Intro to Buffers Practice Problems
Determine the solution with the greatest buffer capacity among the representations of solutions shown below containing one or more of the following components: H2A, KHA, and K2A, where H2A is a weak diprotic acid. Note that K+ ions and water were omitted from the representation for clarity.
Determine what happens to the pH (increases, decreases, or remains the same) of a buffer solution containing equal concentrations of C6H7O7− and C6H6O72−on the addition of each of the following:
Identify a mixture from the following solution pairs where equal volumes of the two solutions, when mixed together, will lead to the formation of a buffer.
a) 0.13 M HCl with 0.15 M NaOH
b) 0.13 M H2SO4 with 0.13 M Na2CO3
c) 0.26 M HCN with 0.13 M KOH
d) 0.10 M NH4Cl with 0.20 M HCl
The following diagram represents four different solutions that may contain any of these four species in varying concentrations: H2A, KHA, and K2A.
K+ ions are not shown for clarity. Identify which of these solutions is a buffer.
Identify which of the following two solutions has a higher buffer capacity:
a) 150.0 mL of 0.15 M KCN-0.15 M HCN
b) 150.0 mL of 0.15 M C2H3O2-0.10 M C2H4O2
Explain your choice.
Consider two buffers that are prepared by using an equal number of moles of hypochlorous acid (HClO) and sodium hypochlorite (NaClO), adding it to water to make a 2.00 L of buffer solution. In Buffer X, 1.50 mol each of hypochlorous acid and sodium hypochlorite was used. In Buffer Y, 0.150 mol of each was used. Identify the buffer that would have a higher buffering capacity.
The image below shows a buffer with equal amounts of weak acid, HA, and its conjugate base, A–. If the column heights pertain to the amount of component present, select the image showing a scenario that does not occur because of the addition of a base or an acid
The image below shows a buffer with equal amounts of weak acid, HA, and its conjugate base, A–. If the column heights pertain to the amount of component present, select the image that shows the buffer after a strong acid is added.
The flask on the left has a mixture of 0.2 M hypochlorous acid and 0.2 M potassium hypochlorite with bromthymol blue. The flask on the right has a solution of 0.2 M hypochlorous acid with bromthymol blue. Choose the solution that is able to resist drastic changes in pH even if small amounts of KOH are added. Explain. (Ka hypochlorous acid = 2.9×10–8)
A particular system uses a buffer consisting of HSO4- and SO42-. Can this be replaced with a buffer consisting of H2SO4 and HSO4-? Briefly explain.
Identify whether the solution of the following compounds results in a buffer:
250.0 mL 0.25 M CHCOOH and 105.0 mL of 0.050 M HCl
Write equations showing how a buffer containing CH3NH2 and CH3NH3Br neutralizes acid and base.
Does the addition of 2.56 g HCl exceed the capacity of the 250.0 mL buffer solution that is 0.125 M of HF and 0.175 M of NaF.
Does the addition of 1.5 g of KOH exceed the capacity of the 250.0 mL buffer solution that is 0.125 M of HF and 0.175 M of NaF.
Select the statements that correctly describe buffers.
a) The pH of a buffer solution does not change significantly when any amount of a strong acid is added.
b) An acid added to the buffer solution reacts with the weak base of the buffer.
c) The pH of a buffer solution is determined by the ratio of the concentration of conjugate base to the concentration of strong acid.
d) A buffer is generally made up of a weak acid and its conjugate base.
e) The Ka of a buffer does not change when any amount of an acid is added to the buffer solution.
A buffer contains significant amounts of acetic acid and sodium acetate.
a. Write a molecular equation showing how this buffer neutralizes added acid HNO3.
b. Write a molecular equation showing how this buffer neutralizes added base KOH.
How would you prepare a buffer solution made of HC7H5O2 and NaC7H5O2 with a pH of 4.10? (Ka HC7H5O2 = 6.5×10−5)
Calculate the pH of a buffer consisting of 0.130 mol of propanoic acid and 0.160 mol of sodium propanoate in 1.00 L. The pKa of propanoic acid is 4.87.
0.125 M HCN and 0.140 M NaCN are mixed to produce a 125.0 mL buffer solution. Determine the pH of the solution if 125.0 mg of HCl was added. (Ka HCN = 4.90×10−10)
0.125 M HCN and 0.140 M NaCN are mixed to produce a 125.0 mL buffer solution. Determine the initial pH of the solution. (Ka HCN = 4.90×10−10)
Determine the pH for the following solution: 25.0 mL 0.17 M HIO with 35.0 mL 0.19 M NaIO (Ka HIO = 2.3×10−11)
A 0.275 L sample of HBr gas was dissolved in a solution at 25°C and 730 mmHg. The solution initially contained 1.740 g of Na2HPO4 and 0.125 L of water. Calculate the pH of the resulting solution if the pKa of dihydrogen phosphate is 7.21.
