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albada
02-25-2012, 12:48 PM
If you use borax/boric acid, plan ahead, borax just doesn't want to dissolve fast, even at 40C :( And I mean many hours, not minutes ...

I've dissolved Borax at room temperature, and it only took a few minutes to dissolve. The Borax available here in the USA is sodium tetraborate decahydrate. Is that what you've got? Anyway, I dissolved it at up to 9 g/L; perhaps you are pushing it higher, closer to the solubility limit.

I've resisted using Borax because of the ten water molecules on it. I'm designing a concentrate, and I'm unhappy enough with having two water molecules on some sodium metaborate. With five times as much water in it, my instructions will have to say, "Keep the concentrate hot until the steam stops forming condensation on your safety glasses." :)

Although, maybe that extra water will steam-out quickly enough and won't be an issue. I might try it.

The concentrate formula I'm narrowing-in on looks like:



Propylene glycol (80-90C) ....... 25 ml
Sodium metaborate 4 mol ....... 3.5 g
Ascorbic acid ........................... 6 g
Boric acid ................................ 4.2 g
Phenidone ............................... 0.08 g
Propylene glycol to .................. 33 ml (1+29 dilution)

pH = 8.0. One liter of dev contains 66 g of sodium sulfite and 33 ml concentrate (1+29).
Time for TMY-2 at 20C is 10.5 minutes.


Today I'll test this formula using citric acid instead of boric. Any comments/critiques of this approach are welcome.

Mark Overton

Photo Engineer
02-25-2012, 01:18 PM
Well, Phosphates are not very good buffers and the only evidence I can offer is the fact that EK does not use them in any developers except those at pH 11 and higher! Borax will dissolve quickly or slowly depending on crystal habit and water of hydration. That is why we are seeing differences here.

Actually, the best alkali for this pH would be TEA, but we have already seen the problems with differences there. I have used the Formulary TEA now for over 10 years with no problem. I talked to them about this and it is Photograde and is used by many other companies for photo products. There is a purer grade with a big jump in price.

You can force an alkali to work in a lower or higher range than what it wants to buffer in. It will work. BTDT. It just can move on you but if the developer is designed to work with this situation there is no problem.

An added thought. You have Ascorbate / Ascorbic acid there and that is a secondary buffer all in itself. If you use HQ then you have the HQ anion present along with HQ which is an acid. This forms a secondary buffer. It depends on the amounts you use.

PE

Rudeofus
02-25-2012, 01:19 PM
I've dissolved Borax at room temperature, and it only took a few minutes to dissolve. The Borax available here in the USA is sodium tetraborate decahydrate. Is that what you've got? Anyway, I dissolved it at up to 9 g/L; perhaps you are pushing it higher, closer to the solubility limit.
Ron recommended much higher concentrations (>25g/l) and I am still stuck trying to dissolve 35g/l in a solution which already contains 40g/l sulfite. If this doesn't work out soon, I'll incrementally add more water and sulfite stock to make it dissolve at last.


I've resisted using Borax because of the ten water molecules on it. I'm designing a concentrate, and I'm unhappy enough with having two water molecules on some sodium metaborate. With five times as much water in it, my instructions will have to say, "Keep the concentrate hot until the steam stops forming condensation on your safety glasses." :)
Don't forget that while you have 10 water molecules per molar unit of Borax, you also have four times the boron, which means you end up with a much more favourable water ratio for borax, depending on metaborate hydration, of course.

Photo Engineer
02-25-2012, 01:31 PM
If this will not dissolve, then dilute it slowly until it is fully dissolved at room temperature. Then you know the maximum concentration that you can use. You can then back calculate the concentration and then make up that "proper" solution again. This will tell you the maximum concentration of your concentrate and of your working solution if you choose to dilute it.

