Essay on Chem 2c Post Lab

(.1063 KIO31) (1 mol KIO214 g) x 6 mol S2O41 mol KIO= .00298.04150= .259 M

The number of moles of ascorbic acid neutralized in the experiment is given by multiplying the moles of NaOH used (0.00235128 mol) by the unit conversion coefficient (1 mol ascorbic acid/1 mol NaOH).

First I will set up the apparatus as show above. I will add 1.5 grams

sodium thiosulphate + hydrochloric acid = sodium chloride (salt) + sulphur + water + sulphur dioxide

The results showed the molarity of the NaOH solution. This experiment was completed twice and a new average molarity

9. Each of the other solutions is already approximately 0.1 M. With these solutions you can pour a small

3.6.3. 2, 4 – D (2, 4–Dichloro phenoxy acetic acid) stock solution (1mg/ml): 10.0mg of 2.4-D being weighed and dissolved completely in 1N NaOH to a final total volume

*Note: A recording mistake was made in the lab notebook, where the final product was supposed to be 0.05g (calculations above) and not 0.55g (calculations below).

For every two moles of thiosulphate, one mole of iodine is needed so the amount of iodine mole needed is half the amount of thiosulphate moles that are needed. Starch is used as the indicator for the end point because the iodine and water is such a faint yellow color that the starch will combine with any of the iodine that is still present and turn it a blue color and the end point for the titration is than easier to notice.

Then copy trial 1 except using 1.5% hydrogen peroxide instead of 3% hydrogen peroxide. Trial 3 12. Prepare .025 KI by mixing 2 mL of distilled water with 2 mL 0.5M KI. Then copy trial 1 except using .025 KI instead of 0.5M KI.

Molarity of the original bleach = [(Average molarity of ClO- in diluted bleach) × (0.100 L)] ÷ (Volume of original bleach)

The purpose of this two-part experiment was to “determine the mass percent of sodium hypochlorite in a commercial chlorine bleach sample” (lab website). Titration is used in this lab twice. In Part one, students will use the solution of KIO3 primary standard and use this solution to standardize a solution containing sodium thiosulfate. When the KI is added to the solution, it forms a brown/yellow color. Then students used titration to add Na2SO3 to make the color go from brown/yellow to colorless. In the second part of the lab, students used the standardized Na2SO3 solution to determine the mass percent of sodium hypochlorite in the commercial chlorine bleach. Students used titration to the standard solution of Na2SO3 to

A way to answer this question is to do at least three experiments, in order to get an overall view of the numbers you need, then to use a calculation device that may be able calculate how much potassium iodate you need, in order to get the results you want. Some materials you need would be potassium iodate (an essential component), water, a magnetic mixer, beakers, sodium meta-bisulfate solution, graduated cylinders, a stop watch, and starch solution. The starch solution and sodium meta-bisulfate is used as Solution B, the solution whose volume will never change throughout the experiment, due to the need to keep a constant variable throughout the experiment.

Fill in the following table, using the value of a obtained from the simulation. Record and calculate data to the nearest three decimal places. The masses in the table are in kg, but the input data are in grams, so make the necessary conversions. The first row has been completed for you.

The rough draft of the results section should include all tables used in your experiment. All values within the tables should be in numerical form and contain units. For instance, if measuring the amount of chloride in water you should report as 2 mg/L or 0 mg/L not as two or none.