Weakening of the Camden tablets due to atmospheric exposure comes eventually through combination with water vapor and CO2, which are always present, at least when and where grapes are grown. A molecule of metabisulfite ( also called pyrosulfite) is converted to 2 molecules of acid sulfite (bisulfite) by 1 molecule of water. 2 molecules of bisulfite are converted by 1 molecule of CO2 to one molecule of H2CO2, a gas, and 1 molecule of sodium sulfite, which removes the acidity of the tablets as time passes under exposure to the atmosphere. It has nothing to do with whether it is meta- or bisulfite. The same fate awaits either.
SO2 in water is a solution of sulfurous acid. It forms when either bisulfite or metabisulfite is dissolved in water.
SO2 in water is a solution of sulfurous acid. It forms when either bisulfite or metabisulfite is dissolved in water.
The salts of Sulphurous acid H2SO3 are theoretically formed by the aqueous dissolution of gaseous SO2. However there is strong spectral evidence that very little of the free acid exists in solution, instead dissolved SO2 forms a clathrate in which the SO2 is trapped inside a water shell without chemical bond formation - in a manner analogous to other gas hydrates. (Cotton and Wilkinson,1966). This structure is usually represented as SO2.7H2O and although there is discussion about the exact form in which it exists.
When you start delving deeply into the chemistry you start to find that what appear at first to be simple equilibrium equations are quite complex. Certainly all references state quite clearly that both Sodium & Potassium Metabisulphite are the principle SO2 generating salts, and preferable to the Bisulphite and Sulphite, none say why.
BASF in their data-sheets for Photo-grade Sodium Metabisulphite give a figure of 66g/100g for potential SO2, while for BASF's 40% aqeous Sodium Bisulphite it's 23g/100g (57.5g/100g equivalent)
The basic equilibrium reaction between Metabibulphite - Bisulphite - SO2 & Water is what's most imortant. It's clear that in commercial applications that you start with Metabisulphite because of it immediate properties of liberating SO2 in solution.
It solution it dissolutes into Bisuplhite, in an alkaline solution the Bisulphite equilibrium is towards Sulphite and in an acidic solution towards SO2.
Certainly the food industry prefers adding Metabisuphite as a preservative in most cases, although Bisulphite is added to fruit juices which are naturally acidic.
Out of curiosity, and to change the subject a bit, what would happen to one if he or she drank a glass of developer, and why? What about fixer?
I have tried a drop or two of developer and it was actually not half bad. Fixer seems much worse, however. I've never tried that. Guess it's the animal instincts saying: "Enough is enough with the tasting of the chemicals!"
If you put the metabisulfite in water, it is not so easily attacked by atmospheric CO2. It still becomes the equivalent of a solution of bisulfite. Surely the equilibrium point can vary with time, temperature, but when it is possible for a gas to form, especially if the gas cannot be dissolved by the amount of water present, the equilibrium is shifted. The addition of one molecule of water to one molecule of metabisulfite will produce 2 molecules of bisulfite. The interaction of 1 molecule of CO2 with 2 molecules of sodium bisulfite will produce 1 molecule of sodium sulfite and 1 molecule of H2CO3 which is not very soluble in water, but there is no water anyway except that in the atmosphere, so the Na2SO3 is left on the surface of the Camden tablet. Whether it then protects the inner layers from being attacked, I do not know. That is a problem in physical chemistry I think, that could be solved by experiment.
If you keep the metabisulfite dry, you need not worry about CO2. OTOH, if you use the bisulfite, it can be attacked directly by CO2 in a perfectly dry atmosphere, leaving sodium sulfite. So it would seem that Camden tablets should be kept in a closed container with silica gel dessicatant. You might still lose SO2, leaving sodium sulfite, but you should be able to tell by the smell. I have a friend nearby who grows grapes and makes fine wines, who uses sulfur candles to sterilize bottles, but only because required by law IIRC.
These are IMHO possibilities. If there is a way to show that they do not occur, let us know.
Haist says that metabisulphite and sulphite both have good preservative quantities. IIRC - and that's a matter of debate - it takes about 80% of the meta to equal the, um, naked sulphite for a given oxygen scavenging ability.
Of course, one results in an acidic solution, the other is alkaline. Setting aside cost differential, use what you want for pH. The reality, of course, is that meta costs many, many times the plain sulphite. The meta is often used to reduce the pH of an otherwise sulfite only mixture.
Haist says that metabisulphite and sulphite both have good preservative quantities. IIRC - and that's a matter of debate - it takes about 80% of the meta to equal the, um, naked sulphite for a given oxygen scavenging ability.
Of course, one results in an acidic solution, the other is alkaline. Setting aside cost differential, use what you want for pH. The reality, of course, is that meta costs many, many times the plain sulphite. The meta is often used to reduce the pH of an otherwise sulfite only mixture.
The argument here was whether one could use sodium bisulfite (acid sulfite) when a recipe called for sodium metabisulfate (pyrosulfite). 1 molecular weight of sodium pyrosulfite + 1 moleculer weight of water is equivalent to 2 moleculer weights of sodium acid sulfite. If you want to make a solution in water equivalent to 10% bisulfite by weight, then 100 grams of sodium bisulfite or 91.35 grams of the metabisulfite in water to make 1 liter of solution will do the job. I contend there is no chemical test to distinguish one solution from the other if they are made with equally pure ingredients.
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