For a general salt, AmBn, the equation would be: The equilibrium constant for such a salt would be: The solubility product expression matches the coefficients in the equilibrium equation, and that the solid is not included in the Ksp equation. In this experiment, you will determine the Ksp for the sparingly soluble salt potassium hydrogen tartrate (also called acid tartarate or bitartrate). It dissolves in water to give potassium ions and hydrogentartrate ions according to the following equation: KHC4H4O6 K+ + HC4H4O6- The solubility product is
In order to synthesize this compound, measure out 1 gram of NaC2H3O2.Weighed 1 gram of NaC2H3O2 and mixed it with ionized water. Boiled 12 mL of 1.0M Acetic Acid added into a beaker containing the sodium carbonate on a hot plate until all the liquid is evaporated leaving solid precipitate form inside of the beaker. Weigh the precipitate form and mix it with water until fully dissolved. Weigh the precipitate remaining. A flame test and a halide ion test were performed.
Standardization of NaOH solution Accurately weigh out a sample of approximately 0.3-0.4 g of primary standard potassium hydrogen phthalate, KHPh, which has been previously dried at 120°C. Do not use more than 0.4 g. To obtain an accurate mass, weigh the sample on weighing paper, slide it into a clean (but not necessarily dry) 250 mL Erlenmeyer flask and reweigh the paper to account for any KHPh that may remain on it. Dissolve the KHPh sample in about 50 mL of CO2-free water and add 2-3 drops of 0.1% phenolphthalein indicator. Begin adding the approximately 0.1 M sodium hydroxide solution from the buret while continuously swirling the flask contents. Do not open the stopcock completely.
Eventually using the NaOH and the acid’s consumed moles, the equivalent mass will be determined. Procedure: Part 2: Obtain 45mL of NaOH, and then weigh 0.3-0.4g of the unknown acid (KH2PO4). Dissolve the acid into 20.00mL water. Record the buret readings, and slowly titrate the NaOH into
Absorbance of the solution was measured at 238 nm. Drug content was calculated by the formula absorbance/ slope* dilution factor. Drug loading Microsphers equivalent to 100 mg of DLTZH/ND taken and washed with 3 x 10 ml of methanol, which removes the free unloaded drug. Filtrate was diluted suitably to Beer’s concentration range and free drug concentration was determined spectrophotometrically at 238 nm and 236 nm for DLTZH and ND respectively. Further microspheres were crushed, dissolved in methanol and made up to 100 ml, which was further diluted suitably to Beer’s rangeand drug concentration was determined.
These tests are essential to the recrystallization process because if the selected solvent is inappropriate, the results from the process will not be pure. Choosing the solvent with a similar structure to that of the solute is crucial because it will dissolve more solute than with a solvent of different structure. Also, the boiling point is an important part of process when determining an appropriate solvent. If a solvent has a boiling point below 100⁰C, it is not right solvent due to loss of the compound being recrystallized, which is soluble in the cold solvent, resulting in the crystals are not forming back. The solute should have a maximum solubility in the hot solvent and a minimum solubility in the cold solvent to be at purest state.
200ml of water was then added to the filtrate in a 500ml beaker with constant stirring. White solid was formed in the process of addition and the solution was then left undisturbed in an ice bath for 10minutes. Once most of the solids had settled at the base of the beaker, the solution was decanted. 10ml of ethanol was added to the remaining suspension and was transferred in a clean centrifuge and centrifuged for 2minutes at 6000rpm. After the first centrifugation, the supernatant was discarded and the residue was washed by adding 14ml of ethanol.
As it was done in the Experiment A, 20 drops of 0.2 M acetic acid and 10 drops of 2% starch solution was mixed well with the juice solution. Before adding the iodine solution, the initial reading of the burette was taken. Then, the titration was started using the iodine solution into the burette with continuous swirling of the flask slowly and carefully. Once the color change started to appear, titration was stopped and final burette reading was recorded. Finally, the amount of vitamin C in the mandarin orange was calculated by using the standardization factor and used iodine solution.
In this experiment iodometric titration is done using an unknown concentration of copper(II) sulfate in a 100ml volumetric flask (A43). The solution had a light blue colouration after it was topped up to the graduation mark, which is expected of a copper(II) solution. From the stoichiometric calculations it is possible to deduce that the number of moles of thiosulfate ion is the same the number of moles of copper(II) ion by looking at reaction 1&2. Before commencement of the experiment the apparatus used were conditioned with D.I. water to remove any existing particles or compounds that could have affected the titration results.
Hence, the amount of ethanoic acid to neutralize sodium hydroxide is lesser too. Distilled water can be poured in the pipette to swill out the remaining sodium hydroxide. Distilled water will not affect the amount of moles in sodium hydroxide, but it will remove the remaining NaOH in the pipette and transfer all in the conical