# Trouble making alkaline water

1. Jan 29, 2017

### barryj

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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2. Jan 29, 2017

### symbolipoint

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.

3. Jan 29, 2017

### Ygggdrasil

Second,
This is incorrect. You have to consider the equilibrium constant of the reaction between bicarbonate and water.

4. Jan 30, 2017

### barryj

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.

5. Jan 30, 2017

### Ygggdrasil

Again, write out the relevant chemical equation. Dissociation of Na+ from HCO3- is not what produces OH-.

6. Jan 30, 2017

### barryj

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

7. Jan 30, 2017

### barryj

Also, I read that Kb = 2.08E-4. So only a small amount of the NaHCO3 will dissociate, yes?

8. Jan 30, 2017

### barryj

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?

9. Jan 30, 2017

### Staff: Mentor

10. Jan 30, 2017

### barryj

I calculated the pH to be 12 based on the erroneous fact that the NaHCO3 would totally dissociate which it doesn't.

11. Jan 30, 2017

### Staff: Mentor

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.

12. Jan 30, 2017

### barryj

Then the HCO3 disassociates into OH and CO2 and then I have the OH I need.

13. Jan 30, 2017

### Staff: Mentor

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.

14. Jan 30, 2017

### barryj

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. Feb 1, 2017

### Kevin McHugh

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