O2]−) is one of the carboxylate family. It is the conjugate base of acetic acid. Above a pH of 5.5, acetic acid converts to acetate:

CH3COOH ⇌ CH3COO− + H+

Seems like dropping the pH of treated low wines to or around 7 after ester hydrolyzation shouldn't leave you with CH3COOH.

[Ferrum]

Sorry, CH3CO- was a typo, it is supposed to be CH3COO-. Note that acetate is not ethyl acetate. Acetate is the conjugated base of acetic acid and is a non-volatile anion, while ethyl acetate is the highly volatile ester we all know. But since we’re dealing with high pH hopefully the ethyl acetate is mostly gone and turned into acetate and the lovely ethanol.

It sure is an acedemic exercise. That's definitely not always the best way to look at it, but when there's people that run at between pH 7 and pH 11 with success I'm just trying to think of what would theoretically be superior. In my opinion it’s always useful and fun to know what’s chemically going on, even if that is just a tiny part of making an actual drinkable brew.

To say above pH 5.5 all acetic acid becomes acetate is not right. It is equilibrium thus you can't express it in absolutes. What they may have meant to say is that above pH 5.5 most of the molecules exist as acetate and below 5.5 most exist as acid. Since pH is a logarithmic scale that would mean that at pH 6.5 there would be ~5% acid and at pH 7.5 there would be ~0.5% acid. If you're interested in the subject you can look up 'isoelectric point', theres graphs there that shows in which proportions proteins that can donate multiple H+ ions exist based on pH level. We’re only dealing with 2 species here, RCOO- and RCOOH, but the distribution depending on pH will be similar.

About the soluability, I may have been wrong about the baking soda. However, I use washing soda without any problems. Maybe a tablespoon or two of soda on 5 L of wines of 30-40% is dissolved well with a small stir. One thing to note is that if theres ~2g of acetic acid in your wines and you add 1g of caustic soda you will end up with a very mild pH and lots of acetic acid present. If you added 5g of washing soda you would not have this problem and end up with a higher pH and basicly only acetate ions (you should calculate this with moles instead of grams but it's just an example to keep it simple). I would not know if this is a realistic figure though.

[der wo]

One thing to note is that if theres ~2g of acetic acid in your wines and you add 1g of caustic soda you will end up with a very mild pH and lots of acetic acid present. If you added 5g of washing soda you would not have this problem and end up with a higher pH and basicly only acetate ions.

Say I have two bottles with identical low wines and I rise both pHs to exact 11, one with washing soda the other with caustic soda, does the low wines with the caustic soda have higher amounts of ethyl acetate and volatile acetic acid?

[Ferrum]

No, in theory they will have the same. I personally don't have a pH meter though so I'd rather be safe by adding lots of a weaker base than be unsure by adding a small amount of strong base.

[Bagasso]

Sorry, CH3CO- was a typo, it is supposed to be CH3COO-. Note that acetate is not ethyl acetate. Acetate is the conjugated base of acetic acid and is a non-volatile anion,

This is where I'm confused, conjugated base sound like "connected" to me but it was just argued that things in solution are not connected.

It sure is an acedemic exercise. That's definitely not always the best way to look at it, but when there's people that run at between pH 7 and pH 11 with success I'm just trying to think of what would theoretically be superior. In my opinion it’s always useful and fun to know what’s chemically going on, even if that is just a tiny part of making an actual drinkable brew.

Sure but pH 7 to pH 11 leaves a lot of room for people to run at a pH that they feel confortable at. Maybe it would make them feel more comfortable running at 11 pH if you tell them that people are drinking alkaline water up to 10 pH like regular water, no affiliation, not a recommendation, just an example.

Since pH is a logarithmic scale that would mean that at pH 6.5 there would be ~5% acid and at pH 7.5 there would be ~0.5% acid.

5% and 0.5% of what though?

A finished wash might have a pH as low as 3.5 and, under normal conditions, there is not 5% acetic acid in there even then, maybe not even 0.5%. It just seems like the academic exercise uses a standard acetic acid solution which makes it an apples and oranges comparison.

[Ferrum]

Conjugated base basicly means the base that is left behind when an acid loses its proton. For CH3COOH that is CH3COO-. If you dissolve a weak acid in water it will always be in equilibrium with its conjugate base. The equilibrium depends on the pH.

What I meant to say with the 5% and 0.5% is that 5% would be in acid form (CH3COOH) and 95% would be in conjugated base form (CH3COO-) at pH = 6.5. Likewise, 0.5% would be in acid form and 99.5% in base form at pH = 7.5. This tells nothing about the amount of vinegar present, that depends on how much was present in your brew as ethlyl acetate or vinegar. Hopefully it is not 5% no, then you may have a serious fruit fly invasion. Most of the acid in your wash should probably come from carbonic acid, citric acid, lactic acid etc. Instead, this figure tells you whether the vinegar that is present, is in the harmless non-volatile conjugate base form, or in the not so tasty volatile acid form.

I think it would be fine dealing with pH 11 or pH 12, but handling caustic soda is rather dangerous, so I woulnt be too comfortable recommending going stronger than washing soda to get a high pH.

