When Adding a base lowers pH

[wolf_merker]

Friends,

I have spent some time doing acid-base titrations in non-aqueous solutions. Recently I discovered a very strange phenomenon that I cannot make much sense of. The system is a combination of three organic but somewhat polar solvents (they mix with water, but NaOH is not soluble in the mixture). Dissolved in the solution is a classic pH indicator at about 20% by mass (fully soluble). The solution is then titrated with NaOH predissolved in water at 1% by mass. Reults: ph starts at 4.2 and increases with base additions (despite NaOH not being soluble); upon reaching pH 7.2 the next addition of base sends the solution crashing to pH 3.2. Additional base added does not change the pH much at this point but keeps decreasing it slightly. The solution color (indicator color) is consistent with these measurements - it goes back to its original color when the pH plummets.

Has anyone ever encountered something like this? Any ideas what could cause it?

 

Wolf

[John Cuthber]

pH is badly defined in non aqueous solutions. Also the pH range of indicators is solvent dependent.

It's weird to have an indicatior present in such large quantities too.

How are you measuring pH?

[wolf_merker]

A couple of things: The indicator is present at the quantity because it is used in an electrochemically active ink at that ratio in my work. Basically, it is functioning as a weak acid solution and as an indicator at the same time. The funny thing is that I have had several solutions display this strange behavior now - reproducibly too.

the pH is being measured with a glass pH probe. This has given us some grief in the past, but we have found one that works reasonably well. At least we can do reproducible titrations in various solvents and get normal (or at least logical) titrations curves. the pH plummeting in the bulk solution seems to be real because both the probe and the "indicator" are telling us the same thing.

When examining the solution closer under a microscope between base additions, you can see dark particles (corresponding to indicator color in base) that dissolve upon v e r y vigorous mixing. This mixing seems to be able to induce the pH plummeting at a slightly earlier point in the titration than would happen with normal mixing and adding.

[Bignose]

that dissolve upon v e r y vigorous mixing.

 

Please note that this is just a shot in a dark and may be very off base, but the phrase I quoted here caught my attention. It made me want to bring up at least a question or too -- in that is the temperature variation being taken into account? If the mixture has to be so vigorously mixed, there is going to be a temperature rise from the viscous dissipation in the fluid. Then, if you had another fluid, which may not be at the same temperature, the temperature change can change the pH, maybe even more so than the basic nature of the new fluid?

 

Or, is the solution unstable? That is, the mixing overcomes the stability of the droplets of the indicator by breaking the droplets, but the addition of a new fluid again causes the indicator to go back to droplets and thus not working like a true indicator should because it is re-coalescing?

 

Like I said, these are just guesses, and may be way off.

[CaptainPanic]

Can it be that the addition of water is the reason? With the NaOH, you also add plenty of water, I think?

 

I can imagine that when there is more water present, more of your (organic) solvents will release a proton... is any of those solvents an organic acids perhaps? (Like acetic acid?)

 

The protons need a "host" (a water molecule) to form the H3O+ ion.

 

I also want to ask if there is any phase separation? It is not uncommon that a homogeneous mixture forms two phases when more water is added. If there is a phase separation, that could be a good explanation of a changing pH value... because all the concentrations will also suddenly change.

[wolf_merker]

Alright, let's give this a shot. The temperature difference is noticeable after mixing. We use a machine called a speedmixer; with the sample size in question and it's lack of viscosity of this particular solution the difference might be about 10 degrees C. My guess is that the dark particles/droplets are NaOH with a tiny bit of indicator. Becuase of the solubility problems with NaOH in our solvent system, we are going to try another base that has been ordered. Perhaps it will show normal titration behavior and this is all just a strange pH depended solubility thing.

 

As for the question of water. This is one we have been considering as well. By weight, we add 50% NaOH in water. So we're not adding much, but in a solution lacking water the effect on the pH probe could be significant. Strangely, if it were reading the water phase (not obvious as a second phase in the solution, but seemingly well dispersed) one would intuitively think it would immediately start reading a very high pH upon addition.

None of the solvents are organic acids. One of them is, however, a protophile. The other two should not affect the pH by direct H+ release or uptake. Again, no real phase separation is noticeable, and the "particles" are solid as seen under a microscope.

