Problem 2: Your facilitator asked you to prepare 450 mL of a buffer solution having a pH of 10.45 composed of Ammonia and Ammonium Chloride. Considering a final salt concentration of 0.123 M, determine the following. pkg = 3.69 Atomic weights: N: 14.01 H: 1.01 CI: 35.45 O: 15.99 63. What is the volume of Ammonia, in mL, needed to prepare the buffer solution? 64. What is the ratio of Ammonium chloride and Ammonia? 65. What is the weight, in grams, of Ammonium chloride? 66. What is the concentration of the Hydroxyl ion? 67. Determine the kb value. 68. Determine the mmoles of the salt present in the solution.

Problem 2: Your facilitator asked you to prepare 450 mL of a buffer solution having a pH of 10.45 composed of Ammonia and Ammonium Chloride. Considering a final salt concentration of 0.123 M, determine the following. pkg = 3.69 Atomic weights: N: 14.01 H: 1.01 CI: 35.45 O: 15.99 63. What is the volume of Ammonia, in mL, needed to prepare the buffer solution? 64. What is the ratio of Ammonium chloride and Ammonia? 65. What is the weight, in grams, of Ammonium chloride? 66. What is the concentration of the Hydroxyl ion? 67. Determine the kb value. 68. Determine the mmoles of the salt present in the solution.

Chemistry: Principles and Reactions

8th Edition

ISBN:9781305079373

Author:William L. Masterton, Cecile N. Hurley

Publisher:William L. Masterton, Cecile N. Hurley

Chapter14: Equilibria In Acid-base Solutions

Section: Chapter Questions

Problem 75QAP: A diprotic acid, H2B(MM=126g/moL), is determined to be a hydrate, H2B xH2O. A 10.00-g sample of this...

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A diprotic acid, H2B(MM=126g/moL), is determined to be a hydrate, H2B xH2O. A 10.00-g sample of this hydrate is dissolved in enough water to make 150.0 mL of solution. Twenty-five milliliters of this solution requires 48.5 mL of 0.425 M NaOH to reach the equivalence point. What is x?
A friend asks the following: Consider a buffered solution made up of the weak acid HA and its salt NaA. If a strong base like NaOH is added, the HA reacts with the OH to form A. Thus the amount of acid (HA) is decreased, and the amount of base (A) is increased. Analogously, adding HCI to the buffered solution forms more of the acid (HA) by reacting with the base (A). Thus how can we claim that a buffered solution resists changes in the pH of the solution? How would you explain buffering to this friend?
Give typed solution only
Be sure to answer all parts. Scenes A to D represent tiny portions of 0.10 M aqueous solutions of a weak acid HA (red and blue; K-4.4 x 10-5), its conjugate base A (red), or a mixture of the two (only these species are shown): (a) Which scene(s) show(s) a buffer? Scene A Scene B Scene Scene D (b) What is the pH of each solution? Scene A: 4.36 Scene B: Scene C: Scene D: 4.52 B (c) Arrange the scenes (by letter) in sequence, assuming that they represent stages in a weak acid-strong base titration.
#6 Calculate the pH during the titration of 40.00 mL of 0.1000 M HCl with 0.1000 M NaOH solution after the following additions of base: (a) 23.00 mL pH = (b) 39.40 mL pH = (c) 54.00 mL pH =
Show All Work. Please 4) Determine whether aqueous solutions of the following salts are acidic, basic, or neutral: (a) FeCl3 (b) K2CO3 (c) NH4Br (d) KClO4 5) Calculate the concentration of each species present in a 0.050-M solution of H2S. (Ka1 = 1.1 x 10^-7 and Ka2 = 6.2 x 10^-8) 6) Calculate the pH of a buffer solution prepared from 0.155 mol of phosphoric acid, (Ka1= 7.52 x 10^-3 just use Ka1) 0.250 mole of KH2PO4, and enough water to make 0.500 L of solution. 7) Calculate the pH at the following points in a titration of 40 mL (0.040 L) of 0.100 M barbituric acid (Ka = 9.8 × 10−5) with 0.100 M KOH. (a) no KOH added (b) 20 mL of KOH solution added (c) 39 mL of KOH solution added (d) 40 mL of KOH solution added (e) 41 mL of KOH solution added
Question 5. Listen Consider the following buffer system: 0.71 M H3PO4 and 0.64 M H₂PO4-. What is the pH of such a system (K₂ = 7.5 x 10-³)? Your Answer: Answer
3
O ADVANCED MATERIAL 3/5 Calculating the pH of a buffer A solution is prepared at 25 °C that is initially 0.42M in ammonia (NH,), a weak base with K, = 1.8 x 10°, and 0.018 M in ammonium bromide (NH,Br). Calculate the pH of the solution. Round your answer to 2 decimal places. pH =
How many moles of B (weak base) would need to be added to 507 mL of 0.28 M BHCl (the acidic salt) to produce a solution with a pH equal to 9.05? Report your answer to the thousandths place. The Kb of the weak base (B) is 4.7 × 10-6 Please explain each step, i want to learn how to solve it myself!
100 mL of 0.10 M NH3 (Kb = 1.8 x 10-5) are mixed with 100 mL of 0.10 M HCN (Ka = 4.9 × 10-10). Which statement is true? (3) The resulting solution is neutral O (2) The resulting solution is basic O (4) No answer is correct as not enough information is given. O (1) The resulting solution is acidic E D Moving to the next question prevents changes to this answer. 4 R F % S T G 6. & 7 H C 8 00 K M Question 2 of 37
Part A What is the pH of a buffer prepared by adding 0.708 mol of the weak acid HA to 0.608 mol of Na.A in 2.00 L of solution? The dissociation constant K of HA is 5.66 x 10-7. Express the pH numerically to three decimal places. ▸ View Available Hint(s) pH = 6.181 Submit Previous Answers Correct Since both the acid and base exist in the same volume, we can skip the concentration calculations and use the number of moles in the Henderson- Hasselbalch equation to calculate the pH. The answer will be the same. Part B What is the pH after 0.150 mol of HCl is added to the buffer from Part A? Assume no volume change on the addition of the acid. Express the pH numerically to three decimal places. ▸ View Available Hint(s) 195| ΑΣΦ pH= 6.124 Submit Previous Answers Request Answer * Incorrect; Try Again; 5 attempts remaining