(Official Hive Translator)
Two OTC prep's of dithionite salts.
(Rated as: excellent)
Here's a nice little collection of proc's related to the manufacture of dithionites, used as reducing agents. It's always been my impression that these agents are undeservedly overlooked here at the Hive - well, who knows maybee one day their time will come (e.g., for selectively reducing a-oximino-aceto/propiophenones to cathinones - and who knows what else).
I know that Na2S2O4 is fairly easy to get OTC for many bees - OTOH, in some parts of the world (like mine) its use is strictly limited to chem labs.
1. Preparation of zinc dithionite - as submitted to Hyperlab by Chemister.
Into an Erlenmeier flask there's placed 5-10g Zn powder, 50 mls abs. EtOH; the flask is stoppered and thru a tube extending almost to the bottom, SO2, dried w/conc H2SO4, is bubbled into the rxn. The flask is also equipped w/an outlet tube, the other end of which is immersed into mercury.
The rxn is over when all the Zn is gone. The crystalline precipitate is filtered and dried in an exicator (preferrably in vacuo) over H2SO4.
As dry crystals, ZnS2O4 is relatively stable to air. Its aqueous solution is a very strong reducing agent.
(A question for everyone - does anyone know if there is a substantial difference in 'reducing power' between Na and Zn dithionite? Any experimental examples using the latter?)
2. Preparation of sodium dithionite from Na formate and SO2 - from Patent US3947559.
Note - in the 'classical' method the reaction is carried under superatmospheric pressures - albeit not very high, like 1,5-3 atm, yet very inconvenient for the kitchen bees. This one describes such a prep'n under atmospheric conditions - the main point of the patent is usage of methoxyethanol as solvt, but it appears that decent results are obtained using EtOH and esp. EtOH/Et(OH)2 system.
a) With EtOH as solvent.
A slurry of 160g (2.35 moles) sodium formate, 240g ethanol (74 DEGo.p. industrial methylated spirit) and 120g water was charged into the reaction vessel and heated up to 75 C with stirring under reflux in an inert atmosphere (N2). To the stirred slurry was added a solution of 60g (1.5 moles) sodium hydroxide in 70g water concurrently with a solution of 200g (3.13 moles) sulphur dioxide in 610g ethanol. The sulphur dioxide solution was added at a uniform rate over 1 hour and the alkali at such a rate that the first 30% was added 21/2 times as fast as the remaining 70%, the total time of addition being 1 hour. The temperature was kept at 75 C at all times. A sticky solid was formed initially but this quickly dispersed to give a normal-looking white solid. The mixture was stirred at 75 C for 2,5 hours after addition was complete and then filtered off under an inert atmosphere, washed with 200g methanol at a temperature above 60 C and finally dried. The white solid obtained weighed 211.9g and contained 70.7% sodium dithionite, Na2 S2 O4. The yield as Na2 S2 O4 was 58.6% calculated on the sulphur dioxide consumed.
b) With EtOH/ethyleneglycol as solvent.
A slurry of 160g (2.35 moles) sodium formate, 180g ethanol (74 DEGo.p. industrial methylated spirit) 60g ethylene glycol and 120g water was charged into the reaction vessel and heated up to 80 DEGC with stirring under reflux in an inert atmosphere (N2). To the stirred slurry was added a solution of 60g (1.5 moles) sodium hydroxide in 70g water concurrently with a solution of 200g (3.13 moles) sulphur dioxide in a mixture of 457g ethanol and 153g ethylene glycol. The sulphur dioxide solution was added at a uniform rate over 1 hour and the alkali at such a rate that the first 30% was added 21/2 times as fast as the remaining 70%, the total time of addition being 1 hour. When addition was complete the product was stirred at 80 DEGC for 21/2 hours, filtered off under an inert atmosphere, washed with 200g methanol at a temperature above 60 DEGC and then dried. The white solid obtained weighed 212.7g and contained 82.3% sodium dithionite, Na2 S2 O4. The yield as Na2 S2 O4 was 67.0% calculated on sulphur dioxide consumed.