 

I'll keep you updated with the progress when the new base arrives and we repeat the titration with a soluble base.

 

Here's another observation in the mean time: When you titrate the solution up and its pH drops down after a certain base addition or additional mixing, the phenomenon only happens once. Upon further titration, everything follows a logical progression with increasing pH at various rates until an appropriate endpoint is reached.

[CaptainPanic]

You're confusing me a little with your choice of words:

 

You say that there's no obvious second phase, but it is well disperded.

"Well dispersed" means to me: "2 phases, in very tiny droplets"...

If you mean that there is only 1 phase, "well mixed" or "homogeneous" would be a better choice of words.

 

Also, I am confused about the concentration of NaOH: the solution of NaOH in water is 1% wt NaOH? (In your last post, you suggest it is 50% wt NaOH?)

 

Finally, I am confused about the particles... in your first post you write that the indicator is fully soluble... so what are the particles?

[wolf_merker]

I'm sorry for the confusion.

There is no obvious liquid separated phase. There are, however, some solid particles. I do not know if they are pure solid, or solid particles in water droplets. My choice of words was unfortunate. What I meant was that a small amt of solvated water could influence the probe more that the bulk solution as a whole.

I'm not sure where you see 1% NaOH... We are adding 50% NaOH 50% water.

 

About the indicator: It is fully soluble. The droplets are the color of the indicator in a very basic solution. So while it is soluble in the solvents, some seems to be involved in the "droplets". I doubt it is part of the solid, but more along the lines of a thin dissolved very basic layer around the droplets/particles. The solid is most likely NaOH which is not soluble in the solvent system. Obviously not a good choice for a base, but it is somehow able to still alter the pH and produce titration curves in the system. Other very similar solutions have titrated just fine in the way, but 2-3 solutions have exhibited the strange pH drop.

[CaptainPanic]

I still think that you have some kind of phase separation, or perhaps the opposite: two phases combining into one, because those are typically processes that can suddenly happen. One droplet extra can be enough to create a new phase.

 

You describe a very sudden pH decrease from 7 to 3, which is strange, since you keep adding NaOH. But you are changing the polarity of the system a lot by adding the 50-50 NaOH-H2O solution.

 

What happens to the pH when you dilute the whole thing with lots of water (like 5x the total volume of water)? By adding lots of water you make sure that everything is dissolved, and you can make a decent measurement of the pH.

[wolf_merker]

Well guys, here is the update. I repeated the titration in water. The "indicator" (used here to make a weak acid solution) exhibits the same behavior. It's quite a shocker. I titrate my aqueous solution with tetramethylammoniumhydroxide (instead of NaOH) and the only change is where this anomaly happens. I titrate up to about 5.8 and measure the pH. While the pH electrode is in the solution it suddenly registers a massive shift in pH to below where I started. Again, the color of the solution agrees with the electrode, but now I have no insoluble species at all. pH progression is as follows. 2.4-->5.8 with base additions. Then 5.8-->1.5 spontaneously, suddenly and with no acid added at all. With further titration the solution behaves normally.

[battousai]

In a previous post, CaptainPanic stated that the addition of water could be causing your strange results. To prove/discount this possibility, I would repeat the original experiment but with a solution of NaOH with a lower and then higher H2O content, and compare the results. Comparing to the original titration, if the pH of the the solution drops when a smaller quantity of NaOH solution of a higher H2O concentration is added, and if the pH drops after a larger quantity of NaOH solution of the lower H2O concentration is added, then water is most likely the problem.

 

In addition I would suggest retrying the original titration, but by using a base that does not need to be in aqueous solution. If you can get the temperature low enough, try using liquid ammonia, or, if you can, try Lithium diisopropylamide. According to wikipedia it will dissolve in non polar organic solvents. From wikipedia as well, it says that n-Butyllithium dissolves in diethyl ether, which is a polar organic solvent, so that might work.

 

Finally, maybe one of the weak bases: Alanine, Methylamine, or Pyridine; mentioned here, might work in your solution. I'd highly suggest retrying the original titration without the use of any water, so if you feel that it's worth it, give it a try.