Donald Qualls said:
In general, unless there's a buffering agent present, the pH of a solution is determined by the strongest alkali present -- so if you have both sulfite and borax, the borax determines the pH and we use the shorthand of saying the borax acts as the accelerator or alkali in that developer. 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 (one might do this with very low levels of carbonate to produce a graded developer, which will exhaust from high activity to low via neutralization of the tiny amount of carbonate, allowing pH to drop to that provided by the borax -- though I've never heard of a developer that does this). Sulfite, in most developers, acts only as a preservative (and silver solvent, if there's enough and the developing time is long enough), because there is a stronger alkali present in sufficient quantity not to be neutralized during development.
I don't know what you mean by "in general", but this is not really right.
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.
) over or under exposed the film. The second bath for me is sodium metaborate, which I think Ole has identified as Borax and which I seem to remember also being called "washing soda"? The idea (short and simple) is that a short bath in "A" gets the developer into the emulsion without any real developing happening, then the activator in "B" gets the developing going, but the absorbed developer is soon exhausted in the highlights, so they don't get all burned out, while it keeps working in the shadows and you end up with good shadow detail.
