Experiments with Metol and ascorbic acid.
The factors affecting developer activity are temperature, pH,
concentration, products of aerial oxidation, products of silver halide
reduction, and the orientations of the heavenly bodies.
According to some experimental results reported in "The Theory of the
Photographic Process" as little as 0.05 moles (6.3 grams) of sodium sulfite in
a liter of Metol developer will double the rate of development that is
obtained when no sulfite is present, pH being held at 8.7 in both cases.
It is explained that oxidation products of Metol restrain development, and
that sulfite counteracts these products by forming the sulfonate.
Ascorbic acid will do the same degree of acceleration without sulfite, also at pH 8.7. In this case the oxidized Metol is reduced back to its original form. The ascorbate is itself oxidized and pH is decreased in the process. The decrease in amount of active Metol in the one case or the decrease of pH in the other may not cause much difference in capacity and storage longevity between a Metol-sulfite and a metol-ascorbate developer. However, it would seem that a well buffered working solution might tip the balance in favor of the ascorbate.
I will explore first a 0.05 molar concentration of Metol, which is
approximately that of D-23, together with a 0.05 molar concentration of
ascorbic acid. It remains to formulate the alkali needed to make the pH less
than that at which an ascorbate becomes a developer and great enough to ensure activity of Metol.
I calculate that 0.1 moles of sodium hydroxide will neutralize the 0.05
moles of sulfuric acid attached to the Metol base plus the 0.05 moles of
ascorbic acid, leaving 0.05 moles each of sodium sulfate and 4-(methyl-
amino)phenol and 0.05 moles of sodium ascorbate. The net pH thus far is still
not alkaline enough to develop film in a reasonable time, nor does it have the
necessary buffering action to keep the pH from going lower yet. A few grams of borax should bring the pH up to about 9.2 but even 20 grams will have little
more effect on pH and should give a cushion against reduction of pH during
The resulting recipe is:
Ascorbic acid.....8.8 g
Sodium hydroxide..4 g
Water to 1 liter.
The concentration of Metol is about 13% greater than in D-23 and the
concentration of sodium sulfite is nil.
The pH of this mixture, which I measure with test strips to be about 9, is
below that at which the ascorbate is a developer of any consequence. The
initial activity of the mixture is about that which could have been obtained
with sulfite in place of the ascorbate, so it does not seem that synergism
between Metol and ascorbate is the explanation.
The ratio of borax to caustic in this formula is quite close to the
effective ratio in sodium metaborate. 14.5 grams of sodium hydroxide and 69
grams of borax in a liter make a solution that is often used as a substitute
for 10% sodium metaborate. 276 ml of such a solution, or 27.6 grams of Kodalk if you prefer, will contain the equivalent of 4 grams of sodium hydroxide and about 20 grams of borax. It may be easier to get sodium metaborate than sodium hydroxide through the UPS. The recipe then becomes:
Ascorbic acid......8.8 g
Sodium metaborate..27.6 g or 276 ml of 10% sodium metaborate solution.
Water to 1 liter.
All ingredients are quickly dissolved in room temperature water. HP5+
developed for 8 minutes at 68 F gives normal contrast, showing that the
solution could be diluted. In fact, diluting with an equal part of water
increased development time by only 25%, probably because the pH of this
solution does not change much with dilution. 125 ml diluted with 125 ml of
water did a 36 exposure roll of HP5+ to normal contrast in 10 minutes, which
means of course that a liter of the above formula will do 16 standard rolls
without reuse. As a matter of interest, twice the above recipe can be squeezed into one liter.
By the way, don't be confused by the fact that there exist 4 mol and 8 mol
metaborates. A gram of one has the same number of atoms of sodium, boron,
oxygen and hydrogen as a gram of the other. At 53.6 C, without losing any
water, Na2B2O4.8H2O becomes NaBO2.4H2O which is stable to 105 C. In other words, the distinction between 4 mol and 8 mol sodium metaborate is academic, not practical when we specify solution strength in terms of weight per unit volume. You may see this for yourself at www.borax.com.
I am attaching a scan of a 10x print from FP4+ developed 8 minutes, 68 F. This is not of very high resolution, but may serve to illustrate gradations. I have also attached a higher resolution scan of a small part of the same print. If you print the high res. scan to 6x9 inches, you will see about a 50X magnification of that part. I have no idea how these will show up, but you can imagine that the originals are better.
As it turned out, the attachments showed up at the end of the thread on superadditivity started by PE. My fault.
Last edited by gainer; 06-03-2007 at 11:23 AM. Click to view previous post history.
Just a couple of quick comments on volume vs. mass measurements: It is well known among ammunition reloaders that the precision and accuracy of volume measurements (with respect to delivering a specified mass) can be quite sufficient for routine measurement of powder charges for reloading of ammunition. Interestingly, the precision required for powder measurements for reloading of ammunition are far more demanding than those required to make film developers, or at least that is my assertion. (The reason for this assertion is that chamber pressures are a very non-linear function of the powder charge, and both accuracy of the load and the safety of the load are strongly dependent on chamber pressures.)
