Borax Project

Murray Kelly said:

And presumably glycerol doesn't affect the buffering ability of borax like glycol?
At 1:50 my scratchings on an envelope indicate there's 5g Borax - 0.4g Phenidone - 0.4g AA per litre (working). My maths has been known to err in the past but is there a decimal place wrong in there?
Long ago I made up some PC-glycerol and it appears to keep just fine. I had no glycol.
Murray

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I don't know. If you made a solution in water with 5g Borax - 0.4g Phenidone - 0.4g AA I think you would find the usual 9.5 or so borax pH. This working solution is much lower, and the only reason I can see for it is the glycerol. I'm working on a second batch with much less heating. I'll let you know.

gainer said:

Here is an interesting fact. 526 grams of borax decahydrate will dissolve in a liter of glycerol. I dissolved 120 grams in 500 ml, added 10 grams each of Phenidone and ascorbic acid, used 1+50 dilution of this solution to develop Arista Premium 400 for 16 minutes at 70F with excellent results. I think this might be a good stand developer. The pH of the working solution was about 8.5. The borax brings with it the water of crystallization, but at most that would amount to about 56 grams of water. If you have time for such frivolities, it's fun to play with. Even with heat it seems it will not all dissolve, but give it some time.

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Is it a stable solution?..Evan Clarke

eclarke said:

Is it a stable solution?..Evan Clarke

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Seems to be quite stable once it all dissolves. The first batch has been sitting at room temp for a few days with no settling or discoloration. Even diluted 1+25 it's quite slow. I'm about to go to my chamber of horrors to try some more stand development.

I have learned that there is a glycerol borate. This is probably (I'm guesing, of course) what causes the sudden apparent change in solubility of borax in glycol as temperature reaches a certain point. The glycerol is no longer just a solvent. Is that good or bad or indifferent? Its formula is (C3H5BO3)n.

I am enjoying this and will try some of this but my interest is that I have a couple friends who are just dabbling in homebrew and want to make D76 but don't have good stirring capability. A water solution probably isn't consistent enough for them but this might be the ticket. I could just make them each a half liter of glycerol/borax. What's the pH of the glycerol/borax?..Thanks..Evan Clarke

If you can stir a teaspoon of sugar into a cup of tea (the kind you drink) you have all the stirring capability you need to mke D-76. Now for the chemicals, assuming you have a scale or balance:

add 2 grams of Metol to a cup of water. It will dissolve right away. In a separate container, mix 5 grams of hydroquinone with 25 grams of sodium sulfite in a cup of water. Add this solution to the first in a container large nough to hold a liter. Bring the volume to 750 ml with water. Add 75 grams of sodium sulfite and 2 grams of borax and stir with a stainless steel or plastic spoon until it dissolves. Add water to make 1 liter. Voila! D-76 home brewed!

Maybe that will help them, I have good heated stirring equipment and just mix it in order and have no problems and like some other developers better. I will pass this routine on to them, thanks..Evan Clarke

Stand developed 30 minutes.

The attached photo is a scan from an 11 power photo print of a negative stand developed for 30 minutes in 1+25 dilution.

gainer said:

Here is an interesting fact. 526 grams of borax decahydrate will dissolve in a liter of glycerol.

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I think if you imagine glycerol as a chain of 3 carbon atoms linked together, with each carbon atom having one hydroxide group connected them you can see why.

The hydroxyl groups are hydrophyllic - that's what accounts for the solubility in water of glycerine and for its hygroscopic nature. I suspect it's also what's responsible for its ability to dissolve numerous inorganic salts. It's kind of like having a molecule that's half water, and half organics.

gainer said:

It turns out that the high solubility of borax in glycol is accompanied by formation of boronate esters that will be completely useless as bases. This is most likely to happen when heat is used to hasten solution.

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My understanding is that you will not make esters of borax this way, it merely makes a glycerol-borax complex. To make an ester, you'll need to drive off the water to push the reaction along. That means lots of heat - pretty much to the point of boiling. About 240°-280° C - now that's hot! And then at that point, you've made an glycerol boriborate ester, which will be useless for your purposes.

gainer said:

I have learned that there is a glycerol borate. This is probably (I'm guesing, of course) what causes the sudden apparent change in solubility of borax in glycol as temperature reaches a certain point. The glycerol is no longer just a solvent. Is that good or bad or indifferent? Its formula is (C3H5BO3)n.