|Re: (A question for everyone - does anyone...||Bookmark|
(Rated as: excellent)
Here's the procedure from Inorganic Syntheses
Anhydrous sodium sulfite and pyrosulfite, free from sulfate, may be prepared by the method of Foerster. Solutions of the salts are readily oxidized by air. Therefore, the product must not be exposed to air until it is thoroughly dried if a high degree of purity is to be attained. Moist crystals of the pyrosulfite decompose slowly even in the absence of air to form the sulfate and sulfur. An odor of sulfur dioxide over this product will indicate that drying was not complete. Both salts may be kept in a desiccator in a hydrogen atmosphere for several months without showing a test for more than a trace of sulfate.
a. Anhydrous Sodium Sulfite. The preparation is best carried out in a widemouthed Erlenmeyer flask arranged as shown. The sulfur dioxide inlet tube is a glass T through one arm of which a small rod is inserted to push off the crystals, which soon obstruct the flow of gas.
One hundred seventy-five grams of sodium hydroxide is dissolved in 500 ml. of freshly boiled distilled water contained in the flask. A slow flow of hydrogen is first started, and sulfur dioxide is then added rapidly. The reaction is exothermic, and the temperature soon reaches the boiling point, where it remains until the hydroxide is converted to the sulfite. At this point the solution turns slightly yellow, and the temperature begins to fall. The flow of sulfur dioxide is now stopped. The sodium sulfite crystals are filtered from the hot solution through a sintered-glass funnel placed in a vacuum desiccator as shown.
If a funnel with sintered-disk is not available, a Bizchner funnel may be equipped with a porous Filtros plate* sealed in with Insalute cement. t Filter paper cannot be used because of the swelling action of the hot solution on the fibers. Gentle suction may be applied to withdraw the liquid, but care must be taken to allow no air to pass through the crystals. When all the mixture is in the funnel, the top of the desiccator is replaced, the stopcock on the vacuum line to the funnel is closed, and the desiccator is flushed out with hydrogen several times. The filtration is then completed by applying vacuum to the funnel. A slow stream of
dry hydrogen is allowed to pass into the desiccator and through the funnel until the crystals are thoroughly dry. The desiccator should not be opened until this condition is reached. This operation requires 24 hours, or more, depending upon the thickness of the layer of crystals in the funnel and the rate of flow of the hydrogen. The product should be kept in a tightly stoppered hydrogen-filled bottle. Yield approximately 225 g. Anal. Calcd. for Na2SO3: Na20, 49.2; S02, 50.8. Found Na20, 49.0; S02, 50.2; sulfates, none. If the solution is cooled to a temperature below 33.4° before filtering, the hydrate Na2SO3.7H2O is obtained in pure form.
b. Sodium Pyrosulfite. The pyrosulfite is prepared in the same apparatus. Two hundred twenty-five grams of sodium hydroxide is dissolved in 500 ml. of boiled distilled water. Hydrogen is passed through the solution to exclude air, and sulfur dioxide is added until the sulfite, which is first formed, is completely dissolved. Crystallization of the pyrosulfite is quite slow, and the flask must be shaken continuously to avoid caking. The crystals are filtered from the cooled solution in the apparatus described above. Yield about 200 g. Anal. Calcd. for Na2S205: Na20, 32.6; S02, 67.4. Found: Na20, 32.9; S02, 66.6; sulfates, none.
Sodium hydrogen sulfite, NaHS03, does not exist as a solid. Foerster found evidence based on freezing-point lowering that in solution no pyrosulfite ions are present, only hydrogen sulfite ions
S205 + H20--+ 2HSO3
The solubility diagram shows that the anhydrous sulfite decreases in solubility as the temperature is raised. It is less soluble than the pyrosulfite; consequently, it separates from solutions in which the ratio of S02 to Na20 varies considerably from that of the pure salt. Crystals of the 7-hydrate are less suitable for laboratory work than the anhydrous sulfite because of their high vapor pressure of water. This causes efflorescence and rapid oxidation in air. Solutions in equilibrium with the pyrosulfite have high vapor pressures of S02 at temperatures above 70°.3 Evaporation of these solutions to increase the yield of crystals results in the loss of S02 unless a stream of this gas is passed through the solution. If the ratio of S02 to Na20 falls much below 1.8, coprecipitation of the normal sulfite is apt to occur. The equilibrium vapor-pressure measurements show that the temperature coefficient of the vapor pressure of water over these solutions is about the same as that of sulfur dioxide. Little is to be gained, therefore, by evaporation under a vacuum.
