Determination of ΔH and ΔS Using the plot of lnKsp versus 1/T, the slope of the best-fit equation and the y-intercept was found. Using van’t Hoff equation, the change in enthalpy ΔH and entropy ΔS of the dissolution reaction was found to be 44 ± 1.3 kJ K-1 mol-1 and 89 ± 4 J K-1 mol-1 respectively. Using van’t Hoff equation, lnK = - (ΔH°reaction)/RT + (ΔS°reaction)/R where ΔH°reaction and ΔS°reaction ¬are the standard enthalpy and entropy change of the reaction respectively. From Figure 1, the gradient and y-intercept was obtained as shown in Table 4 in the Appendix. The enthalpy change and entropy change was then calculated.
In this experiment, it was possible to produce the major products from bromination of acetanilide and aniline. 0.075g of 4-bromoacetanilide and 0.156g of 2,4,6-tribromoanilne were collected from bromination of 0.07g acetanilide and 0.05g aniline with the percent yield of 67.57% and 88.1% respectively. At the end of the experiment, to prove the formation of the major products, melting point of the products were measured. The melting point of the product from the bromination of acetanilide was 164.8-168.50c, which is in the range of the melting point of 4-bromoacetanilide, 165-1690c, as reported on the Chemical Book, CAS Database List (chemicalbook.com). The melting point of the product from the bromination of aniline was 119.8-121.90c, which is in the range of the melting point of 2,4,6-tribromoaniline, 120-1220c, as indicated on PubChem, Open Chemistry Database (pubchem.ncbi.nlm.nih.gov).
The values of K for solutions 1-5 and U were 4.39E4, 4.53E4, 4.23E4, 4.70E4, 6.35E4, and 4.03E4 respectively. The average K for the lab was found to be 4.71E4 and the standard deviation was 8.302E3. The range then that all experimental values of K must fall under is 3.05E4 to 6.37E4. All experimental values of K for this trial fell within the range. Therefore, K can be determined a true constant.
This solution was diluted with diluents to gae a concentration of 0.1 mg/ml solution each of Amoxicillin trihydrate. The HPLC method was applied to the solutions and the results obtained were shown in table 4.6.11. System suitability solution: 25.0 µg/mL each of of USP Amoxicillin RS in Diluent. Precision
Since the solvent is nonpolar, we would expect carotene to have the lowest Rf, then xanthophylls, and chlorophylls would have the highest. As discussed before, this is because chlorophylls are the most polar of the three and carotenes are the least polar of the three. Based on my results, I would say that the packing and running of the column was effective in separating the pigments as there were quite different Rf values for the original than for the other two colors and the band for the yellow was nothing like the bands for the green. There were no limitations in procedure which negatively affected the end results, so there is nothing I would change to it going
CHAPTER 6 RESULTS AND DISCUSSION 6.1. INTRODUCTION The experiment gave the knowledge about various things and various factors played their significance role in it. The experiment stated the Chromium removal and for that we had drawn a calibration curve (graph 6.1) between Absorbance on y axis and concentration on x axis through the table 6.1 as given below. To make calibration curve, we needed the absorbance of the Chromium solution which we got from atomic absorption spectrophotometer (AAS). For calculating % of Chromium removal we have, (C0 – C1) ÷ C0 × 100 Initial concentration (before adsorption) =C0 Final concentration (after adsorption) = C1 So the average efficiency or % Chromium removal = 52.575% The factor analyses are – 1.
The calculated value was 1.6 x 10^-5. Conclusions The resulting Ka of the acetic acid from this experiment’s calculations was consistent with the experimental results. The experimental percent of CH3COOH was calculated at 1.6 x 10^-5, while the actual value was 1.8 x 10^-5. The calculated value is much lower because the pH read from the graph at half of the equivalence was higher than the actual value. To have gotten a 0% error between the experimental and actual value for CH3COOH, the pH would have been measured at about 4.75, which is slightly more acidic than 4.80.
Plot of ln[Aspirin] against time The plot in Fig. 4 shows a strong linear relationship between ln[aspirin] and time. The calculated r2 value is 0.99236 while the gradient is -6.94931 x 10-5. The standard deviation of the regression line is as small as 2.73×10-6, as expected from the properties of ln values. The gradient, which is essentially the change in concentration divided by the change in time, follows the differential rate law for first order reactions where: Rate = (-d[aspirin])/dt=
Upon finding the actual concentrations of salicylic acid, concentration of aspirin in the flask at various times can be found using the equation [aspirin]t = [aspirin]0 – [salicylic acid], since at constant volume, number of moles of initial aspirin decrease to form salicylic acid. Initial concentration of aspirin formed as follows: [aspirin]0 = 0.212g / (180.157gmol-1 * 50/1000 L) = 0.0235 mol L-1. Thus using the first test as sample, [aspirin]t = 0.0235 – 9.981*10-4 = 0.0225 mol L-1. To find the rate constant, we will need to log the value of [aspirin]t and plot it against time to find the rate constant. Figure 1 shows the diluted and actual concentrations of salicylic acid, the concentration and log value of aspirin at various times.
pH of the Amylase/starch will be kept the same. Uncontrolled Environmental conditions Atmospheric conditions The controlled variable Concentration of amylase was kept under control by measuring the amount of amylase used and also it was made sure the percentage of amylase used was 1%. The Amount of amylase/starch used were kept to 5cm3 at all times. Materials needed Beakers Bunsen burner Test tube Thermometer Stopwatch Test plate Glass rod Starch Amylase solution Water bath Iodine solution. Test tube holder Labels Marker Procedure First 5 test tubes were taken and labeled with numbers from 1 to