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OK - to make esters, you usually have to have an acid (often an organic acid), an alcohol, and a catalyst. For example, the esterification of acetic acid in excess ethanol (which also will act as the solvent for the reaction) will take place in the presence of concentrated sulfuric acid (the catalyst) and heat. This results in an ester, ethyl acetate.

gainer said:

It explains why I thought my pH meter had gone bonkers.

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The reason you are getting lower than you expected pH values for borax dissolved in glycerol is that you are releasing boric acid when the glycerol-borax complex is formed.

Glycerol forms a complex with the borax, forming monomeric or dimeric complexes with the glycerol. In borax, the sodium atoms are surrounded by borons atoms connected to other boron atoms with hydroxides bonded to the borons as well. The glycerol complex liberates som eof those boron atoms to form boric acid as it complexes. This free boric acid is what it lowering the pH of your solutions.

In fact, I once used this reaction as the basis to determine the amount of boron (I was looking for borax) in an industrial product that I was hired to deconstruct - I think there was a bit of industrial espionage going on with that project... Anyway, you can titrate a solution of borax (which is basic) with sodium hydroxide (also basic), using pH 7 as the endpoint.

Think about that for a second - you have a solution of borax with a pH of 10 or so, and you titrate it with sodium hydroxide with a pH of 13 or so - how do you get an endpoint of pH 7? (Isn't chemistry fun!?)

The simple way is to add an excess of mannitol (a sugar alcohol with a structure rather similar to glycerol) or glycerol to react with the borax. The borax releases boric acid as a result of the excess glycerol or mannitol. (Neither glycerol or mannitol have much effect on the pH of the solution.) You add some methyl orange or phenolphthalein indicators and titrate to the endpoint of the indicator. The amount of sodium hydroxide used in the titration is proportional to the amount of boric acid liberated in the complexing reaction. You can then calculate how much borax was present in the original sample.

I can list the chemical reaction equation if you like.

So I'd say by dissolving borax into glycerine, you are liberating boric acid and consuming your borax. Probably not what you are really trying to accomplish by using the glycerol as a "solvent".

gainer said:

The onset of my interest in purifying borax was instigated by the assertion, made by PE and Kirk Keyes, that the 20 Mule Team borax from the supermarket is not pure enough for photo work, even though many of us have been doing so for years.

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To set the record straight, I suggested that 20 Mule Team may contain impurities that would interfere with photgraphic uses, and that photo-grade borax gives one some assurance that your reagents are suitable for photographic uses.

Kirk Keyes said:

To set the record straight, I suggested that 20 Mule Team may contain impurities that would interfere with photgraphic uses, and that photo-grade borax gives one some assurance that your reagents are suitable for photographic uses.

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I have refrained from posting in this thread until now. I feel that I have to clarify my position as well.

Kirk not only has his chemistry straight, he has also stated in this quoted post my own thoughts on this matter.

The most common impurities are halide salts and insoluable colloidal salts. Patrick has alluded to the latter in the past saying that some of his Borax solutions are cloudy and the cloudy suspension cannot be filtered out.

This whole thing may be summarized by saying that Patrick is trying to prove his point, that he can compound usable developers. That is conceded, but due to lack of full checks, we don't know how they compare in detail to off the shelf developers. In fact, in some test comparisons shown in other threads, I (and others) have given the tip of the hat to the commercial (albeit modified) developer.

I am not a borax chemist, but in private notes, I have exchanged similar, but less specific informaiton with Patrick discoraging him from the view that an ester or other type of compound is being formed. This was obvious to me, but the details were unclear and I had no text references for help. I thank Kirk for giving us the complete story on this.

PE

interesting! up until a couple of months ago when another member very generously gave me about twenty pounds of old photo chems, I had always used 20 Mule Team Borax from the grocery store in my home brew ....

That stash of chemistry included a one pound jar of Kodak Borax - which I have been using ever since.

I too noticed the cloudiness with the grocery store Borax...and now that it is mentioned, I have not noticed it with the Kodak Borax.


Very interesting. I think I'll leave the 20 Mule team in the laundry room from now on.

I'm no "borax chemist" myself, but as an analytical chemist, you can be exposed to a lot of different analytical techniques that are sometimes forgotten in the modern world of chemistry. Nowadays, to measure boron content, an ICP (Inductively Coupled Plasma atomic absorbtion spectrophotometer) or even an ICP-MS (one with a mass selective spectrometer attached) would be used. 50 to 20 years ago, an AA (atomic absorption spectrophotometer) would have been used. The titration with glycerine was used from about 100 years ago until labs could afford AA's in the 50s and 60s.

