Plz give both answer A buffer solution is made that is 0.482 M inHNO2 and 0.482 M inKNO₂.IfKa forHNO2 is4.50 x 10-4, what is the pH of the buffersolution?pHWrite the net ionic equation for thereaction that occurs when 0.130 mol HI isadded to 1.00 L of the buffer solution.(Use the lowest possible coefficients.Omit states of matter. Use H3O+ insteadof H+)+A buffer solution is made that is 0.396 M inH₂CO3 and 0.396 M inKHCO3.IfKal forH₂CO3 is4.20 x 10-7, what is the pH of the buffersolution?pH =Write the net ionic equation for thereaction that occurs when 0.082 mol HNO3is added to 1.00 L of the buffer solution.(Use the lowest possible coefficients.Omit states of matter. Use H3O+ insteadof H+)

A buffer solution is made that is 0.482 M in HNO2 and 0.482 M in KNO2. If Ka for HNO2 is 4.50 x 10-4, what is the pH of the buffer solution? pH = Write the net ionic equation for the reaction that occurs when 0.130 mol HI is added to 1.00 L of the buffer solution. (Use the lowest possible coefficients. Omit states of matter. Use H3O+ instead of H+) pH + A buffer solution is made that is 0.396 M in H₂CO3 and 0.396 M in KHCO3. If Kal for H₂CO3 is 4.20 x 10-7, what is the pH of the buffer solution? = + Write the net ionic equation for the reaction that occurs when 0.082 mol HNO3 is added to 1.00 L of the buffer solution. (Use the lowest possible coefficients. Omit states of matter. Use H30+ instead of H+)

A buffer solution is made that is 0.482 M in HNO2 and 0.482 M in KNO2. If Ka for HNO2 is 4.50 x 10-4, what is the pH of the buffer solution? pH = Write the net ionic equation for the reaction that occurs when 0.130 mol HI is added to 1.00 L of the buffer solution. (Use the lowest possible coefficients. Omit states of matter. Use H3O+ instead of H+) pH + A buffer solution is made that is 0.396 M in H₂CO3 and 0.396 M in KHCO3. If Kal for H₂CO3 is 4.20 x 10-7, what is the pH of the buffer solution? = + Write the net ionic equation for the reaction that occurs when 0.082 mol HNO3 is added to 1.00 L of the buffer solution. (Use the lowest possible coefficients. Omit states of matter. Use H30+ instead of H+)

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter15: Acid-base Equilibria

Section: Chapter Questions

Problem 2RQ: Define a buffer solution. What makes up a buffer solution? How do buffers absorb added H+ or OH with...

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Sketch the titration curve for a weak acid titrated by a strong base. When performing calculations concerning weak acidstrong base titrations, the general two-slep procedure is to solve a stoichiometry problem first, then to solve an equilibrium problem to determine the pH. What reaction takes place in the stoichiometry part of the problem? What is assumed about this reaction? At the various points in your titration curve, list the major species present after the strong base (NaOH, for example) reacts to completion with the weak acid, HA. What equilibrium problem would you solve at the various points in your titration curve to calculate the pH? Why is pH 7.0 at the equivalence point of a weak acid-strong base titration? Does the pH at the halfway point to equivalence have to be less than 7.0? What does the pH at the halfway point equal? Compare and contrast the titration curves for a strong acidstrong base titration and a weak acidstrong base titration.
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.
A good buffer generally contains relatively equal concentrations of weak acid and conjugate base. If you wanted to buffer a solution at pH = 4.00 or pH = 10.00, how would you decide which weak acidconjugate base or weak baseconjugate acid pair to use? The second characteristic of a good buffer is good buffering capacity. What is the capacity of a buffer? How do the following buffers differ in capacity? How do they differ in pH? 0.01 M acetic acid/0.01 M sodium acetate 0.1 M acetic acid/0.1 M sodium acetate 1.0 M acetic acid/1.0 M sodium acetate
A buffer solution is prepared by adding 5.50 g of ammonium chloride and 0.0188 mol of ammonia to enough water to make 155 mL of solution. (a) What is the pH of the buffer? (b) If enough water is added to double the volume, what is the pH of the solution?
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?
Define a buffer solution. What makes up a buffer solution? How do buffers absorb added H+ or OH with little pH change? Is it necessary that the concentrations of the weak acid and the weak base in a buffered solution be equal? Explain. What is the pH of a buffer when the weak acid and conjugate base concentrations are equal? A buffer generally contains a weak acid and its weak conjugate base, or a weak base and its weak conjugate acid, in water. You can solve for the pH by setting up the equilibrium problem using the K.a reaction of the weak acid or the Kb reaction of the conjugate base. Both reactions give the same answer for the pH of the solution. Explain. A third method that can be used to solve for the pH of a buffet solution is the HendersonHasselbalch equation. What is the HendersonHasselbalch equation? What assumptions are made when using this equation?
When 40.00 mL of a weak monoprotic acid solution is titrated with 0.100-M NaOH, the equivalence point is reached when 35.00 mL base has been added. After 20.00 mL NaOH solution has been added, the titration mixture has a pH of 5.75. Calculate the ionization constant of the acid.
Sketch two pH curves, one for the titration of a weak acid with a strong base and one for a strong acid with a strong base. How are they similar? How are they different? Account for the similarities and the differences.
What mass of NH4Cl must be added to 0.750 L of a 0.100-M solution of NH3 to give a buffer solution with a pH of 9.26? (Him: Assume a negligible change in volume as the solid is added.)
A solution consisting of 25.00 g NH4Cl in 178 mL of water is titrated with 0.114 M KOH. (a) How many milliliters of KOH are required to reach the equivalence point? (b) Calculate [Cl-], [K+], [NH3], and [OH-] at the equivalence point. (Assume that volumes are additive.) (c) What is the pH at the equivalence point?
A buffer is prepared by dissolving 0.0250 mol of sodium nitrite, NaNO2, in 250.0 mL of 0.0410 M nitrous acid, HNO2. Assume no volume change after HNO2 is dissolved. Calculate the pH of this buffer.
Calculate the pH of a buffer solution prepared from 0.155 mol of phosphoric acid, 0.250 mole of KH2PO4, and enough water to make 0.500 L of solution.