UNIT 3: CHEMICAL EQUILIBRIUM (TEXT: Chap 14-pg 627 & Chap 18 pg )

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1 UNIT 3: CHEMICAL EQUILIBRIUM (TEXT: Chap 14-pg 627 & Chap 18 pg ) *Remedial questions on Concentration of Solutions (3.10 pg ) 3:1. ATTEMPT QUESTIONS a) b) c) d) on Page 146 3:2. What volume of 12 M HCl must be added to 1.0 L of 0.1 M HCl so that the resulting sol'n will be 0.5 M? [ Hint : take "x" litres of 12 M HCl] (34.8 ml) 3:3. 15 ml of 0.5 M NaCl & 60 ml of 0.5 M KCl are mixed together. Find the concentrations of each ion in the sol'n. ( [Na + ]=0.1, [K + ]=0.4, [Cl ] = 0.5 M ) 3.4. Practice questions on working out equilibrium concentrations a) A + B AB b) A + B AB c) A 2 + B 2 2 AB d) A 2 + B 2 2AB [in] [in] [in] [in] [Eq] [Eq] [Eq] [Eq] e) A 2 + B 2 2 AB f) A 2 + B 2 2 AB g) 2A + B A 2 B h) 2A + B A 2 B [in] [in] [in] [in] [Eq] [Eq] [Eq] [Eq] i) 2A + B 2 2 AB j) 2A + B 2 2 AB k) 2AB 3 A 2 + 3B 2 l) 2AB 3 A 2 + 3B 2 [in] [in] [in] [in] [Eq] [Eq] [Eq] [Eq] m) 2AB 3 A 2 + 3B 2 n) 2AB 3 A 2 + 3B 2 o) 2AB 3 A 2 + 3B 2 p) 2AB 3 A 2 + 3B 2 [in] [in] [in] [in] [Eq] [Eq] [Eq] [Eq] *LeChatelier's Principle. (14.8 pg ) 3.5. ATTEMPT QUESTIONS a) 15 (pg 660) b) c) d) e) on Page 661 f) (pg 664) 3:6. SO 2 (g) + ½O 2 (g) SO 3 (g) kj Discuss the conditions that would favour a high eq'm conc of SO 3 (g). 3:7. Explain,using LeChatelier's Principle, why increasing the pressure on boiling water causes it to stop boiling until a higher temperature is reached. What kitchen utencil relies on this notion? H 2 O (l) H 2 O (g) H rx = + 44 kj 3:8. MnCl 2 has a fairly low solubility in water ( 2.6 g/100 ml). Using LeChatelier's Princilple would you expect the solubility to be any better in a solution of 1 M NaCl.(write the chemical equations out) ( known as the Common Ion Effect ) *The Law of Chemical Equilibrium (14.3 pg ) 3.9. Attempt question a) 14.8 b) c) d) e) f) g) on page :10. The following system is at eq'm : NH 4 NO 3 (s) 2 H 2 O (g) + N 2 O (g) a) Write the Ke expression b) If the [ H 2 O ] was tripled what would happen to [ N 2 O ]? c) If the [ N 2 O ] was reduced by 16 X what would happen to [ H 2 O ]?