Because of its high solubility and the small effect of temperature, as shown in Fig. 21, potassium sulfite is more difficult to prepare than sodium sulfite. Foerster and earlier workers found it necessary to proceed from a solution of the bisulfite. The pyrosulfite may be crystallized readily from water. Both salts oxidize rapidly in air when moist but are stable when dry.
a. Anhydrous Potassium Sulfite. One hundred grams of pure potassium hydroxide is dissolved in 200 ml. of freshly boiled distilled water (same apparatus as used for sodium sulfite). Sulfur dioxide is added in the presence of a stream of hydrogen until a sample of the solution is just acid to bromocresol green indicator. An equal quantity of potassium hydroxide dissolved in about 100 ml. of water is then added to the solution. Evaporation is carried out at the boiling point at atmospheric pressure in the presence of a slow stream of hydrogen. The crystals are filtered and handled in the manner described for sodium sulfite. Washing is undesirable because of the high solubility and, furthermore, is unnecessary when the crystals have the same composition as the dissolved salt. Yield about 200 g.
b. Potassium Pyrosulfite. The pyrosulfite is precipitated when a solution of 30 per cent potassium hydroxide is saturated with sulfur dioxide and cooled to room temperature. Filtration and drying in the absence of air are necessary to prevent oxidation.
Potassium sulfite forms no hydrates, but the pyrosulfite
forms the compound K2S205.23H2O, which is probably the double salt K2S205.4KHSO3. In solution, the pyrosulfite
is converted to the hydrogen sulfite.
Because of the small solubility of the pyrosulfite, this compound crystallizes from solutions containing as little as 1.2 mols SO2 per mol K20. Excess SO2, therefore, is to be avoided when the pure sulfite is being prepared. Like the sodium compounds, the hydrogen sulfite solutions have high vapor pressures of SO2 at the boiling point,3 but the loss of SO2 during evaporation can never reduce the ratio of SO2 to K20 to a point where the pyrosulfite will not crystallize in pure form. This indicates that the mother liquor from the pyrosulfite precipitation can be evaporated to increase the yield.
1. FORESTER, BROSCHE, and NORBERG-SCHULTZ: Z. physik. Chem., 110, 435 (1924).
2. "International Critical Tables," Vol. 4, p. 236, McGraw-Hill Book Company, Inc., New York, 1928.
3. JOHNSTONE, READ, and BLANKMEYER: Ind. Eng. Chem., 30, 101 (1938)
Chemistry is our Covalent Bond
(Rated as: excellent)
Dithionates may be prepared by the electrolytic or chemical oxidation of sulfurous acid and sulfites. Chemical oxidation is accomplished by the action of chlorine, iodine, hydrogen peroxide, or oxygen in an acid medium; by the action of chromates or permanganates in neutral solution or by the action of metallic ions, or metallic oxides, which are readily reduced. All these procedures give low yields except the method that involves the use of metallic oxides. Various oxides have been employed, but pyrolusite has been found to be most satisfactory, giving yields that vary from 65 per cent to 85 per cent* of the theoretical value. Specific directions are given for the preparation of the calcium, barium, and sodium salts, but the general method may be modified suitably for the production of other dithionates.
Mn02 + 2S02 --> MnS206
MnS206 + Ca(OH)2 --> CaS206 + Mn(OH)2
A 1 L, three-necked, round-bottomed flask is fitted with a thermometer, a mechanical agitator, and a sulfur dioxide
delivery tube, which reaches to the bottom of the flask. The flask is placed in an ice bath. Five hundred milliliters of water is placed in the flask and saturated with sulfur dioxide at a temperature below 10°. Pyrolusite is then added in 1 to 2g. portions until 80 g. has been introduced. The solution is kept saturated with sulfur dioxide during the process, and rapid agitation is maintained continuously. The temperature is not allowed to rise above 10°. Each small portion of added ore is allowed to react completely before more is added. This is evidenced by the change in color of the suspension from black to brown.
