Lab Q: Chemical Equilibrium: Le Chatelier s Principle

Transcription

1 Lab Q: Chemical Equilibrium: Le Chatelier s Principle Poppy Quinlan Partner: Katie Frese February 11, 2014 CHEM 123 L10 TA: Katie Nguyen Lab Performed: January 28, 2014

2 Introduction: Chemical equilibrium occurs when the forward rate of a chemical reaction is equal to the reverse rate of the chemical reaction. This is called dynamic equilibrium. 1 When a stress causes a change in equilibrium, the reaction will naturally oppose the original stress to return to equilibrium. This was discovered in 1888 by Le Chatelier and therefore named Le Chatelier s Principle. 2 This stress can be caused by a change in temperature, a change in concentration, or a change in pressure. During this lab, two of these stresses were evaluated, Part A and B examined changes in concentration, while Part C examined changes in concentration and temperature. For Part A, the reactions below were used to examine the effects of Le Chatelier s Principle due to changes in concentration: 2CrO4 2- (aq) + 2H + (aq) Cr2O7 2- (aq) + H2O (I) Fe 3+ + SCN - [Fe(SCN)] 2+ (II) In Part B of this lab, the following reaction of sparingly soluble salts was evaluated: CaSO4(s) Ca 2+ (aq) + SO4 2- (aq) (III) The saturated solution of these salts were at equilibrium between two phases (solid and aqueous). 3 The changes in equilibrium shifts due to concentration can be examined by the formation of solid when the reaction favors the left. Lastly, for Part C, the reaction: [Co(H2O)6] 2+ (aq) + 4Cl - (aq) [CoCl4] 2- (aq) + 6 H2O (IV) is observed when changes in concentrations and temperatures effect the equilibrium of the reaction.

3 Le Chatelier s Principle can be very helpful, and allows for qualitative statements about an equilibrium reaction. 4 For the reaction: A + B C + D (V) If the concentration is increasing for A or B it causes a shift to the right. If the concentration is increasing for C or D it causes a shift to the left. If the concentration is decreasing for A or B it causes a shift to the left. If the concentration is decreasing for C or D it causes a shift to the right. If the temperature is increased there is a shift towards the endothermic reaction. 4 If the temperature is decreased there is a shift towards the exothermic reaction. 4 If the pressure is increased there is a shift towards the side which produces fewer moles of gas. 4 If the pressure is decreased there is a shift towards the side which produces more moles of gas. 4 Procedure: The procedure from the First Year Chemistry Lab Manual for Experiment Q 1 was followed with the alterations specified below. Starting with Part A-2, 5 drops of 0.1M Fe(NO3)3, 5 drops of 0.1M KSCN, and 10mL of H2O were added to a 50mL beaker. Using this mixture, the solution was added to five separate test tubes using a pipette of unknown volume to add 1mL of solution in to a 10mL graduated cylinder and transferred to the test tubes. To test tube B, 2 drops of 0.1M Fe(NO3)3 were added and compared to test tube A, which was used as a colour reference. Next, 2 drops of 0.1M KSCN was added to test tube C, 2 drops of 0.1M KNO3 was added to test tube D, and 1mL of H2O was added to test tube E. These test tubes were then compared to the colour of test tube A.

4 For part B, Lab Q Part B Procedure handout was used in place of the First Year Chemistry Lab Manual 3 for Experiment Q with the alterations as specified. Using a medium sized test tube, 2.0mL of 1.00M CaCl2 and 2.0mL of 1.00M Na2SO4 were combined from preset dispensettes. With a glass rod, the solution was thoroughly mixed for a little over three minutes. After the solution was mixed, the test tube was centrifuged for two minutes. Using a graduated plastic transfer pipette, 1mL of the liquid on top was transferred in to test tubes B and C. For the remaining test tube A, there was not enough liquid to add 1mL and so the solution was centrifuged a second time for another two minutes as the solution had become mixed. After the second time the solution was mixed, a little under 1mL of the liquid was added to test tube A. Next, 1mL of 1M Na2SO4 was added to test tube A, mixed with a glass stir rod, and let stand for fourteen minutes. While test tube A was standing, 1mL of 1M CaCl2 was added to test tube B, mixed, and let stand for five minutes, and 1mL of 1M NaCl was added to test tube C, mixed, and let stand for fifteen minutes. Once the solutions had been left to stand for the appropriate length of time, it was noted if a precipitate formed or not. Subsequently, 6M HCl was added dropwise to test tubes A and B, which had formed precipitates. In the fumehood for part C, 1.0mL of 1.0M CoCl2 was added to two medium sized test tubes, with test tube A serving as a colour reference. To test tube B, 1.0mL of concentrated HCl was added and observations were recorded. To the same test tube, 1.0mL of H2O was added and observations were recorded. Test tube B was then heated up in a beaker of boiling water for approximately one minute and then placed in ice for about the same amount of time. Observations on colour changes were noted.

