The freezing point of p-xylene was calculated as 13.29C after averaging the data that appeared on Graph 1 once the temperature leveled off. With this value, the Tf for each trial was able to be calculated through Equation 1, which led to Kf being calculated in Equation 2. Both equations were able to be used given that the measurements were in terms of molality, which is not temperature dependent. After completing calculations, the average Kf of the three trials of the p-xylene and toluene solution was computed as as 4.56(C/m) as shown in Table 1, however, the theoretical value was slightly lower than calculated, 4.3(C/m). This resulted in a 6.04% error as shown in Equation 5. Possible causes of error could have resulted from adding too much …show more content…
Given the 6.04% error resulting from Kf, the calculated results were not consistent with what was expected, however, with the theoretical Kf value the results were as predicted. Potential sources of this higher Kf could have resulted from combining too much or too little solvent or solute, which would potentially raise the Kf, or contamination of the substances would impact the intermolecular forces at play. If this lab were repeated, assuring the measurements of the solute and solvent were accurate would provide more exact results, given that freezing point depression is a colligative property. Another potential way to receive a more accurate Kf would include more trials to have an average closer to the theoretical value. The freezing point depression lab centralized on quantitatively finding freezing points and freezing point depression, as well as using calculations to determine Kf and the molar masses of unknown substances. Freezing point depression describes the decrease that appears in the temperature at which a solvent freeze’s when a solute is dissolved in it. Given that the nature of a substance dictates physical properties, a pure substance is expected to have a constant freezing point at a fixed pressure. However, when another substance is added to create a homogenous mixture, …show more content…
The mass of 2mL of the solvent, p-xylene, was measured as well as the mass of 10 drops of toluene, the solute. The temperature of the solution rose to the freezing point after supercooling, then continued to drop as the solution froze, versus stabilizing as the pure p-xylene did. The maximum temperature obtained after supercooling was recorded as the freezing point of the solution. This process was repeated three times, each with a new test tube and the same beaker and scale for measuring the masses, then the Tf was calculated for each trial, as well as the average Kf. For the final portion of the lab, the same procedure as above was followed with the substation of 10 drops of unknown solutes instead of toluene in the p-xylene. Each unknown, A, C, and D, had one trial each. With the experimentally gathered data, the molar masses were then computed and compared to the given compound molar masses to
The theoretical weight was 599.6 mg. This yields a percent yield of 3.7%. Table 1 also illustrates the experimental melting point of 99.3-102.1◦C. A melting point that has a range larger than 3◦C is indicative of impurities in the sample. A few possibilities of impurities could have been unreacted norbornene, and water. Evidence that supports that there was unreacted norbornene in the final sample was the fact that the product was a jelly-like structure. Norbornene by itself has a jelly-like structure. However, once norbornene reacted with the acid-catalyst (H2O2), then it should have changed the chemical structure of the molecule and once the solution was brought back down to room temperature, crystals should have formed. Since a jelly-like, or oil-like product was present at the end of the reaction, then this is indicative that there was unreacted norbornene in the sample. The second impurity that may have been present in the final product was water. Instead of adding 3 mL of sodium bicarbonate and then 3 mL of brine, 3 mL of brine was added first and then 3 mL of sodium bicarbonate was added. This experimental error caused excess aqueous solution to be added to the diethyl ether. Since excess water was added to the final product, about 4x the amount of anhydrous sodium sulfate was needed in order to remove the water from the product. This was another indication that there was too much water in
...s the change in the temperature of both of these batches, 6°C for the pure, and 13°C for the crude. In this final sub-section of the Characterization of Aspirin, the values of absorbance were recorded. Initially, 0.0566 grams and 0.0590 grams of pure and crude Aspirin respectively were obtained and each individually placed into beakers (400 milliliter) and had 250.0 milliliters of distilled water added to them. From each beaker, a tiny amount of the just dissolved solutions was transferred to a cuvette, one cuvette for each type of aspirin. Each cuvette was placed into the ultraviolent spectroscopy mechanism which was connected to a computer and absorbance spectrum values were obtained at 298 nm (Figure 5) (0.1987 pure aspirin, and 0.9549 crude aspirin).
Craig, D. Q. (2002). Pharmaceutical Applications of Micro-Thermal Analysis. Journal of Pharmaceutical Science, 91(5), 1201-1213.