[Bagasso]

Conjugated base basicly means the base that is left behind when an acid loses its proton. For CH3COOH that is CH3COO-. If you dissolve a weak acid in water it will always be in equilibrium with its conjugate base. The equilibrium depends on the pH.

Like I said, just another wordy way of saying, if you add a base it will form a salt, even if it isn't "technically" correct.

What I meant to say with the 5% and 0.5% is that 5% would be in acid form (CH3COOH) and 95% would be in conjugated base form (CH3COO-) at pH = 6.5. Likewise, 0.5% would be in acid form and 99.5% in base form at pH = 7.5. This tells nothing about the amount of vinegar present, that depends on how much was present in your brew as ethlyl acetate or vinegar. Hopefully it is not 5% no, then you may have a serious fruit fly invasion. Most of the acid in your wash should probably come from carbonic acid, citric acid, lactic acid etc. Instead, this figure tells you whether the vinegar that is present, is in the harmless non-volatile conjugate base form, or in the not so tasty volatile acid form.

It isn't about what you meant to say but what it really means. You are throwing around amounts without really knowing if they are correct. Do you have a source or are your numbers just a baseless example.

What I posted says that at 5.5 pH CH3COOH turns into CH3COO- which seems to imply that most if not all CH3COOH at that pH is CH3COO- and it would get closer to 100% with every 0.1 above that even if it never reaches it.

I think it would be fine dealing with pH 11 or pH 12, but handling caustic soda is rather dangerous, so I woulnt be too comfortable recommending going stronger than washing soda to get a high pH.

It isn't about what you feel is fine or not. It is about someone else reading our posts, like HDBN, and thinking running a still at 11 pH sounds very high?

I know it sounds like I'm being an a-hole but it isn't my intention. It is just that while chemists might have a reason for splitting hairs to the fourth decimal point, we don't need that precision. There also seems to be no reason to advise people not to drop the pH of their treated low wines to 7 or any point above that that they feel safe with.

[der wo]

It isn't about what you feel is fine or not. It is about someone else reading our posts, like HDBN, and thinking running a still at 11 pH sounds very high?
For example beton has pH 12.5. Soap has 9.5. Does pH 11 still sound dangerous? And what do you think about cleaning runs with diluted vinegar, which has pH 2.5?
BTW causic soda is used in the food industry for producing pretzel. But yes, don't get it into your eyes. Same like vinegar, high abv, flux or citric acid: Don't get it into your eyes.

I know it sounds like I'm being an a-hole but it isn't my intention. It is just that while chemists might have a reason for splitting hairs to the fourth decimal point, we don't need that precision. There also seems to be no reason to advise people not to drop the pH of their treated low wines to 7 or any point above that that they feel safe with.
And he wrote a few times, that you will turn the sodium acetate partially back to acetic acid regardless what acid you use to get back to pH 7. Partially, because you go back only to 7, not to for example 4, where you perhaps started. And acetic acid turns partially to ethyl acetate again during distillation. Sounds absolutely right for me.

[Bagasso]

For example beton has pH 12.5. Soap has 9.5. Does pH 11 still sound dangerous? And what do you think about cleaning runs with diluted vinegar, which has pH 2.5?
BTW causic soda is used in the food industry for producing pretzel. But yes, don't get it into your eyes. Same like vinegar, high abv, flux or citric acid: Don't get it into your eyes.

Sure, that is why I linked 10 pH drinking water. Still It is perfectly fine for someone to drop the pH to 7 or 8 if they prefer.

And he wrote a few times, that you will turn the sodium acetate partially back to acetic acid regardless what acid you use to get back to pH 7. Partially, because you go back only to 7, not to for example 4, where you perhaps started. And acetic acid turns partially to ethyl acetate again during distillation. Sounds absolutely right for me.

We can agree that he said it a few times but he didn't cite a source. I posted something that said acetic acid turns to acetate at 5.5 pH. Is that 80%, 90%, 95% or 99%? Nobody knows but that is lower than 7.

All I asked was where his numbers came from or was it something just made up to illustrate the point?

[der wo]

You are right. We would need a qualified experiment. Same low wines brought up to pH 11 for example and then one down to 7 and the other remains at 11.
In theory pH11 should be better. In practice we don't know if it's smellable.

[Ferrum]

I don't know if these numbers were right. It was a response to the statement you quoted from wikipedia which is at least misleading. I just wanted to give a better illustration of reality, based on the number you quoted.

I've now looked further into it gained a better understanding myself. The pKa (acid constant) of acetic acid is 4.76 (wikipedia). The ratio of acid form / base form depends on pH and is given by "pH = PKa + 10log (CH3COO-/CH3COOH )". This means that at pH = pKa, in this case pH = 4.76, there would be 50% in acid form and 50% in base form. However, at pH = 7 you would have approximately 0.5% acid form and 99.5% base form according to the equation (I filled in the pH and pKa and replaced the ratio by x to solve it in wolfram alpha). I would say in practical terms you won't have any vinegar at pH = 7. However, if you accidentally lower the pH to 6 you'd end up with approx 25% in acid form which would most likely significantly affect your brew. My conclusion would be that lowering the pH to 7 would be fine, although I don't see a reason to go this low, and only do it if you can accurately measure pH so you don't end up going lower than 7 (source: wikipedia: Acid dissociation constant (under monoprotic acids paragraph)).