PE

Rudeofus
02-25-2012, 02:22 PM
If I dilute this recalcitrant mix slowly, I'll end up with some odd ball concentration and risk another round of precipitation once I add further components to the mix. Instead I'll add just enough water and sulfite stock to reduce the borax level to 30g/l and 25g/l, sulfite concentration kept constant, of course. If buffering is really an issue, or if the experiment requires that buffering is optimal, I have no problem using 500ml on 135 or 120 rolls. Doubling the amount of soup and leaving the amount of film the same should have the same effect as doubling the buffer concentration (assuming good agitation).

I know, a commercial dev wouldn't get away with 500ml required for 135/120, maybe that's the rule we need to break to advance in performance.

albada
02-26-2012, 01:11 AM
The pH of sodium sulfite drops quickly after mixing.
Initially, 2 grams of sulfite in 30 ml of distilled water has a pH of 9.9 to 10.0.
After about 20 minutes, the pH has dropped to 9.75 or 9.80. I confirmed this with two different batches of sulfite.
The odd thing is, after a developer is completely mixed, I haven't noticed a pH-drop in this time-frame. Perhaps ascorbic acid is protecting the sulfite. :(

Any idea what's going on here?
Does aerial oxidation of sulfite occur that quickly? Does it produce an acid which pulls down the pH?

That drop of 0.1 to 0.2 is enough to have a noticeable effect on activity. I'm tempted to mix in just a little sulfite at the start, add everything else, then add the remainder of the sulfite last.

Potassium sulfite is looking better than ever...

Mark Overton

Photo Engineer
02-26-2012, 11:21 AM
Sulfite can oxidize to Sulfate which is a stronger acid and which can lower the pH. I am uncertain as to how fast this can take place though. That sounds to rapid to me. Also, it is common, when mixing developers, to add some or all Sulfite first to protect the developing agent which is added to the mix.

PE

Rudeofus
02-26-2012, 12:42 PM
I'd like to put some perspective of what the pH difference reported by Mark would mean if it came only from conversion of sulfite to sulfate. If we use the pH calculator (http://www.webqc.org/phsolver.php) suggested by him, we use the following constant: 2g of anhydrous Na2SO3 in 30ml water is 1000/30 * 2/126 = 0.529 moles/liter. The following constants can be obtained from their respective wikipedia pages:

H2SO3: pKa1 = 1.857, pKa2 = 7.172
H2SO4: pKa1 = -3, pKa2 = 1.99
NaOH: pKb = 1

Using a concentration of 0.529 mol/l H2SO3 and 2*0.529 mol/l = 1.0582 mol/l NaOH yields a pH of 9.92, which is so close to Mark's measurement that I would be suspicious if a lab worker presented me that number. Now we can increase the concentration of H2SO4 at the expense of H2SO3 and see, how pH value changes.

If 1/4 of H2SO3 is converted to H2SO4, pH value lowers to 9.85.
If 2/4 of H2SO3 is converted to H2SO4, pH value lowers to 9.76.
If 3/4 of H2SO3 is converted to H2SO4, pH value lowers to 9.61.
If 4/4 of H2SO3 is converted to H2SO4, pH value lowers to 7.33.

Now, just for the sake of experiment, I simulated addition of H2CO3 to the mix, note that CO2 is part of air and will always dissolve in water to some extent.

Assuming that none of the H2SO3 is converted to H2SO4, the addition of

0.00001 mol/l H2CO3 lowers pH to 9.91
0.0001 mol/l H2CO3 lowers pH to 9.89
0.001 mol/l H2CO3 lowers pH to 9.66
0.01 mol/l H2CO3 lowers pH to 8.87

Note that 0.001 mol/l H2CO3 would mean only 0.044 g/l of CO2 dissolved in water. Total solubility of CO2 is, according to wikipedia again, 1.45 g/l.

From all this I would conclude that aerial CO2 going into the solution would be a more likely explanation of the pH shift than oxidation of sulfite.

Photo Engineer
02-26-2012, 01:31 PM
Well, good answer. I discounted CO2 assuming that tap water or DW from the bottle would be saturated with air and thus CO2. Maybe it is a little of both, conversion to SO4 and absorption of CO2. IDK.

