An exothermic response is one that includes a negative ΔH value while an endothermic reaction is one that includes a positive ΔH value. Based on the results, the dissolution of potassium hydrogen tartrate has a ΔH value of 3. 89 *104. Therefore as 3. 89 *104 > 0, the dissolution of potassium hydrogen tartrate is endothermic, i. e. high temperature must be added to dissolve the sodium in water. This is further reinforced by the negatively sloped graph above. It can be inferred from the graph that the bigger the temperature, the greater the solubility of potassium hydrogen tartrate.
The spontaneity of a reaction depends on the change in enthalpy (ΔH) and entropy (ΔS), as well as on the absolute temps. The change in the Gibbs free energy (ΔG) can be used to determine if a response is spontaneous or not. Displayed by the equation ΔG = ΔH - TΔS, when ΔG is negative, an activity proceeds spontaneously in the in advance path. When ΔG is positive, the procedure proceeds spontaneously backwards. When ΔG is zero, the process is within equilibrium, with no net change occurring as time passes.
Therefore as the ΔG determined at both 10. 0C and 50. 0C are positive, it can be deduced that the dissolution of potassium hydrogen tartrate is non-spontaneous in the front direction at the conditions examined (10. 0C to 50. 0C). The procedure however proceeds spontaneously backwards at the mentioned temps. In fact, based on the results, the dissolution is only going to be spontaneous at 513. 6K (ΔH / ΔS) and above.
Although both entropy and enthalpy are functions of temp, this experiment assumes that ΔH and ΔS do not change significantly over the number of conditions used. This assumption is valid over relatively small amounts. In this test, the various measurements take place within a little selection of 40K. Therefore it is safe to suppose that the ideals of ΔH and ΔS are relatively invariant over the small changes in temperature.
Many reasons could cause the experimental value to disagree with the literature value. Probably the most significant source of mistake would be the shortcoming to keep the respective temperature during the sluggish step of filtering the salt solution after heating up/cooling. Gravity purification was implemented in this experiment, and although the filter papers were fluted to encourage immediate filtration, the procedure still spanned several minutes. During this time, the temperatures of the salt solution could have easily deviated from the required temperature towards room temperatures. This could cause undesirable recrystallization/dissolution of the salt, thereby, impacting the molarity of the filtrate. This restriction could be defeat by using vacuum filtration to reduce such problems.
Besides the possible undesired recrystallization/dissolution before filtering was over, recrystallization could also arise in the filtrate before the 25mL aliquots were obtained. This might only connect with the samples that were experimented at temperatures above room temps. As the filtrate cools, the salt would recrystalise, triggering a big change in homogeneity and molarity of the filtrate. But the protocol was to immediately aliquot 25mL after filtration, enough time for the filtrate to reach at least 25mL was sufficient enough for significant chilling of the answer, with the large surface area of the answer filtering dropwise and the connection with the cool conical flask. A possible solution to the problem is to heating the filtrate again after filtration before aliquoting to ensure all the salts that could have recrystalised dissolve. However, the heating up should be soothing to avoid significant lost of solvent preventing change in molarity.
Another source of error would be the assumption that the molarity of the given NaOH solution is accurate. Alkaline alternatives such as NaOH absorb skin tightening and from the atmosphere according to the response: CO2 + 2OH- ↔ + CO32- + H2O. Since hydroxide ion is consumed by this reaction, the focus of the sodium hydroxide solution will be evolved. Therefore the specific amount of the NaOH solution may not be the worthiness that that which was mentioned on the container, especially if the perfect solution is was prepared long before the experiment was conducted. Standardization of the NaOH solution should be done just before the test is conducted.
Although in this test, the molarity of the provided NaOH solution was assumed to be appropriate no further standardization was done, safeguards were taken to protect the solution from the skin tightening and that is definitely present in the atmosphere. As during titration, the NaOH solution in the buret will come in contact with air, the buret used was ready for use only when it was needed, and fresh sodium hydroxide should be added if it.
The primary steps of the task was to obtain roughly 200ml of NaOH solution from the stock container. To obtain higher precision though, the NaOH taken from the stock container shouldn't be more than what's necessary for one titration. More NaOH solution should be taken from the stock bottle when needed.
Other safety measures were also used during the experiment to reduce contaminants. Equipment were scrupulously cleaned out and rinsed with solutions that these were to contain before use. The test was also cautiously done to prevent loss of materials through spillage, splashing or splattering.
Conclusion
This experiment efficiently demonstrated the relationships between status functions, including entropy and enthalpy, free energy, spontaneity, and equilibrium constants. Since ΔH and ΔS were both positive, the dissolution of potassium hydrogen tartrate was spontaneous at high temperatures. Which means that the potassium hydrogen tartrate needed energy from the surroundings to dissolve. The change process however, is spontaneous.
A significant percentage error was obtained by looking at of the obtained solubility product frequent with the books value. Although precautions were taken up to ensure accuracy, such an error proves that solubility product constants are extremely difficult to acquire experimentally because of limitations of the experiment and the necessity to identify all chemical varieties and processes within the substance system used to get the values.