In point of fact, the borax that I used in these later experimental developers was the photo grade from Photographers' Formulary. My main interest at this time has shifted from borax to glycerol. I had been using it as a solvent at time when I had no propylene glycol at hand. At room temperature, it seems to serve as well except for the higher viscosity for the preparation of higly concentrated stocks like PC-Glycol. It occurred to me that the high solubility of borax in gylcerol would allow a single solution concentrate of Phenidone, ascorbic acid and borax in glycerol. I am pretty sure that in the process of heating the mixture for faster solution to between 250 and 300 F I did form the glycerol borate. The pH of the working solution was considerably lower than I would have gotten from a water only solution of the same small amounts of Phenidone, ascorbic acid and borax.

The 1+25 working solution of the glycerol stock seems to have buffering properties, as the addition of fairly large amounts of borax produces little change of pH, which stays below that of a 0.1N borax solution. The resulting developer is quite good, but slow.

There is also a glycerol sulfate, which gave me the idea test the solubility of Metol in glycerol. I have dissolved, with heating, 10 grams each of Metol and ascorbic acid in 180 ml of drugstore glycerol, which is nearly anhydrous and of medicinal quality.

Photo Engineer said:

I have refrained from posting in this thread until now. I feel that I have to clarify my position as well.

Kirk not only has his chemistry straight, he has also stated in this quoted post my own thoughts on this matter.

The most common impurities are halide salts and insoluable colloidal salts. Patrick has alluded to the latter in the past saying that some of his Borax solutions are cloudy and the cloudy suspension cannot be filtered out.


PE

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The very first thing I did was to prepare a nearly saturated solution of borax in very hot water. After cooling, most of the borax had settled out, leaving only about 47 grams/liter in solution along with most of the soluble and colloidally suspended impurities. Decanting the solution thus removed most of the soluble impurities. There is no point in trying to filter out any impurities, be they truly dissolved or colloidally suspended, if they can be decanted at the cost of only a few percent of the original amount of borax. Furthermore, even the best grade of borax is likely to be a mixture of pentahydrate and decahydrate, so in any case where an accurate assay of Na2B4O7 is required, the standard procedure is to prepare a saturated solution. This solution, if kept at a temperature above that of its formation, will have a known borax content expressed as weight percent of the decahydrate. Thus, a saturated solution formed at 20 C will contain 4.71% sodium tetraborate decahydrate, or 47.1 grams per liter.

I think you will see if you look at my developers, that I have for many years spent most of my effort designing sulfite-free developers. The major thing I have against sulfite is cost and local avilability. It is often the most expensive component, and among the hardest to get. Anyway, it's part of my fun to see what I can do with what I can get in the middle of Wild and Wonderful West Virginia.

gainer said:

The major thing I have against sulfite is cost and local avilability. It is often the most expensive component, and among the hardest to get.

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It's easy to get it at a photo store. At least the photo shops around here.

gainer said:

Thus, a saturated solution formed at 20 C will contain 4.71% sodium tetraborate decahydrate, or 47.1 grams per liter.

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Pat - do you make corrections for when your borax solution is not at exactly 20C?

I hope you have an accurate and precise thermometer. If you are off by 1 degree, you will get a 4.52% solution if you are low, and 4.94% if you are high. That's an error of 5% for just a change of 1 degree Celcius.

I have to say I think just wieghing your borax is a lot simpler and more expedient. And more accurate.

And that's assuming you have time to let your borax solution sit around so that it's truely saturated - not undersaturated and not over saturated.

Kirk, if the saturated solution has been decanted from the undissolved sediment at 20 C, it will contain 47.1 grams of borax per kilogram of solution at any higher temperature. For practical purposes, even if you measure by volume at a higher temperature, the change due to expansion of the liquid will not throw you off as much as the possible variation in weight per mol of Na2B4O7 due uncertainty of the average amount of water of crystallization. How do you ascertain that number? Every time you open the container, the contents are subjected to possible changes in humidity. The 5 pounds I bought from the Formulary has huge lumps, small lumps and powder, which indicates changes in water of crystallization have occurred, but from what to what?

If you are going to be a stickler for precise measurement of borax, your best bet is a saturated solution. I think the best way to make it is to start from saturation at a higher temperature and to let it cool to the desired temperature. If it takes a while, so be it. as the saying goes, "Think ahead."