* Wide variations in yield may be attributed to such factors as the fineness of subdivision of the pyrolusite, the temperature maintained during reaction, and the concentration of the sulfurous acid. It should also be borne in mind that most commercial samples of pyrolusite are quite impure. Since only the manganese in the form of a manganese(IV) effects the conversion into dithionate, it is necessary to analyze the ore for its available oxygen content. The ore used both originally and in checking was a high-grade Java pyrolusite containing 16.5 per cent available oxygen, as compared with 18.4 per cent required for pure manganese(IV) oxide.
After all the ore has been added (2.5 to 3 hours), the agitation is continued until there is no further change in color of the slurry (light brown). To draw off excess sulfur dioxide, the slurry is placed under vacuum and warmed gently until a temperature of 40° is reached. The gelatinous residue is removed by filtration and washed with warm water.
The filtrate (at 25 to 40°) is stirred continuously while 80 g. of calcium hydroxide (slaked lime) is added over a period of about .5 hour. Stirring is continued for .5 hour longer. The slurry is then heated to 65 to 75° and stirred vigorously while additional hydroxide is added until the slurry is strongly alkaline* toward litmus paper, and stirring is continued for .5 hours longer. An excess of lime does no harm other than to add to the bulk of the solids to be filtered out and washed.
The hot slurry is filtered by suction. The filter cake is washed with 300 ml. of water saturated with calcium hydroxide at 65°. The filter cake is slurried in 400 ml. of water saturated with calcium hydroxide at 65° and refiltered. The residue is washed with 200 ml. of hot water saturated with calcium hydroxide. The filtrates and washings are combined. If the solution now does not show strong alkalinity to litmus paper, a little more lime should be added and the solution refiltered.
Carbon dioxide is passed into the solution to remove the excess of calcium hydroxide as calcium carbonate, which is then removed by filtration.
The calcium dithionate is recovered by concentrating the solution on a steam bath and cooling. Successive crops of crystals are obtained by further concentration and cooling. When the volume of mother liquor has been reduced to 50 ml., the remaining calcium salt is thrown down by the addition of 75 ml. of ethyl alcohol. The crystal crops are not washed but sucked as dry as possible on a Biichner funnel. The crops of crystals are combined and spread out to dry at room temperature.
In a preparation starting with 80 g. of Java pyrolusite,* 194 g. of calcium dithionate, CaS206.4112O, was obtained. Yield 86 per cent.
* An excess over the calculated quantity of calcium hydroxide is required because of the presence in the crude manganese dithionate solution of such impurities as iron, aluminum, and free sulfurous acid.
B. BARIUM DITHIONATE
Mn02 + 2S02 --> MnS206
MnS206 + Ba(OH)2 --> BaS2O6 + Mn(OH)2
The procedure used in the preparation of calcium dithionate is followed except that barium hydroxide is employed instead of calcium hydroxide. For the amounts specified for calcium dithionate, the initial quantity of barium hydroxide, Ba(OH)2-8H2O, should be about 160 g. Because of the greater solubility of barium hydroxide, greater care must be exercised in arriving at the end point (strong alkalinity to litmus paper).
Using 80 g. of 90 per cent pyrolusite, a yield of 203 g. (73 per cent) of barium dithionate, BaS206.2H2O, was obtained.
