A buffer solution made fromHNO2 andKNO2 has apH of 3.72. IfPK forpkaHNO2 is 3.35, what is the[NO2]in the buffer?[HNO2][NO2][HNO2] ThepK value forH2CO3 is 6.38. What mole ratio ofNaHCO3 toH2CO3 is needed to prepare a buffer with a pH of 6.09?[HCO3][H2CO3]

Step 1: Buffer solutionsAqueous solutions containing a mixture of a weak acid and its conjugate base…

Answered: A buffer solution made from HNO2 and…

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 3ALQ: Mixing together solutions of acetic acid and sodium hydroxide can make a buffered solution. Explain....

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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?
Each symbol in the box below represents a mole of a component in one liter of a buffer solution; represents the anion (X-), = the weak acid (HX), = H+, and =OH. Water molecules and the few H+ and OH- ions from the dissociation of HX and X- are not shown. The box contains 10 mol of a weak acid, , in a liter of solution. Show what happens upon (a) the addition of 2 mol of OH- (2 ). (b) the addition of 5 mol of OH- (5 ). (c) the addition of 10 mol of OH- (10 ). (d) the addition of 12 mol of OH- (12 ). Which addition (a)-(d) represents neutralization halfway to the equivalence point?
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.
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 monoprotic organic acid that has a molar mass of 176.1 g/mol is synthesized. Unfortunately, the acid produced is not completely pure. In addition, it is not soluble in water. A chemist weighs a 1.8451-g sample of the impure acid and adds it to 100.0 mL of 0.1050 M NaOH. The acid is soluble in the NaOH solution and reacts to consume most of the NaOH. The amount of excess NaOH is determined by titration: It takes 3.28 mL of 0.0970 M HCl to neutralize the excess NaOH. What is the purity of the original acid, in percent?
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?
The weak base ethanolamine. HOCH2CH2NH2, can be titrated with HCl. HOCH2CH2NH2(aq)+H3O+(aq)HOCH2CH2NH3+(aq)+H2O(l) Assume you have 25.0 mL of a 0.010 M solution of ethanolamine and titrate it with 0.0095 M HCl. (Kb for ethanolamine is 3.2 107.) (a) What is the pH of the ethanolamine solution before the titration begins? (b) What is the pH at the equivalence point? (c) What is the pH at the halfway point of the titration? (d) Which indicator in Figure 17.11 would be the best choice to detect the equivalence point? (e) Calculate the pH of the solution after adding 5.00, 10.0, 20.0, and 30.0 mL of the acid. (f) Combine the information in parts (a), (b), (c), and (e), and plot an approximate titration curve.
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 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
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.
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?
What is the pH of a buffer that is 0.175 M in a weak acid and 0.200 M in the acids conjugate base? The acids ionization constant is 5.7 104.

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A buffer solution made from HNO2 and KNO2 has a pH of 3.72. If PK for pka HNO2 is 3.35, what is the [NO2] in the buffer? [HNO2] [NO2] [HNO2] The pK value for H2CO3 is 6.38. What mole ratio of NaHCO3 to H2CO3 is needed to prepare a buffer with a pH of 6.09? [HCO3] [H2CO3]