The ratio of hydrogen phosphate to dihydrogen phosphate in renal tubular and intracellular fluids is maintained at 1.55 by the body. The phosphate buffer system works in the internal fluids of all cells. Extra hydroxide and hydrogen ions that enter the intracellular fluid are neutralized by this buffer. Calculate the pH of the phosphate buffer system (Ka for dihydrogen phosphate at room temperature is 6.17×10−8).
Calculate the ratio of the volumes of 1.25 M sodium glycolate (NaC2H3O3) and 1.25 M glycolic acid (C2H4O3) necessary to make a buffer of pH 4.25. The pKa of C2H4O3 is 3.83.
Calculate the pH of a buffer solution that is 0.275 M glycolic acid (C2H4O3) and 0.550 M sodium glycolate (NaC2H3O3). The pKa of C2H4O3 is 3.83.
Find the pH of a solution of 0.250 M HF and 0.200 M KF before and after dilution by a factor of 5. The pKa of HF is 3.20.
For a buffer of pH 9.5, which of the following weak acid-conjugate base mixtures would you choose? Explain your choice.
a) CH3COOH and NaCH3COO
b) NH4Cl and NH3
c) HCOOH and NaHCOO
An equation analogous to the Henderson-Hasselbalch equation can be derived, which relates the pOH of the buffer to the pKb of the base component. Which of the following is the correct equation?
A buffer solution with a pH of 7.20 needs to be prepared. If the following 0.20 M solutions are available in the laboratory: HF, NaF, HClO, CH3COOH, NaClO, CH3COONa. Choose a combination that would be best to use to prepare the buffer.
Ka HF = 6.8×10–4
Ka HClO = 3.0×10–8
Ka CH3COOH = 1.8×10–5
A buffer solution is prepared using HF and F–. HF has a pKa of 3.16 and the buffer has a pH of 3.35. Determine which of the following is applicable at pH 3.35 without performing any calculations.
Calculate the concentrations of hydrofluoric acid and sodium fluoride in buffer solution with a pH of 3.57. The solution was measured to have freezing point of -5.50 °C and a density of 1.02 g/mL.
Identify whether to add KOH or HBr to the buffer mixture to adjust the pH to 9.56. The buffer solution of 1-L initially contains 0.500 M in CH3NH2 and 0.500 M in CH3NH3Br. Calculate the mass of the correct reagent needed.
Consider a 250.0 mL solution of 0.20 M propanoic acid. Calculate the mass of sodium propanoate needed to produce a buffer with a pH of 5.11 assuming the volume does not change.
Consider a solution that contains 25 g of CH3COOH and 15 g of NaCH3COO in a 250.0 mL of solution
Calculate the pH using the Henderson–Hasselbalch equation.
If you need to prepare a buffer with a pH of 8.8, which one of the following acid/conjugate base pairs would you choose?
Acetic acid is a weak acid that dissociates into acetate ion and hydronium ions in solutions:
CH3COOH(aq) + H2O(l) ⇌ H3O+(aq) + CH3COO-(aq)
The value of the dissociation constant (Ka) for acetic acid is 1.78×10-5. If you require a buffer of pH 5.25, what would be the ratio of [CH3COO-]/[CH3COOH] that you would use?
A student in an analytical chemistry laboratory needs to prepare a buffer of pH 5.00. She chooses 2-methylpropanoic acid (pKa = 4.84) to prepare the buffer. She has access to 0.100 M 2-methylpropanoic solution and 0.120 M sodium 2-methylpropanoate solution. How much volume (in mL) of each solution does she require to prepare a total of 50.0 mL of her required buffer?
In the titration of 0.2011 g of an unknown monoprotic acid, 25.3 mL of 0.0650 M KOH solution was used to reach the endpoint. When the volume of the base added is 12.65 mL, the pH of the solution is 4.20. Determine the Ka for the acid.
Determine the predominant form of B for the titration of a 25.0 mL 0.15 M solution of a weak monoprotic base, B, with a 0.15 M solution of a strong monoprotic acid, HA, at the equivalence point
Using mathematical proof, identify if the following statement is True or False: At the half-equivalence point of a titration of a weak acid with a strong base, the pH is equal to the pKa for the acid.
At the half-equivalence point, what is the pH of the titration of 15 mL of 0.183 M aniline with 0.15 M HNO3?
