Determine the pH of a buffer solution by constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. The Ka for HNO2 is 6.8 x 104. Complete Parts 1-3 before submitting your answer. 1 2 3 NEXT > A solution is prepared with 0.55 M HNO2 and 0.75 M KNO2. Fill in the ICE table with the appropriate value for each involved species to determine concentrations of all reactants and products. Initial (M) Change (M) Equilibrium (M) HNO2(aq) H₂O(1) H3O+(aq) NO₂ (aq) RESET 0 0.55 0.75 6.8 × 10-4 +x -x +2x -2x 0.55 + x 0.55-x 0.55 + 2x 0.55 -2x 0.75 + x 0.75-x 0.75 + 2x 0.75 -2x MacBook Air Determine the pH of a buffer solution by constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. The Ka for HNO2 is 6.8 x 104. Complete Parts 1-3 before submitting your answer. < PREV 1 2 3 Based on your ICE table (Part 1) and the equilibrium expression for Ka (Part 2), determine the pH of the buffer solution. pH = RESET 0 5.0 × 10-4 0.55 3.30 7.61 10.70 6.39 0.26 MacBook Air

Determine the pH of a buffer solution by constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. The Ka for HNO2 is 6.8 x 104. Complete Parts 1-3 before submitting your answer. 1 2 3 NEXT > A solution is prepared with 0.55 M HNO2 and 0.75 M KNO2. Fill in the ICE table with the appropriate value for each involved species to determine concentrations of all reactants and products. Initial (M) Change (M) Equilibrium (M) HNO2(aq) H₂O(1) H3O+(aq) NO₂ (aq) RESET 0 0.55 0.75 6.8 × 10-4 +x -x +2x -2x 0.55 + x 0.55-x 0.55 + 2x 0.55 -2x 0.75 + x 0.75-x 0.75 + 2x 0.75 -2x MacBook Air Determine the pH of a buffer solution by constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. The Ka for HNO2 is 6.8 x 104. Complete Parts 1-3 before submitting your answer. < PREV 1 2 3 Based on your ICE table (Part 1) and the equilibrium expression for Ka (Part 2), determine the pH of the buffer solution. pH = RESET 0 5.0 × 10-4 0.55 3.30 7.61 10.70 6.39 0.26 MacBook Air

Chemistry: Principles and Practice

3rd Edition

ISBN:9780534420123

Author:Daniel L. Reger, Scott R. Goode, David W. Ball, Edward Mercer

Publisher:Daniel L. Reger, Scott R. Goode, David W. Ball, Edward Mercer

Chapter16: Reactions Between Acids And Bases

Section: Chapter Questions

Problem 16.89QE

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A 10.00-g sample of the ionic compound NaA, where A is the anion of a weak acid, was dissolved in enough water to make 100.0 mL of solution and was then titrated with 0.100 M HCl. After 500.0 mL HCl was added, the pH was 5.00. The experimenter found that 1.00 L of 0.100 M HCl was required to reach the stoichiometric point of the titration. a. What is the molar mass of NaA? b. Calculate the pH of the solution at the stoichiometric point of the titration.
You have a solution of the weak acid HA and add some of the salt NaA to it. What are the major species in the solution? What do you need to know to calculate the pH of the solution, and how would you use this information? How does the pH of the solution of just the HA compare with that of the final mixture? Explain.
Determine the dominant acid-base equilibrium that results when each of the following pairs of solutions is mixed. Indicate the equilibrium by writing 1 for a strong acid, 3 for a weak acid, 4 for an acidic buffer, 7 for a neutral solution, 10 for a basic buffer, 11 for a weak base, and 13 for a strong base. (a) 10.0 mL of 0.15 M NaOH + 15.0 mL of 0.10 M HNO3 (b) 25.0 mL of 0.10 M HCl + 10.0 mL of 0.25 M NH3 (c) 50.0 mL of 0.050 M NaOH + 50.0 mL of 0.10 M NH3 (d) 50.0 mL of 0.10 M NH3 + 50.0 mL of 0.05 M HCl
A buffer is prepared in which the ratio [ H2PO4 ]/[ HPO42 ]is 3.0. (a) What is the pH of this buffer? (b) Enough strong acid is added to convert 15% of HPO42- to H2PO4-. What is the pH of the resulting solution? (c) Enough strong base is added to make the pH 7.00. What is the ratio of [H2PO4-] to [HPO42-] at this point?
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?
A solution is made by diluting 25.0 mL of concentrated HCl (37% by weight; density = 1.19 g/mL) to exactly 500 mL. Calculate the pH of the resulting solution.
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?
Ammonia gas is bubbled into 275 mL of water to make an aqueous solution of ammonia. To prepare a buffer with a pH of 9.56, 15.0 g of NH4Cl are added. How many liters of NH3; at 25C and 0.981 atm should be used to prepare the buffer? Assume no volume changes and ignore the vapor pressure of water.
Ka for formic acid is 1.7 104 at 25C. A buffer is made by mixing 529 mL of 0.465 M formic acid, HCHO2, and 494 mL of 0.524 M sodium formate, NaCHO2. Calculate the pH of this solution at 25C after 110 mL of 0.152 M HCl has been added to this buffer.
Sketch a titration curve for the titration of potassium hydroxide with HCl, both 0.100 M. Identify three regions in which a particular chemical species or system dominates the acid-base equilibria.
You have a solution of the weak acid HA and add some HCl to it. What are the major species in the solution? What do you need to know to calculate the pH of the solution, and how would you use this information? How does the pH of the solution of just the HA compare with that of the final mixture? Explain.
You want to make a buffer with a pH of 10.00 from NH4+/NH3. (a) What must the [ NH4+ ]/[ NH3 ]ratio be? (b) How many moles of NH4Cl must be added to 465 mL of an aqueous solution of 1.24 M NH3 to give this pH? (c) How many milliliters of 0.236 M NH3 must be added to 2.08 g of NH4Cl to give this pH? (d) What volume of 0.499 M NH3 must be added to 395 mL, of 0.109 M NH4Cl to give this pH?