The aspirin crystals were packed into 3 small capillary tubes to ensure that they are compressed so as to prevent any air gaps. Subsequently, the aspirin crystals that are in the 3 capillary tubes are placed into the melting apparatus and the temperature range was recorded. Since the range is quite far from the theoretical value of 140°C, aspirin's purity attained was low due to impurities present. One potential reason is because of the swift cooling. When the aspirin is left to cool, the crystal lattices will form too rapidly which will surround other molecules thus making the aspirin impure. Another reason could be because the recrystallized aspirin has not dry completely and there might me left over solvent that will affect the temperature range of the aspirin.
Another trend in this table which demonstrates this phenomenon is the decreasing FPV of the CHO cells after cooling and freezing/thawing which shows the increasing membrane fluidity. However, compared to the control cells (at 0 mg) the CLC treated cells still showed considerably less membrane fluidity after being cooled.
Specific heat capacity of aqueous solution (taken as water = 4.18 J.g-1.K-1). T = Temperature change (oK). We can thus determine the enthalpy changes of reaction 1 and reaction 2 using the mean (14) of the data obtained. Reaction 1: H = 50 x 4.18 x -2.12.
The temperature probe was placed into the test tube and recorded the temperature of the freezing solution using Logger Pro software. The test tube was held against the inner glass of the ice bath beaker so the test tube was visible to see when the solution froze over. Once the freezing point was measured, the temperature stopped being monitored and the data was recorded. The steps mentioned above for finding the freezing point, also known as ΔTf, was replicated for the 0.0, 0.4, and 0.6 concentrations. To find the freezing point depression, the equation ΔTf = imKf was used. The molality (m) of each solution was then calculated dividing moles of solute by kilograms of solvent, and the Kf value for magnesium chloride is known to be -1.86. Since magnesium chloride breaks down into three ions in deionized water, it was concluded that the Van’t Hoff factor couldn’t exceed three. For better accuracy, the experiment explained above for finding the freezing point depression and Van’t Hoff factor was re-conducted exactly the same to determine more accurate results. Again, the molality of each solution was calculated, and a graph expressing the change in freezing temperature verses molality
Our task was to investigate what the optimum ratio of solute to solvent that will produce the maximum cooling/heating effect?
A Digi-Melt apparatus was used to measure the melting points of the benzoic acid and benzocaine samples which were extracted using the separatory funnel. Based on the observed melting point ranges for each sample, purity could be determined. A tight melting point range that is very close to the reference melting point is considered to be very pure. This is when the range is ~2 ̊C, but is not within 5 ̊C of the reference melting point range
The last part of experiment 5, was learning about specific gravity and temperature. Specific gravity does not have any units, it is unitless. When measuring for the temperature, we used a thermometer to calculate the Celsius of the water, 10% sodium chloride, and isopropyl alcohol. The specific gravity uses a hydrometer to measure the gravity of the liquids. Using the hydrometer, to figure out the measurements we have to look at it from top to bottom. The water for specific gravity was .998 while the temperature of it was 24
The Solubility of Potassium Nitrate Aim To investigate how the solubility of Potassium Nitrate is affected by Temperature. Background Knowledge Potassium Nitrate (KNO3) is an ionic compound. The strong ionic bonds hold the compound in an ionic lattice which gives KNO3 its crystalline structure. These ionic bonds also have other properties which will affect my investigation, I must be aware of these properties for greater accuracy in my method.
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molecules its size it would have a boiling point of -75øC and a freezing point of -125øC4.
On further cooling the χT curve shows a sudden increase to 1.23 cm3.K.mol-1 at T=21 K followed by a sharp decrease down to 0.71 cm3.K.mol-1 at 5 K. The χT maximum de...
...inty between 1.0% (0.1/10.00*100) and 2.13% in the measured volume and 0.1/4.70*100). We also used a digital thermometer that allowed us to read the temperature readings from five degrees celcius to eighty degrees celcius. Since the digital thermometer have an absolute accuracy of plus or minus one degree celcius, it gives a percent uncertainty between 0.125 % (0.1 / 5.00 * 100) and 0.2 % (0.1/ 80.0 * 100). One of the difficulties we faced during the lab is reading the inverted graduated cylinder. To account for the inverse meniscus, we subtracted 0.2 mL from all the volumetric measurements to account for that. Volumetric uncertainty is the most important in determining the accuracy of this experiment since we are constantly checking for the volume throughout the lab. It also is the factor that gives the highest percent uncertainty out of all the instruments used.