C. SODIUM DITHIONATE
Mn02 + 2SO2 -> MnS206
MnS206 + Na2CO3 --> MnCO3 + Na2S206
The crude dithionate solution is prepared in accordance with the instructions given for calcium dithionate. This solution is brought to a temperature of 35 to 40°. With continuous agitation, barium carbonate powder is added a little at a time until further addition does not cause immediate increase in carbon dioxide evolution. Agita tion is continued for 10 minutes, whereupon powdered barium hydroxide is added a little at a time until the slurry is neutral to litmus paper. At this point, a small portion is filtered into a test tube, diluted with an equal amount of water, acidified slightly with hydrochloric acid, and 10 per cent barium chloride solution added. If a precipitate forms, more barium hydroxide (in the form of a saturated solution at 50°) is added to the batch. Addition of barium hydroxide solution and testing for sulfates are repeated, if necessary, until the filtrate is free of sulfate ion. Since any excess of barium hydroxide will be removed in the next step, it is better to add an excess than to use an insufficient quantity, which would leave some sulfates or sulfites in solution.
The slurry is filtered by suction and the filter cake washed with 50 ml. of water at room temperature. To the filtrate warmed to 35° and strongly agitated, about 65 g. of sodium carbonate is added in 1 to 2g. portions. The temperature is brought up to 45°. The slurry is tested a few minutes after each addition with blue litmus paper. When a distinct, permanent but mildly alkaline condition is reached (red litmus changing to a pale, not a deep, blue), the addition of sodium carbonate is stopped. The warm slurry is filtered and the cake washed with 150 ml. of water at 50° made alkaline, to litmus with sodium carbonate. The precipitate is sucked as dry as possible. The filter cake is slurried in 200 ml. of water, made slightly alkaline to litmus with sodium carbonate, and the slurry, heated to 45°, is filtered and washed with 50 ml. of slightly alkalized water at 50°. If the combined filtrates do, not react slightly alkaline to litmus, a little more sodium carbonate is added and the solution refiltered.
Sodium dithionate is recovered by concentrating the solution on a steam bath. Successive crops of crystals are obtained by cooling to about 10°. These crystals are filtered off, sucked as dry as possible, but not washed. A precipitate, which may come down in the first concentration, is removed by filtration from the hot solution before cooling for crystallization. When the volume of mother liquor has been reduced to about 10 ml., it is discarded since most of the dissolved material in it is sodium carbonate. The combined crops of crystals are spread out on absorbent paper to dry at room temperature.
In a preparation starting with 80 g. of Java pyrolusite (90 per cent Mn02), 177 g. (88.5 per cent) of sodium
dithionate, Na2S206-2H2O was obtained.
* The Java pyrolusite contained 16.5 per cent available oxygen (90 per cent MnO2). This material had been ground so that 99.1 per cent passed through 100-mesh, 94.3 per cent through 200-mesh, and 90.7 per cent through 300-mesh.
The solid dithionates are stable at ordinary temperatures but undergo dehydration and decomposition at high temperatures. Above 150° they lose sulfur dioxide, forming sulfates.
In general, the dithionates are very soluble in water. Alkali and alkaline earth dithionates are remarkably stable. Such oxidizing agents as bromine, permanganate, and nitric acid have no effect at ordinary temperatures. In boiling solutions, oxidation to the sulfate takes place slowly. Continued heating with concentrated hydrochloric acid brings about conversion into sulfate with evolution of sulfur dioxide. Such agents as sodium amalgam and zinc in acid solution reduce dithionates to sulfites.
1. ESSIN: Z. Elektrochem., 34, 78 (1928).
GLASSTONE and HICKLING, J. Chem. Soc., 1933, 829. 2. BASSET and HENRY: ibid., 1935, 914. 3. DYMOND and HUGHES: ibid., 71, 314 (1897).
ASHLEY: Chem. News, 94, 223 (1906).
4. ALBU and VON SCHWEINITZ: Ber., 65B, 729 (1932). 5. VON HAUER: J. prakt. Chem., 80, 229 (1860).
MEYER: Ber., 34, 3606 (1901).
SPRING and BOURGEOIS: Bull. soc. shim., [2146, 151 (1885).
VOGEL, Metallboerse, 24, 19 (1929).
Chemistry is our Covalent Bond
Thanks so much lugh! It was very kind of you to put so much time in your posts.
Is there a way to reduce the dithionate anion to the dithionite? Maybe it would be best to try it on the pyrosulphite? In any case, thanks again!