At the half-equivalence point, what is the pH of the titration of 20 mL of 0.178 M lactic acid with 0.150 M NaOH?
Determine the volume of the added acid where the pH = 14 - pKb for the titration of a weak monoprotic base and a strong acid based on the given curve
Determine the volume of the added base where the pH will be equal to the pKa for the titration of a weak monoprotic acid and a strong base based on the given curve
A sample mixture of HCl and H3AsO4, with a volume of 50 mL, was titrated against 0.250 M NaOH. What indicators would be suitable to signal the equivalence points if the first and second equivalence points were reached after 95 mL and 140 mL of the base were consumed, respectively?
A certain liquid has been analyzed and demonstrated a pH change from 6 to 8. Choose the most appropriate indicator to detect this pH change based on the figure below.
The pH of a substance changes from 1 to 3. Determine the most suitable indicator to detect this pH change based on the figure below.
Bromcresol purple is a weak acid and can act as an acid-base indicator. The yellow acidic form is equal in amount to the purple base form when the pH is equal to 6.3. Determine the pKa for bromcresol purple.
The indicator bromthymol blue was added in each of the solutions shown below. Solution A contains 0.020 M hydrocyanic acid. Solution B contains a mixture of 0.020 M hydrocyanic acid and 0.020 M sodium cyanide. Using the figures below as reference, identify the solution with the lower pH.
Drops of bromthymol blue were added to the two solutions shown below. Based on the results, determine if the following statement is true or false: Solution (ii) has a pH of greater than 7.00.
An unknown colorless solution turns red when paramethyl red is added. It turns orange when erythrosine is added. Choose another indicator that will give a more precise value for the pH of the solution
a. Metacresol Purple
c. Quinaldine Red
d. Ethyl Red
An unknown solution turns yellow when bromocresol green is added. It turns red when propyl red is added. Which information can be established about the solution?
An unknown solution turns blue when methyl green is added. It also turns blue when congo red is added. Which range of pH values could the solution be?
Using the figure below, determine which indicator is appropriate for the titration of C17H19NO3 with a strong acid.
Using the figure below, determine which indicator is appropriate for the titration of HCN with a strong base.
Bromophenol blue is an acid-base indicator with a pKa of 4.0. Its color changes from yellow to blue going from its acidic form to its basic form. What color will the solution appear when several drops of this indicator are placed in a 50.0 mL of 0.450 M HBr solution? What pH range will the indicator change color when 0.150 M KOH is slowly added to the solution?
Consider the following situation:
A student was asked to titrate a weak acid with a strong base. The student then started slowly adding the strong base to a fixed amount of weak acid, however, he remembered that he didn't put the indicator. The indicator was supposed to change from colorless to pink when the titration is complete. The student then tried adding a prescribed amount of the indicator to the solution that they already started titrating. The solution did not change color and was still colorless. What should the student do next?
The following is a titration curve for the titration of a weak acid with a strong base. The equivalence point for the titration occurs at pH 9.5. Identify which of the following indicators would be suitable for this particular titration.
A. Ethyl red pKa = 5.42
B. Metacresol purple pKa1 = 1.51, pKa2 = 8.32
C. Phenolphthalein pKa = 9.50
D. Quinaldine red pKa = 2.63
E. Bromothymol blue pKa = 7.0
F. Thymol blue pKa1 = 1.65, pKa2 = 9.20
Two white solids, CaCO3 and BaCO3, are obtained in an analysis. What tests will you perform to determine which is BaCO3?
Provide a method for separating Zn2+ and K+ ions present in a solution by adding only one or two substances.
Based on the following qualitative analysis flowchart, how can you separate Bi3+ ion and Ba2+ ion contained in a solution?
Based on the quantitative analysis, the addition of H2S can separate Pb2+, Cu2+, and Cd2+ from other cations present in a solution. What is the S2− ion concentration needed to start the precipitation of (i) Pb2+, (ii) Cu2+, and (iii) Cd2+ ions if each of the metal ions has a concentration of 0.020 M? When Cu2+ starts to precipitate, what fraction of Cd2+ is still left in the solution?
Ksp PbS = 7.0×10−29, Ksp CuS = 8.0×10−37, Ksp CdS = 1.0×10−28
Based on the given flowchart, what is the difference between the group 2 sulfide precipitates and group 3 sulfide precipitates?
A solution contains both 0.015 M Mn2+ and 0.015 M Ni2+. If K2CO3 is added dropwise to the solution, what is the molarity of CO32– when the second cation starts to precipitate? (Ksp MnCO3 = 2.24×10–11, Ksp NiCO3 = 1.42×10–7)
Potassium carbonate is used to precipitate the cations in a solution that contains 0.025 M Mn2+ and 0.050 M Zn2+. When the second cation starts to precipitate, what would be the remaining concentration of the cation that precipitates first?