5 Finally, for part A-1, 2mL of 1M K2CrO4 was added to two medium sized test tubes, test tubes A and B. To test tube A, 2 drops of 6M HNO3 were added and the colour change was observed, and then 2 drops of 6M NaOH were added, which reversed the colour change. To test tube B, 2 drops of 1M KCl were added and the colour change (or lack of) was observed. Lastly, all equilibrium shifts were explained and all calculations for part B were calculated. Data and Calculations: For Data and Calculations, see attached Lab Q report sheets. Observations: Ions/Solutions Colour/Precipitate formed CrO4 2- Cr2O7 2- K2CrO4 + HNO3 K2CrO4 + HNO3 + NaOH K2CrO4 + KCl Fe 3+ SCN - Fe(SCN) 2+ Fe(NO3)3 + KSCN Fe(NO3)3 + KSCN + additional Fe(NO3)3 Fe(NO3)3 + KSCN + additional KSCN Yellow/Pale Hay Peach Orange Orange Yellow/Pale Hay Yellow/Pale Hay Light Orange/Yellow Colourless Transparent Popsicle Orange Orange Very Dark Orange Fairly Dark Orange

6 Ions/Solutions Fe(NO3)3 + KSCN + KNO3 Fe(NO3)3 + KSCN + H2O CaCl2 + Na2SO4 CaSO4 + NaCl CaSO4 + Na2SO4 CaSO4 + CaCl2 CaSO4 + NaCl CaSO4 + Na2SO4 + HCl CaSO4 + CaCl2 + HCl CoCl2 CoCl2 + HCl CoCl2 + HCl + H2O CoCl2 + HCl + H2O + heat CoCl2 + HCl + H2O - heat Colour/Precipitate formed Fairly Light Orange (due to dilution) Light Orange White/Cloudy Colourless on the top, White Precipitate on the bottom Precipitate formed Precipitate formed No precipitate formed Precipitate went away Precipitate went away Magenta Dark Blue/Purple Pink/Purple Dark Blue/Purple Very Pink/Reddish

7 Discussion: For Part A-1, the Chromate/Dichromate equilibrium, equation (I), due to changing concentrations was observed. Initially, the colour of the reactants was a yellow/pale hay colour due to the CrO4 2- ion. With the addition of 2 drops of 6M HNO3 to test tube A the colour changed to a peach orange. This is due to the increase in concentration of the hydrogen ion causing an equilibrium shift to the right. Next, after adding 2 drops of 6M NaOH, the colour changed back to the original yellow/pale hay colour. The solution changed back to the original colour because the hydroxide ions reacted with the hydrogen ions to form water, decreasing the concentration of the hydrogen, which drives the equilibrium to the left. In test tube B, 2 drops of 1M KCl were added to the 1M K2CrO4 but no colour change was observed due to the fact that neither the potassium ions nor the chlorine ions reacted with the solution to cause an equilibrium shift. Next, the Ferric Thiocyanate complex ion formation was examined in part A-2. A solution containing the reaction in equation (II) was added to five test tubes (A-E) and colour changes due to the shifts in concentrations were observed. Originally, the Fe 3+ ion was a light orange/yellow and the SCN - ion was colourless. Test tube A was used as a colour reference and was a transparent popsicle orange colour. To test tube B, the addition of 2 drops of 0.1M Fe(NO3)3 caused the solution to turn from a transparent popsicle orange to a very dark orange. This is because the concentration of Fe 3+ increases and therefore there is an equilibrium shift to the right. A similar reaction occurs to test tube C from the 2 drops of 0.1M KSCN. The reaction produces a darker orange than the colour of the reference colour in test tube A. This is because there is an increase in concentration of the SCN - ion causing the reaction to shift to the right. In test tube D, there was no colour change and therefore no shift in equilibrium because the addition