I did not continue my formal study of chemical engineering, but changed to aeronautical. I am an engineer by nature, as anyone who knew me during my years at NACA-NASA would testify. As one of our instructors in engineering was fond of saying, "Engineers can do anything."

I think that phrase should be "Engineers will do anything". This does not give it the stamp of approval however, nor does it ensure that their chemistry is correct as Kirk has pointed out.

Imagine building the port on a Mercury capsule (an example given by Patrick in another thread), and it has a 5% variation in size. One would leak (too small), one would crack (too big), and one would be just right! I think that this analogy is apt although the results would be far less draconian in photography. If you had a 5% variation in alkali content, your pictures might be usable but you might wonder why your contrast or sharpness varied as much as it did.

PE

And what makes you think I might have had a 5% variation in alkali content? I told you I used the photo grade borax I got from PF. Is it that bad? Even the stuff I get by processing the household variety is better than 5%.

By the time the Mercury became a project, I had been assigned to do basic human factors research, and to design star charts to be used as backup reentry alignment. What did my education in AE have to do with that? It made me ask questions and search for answers. I no longer have the technical library I could access at Langley Research Center.

Kirk, there is no photo shop within 100 miles of here in any direction. My formulas are not only for my use, but have been used in remote locations in New Zealand. (Jacko Twist, where are you?)

Patrick, the 5% figure came from Kirk's post and suggested the limits of a 1% error in temperature either high or low. That is where it came from, nowhere else. In my experience the Borax from the Formulary is just fine.

As for all of this work, it seems to me that there are two considerations...

1. How much does this purification cost in terms of time and energy? Is it cost effective? It actually is not apparently cost effective as was described in an earlier thread I believe. I did not make the post nor have I tried to locate it or do the estimation myself. The very proposition of purifying my chemistry before I use it is alien to my thinking and work flow.

2. Should I even worry about borax? It is one of the chemicals that Kodak worked very hard to eliminate in the 60s and 70s and resulted in a whole new family of developers which used carbonate, TEA, or very low levels of borax to achieve superior results. The reason? Borax is toxic to some plants, some bacteria in sewage systems, some insects, and small children. Kodak wanted to become greener. So, I seldom use Borax.

PE

gainer said:

I am pretty sure that in the process of heating the mixture for faster solution to between 250 and 300 F I did form the glycerol borate.

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You may have made the ester that way.

But that's super effing hot - I wouldn't want to have to heat anything up that hot just to dissolve borax in something!

Borax is soluable at the rate of 1 kg / kg of water at 20 deg C according to my handbook. That is pretty darn high and should not need much heat at all to make a solution. I don't intend to try. See my last post. Or, consider what Kirk said. That is HOT.

PE

Photo Engineer said:

Borax is soluable at the rate of 1 kg / kg of water at 20 deg C according to my handbook. That is pretty darn high and should not need much heat at all to make a solution. I don't intend to try. See my last post. Or, consider what Kirk said. That is HOT.

PE

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Water to make tea should be nearly that hot. Are you sure of that number? The chart from www.borax.com says 4.71% at 20 C. We are at cross purposes here. I heated borax-water to get a saturated solution as it cooled. I heated borax-glycerol to make it dissolve faster and ran into a response that appeared to be a chemical reaction at around 250 F. Of course that is hot, but I think a roast coming out of the oven is that hot or hotter. A roast should be seared in a salted cast iron pan before being put in the oven. You want hot? Be a cook for a while.

Patrick;

Water boils at 212 F.

The Merck index lists borates and boric acid being quite soluable in glycerol. It states that 1 gram of Sodium Borate is soluable in 1 ml of water at 20 deg C. That is what I quoted above regarding borates in glycerine which is what you asked me.

A temperature of 250 F is above the BP of water at normal atmospheric pressure. A roast is nominally at 350 deg F coming out of the oven and will burn the pan unless you have kept liquid in the bottom of some sort. Grilling temperatures for meat are above 400 F. This too must be moistened somehow to prevent excess burning of the meat. The internal temperatures either way should reach 160 F or higher depending on the degree of doneness you wish. Most organic chemists love to cook.

I don't discount the formation of an ester, but it would have to take place at higher temperatures, I think, or would require a catalyst. What you observe is entirely in line with Kirk's explanation and not with the formation of an ester.

PE