* Trial 1 36-25.5=10g * Trial 2 36-25=11g * Trial 3 35.5-25=10g 2. Calculate the density of the unknown liquid for each trial. (Divide the mass of the liquid calculated above by the volume of the liquid.) * Trial 1: 10.5/50=0.20g/mL * Trial 2: 11/49=0.20 g/mL
2) Percent recovery for isolation of benzoic acid % Recovery = mass of recovered material _________________________________ x100% mass of starting material = (0.43/1.01) x100% = 42.57% That concludes that the percent recovery is 42,57%. 3) Percent recovery for isolation of hydroquinone dimethyl ether % Recovery = mass of recovered material _________________________________ x100% mass of starting material = (0.16/1.01) x100% = 15.84% That concludes that the percent recovery is 15.84%. Table 2: : Experimental IR peaks compared to literature IR peaks for Benzoic acid Functional groups | Experimental peak (cm-1) | Literature peak (cm-1) | O-H | 3407-2563 | 3400-2564 | C=O | 1689 | 1689 | C-H |
* Trial 1 37.00(g) – 26.10(g) = 10.90(g) * Trial 2 36.70(g) – 26.15(g) = 10.55(g) * Trial 3 36.10(g) – 26.05(g) = 10.05(g) 2. Calculate the density of the unknown liquid for each trial. (Divide the mass of the liquid calculated above by the volume of the liquid.) * Trial
Lab 4: Determination of Percent by Mass of the Composition in a Mixture by Gravimetric Analysis Introduction Thermal gravimetric analysis is used to determine the percent by mass is used to determine the percent by mass of a component in a mixture. When a mixture is heated to an appropriately high temperature, one component in the mixture decomposes to form a gaseous compound. The mass of this particular component is related to the mass of the gaseous compound. In this experiment, the percent by mass of sodium hydrogen carbonate (NaHCO3) and potassium chloride (KCl) in a mixture will be determined. Experimental First, we weighed 2 samples, each has 1 gram of NaHCO3-KCl mixture Second, we put the samples in 2 crucibles (A and B) and weighed them.
Adjust the percent transmittance to 100% 4)our out the water in the cuvet, and fill with 2/3 of the reference solution. Read and record absorbance data. Read from lower concentration to higher concentration. 5)Continue to collect absorbance data for al reference and test solutions 6)Dispose of the contents of the cuvets. Data Tables #1 Reference Solutions for the Calibration Curve Sample [FeSCN2+] Absorbance Reference Solution #1 4x10-5 .2034 Reference Solution #2 6x10-5 .3028 Reference Solution #3 8x10-5 .3915 Reference Solution #4 1.0x10-4 .4908 Reference Solution #5 1.2x10-4 .5768 #2 Test Solutions Temperature - 21.9°C Sample [Fe3+] [SCN-] Absorbance Test Solution #6 1.0x10-3 2.0x10-4 .1002 Test
Rate= k[HSO3-]x[IO3-]y = 0.0005sec-1[HSO3-]1[IO3-]1 Table 2: Time per experiment Experiment | Mixing | 1st trial time (s) | 2nd trial time (s) | Avg time(s) | 1 | #1 and #6 | 37.56 | 33 | 35.28 | 2 | #2 and #7 | 40.54 | 38.53 | 39.535 | 3 | #3 and #8 | 51.08 | 55.79 | 53.435 | 4 | #4 and #9 | 77.00 | 73.50 | 75.25 | 5 | #5 and #10 | 203.00 | 149.19 | 176.095 | Table 3: Calculate moles of HSO3- Test tube numbers being mixed | Volume of HSO3- used | Moles of HSO3-Moles=molarity x liters | #1 and #6 | 10 mL | 0.01 L x 0.001 M = 0.00001 moles | #2 and #7 | 10 mL | 0.01 L x 0.001 M = 0.00001 moles | #3 and #8 | 10 mL | 0.01 L x 0.001 M = 0.00001 moles | #4 and #9 | 10 mL | 0.01 L x 0.001 M = 0.00001 moles | #5 and #10 | 10 mL | 0.01 L x 0.001 M = 0.00001 moles | Table 4: Calculate the rate of IO3- IO3-(aq) + 3HSO3-(aq) I-(aq) + 3SO42-(aq) + 3H+ (aq) Test tube being mixed | Moles of HSO3- | Moles of IO3- reacting | Conc. of IO3- (volume is 20 mL) | Rate of IO3- (conc./avg time) | #1 and #6 | 0.00001 | 3.33 x 10-6 | 1.67 x 10-4 | 1.67 x
Name: Date: Period: Course: CHEM 510 Ideal Law Practice Set 2.0 Directions: Solve the following problems on a separate sheet of paper using the 5-step method to show your work. All answers must be in proper SIG FIGs. Use the Ideal Gas Law (PV=nRT) to solve each of the following. R = 8.31 (L•kPa)/(K•mol) or 0.082 (L•atm)/(K•mol) 1. Calculate the volume 3.00 moles of a gas will occupy at 24.0 °C and 101.3 kPa.
ABSTRACT: The purpose of this experiment was to determine the rate constants and ionic strengths of the series and to prove the Bronsted relation. In order to do so, known amounts of KI, Na2S2O3, KNO3, EDTA, starch and K2S2O8 were mixed up, and placed in the spectrophotometer until the %T reached 60%, and time was recorded. In the first part of the calculations, for flask 1, 2 and 3, the true reaction rate was calculated using the equation k = (1/∆t) x ([S2O32-]/[Iodine][S2O82-]). Which resulted in values of 2.8878765.66 x 10-3 s-1 , 3.159845 x 10-3 s-1, and 3.079703 x 10-3 s-1, these values are all similar to each other since they contain no electrolyte reacting with the persulfate solution. The apparent reaction rate was calculated using the equation, kapp= (1/∆t) x ([S2O32-]/[S2O82-]) which resulted in apparent rate constants of 5.66 x 10-5 s-1, 6.1958 x 10-5 s-1, 6.0356 x 10-5 s-1.
In this titration, three major steps involved were first measure out an amount of known solution, add indicator and three, to titrate using unknown solution. At the first colour change the titration would stop. Abstract The purpose to this lab is to standardize an approximately 1 mol•L-1 solution of hydrochloric acid. There were many steps taken in this titration but the main steps were measuring out the amount of known solution, adding the indicator and finally, titrating the solution until the first colour change. In each trial, the initial reading, final reading and the volume of HCl used was recorded down as quantitative results.