Why is my homemade alkaline water not reaching the expected pH level?

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Discussion Overview

The discussion revolves around the process of creating homemade alkaline water using baking soda (sodium bicarbonate) and distilled water. Participants explore the expected pH levels, the dissociation of baking soda in water, and the chemical reactions involved. The conversation includes theoretical considerations, experimental observations, and calculations related to pH and molarity.

Discussion Character

  • Exploratory
  • Technical explanation
  • Debate/contested
  • Experimental/applied

Main Points Raised

  • One participant describes their method for making alkaline water and calculates a theoretical pH of 12.88 based on their measurements.
  • Another participant points out that a pH of 8 is still considered alkaline, questioning the expectation of a higher pH.
  • Concerns are raised about the assumption that the molarity of hydroxide ions is equal to the molarity of sodium bicarbonate, with a suggestion to consider the equilibrium constant of the reaction.
  • Some participants express confusion about the dissociation of bicarbonate and its role in affecting pH, with differing views on how it behaves in solution.
  • One participant mentions that heating baking soda could convert it to sodium carbonate, which may lead to different dissociation behavior and higher pH levels.
  • Another participant calculates the pH based on the incorrect assumption of complete dissociation of sodium bicarbonate.
  • There is a discussion about the dissociation products of bicarbonate and their impact on pH, with conflicting interpretations of the chemical reactions involved.
  • One participant shares their experimental result after heating sodium bicarbonate, noting a pH increase but expressing uncertainty about the underlying processes.

Areas of Agreement / Disagreement

Participants do not reach a consensus on the expected pH levels or the mechanisms by which baking soda affects pH. There are multiple competing views regarding the dissociation of bicarbonate and the effects of heating sodium bicarbonate.

Contextual Notes

Limitations include assumptions about complete dissociation, the dependence on specific conditions for reactions, and varying interpretations of chemical behavior in solution.

barryj
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I am trying to make my own alkaline water using baking soda and distilled water.
First I measure out 0.018 moles of baking soda as follows
(0.25 tsp)(1ml/.202 tst)(1.2 grams NaHCO3/ml)(1 mole NHCO3/1 gram NaHCO3) = .018 moles NaHCo3
I measure out 0.236 liters of distilled water i.e. 1 cup.
The molarity of the NaHCO3 is not M= 0.018 moles NaHCO3/0.236 liters of water = 0.076
The pOH of the solution is –log(0.076) = 1.12 thefore the pH is 14 – 1.12 = 12.88.
Note: the molarity of [OH] should be the same as the molarity of NaHCO3.
However, when I put ¼ tsp baking soda into as cup of distilled water the pH measures only 8.0.
What is wrong here? Is the baking soda not disassociating?
 

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pOH, negative log of the hydroxide (molarity) concentration.
pH, negative log of the hydronium (molarity) concentration.
(Really, should be "activity" instead of "molarity").

pH of 8, as you measured, IS ALKALINE.
 
First, always start with a balanced chemical equation.

Second,
barryj said:
Note: the molarity of [OH] should be the same as the molarity of NaHCO3.
This is incorrect. You have to consider the equilibrium constant of the reaction between bicarbonate and water.
 
Symbolipoint: Yes 8 is alkaline but I expected around 12 not 8.

Yqqqdrasil: I would have expected the equilibrium constant to be basically one way since NaHCO3 is ionic,
Getting a pH of only 8 seems weird. I will check this however.
 
barryj said:
Yqqqdrasil: I would have expected the equilibrium constant to be basically one way since NaHCO3 is ionic,
Getting a pH of only 8 seems weird. I will check this however.

Again, write out the relevant chemical equation. Dissociation of Na+ from HCO3- is not what produces OH-.
 
NaHCO3 -> Na + HCO3
HCO3 -> OH + CO2
but it must be in water? so NaHCO3 + H2O -> Na + OH + CO2 + H2O
Yes/No? This is more complicated that I first thought
 
Also, I read that Kb = 2.08E-4. So only a small amount of the NaHCO3 will dissociate, yes?
 
I now read that if I heat the baking soda to 400F for one hour, it will convert to Na2CO3 and this will dissociate and my equation will work. Yes?
 
  • #10
I calculated the pH to be 12 based on the erroneous fact that the NaHCO3 would totally dissociate which it doesn't.
 
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  • #11
I am afraid you are misunderstanding why the pH of the solution changes. NaHCO3 is fully dissociated into Na+ and HCO3-, of these only HCO3- is responsible for pH changes - partially because it dissociates further, partially because it hydrolyzes.
 
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  • #12
Then the HCO3 disassociates into OH and CO2 and then I have the OH I need.
 
  • #13
No, HCO3- doesn't dissociate the way you wrote. When it dissociates it produces H+ and CO32-, so if anything it lowers the pH during the dissociation, not makes it rise.

Please read on Bronsted-Lowry theory of acids and bases.
 
  • #14
Here is what I did. I baked some NaHCO3 at 450F for one hour. I read this will turn the NaHCO3 into Na2CO3. I put this, 1/4 tsp like I described in my original post, and the pH went up from around 7 to 10.5. It was not a real carefully controlled experiment, in my kitchen, but the pH did rise and the taste of the Na2CO3 was more bitter than the original NaHCO3. Not sure of the process here but it seems like the heat drove out the H. Now when Na2CO3 is in water it produces NAOH and carbonic acid? the Na OH is a strong alkine and the caerbonic a weak acid? Na2CO3 + H2O -> Na + OH + H CO3
 
  • #15
barryj said:
Symbolipoint: Yes 8 is alkaline but I expected around 12 not 8.

Yqqqdrasil: I would have expected the equilibrium constant to be basically one way since NaHCO3 is ionic,
Getting a pH of only 8 seems weird. I will check this however.

No, sodium bicarb is not that basic in solution, it is a weak base.
 
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