It is well known that a volume measurement of a powder or granular material may give different masses if the physical properties of two different samples of material are different. However, as long as the powder or granular material is fairly uniform within a given sample and from batch to batch the uniformity of the measurement can be quite good. Many household, pharmaceutical or nutraceutical materials are quite uniform from batch to batch, and can be expected to give quite uniform results (even on a batch to batch basis) when measured by volume, as long as no serious changes take place during transport and storage, such as clumping or serious settling. Therefore, once these are calibrated by someone with a scale then someone else with only volume measurement devices (teaspoons, etc.) do a pretty good job of accurate and precise measuring. In some cases (e.g. settling) one can even manipulate a material (e.g. sift of fluff it up) to bring it back to a uniform and reproducible condition, as all home bakers know.
One person in the photo community (I forgot who) has pointed out that volume measurements may even be superior to mass measurements in certain cases. In particular, if a compound absorbs water without a large change in volume then mass measurements will often give a poorer result than volume measurements with regard to measuring the amount of active material in a sample. This is particularly the case for a hygroscopic material that has had significant exposure to atmospheric humidity.
Most scientists or engineers will agree that for a well characterized material a mass measurement using an accurate scale will almost always be more accurate and precise than a volume measurement, even using high quality volumetric devices. However, it is often the case that volumetric devices are "fit for purpose" (i.e. good enough) and are almost always more convenient than mass measurements. This is particularly the case for applications that do not require the utmost precision and accuracy. Most photographic applications probably fit into this category. The trick, of course, is to only use volume measurements on materials that are well suited to this type of measurement. Most uniform powders, granular, or liquid materials would fall into this category, but this certainly does not apply to all materials.
A final note: just in case someone is concerned about whether I am qualified to have an opinion on this matter, I am a professional chemist. I make my living by overseeing research and development of analytical methods and by overseeing some aspects of a routine testing laboratory.
A quick question for my own information: Is the cation critical mainly because of its effect on complexation of silver ions? In particular, ammonia and amines can help solublize silver compounds (e.g. silver halides) by forming so called "complex ions" Neither sodium nor potassium can do this, but ammonium ions (in the form of the non-ionic form ammonia, which exists in equilibrium with the ammonium ion in water solutions) readily forms silver complexes. Would this be the reason that the ammonium form of hypo is a more rapid fixer than the sodium, and why the addition of certain amino-compounds act as grain reducers in developers, e.g. by the mechanism of silver halide solublization?
Originally Posted by Photo Engineer
If sodium and potassium act as "spectator ions" in the fixing process then I would expect the substitution of one for the other to have a relatively small effect. This is not to say that sodium and potassium salts would behave identically, but only that the differences would not be extreme.
Interestingly, high concentrations of chloride can also form complexes with silver ions and solublize silver compounds, which is probably the chemistry behind a discussion I recently read about using salt water as a weak fixing agent.
In answer to your first comment Alan, as you know, powder is classified into types and those types require specific settings on the powder dispenser that is commensurate with the powder size. Also, powder is made in uniform pellets or flakes (photo chemicals are not). Powders are also generally of the same or nearly the same density. Therefore, in the case of powder, you have what is effectively a calibrated variable sized spoon. Also, when rotating the dispenser, you 'cut' individual powder particles to make an even 'spoonfull'. Therefore, the argument becomes unimportant in this case as you cannot (or do not) cut up individual crystals of photographic chemicals to insure a uniform volumetric measurment, but you do in reloading. These subtle details are often missed.
BTW, I reload .223, .243, .45, .9mm and 12 guage. I encounter this quite a bit and have had a lot of time to consider the problem. I use a variety of powders and have verified the accuracy due mainly to uniformity and grain cutting! This differentiates it from the photographic use we are discussing here.
As to your second post, the cation has several functions in increasing fixation or retarding it. One is size which relates to diffusion. Potassium ion is larger than Sodium is larger than Lithium, but Lithium is about the same size as Ammonium. Therefore, where or how does the fact that Ammonium based fixers are faster than Lithium fixers come in?
Well, Ammonium ion is both a cation and a ligand and therefore the size of the Silver Ammonium Thiosulfate complex is reduced considerably by the fact that the ammonium ion becomes part of the complexation cage with the silver ion. It is not an external cation. This tends to speed the rate of both soluabilization and diffusion. The rate of dissolution of silver halide is due to the effective syngergy between ammonium ion and thiosulfate ion.
So, as a matter of actual practice, it is a fact that there is a syngergistic effect going on here. This was not fully understood until work in the 60s and 70s which I took part in. Mixtures of silver halide solvents become synergistic and speed fixation.
See my Super Fix formula posted here for an example. Jay DeFehr posted it from the thread on Photo Net. The entire thing played out there.
Hope this helps.
I should add as a separate footnote that the error in the data in Mees compared to the table in Mees caused some of the confusion among workers early on. The text is correct, the table has several errors which imply that Ammonium Thiosulfate is slower than some other fixers.
