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: 

(a) K₃C₆H₅O₇

(b) NaBr

(c) C₆H₈O₇

(d) HCl 

(e) K₂C₆H₆O₇ 

(f) NaOH

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.

Identify the buffer from the given mixtures 

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. (K_a hypochlorous acid = 2.9×10^–8) 

A particular system uses a buffer consisting of HSO₄^- and SO₄^2-. 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.

Identify which of the following solutions can act as a buffer:

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? (K_a 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 pK_a 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. (K_a 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. (K_a 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 (K_a 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 pK_a 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 (K_a 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 pK_a 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 pK_a 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 pK_b 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.

K_a HF = 6.8×10^–4       

K_a HClO = 3.0×10^–8          

K_a CH₃COOH = 1.8×10^–5         

A buffer solution is prepared using HF and F^–. HF has a pK_a 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 (K_a) 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 (pK_a = 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 K_a 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 pK_a 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 - pK_b 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 pK_a for the titration of a weak monoprotic acid and a strong base based on the given curve

A sample mixture of HCl and H₃AsO_4, 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 pK_a 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

b. Thymolphthalein

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 pK_a 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  pK_a = 5.42

B. Metacresol purple  pK_a1 = 1.51, pK_a2 = 8.32

C. Phenolphthalein  pK_a = 9.50

D. Quinaldine red  pK_a = 2.63

E. Bromothymol blue  pK_a = 7.0 

F. Thymol blue  pK_a1 = 1.65, pK_a2 = 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?

K_sp PbS = 7.0×10⁻²⁹, K_sp CuS = 8.0×10⁻³⁷, K_sp 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 CO₃^2– when the second cation starts to precipitate? (K_sp MnCO₃ = 2.24×10^–11, K_sp 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.

K_sp(AgCl) = 1.8×10^–10
K_sp(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:

K_f for [Cu(en)₂]²⁺ = 1.00×10²⁰

K_sp 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 K_sp for Ca(OH)₂ is 4.7×10^–6 and the K_f for [Ca(P₂O₇)]²⁻ is 4.0×10⁴, determine the equilibrium constant (K_eq) for the reaction.

(b) What is the molar solubility of Ca(OH)₂ in 2.50 M Na₂H₂P₂O₇ solution?

[K_a2 H₂P₂O₇^2– = 2.51×10^–7; K_a3 HP₂O₇^3– = 3.89×10^–10]

What is the molar solubility of AgCN in 1.00 M NaCN solution? Use the following values: K_sp AgCN = 6.0×10^–17 and K_f [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: K_sp AgCl = 1.8×10^–10 and K_f [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 K_sp for Ni(OH)₂ is 5.5×10^–16 and the K_f for [Ni(NH₃)₆]²⁺ is 2.0×10⁸.

What is the net ionic equation for the dissolution of Cu(OH)₂ in NH₃? K_sp for Cu(OH)₂ is 1.60×10^–19 and K_f 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.

K_sp (AgCl) = 1.8×10^–10

K_f ([Ag(SCN)₄]^3–) = 1.2×10¹⁰

The amount of Hg(OH)₂ (K_sp = 3.1×10^–26) dissolved in a solution can be increased by the formation of the [HgCl₄]^2– complex ion (K_f = 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 K_sp of Al(OH)₃ is 1.3×10^–33 and K_f 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 K_f 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)?

Find the molar solubility of ZnS in 0.167 M of LiCN (K_sp ZnS = 2.00×10^–25)

Determine the equilibrium constant for the reaction with their K_sp and K_f values

CuS (s) + 6 CN^- (aq) ⇌ Cu(CN)₄^2- (aq) + S^2- (aq) (K_sp = 2x10^-47, K_f = 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. (K_f Ni(NH₃)₆²⁺ = 1.2x10⁹)

Which one of following statements about copper(I) iodide (CuI, K_sp = 1.10×10^–12) and Copper(II) carbonate (CuCO₃, K_sp = 1.40×10^–10) is incorrect.

Cadmium(II) hydroxide, Cd(OH)₂, is a sparingly soluble salt with a very small solubility product: K_sp = 2.5×10^–14.

Cd(OH)₂ reacts with ammonia (NH₃) to form the following complex:

Cd(OH)_2(s) + 6 NH_3(aq) ⇌ [Cu(NH₃)₆]²⁺_(aq) + 2 OH^-_(aq) ; K_eq = 6.5×10^–9

Calculate the value of formation constant K_f 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: AsO4^3–(aq) or PO₄^3–(aq).

AsO₄^3–(aq) + 2 H₂O(l) + 2 e^– → AsO₂^–(aq) + 4 OH^–(aq)       E°_red = –0.71 V

PO₄^3–(aq) + 2 H₂O(l) + 2 e^– → HPO₃^2–(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 CrO₄^2– 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.

Identify the metals that can be oxidized with Pb²⁺(aq) but not with Cd²⁺(aq).

Arrange the species in order of decreasing ability as a  reducing agent 

Sn⁴⁺, I₂, Ga³⁺ , Ca²⁺