8 of 2 drops of 0.1M KNO3 does not cause a reaction with the solution. Lastly, test tube A was compared to test tube E after the addition of 1mL of water, diluting the solution. When the equilibrium was re-established, the colour of the Fe(SCN) 2+ ion is less intense than the expected half intensity due to the fact that for the equation K=[Fe(SCN) 2+ ]/[Fe 3+ ]*[SCN - ] when affected by the dilution becomes Kdilution=0.5[Fe(SCN) 2+ ]/0.25([Fe 3+ ]*[SCN - ]). Since there is less reactants than products after dilution, there is a shift towards the reactants, causing the colour to be less intense than the expected half intensity. For part B, equation (III) is evaluated. When 2.0mL of 1.00M CaCl2 and 2.0mL of 1.00M of Na2SO4 are mixed a cloudy precipitate is formed in a colourless solution due to the separation of CaSO4 and NaCl. To test tubes A, B, and C 1.0mL of CaSO4 is added. To test tube A, 1mL of 1M Na2SO4 is added, and after around fifteen minutes, a precipitate is observed. This is because the concentration of SO4 2- increases and therefore there is an equilibrium shift to the left, which forms the solid precipitate. In test tube B, 1mL of 1M CaCl2 is added and allowed to rest for five minutes. After five minutes, a visible precipitate forms due to the increase of concentration of Cl 2- causing a shift towards [Fe(SCN)] 2+ (s). When 1mL of 1M NaCl is added to test tube C, after fifteen minutes no precipitate is formed. This is due to the fact that neither of the ions react with the solution to cause a shift in equilibrium. Lastly, when 6M HCl is added dropwise to test tubes A and B, where precipitate had formed, the addition of HCl did not make the precipitate disappear. This is because CaSO4 is nearly completely insoluble in HCl. With the sparingly soluble solid, CaSO4, the Ca 2+ and the SO4 2- are effectively removed from the reaction, forcing the equilibrium shift towards the solid to remain.

9 Lastly, for part C, the equilibrium reaction from equation (IV) was examined. In a test tube, 1.0mL of 1.0M CoCl2 (initially a pink/red) and 1.0mL of concentrated HCl were added which turned the solution a dark blue/purple and therefore indicates a shift to the right because of the increase in concentration of Cl -. To the same test tube 1.0mL of water was added. The solution turns to a pink/purple colour which indicates an equilibrium shift to the left because of the increase in concentration of H2O. Next, the test tube was placed in boiling water which turns the solution a dark blue/purple indicating a shift to the right. When the test tube was put in ice, the solution returned to a pink/purple colour shifting back to the left. These equilibrium shifts indicate that the reaction in endothermic. When there is an equilibrium shift to the right due to heat, the reaction favors the reactants, causing KC to increase. When there is an equilibrium shift to the left due to cooling, the reaction favors the products causing KC to decrease. Using this information and table 1 from the lab manual, 1 it shows that the reaction is in fact endothermic. There are several possible sources of error throughout this lab including contamination of chemicals or equipment, uncertainty in measuring instruments and equipment, or loss of solution when transferring from one piece of equipment to another. Also, for test tube A, in part B there was not enough remaining clear supernatant liquid to fill the test tube with 1mL of the solution with the 1M Na2SO4 and therefore had an estimated mL of liquid instead which could have affected the reaction. Although it may have had an affect, a precipitate still formed, which was as expected. During each reaction that was evaluated all of the equilibrium shifts and colour changes concurred with Le Chatelier s Principle, both visually and mathematically. When a change in the concentration of the reaction occurs, a visual colour change showed the reaction shifts either to

10 the left or the right. Le Chatelier s Principle also helped determine that reaction (IV) was endothermic because as the temperature increased, KC increased and as the temperature decreased, KC decreased (as determined by colour change and which side of the reaction was favored). This lead to the conclusion that the H is positive and therefore endothermic. Also, for part B, Ksp and Q were calculated for test tubes A, B, C, as well as the Ksp of Na2SO4. These calculations proved that the reaction shifted in the same direction as predicted by Le Chatelier s Principle. As all of the reactions in this lab matched the expected equilibrium shifts as predicted by Le Chatelier s Principle, this lab was a success. Conclusion: Using Le Châtélier s Principle, the examination of several chemical equilibriums that experienced either a change in concentration or a change in temperature were observed. The observations observed all concurred with Le Châtélier s Principle therefore making the lab a success.

11 References: [1] First Year Chemistry Manual: Chem 111/113 & 121/123, University of British Columbia: Kelowna, BC, ; p [2] Ihde, J.; Journal of Chemical Education, March 1989, 66(3), /ed066p238?prevSearch=%255BAbstract%253A%2BLe%2BChatelier %25E2%2580%2599s%2Bprinciple%255D&searchHistoryKey= (accessed February 6, 2013) [3] Lab Q Part B Procedure handout [4] Petrucci, R. F.; Herring, F. G.; Madura, J. D.; Bissonnette, C. General Chemistry Principles and Modern Applications,10th ed.; Pearson: Toronto, ON, 2011;