Determine the cation that will precipitate first if potassium carbonate is used to precipitate one of the cations in a solution that contains 0.025 M Mn2+ and 0.050 M Zn2+. Calculate the minimum concentration of K2CO3 that will start the precipitation of the cation.
A solution of two silver salts contains 0.0111 M Cl- and 0.0125 M Br-. Upon addition of a very dilute solution of silver nitrate (AgNO3), a precipitate is formed.
A. Identify the precipitate as AgCl or AgBr.
B. Calculate the minimum concentration of Ag+ at which the precipitate begins to form.
Ksp(AgCl) = 1.8×10–10
Ksp(AgBr) = 5.0×10–13
When KOH(aq) is added to the solution, will the solubility of Ni(OH)2 increase, decrease, or remain the same? Give the balanced net ionic equation for the dissolution process. [Hint: Refer to complex ions formations]
What will happen to the solubility of Fe(OH)3 when KCN(aq) is added to the solution? Provide the balanced net ionic equation for the dissolution process. [Hint: Refer to complex ions formations]
Determine the molar solubility of copper(II) carbonate (CuCO3) in a 0.350 M solution of ethylenediamine (abbreviated as en). Use the following data:
Kf for [Cu(en)2]2+ = 1.00×1020
Ksp for CuCO3 = 2.40×10−10
Pyrophosphate (P2O74−) is added to laundry detergents to prevent the precipitation of insoluble calcium and magnesium salts. Consider a scenario where Na2H2P2O7 is added to the solution and we're trying to dissolve Ca(OH)2 salts because Ca2+ ions form a soluble complex with pyrophosphate.
(a) Given that the Ksp for Ca(OH)2 is 4.7×10–6 and the Kf for [Ca(P2O7)]2− is 4.0×104, determine the equilibrium constant (Keq) for the reaction.
(b) What is the molar solubility of Ca(OH)2 in 2.50 M Na2H2P2O7 solution?
[Ka2 H2P2O72– = 2.51×10–7; Ka3 HP2O73– = 3.89×10–10]
What is the molar solubility of AgCN in 1.00 M NaCN solution? Use the following values: Ksp AgCN = 6.0×10–17 and Kf [Ag(CN)2]– = 3.0×1020.
What is the molar solubility of AgCl in pure water and in 0.20 M Na2S2O3 solution? Use the following values: Ksp AgCl = 1.8×10–10 and Kf [Ag(S2O3)2]3– = 4.7×1013.
Provide the balanced net ionic equation for the dissolution reaction of Ni(OH)2 in aqueous NH3 forming [Ni(NH3)6]2+. Determine the equilibrium constant for the reaction given that the Ksp for Ni(OH)2 is 5.5×10–16 and the Kf for [Ni(NH3)6]2+ is 2.0×108.
What is the net ionic equation for the dissolution of Cu(OH)2 in NH3? Ksp for Cu(OH)2 is 1.60×10–19 and Kf for [Cu(NH3)4]2+ is 5.60×1011. Determine the equilibrium constant for the dissolution reaction.
When 2.50×10–3 mol of CuCl are added to 450.0 mL of a 0.250 M KCN solution, a complex ion with the formula [Cu(CN)2]– is formed. Kf for the complex ion is 1.00×1016. What is the fraction of uncomplexed Cu+ ions in the solution?
Calculate the solubility of AgCl (in mol/L) in 0.15 M KSCN.
Ksp (AgCl) = 1.8×10–10
Kf ([Ag(SCN)4]3–) = 1.2×1010
The amount of Hg(OH)2 (Ksp = 3.1×10–26) dissolved in a solution can be increased by the formation of the [HgCl4]2– complex ion (Kf = 1.1×1016). What is the initial amount of KCl in molarity needed to increase the molar solubility of solid Hg(OH)2 to 1.2×10–3 mol/L when solid Hg(OH)2 is added to a KCl solution?
If Ksp of Al(OH)3 is 1.3×10–33 and Kf of [Al(OH)4]– is 3.0×1033, calculate the amount of OH– in molarity needed to dissolve 0.0780 mol of Al(OH)3 in 1.50 liter of solution
If the Kf of [Co(NH3)6]2+ is 1.30×105, what is the molarity of Co2+(aq) and [Co(NH3)6]2+(aq) at the equilibrium of a solution made by dissolving 1.75 g CoBr2 in 200 mL of 0.750 M NH3(aq)?
