Did you mean aluminum hydroxide? Dan
Originally Posted by Photo Engineer
Does having the "milky" appearance interfer with the functioning, including hardening of the fixer?
Claire (Ms Anne Thrope is in the darkroom)
Aluminum hydroxide can only form at a pH well above 7. As I said your initial solutions is quite acidic due to the presence of the metabisulfite. You are then adding the hardener which contains acetic acid. When solutions of thiosulfate are made too acidic or when a local region of the solution becomes too acidic due to lack of stirring colloidal sulfur will come out of solution.
Advantage of this reaction is taken in the making of hypo alum toner. A solution of sodium thiosulfate and potassium alum are mixed and heated. The alum being acidic itself causes some of the thiosulfate to decompose into various sulfur compounds and colloidal sulfur. Prints are placed in this bath to be toned brown.
BTW, the decomposition of thiosulfate is the basis of a freshman chemistry demonstration called "The Setting Sun". A solution of sodium thiosulfate is set up with a slide projector behind the beaker shining on a movie screen. Acid is slowly added to the solution with stirring and what appears as an image of the sun slowly goes from white to yellow to orange and finally to dark red. This is caused by collodial sulfur scattering the light. As the particle size gets larger the light is scattered more resulting in the color shifting to red.
Gerald another basic chem demo is taking water you boiled red cabbaage in to use as a universal indicator. It changes from clear to pink depending on the ph. At the end of the demo I use to do at grade schools in the SW, I would ask what they thought putting silver nitrate would do to the tap water we experimented with. I was surprised when I was doing the chem demo in Park City Utah, and the solution immediately turned milky and cloudy right out of the tap, without me adding any silver nitrate. Later I talked to one of the City Engineers, and he told me that the mines in the area have effected the water supply. That is when I started using bottled water religiously.
I figured since This thread didn't indicate that it was precipatate, it had to be something similar to the cloudy milkiness you get from that experiment.
The solution can be either too acidic or too alkaline.
If too acidic, either the hypo will sulfurize and form colloidal sulfur with release of sulfur dioxide gas and hydrogen sulfide gas or if too alkaline the alum will precipitate as aluminum hydroxide. It is interesting to note that although the alum would be expected to precipitate at pH 7 or higher, I have observed it to ppt at as low a pH as about 6.0. It probably depends on the other ingredients salting it out. Aluminum salts are 'amphoteric' that is they can behave as either acidic or alkaline salts depending on the phase of the moon.
The use of ammonium or sodium salts has no effect on alum precipitating nor on any propensity to sulfurize. Hypo is hypo (thiosulfate), sulfite is sulfite, and alum is alum regardless of the mix of sodium and ammonium ions present. The overriding factor here is probably pH.
Just FYI, adjusting pH with sulfuric acid vs acetic acid at exactly the same concentration will give different results. The reason is that sulfuric acid is far more acidic at the same concentration than the same amount of acetic acid. Therefore, you are safe adjusting pH with acetic acid up to about 50% concentration whereas you would destroy your fix through sulfurization and release of sulfur dioxide gas at that concentration of sulfuric acid.
The point being that aluminum Sulfite can be quite acidic. It is basically aluminum hydroxide dissolved in sulfuric acid. At least that is one way to look at it. Therefore, it can have a profound effect on pH if it is too concentrated. If the fix is too concentrated and too alkaline, it can suddenly swing the other way and become too alkaline.
This is not a simple situation. Fix chemistry is quite complex.
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"If too acidic, ... the hypo will sulfurize and form colloidal sulfur
with release of sulfur dioxide gas and hydrogen sulfide gas ..."
Both of which are highly soluble in water and may give no clue.
I'm not so sure of H2S formation in that solution.
"... or if too alkaline the alum will precipitate as aluminum hydroxide."
That, at least from my reading, can be less than ph7.
"...and alum is alum... The overriding factor here is probably pH."
And in this case a potassium aluminum sulfate compound. Dan
Dan, the decomposition of hypo often leads to formation of H2S either through oxidation of acidification or both. The decomposition of sulfites to sulfur dioxide takes place with acidification. Both gases are sufficiently insoluable in the acidic hypo that the gas is given off in noticeable amounts.
Both are toxic.
Decomposition of alum hardening fixers at pH values over about 6 will lead to precipitates of aluminum hydroxide.
That fixer will have too high a pH to use alum hardener without an extra means of solubilizing alum. Old (pre WWII) alum hardening fixers have working pH of 4 to 4.5 to avoid this precipitation problem. In such a low pH range, the hardening effect is reduced, and the sulfurization becomes more serious problem after majority of bisulfite ions are oxidized. In modern alum hardening fixer, the solution contains agents that form soluble aluminium complex that also posseses hardening function. Classic work by Crabtree (or might it be Russell, I need to check) was to include boric acid in the fixing bath. There is a new patent assigned to Konica which uses alpha-hydroxyl acids and their derivatives for this function, allowing to formulate stable alum hardening fixer without using borates. With these agents, a practical alum hardening fixer can be formulated in the pH range of 5 to 6, but not much higher. Sludge will form above pH 6, even with boric acid. The hardening effect of alum is maximum at the highest pH that does not precipitate out insoluble aluminium salts. Therefore, good pH buffering is very important in optimal design of alum hardening fixers.
I would recommend you to buy a hardening fixer, or study the history and chemistry of fixing baths from the original research reports (that is, not Anchell and Troop or anything like that). If you want to jump start, I suggest to find a decent Kodak formula that incorporates boric acid in addition to acetic and/or citric acid in the fixing bath as a starting point.
I am pretty sure that the problem you have is not sulfurization. (Comments by Jerry Koch is irrelevant.) This is because sulfrization (acid decomposition of thiosulfate to form elemental sulfur) is very slow and almost insignificant within the pH range you are looking at, as long as bisulfite ions are not fully oxidized. You are seeing immediate precipitation after addition of alum solution. This is exactly what happens when the fixer's pH is too high in absence of borate (or another suitable agent) in the solution.
Originally Posted by Alessandro Serrao
Does it has any effect on the process?
I mixed some fix and it became milky today. It wasn't the good stuff since a shop I get supplies from didn't have my usual. I dipped a leader in to see time to clear. It cleared in a reasonable time... But I swear it looked like I'd soaked my film in a mud bath, it was covered with crud that wouldn't easily rinse off. So I didn't use it, and grabbed a new box of Kodak fix.