It would be great if someone could do an experiment with this. I expect the difference would not be noticeble, but it's great to have the practical results.

The theory interests me because it can also be applied to other situations. For example, I have some low wines of a ferment that got infected during mashing. I'm not really into the funky stuff, but after stripping it, it smells a bit of vomit but also of pineapple. It seems like a lot of propionic acid / butylic acid was formed by bacteriae and partly esterized to ethyl propiate during stripping. It seems reasonable to do the exact opposite of what we are doing here and lower the pH using sulphuric acid and maybe reflux a few hours to strongly promote ester formation. In this case it may be a good idea to bring the pH back to 7 before running so remaining acids are mostly in base (salt) form but the pH is neutral enough to not hydrolyse the esters.

[der wo]

I would add sulphuric acid to the low wines (something like 1ml per liter low wines) and nothing else. Then most acids get esterified. And the remaining acids will be mostly in the tails. If you have many acids in the low wines (due to bacteria for example), the rising acidity of the distillate will determine the point where to cut the tails. So when you esterify many of them, you not only get more esters, but also you can go deeper into the tails. More fruity esters and more depth, only advantages IMO, at least for high aromatic spirits.

We are way off topic. Either start a new thread or jump in in one of my Rum threads, if you want to discuss it further.

[Bagasso]

I don't know if these numbers were right. It was a response to the statement you quoted from wikipedia which is at least misleading. I just wanted to give a better illustration of reality, based on the number you quoted.

That is what I though. Not that there is anything wrong with that but I think it is good to be clear that the number are for illustrative purposes.

I've now looked further into it gained a better understanding myself. The pKa (acid constant) of acetic acid is 4.76 (wikipedia). The ratio of acid form / base form depends on pH and is given by "pH = PKa + 10log (CH3COO-/CH3COOH )". This means that at pH = pKa, in this case pH = 4.76, there would be 50% in acid form and 50% in base form. However, at pH = 7 you would have approximately 0.5% acid form and 99.5% base form according to the equation (I filled in the pH and pKa and replaced the ratio by x to solve it in wolfram alpha).

What do you make of this, taken from this article, because it seems to contradict your math:

Acetic acid is a weak, effectively monoprotic acid in aqueous solution, with a pKa value of 4.8. A 1.0 M solution (about the concentration of domestic vinegar) has a pH of 2.4, indicating that merely 0.4% of the acetic acid molecules are dissociated.

If I'm reading this right it says 99.6% acetic acid is acetate at 2.4 pH or is "dissociated" in this statement a term that refers to something besides the conjugate base?

I actually appreciate the explanation of how things exist in a solution. It actually helped me understand something in regards to extractions and why just going slightly above and below pH 7 isn't enough.

I would say in practical terms you won't have any vinegar at pH = 7. However, if you accidentally lower the pH to 6 you'd end up with approx 25% in acid form which would most likely significantly affect your brew. My conclusion would be that lowering the pH to 7 would be fine, although I don't see a reason to go this low, and only do it if you can accurately measure pH so you don't end up going lower than 7 (source: wikipedia: Acid dissociation constant (under monoprotic acids paragraph)).

I agree but nobody is saying you have to hit 7 on the nose. I was talking about lowering to a place someone might be more comfortable running at, maybe 8 or 9.

I mean if someone is really dead set on running at 7 then they would probably see the purchase of a pH meter as a worthwhile investment. A 5 meter roll of litmus paper is less than $10 on amazon. It would last you years even if you test a batch a week.

[Ferrum]

What do you make of this, taken from this article, because it seems to contradict your math:

They're saying 99.6% is in acetic acid form and not dissociated. That makes sense because it is a weak acid so it only partly dissociates in water.

Using the same formula and their slightly different pKa for acetic acid: 2.4 = 4.8 log(x) yields x = 0.00398 so [CH3COO-]/[CH3COOH] = 0.00398 -> 0.398% is CH3COO- and the rest is CH3COOH. So that is how they obtained the numbers.

The difference is that theyre talking about an acetic acid solution in pure water. Therefore the pH of the solution is dictated only by the acetic acid and it has a low pH. The low pH causes the acid to stay in acid form mostly.

We are talking about an acetic acid solution in a 'buffer' where the pH is determined by all acids/bases we added and only slightly by the acetic acid present. Because we have a high pH the acid dissociates mostly into its conjugate base (CH3COO-)

[Bagasso]

They're saying 99.6% is in acetic acid form and not dissociated. That makes sense because it is a weak acid so it only partly dissociates in water.

Ok, dissociate means to play a part in a reaction but since it is in a solution it reacts but is separate. Something like that, right?

Because we have a high pH the acid dissociates mostly into its conjugate base (CH3COO-)

I think we took the scenic route to get to the same spot, add a base (raise pH) and at some point acetic acid will stop being acetic acid and turn to acetate.

To be sure where you are at buy a meter or some litmus paper.

Or maybe make your own with red cabbage.