Experiment 5 Redox Titratio

KIO3 + 6 HCl + 6 Na2S2O3 → 3 Na2S4O6 + KI +3 H2o + 6 NaCl

1. To titrate a hydrochloric acid solution of “unknown” concentration with standardized 0.5M sodium hydroxide.

Enter the precise mass in GRAMS of the potassium iodate used to prepare your primary standard solution.

Discard the solution in the appropriate container as directed to you by your lab instructor.

mL cylinder to the beaker on the stir plate and empty it into the beaker. Place the pH probe in the beaker and record the pH in the data table. Drag the beaker to the red disposal bucket. Double-click the bottle of NaHCO3 to move it to the Stockroom counter. Repeat steps 5 and 6 for KNO3.

After the twenty minutes elapsed, the flask was cooled to room temperature and then titrated with the remaining NaOH until the colorless solution remained pink. The final volume was then recorded. While solution #1 was heating the same process was repeated with solution#2 and the second burette

Again, label 7 1.5ml tubes 0 thru 6. Place 15μl of each serially diluted extract into its corresponding labeled tube. Next add 465μl of media into each tube. Then 60μl of Alamar blue in each tube. Finally add an additional 60μl of cells (adjusted to 10,000 cells/20 μl). Vortex each tube for 5 seconds. Now, take 3 different samples 190μl samples of concentration 0 and put it in Wells A2, B2, and C2. Repeat this step again by taking 3 more different 190μl samples of concentration 1 and putting it in wells A3, B3, C3. It should be noted that it is important to vortex each 1.5μl tube again be-fore putting it into the 96 well plate. Contin-ue this same procedure consecutively for the re-maining concentrations.

ii. The second part of the titration series involves titration of NaOH with Hydrochloric acid (HCL). Again, three reps of titration and a blank titration have to be completed. A volumetric pipet is used to measure 10.00mL of HCL into three labeled conical flasks. Then the flasks are filled with deionized water until about the 50mL mark. A buret is

These equations can only be carried out and be visible after the iodine has completely reacted with thiosulphate added – two moles of thiosulphate for every mole of iodine. Once all the thiosulphate has been used up in the reaction, the colour will start to appear.

During a titration the pH of the solution will be monitored using a pH meter from that we get a titration curve. The titration curve is then used to determine the equivalent molecular weight and Ka value of the unknown weak acid, from that we are

1. 10 drops of 1 M K2CrO4 was added to the solution and stirred for about 10 minutes.

Add the same amount of Benedict’s solution in each tube and for the ones that need iodine use the same amount.

In order to find a standard solution you must first, add water from a cleaning bottle to dissolve it. Second, you would need to use some of the water to rinse off all the substance off the watch glass. This needs to be done at least two times. Then stir with a glass rod until all the of the solid substance is dissolved, then put the solution to the volumetric flask. The laboratory method is when you drop the titrant into the analyt. (basically titration). The results you get are titration, pH level, and the moles. The equation to find molarity/concentration is: M1 x V1 = M2 V2. In conclusion, titration is very useful in the modern day. It's also pretty

The excess iodine is then back titrated by standard sodium thiosulphate solution, by using a burette and starch solution (used as end point indicator) Similarly, iodine can also be

Redox Titrations involve standard solutions of oxidants and reductants, the titrant. A change in oxidation state is effected by the oxidant or reductant. This change of oxidation state accompanies a change in the indicator, signalling the endpoint of the titration has been reached.