Ghazaldehydrate Lab Report

The two most obvious formation of the precipitate were the combinations with the MgSO4. The MgSO4 and NH3 solution became very opaque and the MgSO4 and Na2CO3 turned from liquid to a full solid white substance. The Na2CO3 and CH3COOH did not have as strong of a reaction, however, the precipitates were able to be visualized with in the clear

10 microliters of the sample is then added and the assay absorption is measured at 340nm. If absorbance was above 1.5, samples were diluted.

0.1 gram of my product from the second trial was weighed in a tray and was then added to a fourth test tube containing 2.0 mL of Iron (III) chloride, which was measured using a 10 mL graduated cylinder, to test for

Heavy precipitate emerged immediately and solution turned white in color; solution then became opaque and turned light, bright blue in color.

The organic layer on the top formed, and was also extracted out through vacuum filtration like the step before. The solid that was produced from the reaction was bright yellow in color. It then was set out to dry for a week, and worked on the next experiment time. After setting the solid aside, the original solution that was still in the separatory funnel, 14 mL of 10% NaOH was added. The aqueous layer was then again separated into a separate flask. This same step was repeated twice more. The aqueous solution that formed was then cooled in an ice-bath. Concentrated HCl was then added until a pH reading of 1 was reached making it an acidic solution. This took 13 mL of HCl to reach the adequate pH. Another precipitate formed from this process, and it was extracted through vacuum filtration. The solid precipitate that was formed was a cloudy white color. It also, was set out to dry for a week until the next experiment time.With the remaining solution in the separatory flask, CaCl2 (a drying agent, which helps dry the precipitate that was filtered out so that no extra liquid will be left over 5), was added. This solution was boiled over a steam bath until no bubbles were present, and left to dry

Unfortunately, achieving this task is very unlikely due to three major unavoidable sources of error. The first one being the adherent property of the precipitate. The precipitate had a tendency to stick to the sides of the beaker and even the bottom of the stirring rod. Even after attempting to wash it off, most was able to be cleaned but a little bit of the precipitate was still stuck and would not come off. This leads to having less than the actual yield.

The many different aqueous solutions involved were potassium nitrate, zinc nitrate, potassium ferricyanide, sodium phosphate, sodium hydroxide, magnesium nitrate, copper (II) nitrate, sodium chloride, iron (III) nitrate, calcium nitrate, sodium sulfate, and sodium carbonate. These different reactants were combined to test whether or not a precipitate, color change, or both would occur in that in that those reactions. The purpose of this lab was to demonstrate what a double replacement reaction as well as a precipitate

A solute and solvent were chosen from the list to start with. Around 10 mg of solid and 0.25 mL of solvent were added to a clean reaction tube. A pipette was useful in transferring the solvent to the tube. The mixture was stirred with a glass stirring rod. If the substance was completely dissolved at room temperature, a few drops of water would then be added to see if the solid precipitated.

From that point, the oxalate solution was added to the iron (III) chloride solution in the 150 mL beaker and stirred. To ensure the highest percent yield, any remaining K2C2O4 was rinsed with 3 mL of DI water and added to the 150 mL beaker. After all solutions were

I prepared a new ferrioxalate solution by adding 1g of into 25mL of distilled water in a 150mL beaker. Then, I added 1g of oxalic acid and placed it on the bottom shelf. 2. I cut out 4 pieces of paper circles with approximately 9cm in diameter. 3.

The purpose of this lab was to become familiar with the three different balances and two different methods used to find the weight and mass of chemicals and compounds in the ChemLab program. The lab was performed by using three different types of balances, and the direct weighing and weighing by difference methods.

The absorbance obtained from the results of each test sample, then put into the equation of calibration curve in accordance with the respective raw substances, after which it would have obtained the value of molar absorptivitas. The molar absorptivitas results obtained are used in the determination of the percentage of the sample levels

Methods and Materials 0.3g methyl benzoate and 0.6ml sulfuric acid were added into a small test tube in an ice bath. A mixture of 0.2ml sulfuric acid and 0.4ml nitric acid was then added to methyl benzoate solution. The contents were mixed thoroughly and then allowed to sit in the ice bath for another 5 minutes. The test tube then put in a drying rack for 15 minutes. After the 15 minutes, the reaction was poured into a 25ml Erlenmeyer flask with 2 small ice chips in it.

By using the average absorbance of every solution prepared here, and knowing the concentration of standard iron added to each, a standard addition curve was constructed with the purpose of determining the original concentration of iron. Figure 3 Illustrates the standard addition curve constructed using the data from Table 1. The line of best fit is also included in the graph, as well as the equation representing it and its coefficient of determination.

The black precipitate was allowed to settle and then the supernatant, the clear liquid that lies above a precipitate, was decanted, or poured carefully off. Then, 200 mL of hot distilled water was added and the precipitate was allowed to settle to repeat the decanting process again.