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Old 09-02-2009, 06:19 AM   #1
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Default Finally got my pH overview article good enough for release

I have been working on a few articles about pH for a while. The intent is to provide more insight on the subject of pH from a brewer's point of view. The first article deals with pH basics: what it is, how it is affected and measured. While there are a lot of formulas that could have been presented in that article I intentionally refrained from presenting them b/c that tends to scare people away. And the ones who understand these formulas will be getting their info from other places anyway

The last thing I had to do tonight was to prepare and take pictures of the red cabbage juice pH series. Until I started looking into this pH stuff I didn't even know that the color of red cabbage juice is pH sensitive. It sure made for a neat experiment and pictures but is not very applicable to brewing practice:

An Overview of pH - German Brewing Techniques

I think that the key to understanding pH in brewing is understanding how pH affects the disassociation of weak acids and bases. I hope that I did a good enough job with that.

I always appreciate feedback

Kai

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Old 09-02-2009, 12:05 PM   #2
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Very good article. Brings back memories of making my own cabbage pH test strips as a kid too!

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Old 09-02-2009, 06:47 PM   #3
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Excellent resource. You might also consider adding a section on why brewers should care about pH, i.e. all of the different effects that it has throughout the brewing process, from mashing to fermentation.

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Old 09-02-2009, 06:51 PM   #4
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My god you keep yourself busy! You are a brewing sciences ANIMAL!

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Old 09-02-2009, 07:15 PM   #5
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Quote:
Originally Posted by stoutaholic View Post
Excellent resource. You might also consider adding a section on why brewers should care about pH, i.e. all of the different effects that it has throughout the brewing process, from mashing to fermentation.
It's coming. Don't worry . This will be part II.

Kai
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Old 09-03-2009, 12:53 PM   #6
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Great start...very much anticipating the future parts!

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Old 09-03-2009, 02:46 PM   #7
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Looks solid so far.

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Old 09-03-2009, 07:50 PM   #8
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I know there's more to the series and I'll try not ask questions that weren't in the scope of each part. But there was something I wasn't sure I fully understood.

Quote:
pH is a measure of the acidity of a solution. It is a measure of the concentration of hydrogen ions (H+; proton). The higher the concentration of hydrogen ions in a solution the more acidc it is and the lower their concentration the more basic it is.
Quote:
A substance that donates H+ to or accepts OH- ions from its environment is called an acid. It lowers the pH. A substance that accepts H+ or donates OH- is called a base and it raises the pH.
So pH is a measurement of H+ ions. Suppose we have a substance accepts or donates OH- ions (implying it does not directly affect the H+ ions)...how does that change the pH? Since pH is a measure of the H+ and not the OH-.

Is it because:
In the case of an OH- acceptor, water molecules disassociate and form more OH- and H+ ions...yielding a resultant increase in H+.
In the case of an OH- donator, the extra OH- joins with some of the existing H+ to form water...yielding a resultant decrease in H+.
???
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Old 09-03-2009, 08:00 PM   #9
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wow, you keep amazing me. now I get to learn more from the master.

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Old 09-03-2009, 10:09 PM   #10
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Quote:
Originally Posted by SpanishCastleAle View Post
So pH is a measurement of H+ ions. Suppose we have a substance accepts or donates OH- ions (implying it does not directly affect the H+ ions)...how does that change the pH? Since pH is a measure of the H+ and not the OH-.

Is it because:
In the case of an OH- acceptor, water molecules disassociate and form more OH- and H+ ions...yielding a resultant increase in H+.
In the case of an OH- donator, the extra OH- joins with some of the existing H+ to form water...yielding a resultant decrease in H+.
???
In the Bronstead-Lowry theory of acids and bases, an acid is a proton donator, and a base is a proton accepter. Under this model, strong bases such as NaOH (Sodium Hydroxide or Lye) dissociate into Sodium ions and OH groups. Weaker bases act similarly, but dissociate less fully. Acids do the same thing, but form H+ and some negatively-charged ion. Here "weak" and "strong" are not measures of pH, but measures of how readily the acid or base dissociates.

The "OH acceptor" that you're asking about is a proton. When a H+ and OH- combine, they form water. Pure water has both OH- and H+ present in very small quantities (in any aqueous solution, the concentration of OH times the concentration of H+ = 1x10^-14 As you might expect, the concentrations of each in pure water are equal at 1x10^-7, which leads to the definition of pH=7 as that of pure water).

When something that "donates" OH- ions (which is equivalent to accepting H+ ions) is added to a system in equilibrium, the equilibrium will shift one way or the other due to the H+ + OH- = H2O reaction. This shift is easy to calculate for strong acids and bases, and less easy for weak acids/bases. This goes into the math that Kaiser is purposefully avoiding (people tend to be intimidated by logarithms).
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