7# "V7!TXTXTXTXTXTf<TTT TTTT4UxTUz UU*UTXUU UUUOUUUUUUChemistry 142 PREPARATION OF SODIUM THIOSULFATE PENTAHYDRATE Introduction In this experiment you will synthesize sodium thiosulfate pentahydrate. This material will be used in the next experiment on kinetics. Sodium thiosulfate pentahydrate, Na2S2O3.5 H2O, is a substance that finds several important uses, both in the chemistry laboratory and in commercial situations. The thiosulfate ion, S2O32-, has a Lewis structure that is comparable to that of the sulfate ion, SO42-, with one of the oxygen atoms of sulfate replaced by a sulfur atom. The presence of the second sulfur atom in the thiosulfate ion makes this species a better reducing agent. The ability of thiosulfate to serve as a reducing agent is the basis for most of its common uses. Uses of Sodium Thiosulfate Standard solutions of sodium thiosulfate are used as the titrant in analyses of samples containing elemental iodine 2S2OA(2-,3) + I2 > A(S4OA(2-,6) + 2I-,tetrathionate ) (1) Such titrations are monitored by watching for the disappearance of the brown color of elemental iodine as the equivalence point is approached. This is usually enhanced by addition of a small amount of starch near the endpoint of the titration, which results in the production of the characteristic blue/black color of the starch/iodine complex. Although not many real samples would be expected to contain elemental iodine itself, iodine can be generated in a sample containing an oxidizing agent by addition of an excess amount of potassium iodide; the half reaction is: 2 I- > I2 + 2 e- (2) Naturally, the quantity of elemental iodine produced is directly related to the quantity of oxidizing agent present in the sample. By titration of the elemental iodine thus produced with standard thiosulfate solution, the quantity of oxidizing agent in the original sample may be determined. Commercially, sodium thiosulfate pentahydrate is used in photography, where the substance is known more commonly as "hypo". Photographic film contains a thin coating of silver bromide (AgBr) spread on plastic. Silver bromide is light sensitive; silver ions are reduced to silver atoms when exposed to light through a camera's lens. The production of elemental silver on the film forms the negative image. When exposed photographic film is developed, any unreacted silver bromide must be removed from the film, or it will also be reduced when the film is exposed to further light. Thiosulfate ion complexes and dissolves silver ions, but is not able to dissolve metallic silver AgBr(s) + 2 S2OA(2-,3) (aq)  Ag(S2O3)A(3-,2) (aq) + Br - (aq) (3) By using sodium thiosulfate to "fix" the film, only the negative image (elemental silver) remains on the film. All unreacted AgBr is washed away. Synthesis Sodium thiosulfate may be prepared by reaction between aqueous sodium sulfite and elemental sulfur. Since elemental sulfur is not soluble in water, a small amount of detergent is added to the system to promote "wetting" of the sulfur. Safety Precautions *Sodium sulfite is toxic and irritating to the skin. Sodium sulfite evolves toxic, noxious sulfur dioxide gas if acidified or heated. All procedures involving heating should be carried out in the hood. *Elemental iodine will stain the skin if spilled. *Silver salts will stain the skin if spilled. The stain is elemental silver and requires several days to wear off. *Use tongs or a towel when handling hot glassware. A. Synthesis of Sodium Thiosulfate 1. Weigh 5 grams of powdered sulfur into a small beaker. Transfer the sulfur to a mortar. Add 5 drops of 50% laboratory detergent and make a paste. Add 5 ml of water to thin out the paste, then use no more than 50 ml of water to transfer the paste to a 400 ml beaker. Weigh 15 grams, to the nearest 0.1 gram, of sodium sulfite into a small beaker. Transfer the sodium sulfite to the sulfur slurry. Stir to dissolve the sodium sulfite. Determine the pH of the mixture with pH paper. 2. Transfer the beaker to a hot plate under the hood, cover with a watch glass, and heat the mixture to a gentle boil for approximately 30 minutes. The presence of the detergent makes this mixture tend to bubble and foam excessively. Keep the heat as low as possible while still maintaining boiling. Stir the mixture frequently while heating to promote the mixing of the powdered sulfur. During the heating period, watch for a subtle change in the appearance and color of the powdered sulfur; the yellow color of the sulfur becomes somewhat lighter as the reaction proceeds. 3. After boiling for 30 minutes, check the pH of the solution with pH paper. The pH of the solution should have dropped to nearly neutral. If it has not, continue heating for an additional 15 minutes. 