Determine the equilibrium constant for the reaction with their Ksp and Kf values
CuS (s) + 6 CN- (aq) ⇌ Cu(CN)42- (aq) + S2- (aq) (Ksp = 2x10-47, Kf = 1x1025)
Determine the concentration of Ni2+ that remains when a solution of 1.3x10-3 M Ni(NO3)2 and 0.173 M NH3 reaches equilibrium. (Kf Ni(NH3)62+ = 1.2x109)
Which one of following statements about copper(I) iodide (CuI, Ksp = 1.10×10–12) and Copper(II) carbonate (CuCO3, Ksp = 1.40×10–10) is incorrect.
Cadmium(II) hydroxide, Cd(OH)2, is a sparingly soluble salt with a very small solubility product: Ksp = 2.5×10–14.
Cd(OH)2 reacts with ammonia (NH3) to form the following complex:
Cd(OH)2(s) + 6 NH3(aq) ⇌ [Cu(NH3)6]2+(aq) + 2 OH-(aq) ; Keq = 6.5×10–9
Calculate the value of formation constant Kf for [Cd(NH3)6]2+.
Given Br–(aq), Ca(s), Hg22+(aq), H2(g), identify the weakest and strongest reducing agent.
Given the following: Cl2(g), Na+(aq), S(s), H2O(l), identify the weakest and strongest oxidizing agent.
Rank the following species by increasing standard reduction potential: Au+(aq), F2(g), ClO–(aq).
Referring to the Standard Reduction Potential Table, predict whether the oxidation of Cu+(aq) by I2(s) will take place under standard-state conditions.
Refer to the following substances: Zn(s), HNO2(aq), HClO(aq), H2C2O4(aq), Bi(s), and IO3–(aq).
(a) Hg2Cl2(s) is capable of oxidizing which substances?
(b) H3AsO3(aq) is capable of reducing which substances?
Refer to the list of standard reduction potentials below:
Identify which of the following species can be reduced by Y.
Use the following data to identify which is the stronger reducing agent: AsO43–(aq) or PO43–(aq).
AsO43–(aq) + 2 H2O(l) + 2 e– → AsO2–(aq) + 4 OH–(aq) E°red = –0.71 V
PO43–(aq) + 2 H2O(l) + 2 e– → HPO32–(aq) + 3 OH–(aq) E°red = –1.05 V
Use the following data to identify which is the stronger reducing agent: Ba(s) or Co(s).
Ba2+(aq) + 2 e– → Ba(s) E°red = –2.90 V
Co2+(aq) + 2 e– → Co(s) E°red = –0.28 V
Under standard conditions, list the following oxidizing agents in an acidic solution in order of increasing strength: Sn2+, MnO4–, N2, Br2, H2SO3.
Br2(l) + 2 e– → 2 Br–(aq) E°red = +1.07 V
H2SO3(aq) + 4 H+(aq) + 4 e– → S(s) + 3 H2O(l) E°red = +0.45 V
MnO4–(aq) + 8 H+(aq) + 5 e– → Mn2+(aq) + 4 H2O(l) E°red = +1.51 V
MnO4–(aq) + 2 H2O(l) + 3 e– → MnO2(s) + 4 OH–(aq) E°red = +0.59 V
N2(g) + 5 H+(aq) + 4 e– → N2H5+(aq) E°red = –0.23 V
N2(g) + 4 H2O(l) + 4 e– → 4 OH–(aq) + N2H4(aq) E°red = –1.16 V
Sn2+(aq) + 2 e– → Sn(s) E°red = –0.14 V
The following table lists a series of hypothetical reactions in aqueous solutions with their standard electrode potentials:
Identify the substance(s) that can oxidize Z2+.
Iridium is an impurity found in a mixture of palladium and silver metals that have undergone electrorefining. The oxidation state of iridium is Ir3+ and its standard reduction potential is as follows:
Ir3+(aq) + 3 e– → Ir(s) E°red = 1.156 V
Using this information, what is likely to happen to iridium impurities during electrorefining? When palladium is oxidized, do the impurities remain at the bottom of the refining bath (unchanged) or do they dissolve as ions? If they dissolve in the solution, will they plate out on the cathode?
Pd2+(aq) + 2 e– → Pd(s) E°red = 0.951
Use the standard reduction potential of CrO42– to determine if it is a reducing or oxidizing agent.
Use the standard reduction potential of Pb(s) to determine if it is a reducing or oxidizing agent.
Use the standard reduction potential of BiO+(aq) to determine if it is a reducing or oxidizing agent.
Consider the following metal cations:
Fe3+, Ni2+, Co2+, Ca2+
Identify which of the following is the best oxidizing agent.