Objective In this experiment, the critical micelle concentration of sodium dodecyl sulfate (SDS) is determined by the conductance method.
Procedure In this experiment, a series of SDS solutions at various concentrations are tested for their conductance at two different temperatures, 25 °C and 50 °C. For detailed procedure, refer to the lab manual (J. F. Wójcik and T. S. Ahmadi, Experimental Physical Chemistry, 2015; p.125-129.).
Data Sodium dodecyl sulfate has a molecular weight of 288.372 g/mol, with a density of 1.01 g/cm³. The melting point of SDS is in the ranges of 204 -205.5 °C. In this experiment, 8.6151 g of SDS was weighed to make a 500-mL 0.06 M solution.
Before any calculations, all the conductance data were baseline corrected with the conductance of pure water at each temperature. To
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According to a literature Domínguez, A, the critical micelle concentration of SDS at 25 °C determined by the conductance method is 0.0080 M. Compared to the literature value, the experimental value has a percent error of 8.25%. The discrepancies between the two values could be a result of slightly different methods, the presence of impurities, or the slight change in concentrations due to evaporations of the solvent. When impurities are present in the solution, they can affect the ability for SDS molecule to form micelles. This can affect the critical micelle concentration. Note that the critical micelle concentration of SDS at 25 °C is lower than at 50 °C. This is most likely due to the fact that at higher temperatures, there is more movement between molecules, which makes them more difficult to come together and form a proper micelle structure. Therefore, more SDS molecules are needed in the solution so that they can form micelle at higher temperature, explaining why the CMC at 50 °C is higher. There are no known literature value for CMC of SDS at 50
METHOD: The following procedure was taken from the 2017 Millsaps College lab manual.1 The experiment was split into two parts, part A and part B. Part A was to find the heat capacity while part B determined the specific heat of an unknown metal. This was the final goal of the lab. To start, a temperature probe had to be connected to a LabQuest2 data collection device. 100.0 mL of deionized had to be added into a Styrofoam cup.
The purpose of this lab was to be able to use physical characteristics to determine the identity of an unknown compound. The data from this experiment classified aluminum as metallic; ascorbic acid, paraffin, palmitic acid, sucrose, graphite, and water as molecular; sodium chloride as ionic. In order to determine this, 3 tests were conducted. The first test was to test the conductivity of each substance at room temperature. In this test, only graphite and aluminum conducted.
For Herbert Run the conductivity level was 687µS/cm. The Turbidity level was 0 FAU and the Nitrate level was 0.02ppm. I accept my hypothesis and reject parts of my hypothesis. I reject that both streams have a high turbidity level. Both streams’ turbidity level is zero.
The quantitative solubility of the unknown compound was determined to be 29/100ml. The known solubility of sodium sulfate is 28.11g/100mL water. Using the found solubility to compare to the known solubility of sodium sulfate. This solution created in the solubility test, the conductivity of the unknown compound was tested using an Ohmmeter to measure the resistance of the solution. Resistance is the measure of a substances ability to conduct
Goals The primary goal of this experiment was to identify an unknown compound by running various tests to determine the qualitative solubility, conductivity, and pH value of the compound. Tests were also performed for the presence of specific cations and anions in the compound. The second goal was to discover the reactivity of the unknown compound by reacting it with different types of substances. The third goal of this project was to calculate the quantitative solubility of the unknown compound in water.
The powder on the filter paper could've fell and this caused it to have a smaller percent purity, percent yield and also cause a lower absorbance and concentration of pure ASA. Another error would be not using a properly dried sample for the pure ASA in part C when making the crystals, this could have cause tye percent yield error. This would make a lower melting point. To prevent this from occurring next time there could be a dry sample that is completely dry and this would not alter the mass of the sample and this would make the solution have a more
This finding could have been due to experimental errors which affected the results. Discussion
A typical micelle in aqueous solution forms an aggregate with the hydrophilic "head" regions in contact with surrounding solvent, sequestering the hydrophobic single-tail regions in the micelle centre. This phase is caused by the
The data table provided below obtained melting point data for crude product, pure product, and mixture of the pure and 4-tert-butylbenzyl. 12. The TLC data obtained is provided in a table below. The TLC data was conducted solely in a 9:1 hexane/ethyl acetate solvent solution as opposed to the 1:1 and pure hexane solution as well. This was due to the lack of time, but as explained in number 7, a very polar solvent (1:1 solution) or non-polar solvent (pure hexane) is not ideal when obtaining
There are multiple points both at 43°C and at 72°C which indicates that liquid was collected at these temperatures. Based on this information, it would appear that two different liquids were present in solution and that one liquid has a boiling point of approximately 43°C and that the other has a boiling point of approximately 72°C. The literature value boiling point for DCM in is reported to be about 40°C and it is about 80°C for cyclohexane. Based on the graph, DCM was collected from 4 ml to 22 ml, thus 18 ml of DCM was collected.
The concentration calculation is based on the absorbtion which depends on the p-nitrophenolate form. This could have contributed to the drop
The reagents used were Diphenylamine reagent which contains concentrated H2SO4. The standard solution used for this test is the deoxyribose standard solution. In the sample, only a faint blue solution appeared, which indicates a small presence of deoxyribose. In test for Phosphate, the standard solution was the Phosphate solution and the reagents used were concentrated H2SO4, concentrated HNO3, 2.5% ammonium molybdate solution.
In this experiment, the amount of water lost in the 0.99 gram sample of hydrated salt was 0.35 grams, meaning that 35.4% of the salt’s mass was water. The unknown salt’s percent water is closest to that of Copper (II) Sulfate Pentahydrate, or CuSO4 ⋅ 5H2O. The percent error from the accepted percent water in CuSO4 ⋅ 5H2O is 1.67%, since the calculated value came out to be 0.6 less than the accepted value of 36.0%.This lab may have had some issues or sources of error, including the possibility of insufficient heating, meaning that some water may not have evaporated, that the scale was uncalibrated, or that the evaporating dish was still hot while being measured. This would have resulted in convection currents pushing up on the plate and making it seem lighter by lifting it up
Introduction The goal of the experiment is to examine how the rate of reaction between Hydrochloric acid and Sodium thiosulphate is affected by altering the concentrations. The concentration of Sodium thiosulfate will be altered by adding deionised water and decreasing the amount of Sodium thiosulphate. Once the Sodium thiosulphate has been tested several times. The effect of concentration on the rate of reaction can be examined in this experiment.
It was calculated and found that the concentration of benzoic acid was higher at 30℃ (0.0308M) than at 20℃