This was corrected in James, not by doing work to extend the knowledge, but rather by removing the table which was in error. Our work was ongoing when Howard James was rewriting M&J and Grant Haist was aware of much of the internal and external work. It was decided to let the Haist statement with external juried references stand as the evidence and to correct the situation regarding cations in fixers.
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More on reloading that I almost forgot.
When you use a powder dispenser, it has a micrometer adjust, and the good reloader checks the charge on a balance while adjusting the micrometer to insure he/she is on the money so to speak. So, in effect, you have a micrometrically adjustable 'teaspoon' which you are expected to calibrate by weight for verification, and which performs optimally by slicing off excess chunks of powder to get an even level amount with each turn of the dispenser.
In addition to this, the manufacturer publishes sets of weight/volume tables in rather largish books.
And finally, the user (reloader) is expected to test the load by firing groups of loads for test purposes.
I'm glad this was brought up, because it highlights the fact that only when you have a uniform material of known density and solid form can you be sure of getting the right amount of powder by volume.
At Kodak, we dispensed solids by weight using strain guages under the container. These guages are highly sensitive to weight and are set to sense the exact weight less the tare value of the container. We do not use volume for solids.
In many situations, we even use weight for liquids.
As noted above, volume changes with the 'fluffing' of a powder, but also varies with temperature in some cases or with humidity. A caked solid has more packing density than a fluffy quantity of the same chemical. It therefore delivers more weight / volume.
I thought I could stay out of this fray, but I have to say this: I did not ever publish a formula for fixer, and if I had it would not have contained borax. You did not respond to my questioning the probability of significant potassium content in technical grade borax. The refinement of raw borax ore by the producers of 20 Mule Team products involves, as you might expect, solution in hot water and crystallization. Although there may be potassium remaining as a soluble impurity among the 1300 ppm maximum allowed in technical grade, it is difficult to see how that fraction of the small amount of borax used in any of the formulae I have seen could have a drastic effect on fixation, except in the psycholigical sense. Even so, it can easily be reduced by a factor of 50 or so in the home laboratory simply by making a saturated solution of the technical grade with at least double the amount needed for saturation, discarding the clear liquid, and making another saturated solution with the sediment from the first. If the solutions are saturated at high temperature, the last solution upon cooling will produce crystals of very pure borax decahydrate. How much did it cost to do this with borax that costs 50 cents a pound at the local supermarket? These are procedures that require little attention.
Flashpoint is the temperature at which an open flame close to the surface of a liquid will ignite the liquid. I tested that with the propylene glycol. It ignited as predicted, but there was no explosion, an open flame was required for ignition, and the burning glycol was easily extinguished.
A fire is achieved when a flammable vapor reaches a hot element that is above the flashpoint of the vapor.
In this case, you heat the liquid until vapor begins to fill the container. The heavy vapor will rise over the edges of the container and drop downwards due to the density compared with the atmosphere, and will ignite when it hits a surface heated at or above the flashpoint.
This can take place over distances of 6 feet or more. I have seen it, and it is much more difficult to contain and extinguish. The sudden ignition of even a moderate amount of vapor can be nearly explosive in force.
Very often, this is seen with gasoline or solvent based paints which can ignite when even opened in a shop with a lit pilot lamp on a water heater or furnace over 6 feet away. I was a consultant for an attorney on such a case many years ago.
In many cases the ratio of air to vapor determines whether you get a fire or an explosion.
IIRC, TF-3 uses Borax for buffer. I don't have my copy of A&T handy, but if so then the Borax would have to be essentially free of Borax. Common salts vary in the Potassium and Sodium ratio around the world so it could depend on the source.
Found A&T and looked up some formulas...
TF2 and TF3 contain Borax. There is Kodalk in Kodak's F6, F10 and F24. F5 contains Boric acid. All, being photo grade are free of Potassium as per the information from Haist.
Where did you get that definition of flashpoint? IIRC the flashpoint of propylene glycol is about the temperature of boiling water. It would be very difficult to use it as an automobile antifreeze would it not?
Originally Posted by Photo Engineer
Wikipedia:The flashpoint of a flammable liquid is the lowest temperature at which it can form an ignitable mixture in air. At this temperature the vapor may cease to burn when the source of ignition is removed. A slightly higher temperature, the fire point, is the temperature at which the vapor continues to burn after being ignited. Neither of these parameters are related to the temperatures of the ignition source or of the burning liquid, which are much higher.
Thhe flash point of propylene glycol is 210 F by the closed cup method and 224 F by the open cup method. You can hold a rod heated to 224 F near the surface of propylene glycol all day long without igniting it. The autoignition point is 700 F. If the glycol were to boil over onto a red hot heating element, there would certainly be a fire. If the vapor from glycol at its flashpoint overflowed the container, there might or might not be a fire, depending on the partial pressure of glycol at the hot surface, which would have to be at least at the autoignition point of the glycol.