Your comment sounds reasonable.

PE

Rudeofus
02-26-2012, 02:07 PM
Well, good answer. I discounted CO2 assuming that tap water or DW from the bottle would be saturated with air and thus CO2. Maybe it is a little of both, conversion to SO4 and absorption of CO2. IDK.
There is a good chance that Mark dissolved the sulfite in hot water or it would have taken forever. After it cooled down, it took in extra air, maybe more slowly than the cool down. After 20 minutes some equilibrium could have been reached.

If 1% of the SO3-- would convert to SO4-- in 20 minutes, SO3-- would not be used as preservative for developers and fixers, which should last at least a few weeks. A 1% decay in 20 minutes would equal a 50% decay after one day and a 99.4% decay after one week. Note that a 1% decay of SO3-- would not be noticeable with the pH meter.

albada
02-26-2012, 05:09 PM
The odd thing is: The water was at room-temperature of 22-23C.
My distilled water has always been kept capped, so there's probably little gas dissolved in it.
I let the beakers sit overnight, and their pH's were 8.65 and 8.72 next morning.
In comparison, a test-developer I also let sit overnight went from 7.91 (at night) to 7.95 next morning. Its pH went up!

Anyway, I think I'll let some distilled water sit for hours to absorb all the gases it can, and then mix some sulfite and measure pH.

Mark Overton


There is a good chance that Mark dissolved the sulfite in hot water or it would have taken forever. After it cooled down, it took in extra air, maybe more slowly than the cool down. After 20 minutes some equilibrium could have been reached.

If 1% of the SO3-- would convert to SO4-- in 20 minutes, SO3-- would not be used as preservative for developers and fixers, which should last at least a few weeks. A 1% decay in 20 minutes would equal a 50% decay after one day and a 99.4% decay after one week. Note that a 1% decay of SO3-- would not be noticeable with the pH meter.

Photo Engineer
02-26-2012, 05:13 PM
DW does pick up gasses as it condenses in the still. It also absorbs gases through the plastic of the bottle!

The results may be an erratic meter. Did you check the calibration before use?

PE

albada
02-26-2012, 07:47 PM
DW does pick up gasses as it condenses in the still. It also absorbs gases through the plastic of the bottle!
The results may be an erratic meter. Did you check the calibration before use?
PE

The electrode is only 2-3 weeks old, and always been stored in storage solution. But I suspected the probe too, which is partly why I mixed a second batch. After measuring the low value in the first batch, the meter immediately popped back up to a high value when placed into the second batch. Then the second batch began drifting down. This afternoon, the pHs of both of those batches measured very low (8.something), but the meter acted fine with a test-brew I made soon afterwards. It was exposed to that brew for half an hour, because I work slowly (meticulousness eats time). Its behavior was steady the whole time. I checked the pH of the 10.01 cal-solution, and the meter measured .03 off. I'm thinking the meter is fine.

I'm wondering if the sulfite oxidizes faster at high pHs. If so, I could add the acid immediately after pouring in the sulfite. Although this would make no difference if Rudi's CO2-theory is correct.

Mark Overton

Rudeofus
02-27-2012, 03:03 AM
I'm wondering if the sulfite oxidizes faster at high pHs. If so, I could add the acid immediately after pouring in the sulfite. Although this would make no difference if Rudi's CO2-theory is correct.

You could try the following: since you don't want a pH of 10 anyway, you could start with a suitable mixture of sulfite and bisulfite to get a pH of 8. The pH value of this mix should stay a lot more stable than sulfite alone, CO2 or not.

albada
02-27-2012, 07:19 PM
You could try the following: since you don't want a pH of 10 anyway, you could start with a suitable mixture of sulfite and bisulfite to get a pH of 8. The pH value of this mix should stay a lot more stable than sulfite alone, CO2 or not.

