We like to think that the second bath of two can only work while the first bath remains in the emulsion. It might be illuminating to use the 2-bath in the normal manner and use the remains of the second bath as a single solution developer for another piece of film. My money says you'll get an image after a while. Might start a trend among the masochists who leave film in weak developer for hours under the guise of stand development.
WOW!!! You guys are insane.. My head hurts after reading all of that... So maybe I should try a divided D23 or D76.... Experiement with the Bath B's, which will cause increase/decrease in contrast and grain??
I don't know what you mean by "in general", but this is not really right.
Originally Posted by Donald Qualls
The pH of a solution is not determined by the strongest alkali present, the pH of the solution is determined by the equilibrium of the hydrogen ions donated (or accepted) by the compounds in the solution. Of course, one must consider relative concentrations of these compounds in the solution, as well as the alkalinity (or acidity) of each compound.
"unless there's a buffering agent present, the pH of a solution is determined by the strongest alkali present" is wrong. You can easily make a solution that contains both sodium sulfite and sodium hydroxide and the pH can be much closer to the pH of a solution that contains a lot of sulfite, as opposed to one that conatins a lot of hydroxide.
Buffering is a measure of a solutions resistance to change it's pH. Chemicals which are pH buffers are weak acids or bases, acids are proton (H+) donors, and bases are proton acceptors. Acids and their conjugate bases are in equilibrium, since equilibria are related to the properties of the reactants and products. So once you put these compounds into solution, they react and the solution comes to equilibria. It really has got nothing to do with "the strongest alkali present", as that statement does not take into account the concentration, i.e. the ratio, and the resulting equilibrium of the compounds present.
And in the case of solutions that conatin both carbonate and borax, they are both are pH buffers, so then what does that mean in reagards to what you said?
So if you have both sulfite and borax, the borax does have a greater alkalinity than the sodium sulfite, and it is even a better buffer, but depending on the amounts of each in solution, the pH may range from around 7 if there is a large excess of sulfite, to above pH 9 for a large excess of borax. You can get any pH you want in between those two values with different ratios of those two compounds.
Or take solutions of carbonate and bicarbonate. A solution that has a high amount of carbonate and a very low amount of bicarbonate, will have a pH of abound 11. As you add more bicarbonate (or acid to convert some of the carbonate into bicarbonate), the pH of the solution will lower. When you have equal amounts of carbonate and bicarbonate in solution, the pH will be about 8.3. Continue on making solutions that have a great excess of bicarbonate over carbonate in them, and you will get a solution that has a pH of around 7.0.
Even though the carbonate is a stronger base (and it has a greater alkalinity for a given amount) than the bicarbonate, it's not which is stronger, it is the ratio of them in the solution that determines the pH.
The same goes for different ratios of borax and carbonate. The pH of a solution containing both borax and carbonate will reach an equilibrium pH - one that is somewhere between the pH of pure and concentrated solutions of each of these compounds.
And therefore, you can't make statments like "If there were (for some reason) both borax and sodium carbonate, the carbonate would act as the alkali because it produces a higher pH than the borax." This statement does not take into accound the amount of the borax or the carbonate, and it completely ignores the fact that any solution of these two compounds will reach an equilibrium.
It's not like the carbonate does all the pH work, and once that gets used up, the pH of the solution jumps down to the pH of the next "strongest alkali", the borax. Doesn't happen like that. (I do note that you do not say it drops in a big step, but you seem to imply that it does.)
The pH of a mixture of carbonate and borax will change continuously as acid is added to it (or even stronger base). There is a continuum of pHs there that can be made, and you should consider a solution that contains both borax and carbonate (or any other compounds) to be a mixture that has a distinct pH based on the equilibria of the concentrations of those two compounds.
Kodak packaged products are now dated so you should not receive old chemicals.
Commercial products produce excellent results and unless you need some special effect not obtainable, I would invest in good glass storage containers instead.
Another thing to consider is metol will age on your shelf. I costed Dektol and D72 in a gallon size and there was no cost savings based on Photographers Formulary costs a few years ago.
Learn to print well before branching into unknown territory.
I've got news for you and Mr. Qualls. Sodium sulfite has
a GREATER alkalinity than borax. One-shoters don't need
borax. The borax is there to absorb OH ions produced
by the oxidation of hydroquinone; ph maintanance.
The ph of bicarbonate and borax are nearly the same;
9 plus a tenth or two depending on the data you read.
The ph of sulfite runs above those two, up to
several tenths, also depending. While on the
subject of depending, who have you two
been reading? Maybe A & T? Dan
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R. Suzuki suggests sodium carbonate and the bicarbonate
Originally Posted by Kirk Keyes
for that purpose.
I believe most developing agents are acid and if the ratio
of the agent and other components is low it can pull the ph
quite low. I worked with an 8 and 80 gram D23. The sulfite
would not pull the ph above 7.8 +/- .1 and I've seen
sulfite peged at a ph of 9.8 by a supplier. Dan
I just checked the pH of a 5g/100ml soultion of sodium bicarbonate, and I got pH 8.0. A freshly made solution of sodium sulfite at 5g/100ml was pH 10.0.
Originally Posted by dancqu
I've never tried to titrate sodium sulfite solutions, so I'm not sure about the claim made here about it having a higher alkalinity than borax.
Originally Posted by jim appleyard
I'll sort through my box of formulas later tonight and try to find it.
Not sure? You say the sulfite measured ph 10.0. I've measured
Originally Posted by Kirk Keyes
borax 1% at 9 and I see via the WWW 9.18 mentioned.
I've measured bicarboate, also 1%, at 9 but think that may be
atypical or an error on my part; 8.2 is indicated. Dan
Perhaps your bicarbonate was either not high quality (i.e. it may have had some carbonate in it), or, as you say, was an error on your part. Can't really speculate without more info.
Originally Posted by dancqu
A 50/50 solution of carbonate/bicarbonate should have a pH of 8.3, as I said below. That fact is routinely used to calculate the concentration of each componant of a solution containing carbonate species. If you have a pH below 8.3, you have more bicarbonate, if the pH is above 8.3, you have more carbonate.
See "Standard Methods for the Examination of Water and Wastewater", any edition, Method 2320, Parts A and B. It's a nationally approved method for determining the alkalinity of a solution.