Introduction For this experiment, we are going to determine the effect of temperature on solubility, to be done in a chemical by dissolving a solute in a definite amount of solution which is saturated. Specifically, the goal of this experiment is to prepare a saturated solution of Na 2 C 2 O 4 in water at different temperatures, determine the effect of temperature in solubility, and to apply Le Chatelier’s Principle. We can do all this by simply titrating a certain amount of standard KMnO 4, and measuring how much KMnO 4 was needed to help Na 2 C 2 O 4 reach chemical equilibrium at certain temperatures. In doing this, we will have fulfilled all the goals of our experiment, as well as being able to determine the Solubility of Na 2 C 2 O 3 at both 20 and 100 C, in addition to determining both experimental and theoretical ∆ H for the reactions. Procedure A. Preparation and Standardization of KMnO 4 Solution: I weighed out about 2.
8 grams of potassium permanganate and dissolved it in about 200 ml of distilled water. I made sure to make all of the permanganate dissolve, of course, by stirring it thoroughly. I then weighed about 2 samples of sodium oxalate (. 47 and. 50 g) then placed them in some 200-ml Erlenmeyer flasks, then and added about 50 ml of distilled water to them as well. I then titrated the potassium permanganate solution after adding about 15 ml of 3 M sulfuric acid to the oxalate and heating the whole solution to about 60-90 C, using a magnetic stirring bar to help stir.
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I kept on slowly titrating at the same temperature until the solution turns very light pink. I then stopped, recorded the amount of KmnO 4 used, then went on. B. Preparation and Analysis of Samples: I weighed out about 12 g of sodium oxalate and then put it into a 250-ml Erlenmeyer flask, along with 150 ml of distilled water. I then stirred it around, to make sure that the solution dissolved completely.
After stirring, I measured its current temperature then allowed it to settle for a while. When it is done, I used a pipette to remove about 5 ml of the clear saturated solution at room temperature from the Erlenmeyer flask, calling it B 1, then I heated the solution to 40 C then removed another 5 ml sample, called it C 1, and then got a cool 5 ml sample of 3 C solution, called it A 1, then heated the solution to 80 C and got one last sample, called it D 1. I then added 50 ml of distilled to the test tubes A 1, B 1, C 1 and D 1, then titrated them with the potassium permanganate solution (making sure to add 15 ml 3 M sulfuric acid in each), then recorded the amount it took for them to turn pink. Data.
Standardization of KMnO 4 solution: Sample g Na 2 C 2 O 4 moles Na 2 C 2 O 4 (x 10^-3) moles KMnO 4 (x 10^-3) ml KMnO 4 used M KMnO 4#1. 47 3. 51 1. 4 15. 5. 090#2.
50 3. 73 1. 5 16. 093 Average M = . 091 B. Solubility of Na 2 C 2 O 4 at different temperatures: in M and g/ L.
Sample and temp (C): ml of KMnO 4 titrated moles KMnO 4 (x 10^-4) moles NaC 2 O 4 dissolved (x 10^-3) M (moles/L) Na 2 C 2 O 4 g Na 2 C 2 O 4 dissolved in sample g / L Na 2 C 2 O 4 A 1 – 3 5. 0 4. 55 1. 14.
228. 152 30. 5 B 1 – 31. 1 6. 0 5. 46 1.
37. 274. 183 36. 6 C 1 – 45 8. 8 8. 01 2.
00. 400. 268 53. 7 D 1 – 80 9.
5 8. 65 2. 16. 432. 290 57. 9 Solubility of Na 2 C 2 O 4 in g / L Experimental Value in g / L Handbook Value in g / LAt 20 C 36 37 At 100 C 73 63.
3 C. Relationship of K and ∆ Sample Average M of Na 2 C 2 O 4 K = [Na^+]^2[C 2 O 4] ln K T K 1 / T K (x 10^-3) A 1. 228. 047 -3.
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PLANNING Investigating the Kinetics of the reaction between Iodide ions and Peroxodisulphate (VI) ions By the use of an Iodine clock reaction I hope to obtain the length of time taken for Iodine ions (in potassium iodide) to react fully with Peroxodisulphate ions (in potassium Peroxodisulphate). I will do three sets of experiments changing first the concentration of iodide ions, then the ...
05 276 3. 62 B 1. 274. 082 -2. 50 304. 1 3.
29 C 1. 400. 256 -1. 36 318 3. 14 D 1. 432.
322 -1. 13 353 2. 83 At 20 C. 276. 084 -2. 48 293 3.
41 At 100 C. 472. 421 -0. 87 373 2. 68 Experimental ∆ H = 13. 9 KJ / mol Theoretical ∆ H = 18.
3 KJ / molDiscussionWith the results from my experiment at hand, it is very obvious that the effect of temperature on solubility is that the higher the temperature of the reaction, the more chemical is dissolved, both in grams and moles, which in turn gives a higher equilibrium constant. Additionally, the higher the temperature of the solution, the more standard KMnO 4 was needed in order to balance out the sodium oxalate. Regarding ∆ H and K, it is apparent that they are proportional to each other, the higher the rate law, the more change in heat there is, which effectively means that the more amount of a substance is used, the greater the amount of reactant dissolved in order to reach chemical equilibrium and / or minimize the amount of change, trying to return back to equilibrium. This in turn gives a greater amount of heat discharged as well, with regards to the amount of temperature.
Conclusion In conclusion, Le Chatelier’s Principle has indeed been proven to be a valid principle regarding chemistry because of the fact that through my experimentation, I have consistently noticed that the greater the amount of reactants and products I had, as well as the higher the temperature, the greater the amount of work and heat was generated (obvious because of a large rate law), which means that there is a greater struggle to reach chemical equilibrium the more substance there is. The heat given off was also proportional to these reactions. Questions 1. When I compared my experimental values and handbook values of the solubility of Na 2 C 2 O 4 at both 20 and 100 C, I have found that my experimental value for 20 C closely matched that of the handbook value (36 g / L and 37 g / L respectively), which was very good, showing that I performed the experiment correctly. Regarding my solubility at 100 C, I found that I was a bit more off my intended value than I would have liked, getting about 73 g / L when my handbook value was at about 63 g / L. I can probably attribute my mistakes to not holding my temperature constant throughout that part of the experiment.
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2. The dissolving process is endothermic because of the fact that as I dissolved the Na 2 C 2 O 4, I noticed the temperature of the solution drop as I titrated the KMnO 4 onto it. I can also attest to this observation by noting that when I touched the beaker, it felt cooler than previous. 3. Yes, the results I obtained are in agreement with this principle because I have already noted in my discussion that the greater the amount of reactants and products, the bigger the rate law will be, which means that there is also a larger ∆ H involved.
All of this points to the fact that there is a greater reaction that tries to make the solution reach equilibrium once disturbed.
