At the second titration, the experiment was exactly the same as the first titration but the H2SO4 solution was used to instead of the HCl solution. The same step as the first experiment was repeated. In the third experiment, the buret had used Ba (OH) 2 thoroughly rinsed three times. 10mL of HCl
Measure the solution by right clicking on the beaker and choose pH meter *Then I had to measure the pH of 0.1 M sodium hydroxide Get a 100 mL beaker from the equipment menu Right click on the beaker, select chemicals, and add 50 mL of 0.1 M sodium hydroxide. Measure the solution by right clicking on the beaker and choose pH meter Part 2: *First add 35 ml of unknown acid a to 100 mL beaker. Select all chemicals from the toolbar in the chemicals section, choose unknown acid a. Put the volume at 35 mL in a new 100 mL beaker. *Then add two drops of phenolphthalein indicator to the beaker by right clicking, choosing indications, and adding 2 drops of phenolphthalein.
Part C: Density of Sodium Chloride (NaCl) Solution, a sample of NaCl was obtained and measured using a 100mL beaker and a 10mL pipet to determine the concentration of the solution. In order to obtain the appropriate result, a calibration graph and density measurement was used to determine the concentration of the sodium chloride solution. In conclusion, based on the water temperature of 21.8°C in part A’s graduated cylinder experiment obtained, it was determined that the average density was .0973g/mL with a percentage error of 2.5%. When graphed the measurement was equal to Y=0.988x. Part B: The graduated pipet’s average density at 22.3 °C was determined to be 0.9785g/mL with a percentage error of 1.89% shows the graduated pipet to be more accurate and precise.
Part B: A similar experiment was performed after the 15% NaI in acetone solution. This time, ten test tubes were marked and we put 2 mL of 1% ethanolic silver nitrate in each test tube. The same corresponding halide was added to the mixtures as the first part of the experiment. Once these halides were added, we observed to see which turned cloudy. The mixtures that didn’t turn cloudy were put into a hot water bath of about 100 C for about a min.
Next, 50 mL of distilled was placed into the 150 mL. Twenty drops of Bromothymol blue were added to the 150 mL beaker solution. The pH was then recorded. Five mL of this solution was transferred into three separate 100 mL beakers. In one of these beakers, 1 mL of HCl was added to the solution, making this the “Yellow” beaker.
The buffer checks out because pH=pKa ± 1.0. Results: Phosphate buffer: The initial pH for the phosphate buffer is 6.83 and it takes 18 drops to equal 1.0 mL of 1.0 M HCl. The volume of HCl required to break the buffer was 1mL, and the known concentration of the HCl 1.0 M The number of moles of HCl consumed by the buffer was 0.02. [pic] Acetate buffer: The target pH of NaCH3CO2 was 5.25 and the actual measured pH for the acetate buffer was 11.21. The number of drops to equal 1Ml of NaOH is 17 drops.
After adding three drops of brymothymol blue indicator, the solution was now a clear blue colour. In the first and fourth trial, the bromothymol blue indicator caused the solution in the Erlenmeyer flask to turn a pale yellow colour from the blue colour that the solution already was after the 10.00mL of acid was added. This indicates that the solution was now acidic. In the second, third and fifth trial, the bromothymol blue indicator caused the solution in the Erlenmeyer flask to turn a pale green colour from the blue colour that the solution already was after the respective amounts of acid was added to the solution. This indicates that the solution was now neutral, or close to a pH of 7.
Add 0.5 ml concentrated HCl and 1.0 ml 15% KI solution. Mixed exactly 1 minute and leave for 5 minutes in a dark place. Add 0.5 ml starch solution, 20 ml distilled water. Mix and titrate with sodium thiosufate solution. Calculate the exact normality of Na2S2O3 knowing that in this chemical reaction 1 gram-equivalent of K2Cr2O7 react with 1 gram-equivalent of Na2S2O3 (1 mole K2Cr2O7 react with 6 moles Na2S2O3).
Brad Hopkins CHM 145 A Shadi Abu-Baker Cheng Guo Alex Hudson 10/10/2011 Experiment #4: Ionization Constants of Weak Acids Qualitative Data: The indicator dye that we used was Blue (1). During the qualitative analysis I was able to use the shades of the buffer solutions after the dye was added to determine an estimate of the pKa of the dye. This is because according to the Henderson-Hasselbalch equation, the pH = pKa when there are equal amounts of the forms of due present. To determine when the amounts were equal, observations were made on the dye coloration. When the KH2PO4 was of a greater volume the resulting dye was yellow in color; when the K2HPO4 was in greater volume the resulting dye color was dark blue in color.