Campden Tablets vs. Hydrogen Peroxide

[TheLastDamnBatch]

Various posts across the internet involve substituting HP for CT or even sometimes using both.

1.) How does hydrogen peroxide remove chlorine/chloramine?
2.) Is it safe to use hydrogen peroxide to remove chlorine/chloramine?
3.) What concentration should be used?
4.) Do hydrogen peroxide and campden tablets chemically react with each other to produce something toxic?

More research has lead to these threads:

Dechlorination using drugstore hydrogen peroxide:
http://hbd.org/discus/messages/40327/41985.html

Peroxide for starter aeration:
http://hbd.org/discus/messages/26895/29408.html?1112143606

 

[WoodlandBrew]

Im actually testing the use of hydrogen peroxide for yeast aeration right now! Nothings new I guess. I haven't heard of using it to remove chlorine.

 

[TimT]

Im actually testing the use of hydrogen peroxide for yeast aeration right now! Nothings new I guess. I haven't heard of using it to remove chlorine.

I am quite sure it is just plain out toxic to the yeast. I thought of it myself before and the research I did found it was purely negative. Too long ago to specifically cite anything though.

 

[WoodlandBrew]

So far my results are agreeing with your research. At the 3% solution strength it's equivalent to 7000ppm of oxygen. That is quite a bit of oxygen stress. Even when diluted down to the equivalent of 11ppm there was significant lag at the start of fermentation. I got the idea while reading a paper on vitality that used 2(HO) to induce oxygen stress, and wondered what it would do in extremely low quantities. 1ml is equivalent to about 1ppm in 5 gallons.

 

[WoodlandBrew]

The effect of hydrogen peroxide, 2(HO) on chlorine ions, Cl-, is that it forms hydrochloric acid. The amount that stays as HCl and the amount that disassociates, dropping the pH, will depend on concentration. If your water has 3ppm of free Chlorine ions, then that will be about 3ppm of HCl. This is a 0.000082 molar solution. -log(0.000082)=4.1pH. That's a little low, but not terrible. And this is assuming full disassociation which is normally reasonable for a strong acid like HCl. Some commercial brewers adjust mash pH with HCl.

 

[WoodlandBrew]

Yes, thank you. The other two assumptions are that all of the free chlorine is converted to HCl, and that in the mash all of the hydrogen ions disassociate. It would be interesting to see a pH measument before and after the addition of 2(HO) to a mash.

 

[Wynne-R]

Drug store peroxide is not pure. It has a stabilizer to keep it from going flat, a few days instead of a few hours.

I don't know what it is, but it tastes bitter.

 

[SynthesisError]

I don't think it's a good idea, for a couple of reasons.
1. Because HOOH is relatively unstable, it would be difficult to add a precise amount without first titrating using a primary standard or otherwise quantifying the actual concentration. This is one of those cases where what the label says (usually 3%) is irrelevant. It IS possible to over-oxygenate.
2. The reaction described in which HOOH + Cl- --> HCl + H2O only happens appreciably at a pH above 7.
3. In order to effectively get the oxygen in solution for aeration, you would need to catalyze the decomposition of the HOOH. A rust-free iron plate with large surface area would work, but could have unintended consequences on flavor if the wort has a low enough pH. Platinum works too, but if you have a couple ounces sitting around, I'd like to borrow some...
4. Combining HOOH with campden (Na2S2O5) triggers a redox reaction producing sulfur dioxide. While it probably wouldn't kill you, your brew would have the aroma of burnt hair or rotten eggs.
5. HOOH is a strong oxidizer, and if any is present when your fermentables are added, you will further improve flavor by adding a sophisticated wet cardboard finish to your brew.