4. When the pH has reached the nearly neutral point, remove from the heat. Allow the mixture to cool to the point where it can be handled easily. Filter the mixture by gravity or suction to remove unreacted sulfur. Collect the filtrate in a clean beaker. Using the beaker calibration marks, note the approximate volume of the solution. If you have 35-40 mL or less, it should not be necessary to reduce the volume further, proceed to step 7. If the volume is substantially greater than 40 mL, proceed with step 5. 5. Transfer the filtrate to an evaporating dish or casserole. Heat the solution gently on a hot plate. Do not rapidly boil the solution, or a portion of the product may be lost to spattering. Continue heating until the volume of the solution has been reduced by approximately 50%. 6. Protect your hands with a towel while transferring the hot contents of the evaporating dish to a clean, dry 250 ml beaker. 7. Put the beaker into an ice-water bath to allow crystals to form. If no crystals have formed by the time the solution has reached room temperature, you probably have a supersaturated solution. Scratch the bottom and side of the beaker to assist nucleation for crystal formation. Addition of a seed crystal may also be helpful. If crystals still do not form, return the solution to the evaporating dish and reduce the volume of the solution by an additional 25%. Cool the solution again to room temperature to allow crystallization. 8. Filter the crystals through a BO(,u)chner funnel under suction. Wash the crystals with about 10 mL of ice cold 1:1 ethanol-water to remove as much detergent as possible. Allow the suction to continue for several minutes to promote drying of the crystals. 9. Remove a small portion of the product for the tests indicated below. Allow the remainder of the product to dry on a watchglass for approximately one hour. Write a balanced equation for the synthesis. Weigh the product and calculate the % yield. Remember to save your product for use in the next laboratory, "Rate of Reaction: Thiosulfate Decomposition". Store your product with a second watch glass inverted over it. B. Tests on Sodium Thiosulfate Place 5 mL of distilled water in a clean test tube. Add 5-6 drops of 0.1 M I2/KI reagent, followed by a small amount of your product. Describe what happens, and write a balanced equation for the reaction. Place 2-3 mL of 0.1 M NaCl in a clean test tube. Add 3-4 drops of 0.1 M silver nitrate. Record your observations. Add a small amount of your product and shake the test tube. Record your observations and write balanced equations for both reactions. Report In your notebook, record all of your observations, write balanced equations for the reactions in Part A and B, calculate the percent yield of your product and answer the following questions. a) Give the Lewis dot structure for the tetrathionate ion, S4O62-. The tetrathionate ion is the product in reaction (1), above. Tetrathionate has the arrangement of atoms:  b) The I2/KI reagent used in Part B contains the triiodide ion, I3-. What is the shape of this ion? c) Write a balanced equation for the production of SO2(g) from an acidic solution of Na2SO3. This reaction is a handy laboratory method for the production of SO2 in air pollution studies. Adapted from James F. Hall, Experimental Chemistry, D.C. Heath, Lexington, Massachusetts, 1986, Exp. 37. Chemistry 142 PREPARATION OF SODIUM THIOSULFATE PENTAHYDRATE PRELABORATORY ASSIGNMENT Name____________________________ 1. Give the Lewis dot structure for the thiosulfate ion and the sulfate ion. 2. Which is the limiting reagent in your synthesis? 3. 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mdl_patdict begin /FontType 3 def /FontMatrix [1 0 0 1 0 0] def /FontBBox [0 0 1 1] def /Encoding StandardEncoding def /CharProcs 27 dict def CharProcs begin /A {8 8 true [8 0 0 -8 0 8] mdl_A_pat imagemask} def /B {8 8 true [8 0 0 -8 0 8] mdl_B_pat imagemask} def /C {8 8 true [8 0 0 -8 0 8] mdl_C_pat imagemask} def /D {8 8 true [8 0 0 -8 0 8] mdl_D_pat imagemask} def /E {8 8 true [8 0 0 -8 0 8] mdl_E_pat imagemask} def /F {8 8 true [8 0 0 -8 0 8] mdl_F_pat imagemask} def end /BuildChar { 1 0 0 0 1 1 setcachedevice exch begin Encoding exch get CharProcs exch get end exec } bd end /mPatFnt mdl_patdict definefont pop /mVcGd F def /vKvcm { /mVcGd F def } bd /mTM matrix def /mBus matrix def /mVcM matrix def /mVS F def /mVO { mVcGd mVS not and { mBus currentmatrix pop mVcM setmatrix /mVS T def } if } bd /mVF { mVS { mBus setmatrix /mVS F def } if } bd /mCrtVirt { 0 begin gS mVF /bbox 4 array def /pbox 4 array def /smx matrix def /tmx matrix def bbox astore pop np :M :L pathbbox pbox astore 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