If you have chemistry questions....

Oregon Panda said:

My wife makes me drink baking soda and citric acid every morning.
It tastes terrible.
Is she trying to kill me?

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I don't get the point, either. Do you mix them first and drink the sodium citrate, or let it bubble up in your stomach and burp for the next hour?

She puts it in water. Says its "good for me". I dunno though.

lol it was more a joke than anything but now I'm not so sure. OMG SQUIG my life is at stake here!

woodsmaneh, that sounds like excessive rambling to me.

You make a buffer by mixing a weak acid with its conjugate base generally speaking.

So a common mix for a buffer would be Citric Acid (weak acid) and sodium citrate.

So, just to clarify:

Citric Acid: H3C6H5O7 (acidic protons first)

Citrate: C6H5O7 -3

Sodium citrate (in this case trisodium citrate) : Na3C6H5O7


Mixing a strong acid with a strong base just produces water and a salt.

Example:

NaOH + HCl = H2O + NaCl (table salt).

SQUIGGS! Enquiring minds have got to know....

Oregon Panda said:

She puts it in water. Says its "good for me". I dunno though.

lol it was more a joke than anything but now I'm not so sure. OMG SQUIG my life is at stake here!

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ttystikk said:

SQUIGGS! Enquiring minds have got to know....

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3NAHCO3 + C6H8O7 --> Na3C6H5O7 + 3CO2 + 3H2O

(sodium bicarbonate) + (citric acid) ---> (sodium citrate) + (carbon dioxide) + (water)

In short, it's a fancy way of increasing your sodium intake and making you burp. I suppose citrate might be beneficial, but you're probably better of getting it from citrus fruit.

TLDR -- Your wife is wrong. If you have any concern that you eat to much sodium already (like most Americans do) you should probably stop drinking this.

Hey squiggs, long time lurker of your thread, first question.
I'm being told a pinch of sugar in a 5-gal bucket will neutralize chloramines. Does this work, and can you tell me how? Thanks!

altitudefarmer said:

Hey squiggs, long time lurker of your thread, first question.
I'm being told a pinch of sugar in a 5-gal bucket will neutralize chloramines. Does this work, and can you tell me how? Thanks!

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It will not.

Sucrose (table sugar) is a non-reducing sugar. It cannot donate hydrogens easily enough to react with the chloramine.

I know that ascorbic acid (vitamin C) works for this and is the industry standard. It's also fairly cheap, especially in bulk.

Glucose MIGHT work for this, and the scheme would look something like this mechanism I just drew up for you.

hope alls been well squiggs!

Actually I linked from the wrong OH group in the initial arrow, it would actually be the O from the dotted cone that would loop around and make the attachment to the aldehydic carbon--but you get the point.

Thanks, Squiggs. I appreciate solid info. Pm if there's anything I could offer in return.

altitudefarmer said:

Thanks, Squiggs. I appreciate solid info. Pm if there's anything I could offer in return.

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Helping is it's own reward, good luck getting the nasties out of your water.

squiggly said:

Helping is it's own reward, good luck getting the nasties out of your water.

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I was the source of the table sugar thing, and I read it in what I thought was a reputable site for gardening. Citric acid is easy to get, that will do fine.

Hey Ty, I've heard that before, brother. I sure wasn't pointing fingers or doubt in your personal word. Much respect.

altitudefarmer said:

Hey Ty, I've heard that before, brother. I sure wasn't pointing fingers or doubt in your personal word. Much respect.

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Thanks brother. Here's the thing about me; I'm a passionate debater for what I believe is right. Today, we had a clear example of someone who knows their shit saying that my information was wrong. Instead of taking it personally, I'll just preach the new gospel of citric acid or the like.

I care about the truth. I care not whose truth it is, I'm just interested in the truth. This is what the steppingstones of progress look like. I'm here to pave that road forward, not argue who laid which bricks- or worse, that the bricks are wrong, in the face of obvious expertise that says otherwise.

I'm stoned x I ramble... but damnit, this is right at the heart of who I am, and why I'm here. Ya feelin' me?

Hey squiggly I have a question on Potassium and sodium silicate r.e. pH and stability/solubility.

Ive been a bit confused with the chemistry because it seems to me that to keep Si soluble and non reacctive in concentrated stock solutions, pH needs to be kept high (>8). But to keep Si soluble and non reactive in working solutions, pH must be kept neutral(<8).

As I understand it, at concentrated levels, without the presence of other ions (e.g. in stock solutions), Silica solubility seems to increase with increasing pH. However at diluted levels, in the presence of ions (e.g. working solution), silica solubility seems to be in that 4-8 pH range again [see thumbnailed graphs]. As I understand it, at high pH (>8) silicic acid turns anionic, making it incredibly soluble and stable with ITSELF at concentrated levels, but seems to be highly reactive to cations in this anionic state (at high pH).

Is this correct or have I misinterpreted the info? Hope you can set things straight.

What causes blue balls? Chemically speaking of course. :)

dizzlekush said:

Is this correct or have I misinterpreted the info? Hope you can set things straight.

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When you're talking about this anionic transformation, it happens in a stepwise fashion--in effect the stuff acts as a weak base. So you start with, at very high pH, SiO3 (2-). As you lower the pH some of this converts to HSiO3 (1-) by abstracting a proton from solution. So you end up with a ratio of HSiO3(1-)/SiO3(2-) which increases with decreasing pH. When the ratio collapses and you end up with 100% SiO3(1-) in solution if you decrease the pH even further you will start a new ratio of H2SiO3(solid)/SiO3(1-) and the stuff will start to precipitate out.

If you look at the charge on magnesium it's 2+, so what's happening here is that in the HSiO3(1-) form the anion can effectively chelate the magnesium (2 equivalents per Mg cation). but once you have SiO3(2-) in solution the anion and cation can start to form strong ionic bonds, which can cause them to precipitate out.

Adding other cations/anions to solution can muddle the predictive process here.

If you look at your second graph you can still have a pH above 8.00 and keep the stuff soluble. When you hit a pH around 8.75 the HSiO3(1-) starts giving up protons and converting to the (2-) form. That's when you'll start to get magnesium salts.


To be clear, this is probably a faulty explanation because Sillicic acid actually acts very strangely in solution and forms complex structures which interact with each other. It really is weird stuff. When you add a bunch of other stuff to it the thermodynamics become totally unworkable from a theoretical point of view--and that's why experimentation becomes necessary.

So while these transformations/ratios I've described might not accurately depict what is actually happening--this is the general thermodynamic form you are looking at from an idealized perspective.

I couldn't find pKa/pKb values for the stuff so it's tough to say at what pH what is happening. I think the general idea here is to keep the concentrations low to avoid all of this mess. The sweet spot for maximum solubility, if you superimpose both of these graphs, appears to be somewhere in the 8.5-8.75 pH range.

Thought I'd share with everyone that I took my first toke in 6 months today. Sweet release :)

squiggly said:

Thought I'd share with everyone that I took my first toke in 6 months today. Sweet release :)

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I know that feeling man, a great one it is. Have you been able to kick the back to back cigarettes yet?. Hope you are well man

I bet ya got seriously high after a 6 month break, huh? I took a year off once....and the first bowl was a 2 hour spatial/visual trip that was very intense at times. Totally blew me away.