 

[WoodlandBrew]

Some good food for thought there Synthesis,

1. I'm not sure that a precise amount is necessary. For aeration the control most brewers have (without a DO meter ) is already imprecise. For removing chlorine it only takes a ml to cancel everything, and going over has no real adverse side effects. But if you really do want to know what the concentration is there are other ways to get a good idea without titration. Such as relative density with a hydrometer or by refraction index with a refractometer.
2. In the post that the OP linked there is evidence that the removal of chlorine worked with tap water at room temperature. Are you saying it will separate back out?
3. I'll have to look at that. Hydrogen peroxide is used for inducing oxygen stress to yeast. (Information below, It's a good read)
4. It's good to know about the side effects of using both 2(HO) and Campden together, but I don't think anyone was suggesting that.
5. "sophisticated wet cardboard" ha ha. The intent would be to add an equivalent amount to using pure O2 as many people have success with.

Vitality/Viability Assays
FRITHJOF THIELE and WERNER BACK
Radedeberger Gruppe KG, Darmstadter Landstr. 185,60598 Frankfurt, Germany
Lehrstuhl fijr Technologie der Brauerei I, Technische Universitat MiinchenWeihenstephan,
Weihenstephaner Steig 20, 85354 Freising, Germany.
Corresponding author. Tel: +49 (0) 69 6065 157; fax: +49 (0) 69 6065 99157.
E-mail: [email protected]

quote from the paper showing the use of 2(HO) for oxidative stress:

• Ethanol stress: I and 2 hours at 36°C in 10% (w/w) ethanol
• Starvation stress: 1 and 3 days under deionized water at 25°C
• Oxidative stress: 1 hour in 0.1 % and 1.0% (w/w) H202 at 25°C
• Osmotic stress: 2 days in 15% and 20% sorbitol solution at 25°C

 

[Kaiser]

The problem that I see with Hydrogen peroxide in the wort (if you can find it food grade) is that it is a very strong oxidizer and it can oxidize wort compounds more readily than you can with O2. So if Hydrogen Peroxide is suggested as a wort aeration replacement one has to do extensive field tries to determine that it is not detrimental to beer quality or beer flavor stability.

Kai

 

[SynthesisError]

Hydrometers and refractometers aren't precise enough when your eventual target is on the order of magnitude of ppm. I suppose if you had a good analytical balance and a precise pipet, you could devise a somewhat accurate relative density determination by comparing your stock HOOH solution with distilled water.

The link posted would be more helpful if he had included a local water profile. Also, colorimetry varies in its applications; I have no idea what indicator he used, but it is easy to interfere with most indicators by simply adjusting the pH, whether the chlorides were consumed or not. The experiment should have included a NaOH buffer.

Also, free chlorine primarily exists in water as hypochlorite ion (OCl-), not as free radical Cl-. This is a dissociation product of HOCl, hypochlorous acid, which is much weaker than HCl. At a pH of less than 7 (excess H+ in solution), common ion effect shifts the equilibrium of the dissociation towards the reactant side, meaning that there is very little chlorine available for reaction with HOOH. If the pH in the system is adjusted, say with CaCO3 (like in hard tap water), the acid tends to dissociate more because the free OH in solution tends to pull the hydrogen off of the acid.

 

[WoodlandBrew]

Hydrometers and refractometers aren't precise enough when your eventual target is on the order of magnitude of ppm. I suppose if you had a good analytical balance and a precise pipet, you could devise a somewhat accurate relative density determination by comparing your stock HOOH solution with distilled water.
[/QUOTE]
I agree trying to use a hydrometer or refractometer to measure a 0.01% 2(HO) solution would not work very well. Using it to measure the 3% solution to see how close it might be to 3% was the idea.

Also, free chlorine primarily exists in water in the water as hypochlorite ion (OCl-), not as free radical Cl-. This is a dissociation product of HOCl, hypochlorous acid, which is much weaker than HCl. At a pH of less than 7 (excess H+ in solution), common ion effect shifts the equilibrium of the dissociation towards the reactant side, meaning that there is very little chlorine available for reaction with HOOH. If the pH in the system is adjusted, say with CaCO3 (like in most tap water), the acid tends to dissociate more because the free OH in solution tends to pull the hydrogen off of the acid.

Good to know. Thanks.