That's a good idea. I'm using boric acid instead of metabisulfite, but the same principle applies.
Still thinking about this: If pH actually drops, then I have a hard time believing that D-23 users haven't noticed this over the decades. Yet, the meter's acting fine in all other respects. Strange.

BTW, that formula I posted a few postings ago here (http://www.apug.org/forums/viewpost.php?p=1308010) works well, but is slightly grainier than XTOL. Don't know why.

Mark Overton

albada
03-09-2012, 01:06 AM
Well, I got another mystifying result that more knowledgeable folks might want to remark on. I developed a roll of Tri-X in a stainless steel tank using 225 ml of the following one-litre formula:



Sodium sulfite
48 g


Ascorbic acid
7.4 g


Sodium metaborate
4.2 g


Dimezone S
0.11 g



Target pH = 8.00. Tri-X: 10.5 minutes at 20C. Tmax-400: 13 minutes.



The roll came out rather thin. I'm not disappointed because my electronic enlarger likes it and the grain is surprisingly fine. However, test-strips dev'd in the same formula come out with normal density. The developer was clear when poured out, so it wasn't exhausted. Any ideas why the roll was thin?

Another question: Is there any reason to think the formula below would give better or worse image-quality than the one above?:



Sodium sulfite

48 g


Ascorbic acid

5 g


Sodium metaborate

3.1 g


Boric acid

3.2 g


Dimezone S

0.11 g




The only difference is the addition of boric acid and reduction/adjustment of the other acid and alkali. Same pH of 8.00 and same dev-times. My buffering tests show that both formulas have the same buffering strength: A 15% change in the sulfite causes a pH-shift of 0.05.

BTW, I learned how to create tables in apug. :)

Mark Overton

Photo Engineer
03-09-2012, 09:51 AM
Those two developers should be similar, but then chemicals don't know chemistry! :D

It is pretty much T&E (Trial and Error)

PE

albada
03-12-2012, 04:16 PM
Does developing a roll of film destroy a certain number of milligrams of developer? I'm still investigating why test-strips are denser than a roll. That happens with both the Dimezone S formula above, and the Phenidone formula below:



Sodium sulfite
48 g


Ascorbic acid
7.4 g


Sodium metaborate
4.2 g


Phenidone
0.05 g



Target pH = 8.00. For Tri-X: 15.5 minutes at 20C (came out thin).

Here was my experiment:

Mix the developer.
Dev a test-strip. Leader-density = 2.70.
Dev a roll of Tri-X. Leader-density = 2.60.
Dev a test-strip. Leader-density = 2.46.

So developing a roll weakened the developer considerably. Yet the dev was clear when poured out, so it wasn't exhausted, at least not exhaustion that makes it turn yellow or orange. I'm thinking that the culprit is simply that developing a roll destroys some of the Phenidone. Based on the numbers above, and knowing I used 220 ml in the tank, I'd guess around 7 milligrams of Phenidone was destroyed (out of the original 11 mg). Sound plausible?

Mark Overton

albada
03-12-2012, 04:31 PM
I made an H&D curve.

I bought a densitometer, thanks to Greg Davis who advertized it on apug's classifieds. And I got a 21-step Stouffer wedge. Put those two together, and you can make H&D curves. Here's my first:
47706
This is taken from the roll developed in the prior posting.
Notice that it's upswept. The film was underdeveloped some, which might be responsible for the long toe.
But I'm still proud of my H&D curve. The X-axis is simply the step-number, but that can be converted to the proper lux-seconds.
BTW, H&D stands for Hurter-Driffield, its inventors. But I like to think it stands for Helios-Density, where Helios is the Greek sun-god, representing illumination on the X-axis, and where Density is the Y-axis.

Mark Overton

Alan Johnson
03-12-2012, 05:16 PM
It is usually considered that it is the ascorbate which is used up, as reported by Gainer:
http://unblinkingeye.com/Articles/Synergism/synergism.html