2 *Calculating the equilibrium constant ( Ke or Kc) (14.9 pg ) Attempt questions a) b) c) d) e) on page 663 in your text 3:12. Consider this eq'm : PCl 5 (g) + H 2 (g) 2 HCl (g) + PCl 3 (g) in a 2 L vessel, at eq'm, there were 2 mol of PCl 5 (g), 2 mol of H 2 (g),6 mol of HCl and 2 mol of PCl 3 (g) a) Calculate Ke for this rx. b) Calculate Ke for the reverse rx. 3:13. A rx'n occurs according to this equation : A (g) + 2 B (g) 2 C (g) * 3 mol of A and 5 mol of B are reacted in a 3 L container. After equilibrium had been reached 1 mol of A remained in the container. Calculate Ke for the rx. 3:14. Attempt questions a) b) c) d) on page 663 *Using the equilibrium constant ( Ke or Kc) to calculate eq m concentrations (14.9 pg ) 3:15. Given : Pb 2+ (aq) + Sn (s) Sn 2+ (aq) + Pb (s) ; Ke = 1.2 * If 0.1 mol of lead nitrate is dissolved in 200 ml of water and a chunk of solid tin added, find the eq'm concentrations of both metallic ions. 3:16. Attempt questions a) b) c) on page 663 ( all perfect squares) 3:17. For the reaction : CO 2 (g) + H 2 (g) CO (g) + H 2 O (g) Ke = 0.02 Suppose that 22 g of CO 2 and 2 g of H 2 are placed in a 2.0 L container. Find the eq'm concentrations of all species and plot a rough conc - time graph. 3:18. Attempt questions a) b) on page 664 ( use the quadratic eq n) 3:19. The following system was allowed to reach eq'm : SO 3 (g) + NO (g) NO 2 (g) + SO 2 (g) The 1st eq'm concentrations were : [NO 2 ] eq m = 0.8 M, [SO 2 ] = 0.12 M, [NO] = 0.4 M, & [SO 3 ] = 0.48 M * NOW an additional 0.52 M SO 3 was added to the system, determine the 2nd eq'm concentrations of all species at this new eq'm state and plot a rough conc - time graph. *Free Energy ( G ) and Equilibrium (18.8 pg 818) Attempt questions a) b) c) d) on page 840 in your text 3:21. Using your data sheet calculate G and then Ke for the reaction 2 NO (g) + 2 CO (g) N 2 (g) + 2 CO 2 (g) at 500 C. 3:22. Consider this reaction : CO (g) + 1/2 O 2 (g) CO 2 (g) a) Calculate G and Ke for the rx at 300 C b) Calculate the temp at which G =0. What is Ke at this temp? c) Discuss the temp range of spontaneity for the rx

3 Chemical Equilibrium: Finding a Constant, Ke ( Kc ) The purpose of this lab is to experimentally determine the equilibrium constant, Kc, for the following chemical reaction: Fe 3+ aq) + SCN - (aq) FeSCN 2+ (aq) iron(iii) thiocyanate ferrocyanate ion When Fe 3+ and SCN - are combined, equilibrium is established between these two ions and the FeSCN 2+ ion. In order to calculate Kc for the reaction, it is necessary to know the concentrations of all ions at equilibrium: [FeSCN 2+ ]eq, [SCN - ]eq, and [Fe 3+ ]eq. You will prepare four equilibrium systems containing different concentrations of these three ions. The equilibrium concentrations of the three ions will then be experimentally determined. These values will be substituted into the equilibrium constant expression to see if Kc is indeed constant. In order to determine [FeSCN 2+ ] eq, you will use the colorimeter shown in Figure 1. The FeSCN 2+ ion produces solutions with a red color. Because the red solutions absorb blue light very well, the blue LED setting on the colorimeter is used. The computer-interfaced colorimeter measures the amount of blue light absorbed by the colored solutions (absorbance, A). By comparing the absorbance of each equilibrium system, Aeq, to the absorbance of a standard solution, Astd, you can determine [FeSCN 2+ ]eq. The standard solution has a known FeSCN 2+ concentration. To prepare the standard solution, a very large concentration of Fe 3+ will be added to a small initial concentration of SCN - (hereafter referred to as [SCN - ]i. The [Fe 3+ ] in the standard solution is 100 times larger than [Fe 3+ ] in the equilibrium mixtures. According to LeChatelier's principle, this high concentration forces the reaction far to the right, using up nearly 100% of the SCN - ions. According to the balanced equation, for every one mole of SCN - reacted, one mole of FeSCN 2+ is produced. Thus [FeSCN 2+ ]std is assumed to be equal to [SCN - ]i. Assuming [FeSCN 2+ ] and absorbance are related directly (Beer's Law), the concentration of FeSCN 2+ for any of the equilibrium systems can be found by: [FeSCN 2+ ]eq = A eq Astd X [FeSCN2+ ]std Knowing the [FeSCN 2+ ] eq allows you to determine the concentrations of the other two ions at equilibrium. For each mole of FeSCN 2+ ions produced, one less mole of Fe 3+ ions will be found in the solution (see the 1:1 ratio of coefficients in the equation on the previous page). The [Fe 3+ ] can be determined by: [Fe 3+ ]eq = [Fe 3+ ]i [FeSCN 2+ ]eq Because one mole of SCN - is used up for each mole of FeSCN 2+ ions produced, [SCN - ] eq can be determined by: [SCN - ]eq = [SCN - ]i [FeSCN 2+ ]eq Knowing the values of [Fe 3+ ] eq, [SCN - ] eq, and [FeSCN 2+ ] eq, you can now calculate the value of K c, the equilibrium constant. PROCEDURE 1. Pipet 5.0 ml of M Fe(NO 3 ) 3 into each of 4 labeled test tubes. CAUTION: Fe(NO 3 ) 3 solutions in this experiment are prepared in 1.0 M HNO 3 and should be handled with care. 2. Pipet 2, 3, 4 and 5 ml of M KSCN into tubes 1-4, respectively. Then pipet 3, 2, 1 and 0 ml of distilled water into tubes 1-4, respectively, to bring the total volume of each test tube to 10 ml. Mix each solution thoroughly with a stirring rod. Be sure to clean and dry the stirring rod after each mixing. Measure and record the temperature of one of the above solutions to use as the temperature for the equilibrium constant, Kc. Volumes added to each test tube are summarized below: Test Tube Number Fe(NO 3 ) 3 (ml) KSCN (ml) H 2 O (ml) Prepare a standard solution of FeSCN 2+ by pipetting 18 ml of M Fe(NO 3 ) 3 into a test tube labeled 5. Pipet 2 ml of M KSCN into the tube. Stir thoroughly.

4 3. You are now ready to calibrate the colorimeter. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use a colorimeter cuvette, remember: All cuvettes should be wiped clean and dry on the outside with a tissue. Handle cuvettes only by the top edge of the ribbed sides. All solutions should be free of bubbles. Always position the cuvette with its reference mark facing toward the white reference mark to the right of the cuvette slot on the colorimeter. Choose Calibrate from the Experiment menu and then click Perform Now. Place the blank cuvette in the cuvette slot of the colorimeter and close the lid. Turn the wavelength knob of the colorimeter to the 0% T position. In this position, the light source is turned off, so no light is received by the photocell. Type 0 in the % edit box. When the displayed voltage reading for Input 1 stabilizes, click Keep. For Reading 2, turn the wavelength knob of the colorimeter to the Blue LED position (470 nm). In this position, the colorimeter is calibrated to show 100% of the blue light being transmitted through the blank cuvette. Type 100 in the % edit box. When the displayed voltage reading for Input 1 stabilizes, click Keep., then click OK. 4. You are now ready to collect absorbance data for the four equilibrium systems and the standard solution. Click Collect. Empty the water from the cuvette and rinse it twice with ~1-mL portions of the Test Tube 1 solution. Then fill the cuvette 3/4 full with the solution from Test Tube 1. Wipe the outside of the cuvette with a tissue and then place the cuvette in the colorimeter. After closing the lid, wait for the absorbance value displayed in the Meter window to stabilize. Then click Keep, type 1 (the trial number) in edit box, and press the ENTER key. 5. Discard the cuvette contents & rinse the cuvette twice with the Test Tube 2 solution and fill the cuvette 3/4 full. Follow the procedure in Step 6 to find and record the absorbance of this solution. Type 2 in the edit box and press ENTER. 6. Repeat the Step 6 procedure to find the absorbance of the solutions in Test Tubes 3, 4, and 5 (the standard solution). 7. From the Table window, record the absorbance values for each of the five trials in your Data and Calculations table. PROCESSING THE DATA 1. In the Data and Calculations table, write the K c expression for the reaction. 2. Calculate the initial concentration of Fe 3+, based on the dilution that results from adding KSCN solution and water to the original M Fe(NO 3 ) 3 solution. See Step 2 of the procedure for the volume of each substance used in Trials 1-4. Calculate [Fe 3+ ] i using the equation: [Fe 3+ ]i = Fe(NO 3)3 ml X ( M) This should be the same for all four test tubes. total ml 3. Calculate the initial concentration of SCN -, based on its dilution by Fe(NO 3 ) 3 and water: [SCN - KSCN ml ]i = X ( M) total ml In Test Tube 1, [SCN - ]i = (2 ml / 10 ml)(.0020 M) = M.. 4. [FeSCN 2+ ] eq is calculated using the formula: [FeSCN 2+ ]eq = A eq Astd X [FeSCN2+ ]std where A eq and A std are the absorbance values for the equilibrium and standard test tubes, respectively, and [FeSCN 2+ ] std = (1/10)(0.0020) = M. Calculate [FeSCN 2+ ] eq for each of the four trials. 5. [Fe 3+ ] eq : Calculate the concentration of Fe 3+ at equilibrium for Trials 1-4 using the equation: [Fe 3+ ]eq = [Fe 3+ ]i [FeSCN 2+ ]eq 6. [SCN - ] eq : Calculate the concentration of SCN - at equilibrium for Trials 1-4 using the equation: [SCN - ]eq = [SCN - ]i [FeSCN 2+ ]eq 7. Calculate Kc for Trials 1-4. Be sure to show the Kc expression and the values substituted in for each of these calculations. 8. Using your four calculated K c values, determine an average value for K c. How constant were your K c values?

5 DATA and CALCULATIONS ( show all your work in the space provided) Absorbance Trial 1 Trial 2 Trial 3 Trial 4 Absorbance of standard (Trial 5) Temperature C K c expression K c = [Fe 3+ ] i [SCN - ] i [FeSCN 2+ ] eq [Fe 3+ ] eq [SCN - ] eq K c value Average of K c values K c = at C

6 Assignment:EQUILIBRIUM(1) 1.For the system 2 SO 2 (g) + O 2 (g) 2 SO 3 (g) H is negative for the production of SO 3 Assume that one has an equilibrium mixture of these substances. Predict the effect of each of the following changes on the number of moles of SO 3 and on the value of the equilibrium constant. Briefly account for each of your predictions. (a) Decreasing the volume of the system. (b) Adding oxygen to the equilibrium mixture. (c) Raising the temperature of the system. 2. Sulfuryl chloride, SO 2 Cl 2, is a highly reactive gaseous compound. When heated, it decomposes as follows: SO 2 Cl 2 (g) SO 2 (g) + Cl 2 (g). This decomposition is endothermic. A sample of grams of SO 2 Cl 2 is placed in an evacuated 1.00 litre bulb and the temperature is raised to 375K. (a) What would be the pressure of this gas (in kpa) in the bulb if no dissociation of the SO 2 Cl 2 (g) occurred? (use PV=nRT) (b) When the system has come to equilibrium at 375K, the total pressure in the bulb (the sum of all the partial pressures of all the gases present) is found to be 145 kpa. Calculate the partial pressures of SO 2, Cl 2, and SO 2 Cl 2 at equilibrium at 375K. (recall Daltons Law of partial pressures) (c) Write the equilibrium constant expression for the decomposition of SO 2 Cl 2 (g) at 375K. Calculate the value of the equilibrium constant (K P ) for this rx. (d) If the temperature were raised to 500K, what effect would this have on the equilibrium constant? Explain briefly. 3. CO 2 (g) + H 2 (g) H 2 O (g) + CO (g) When H 2 (g) is mixed with CO 2 (g) at 2000 K, equilibrium is achieved according to the equation above. In one experiment, the following equilibrium concentrations were measured. [H 2 ] = 0.20 mol/l [CO 2 ]= 0.30 mol/l [H 2 O] = [CO] = 0.55 mol/l (a) Using the equilibrium concentrations given, calculate the value of K e for the reaction. (b) When the system is cooled from 2000 K to a lower temperature, 30.0 percent of the CO(g) is converted back to CO 2 (g). Calculate the value of K e at this lower temperature. (c) In a different experiment, 0.50 mole of H 2 (g) is mixed with 0.50 mole of CO 2 (g) in a 3.0-liter reaction vessel at 2000 K. Calculate the equilibrium concentration, in moles per liter, of CO(g) at this temperature. 4. a) Do question on page 665 b) Draw a rough conc vs time graph to represent the changes that took place in this question. 5. NO (g) + 1/2 O 2 (g) NO 2 (g) a) Calculate G and Ke for this 200 C b) Calculate the temperature at which G=0 or Ke=1 and comment on the spontaneity of this reaction and its temperature range