Question 1

Which one of these combinations would give a buffer that would be most effective in the pH range 3.0 tO₄.0?

Accepted Answer

The most effective buffer in the pH range 3.0 to 10.0 would have a pKa value close to the desired pH range. The pKa value of HF is 3.17, which is very close to the desired pH range. NaF is used as the conjugate base. Therefore, the buffer system of 0.10 M HF and 0.10 M NaF would be most effective in this pH range.

Question 2

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions is the pH greater than the p K _a of the weak acid?

Accepted Answer

When [HX] < [NaX], the concentration of the conjugate base (X-) is higher than the concentration of the weak acid (HX), which results in a pH greater than the pKa of the weak acid. This is because the higher concentration of the conjugate base will better neutralize added H+, leading to a more basic solution.

Question 3

What is the pH of a buffer solution that consists of 0.83 M HBrO and 0.53 M KBrO? K _a(HBrO) = 2.3×10^{- 9}

Accepted Answer

The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]), where [A-] is the concentration of the base form (KBrO in this case), [HA] is the concentration of the acid form (HBrO), and pKa is the negative logarithm of the acid dissociation constant (Ka). Given Ka = 2.3×10^-9, pKa = -log(2.3×10^-9) = 8.64. Plugging in the given concentrations, pH = 8.64 + log(0.53/0.83) = 8.64 - 0.19 = 8.45.

Question 4

A solution of HCOOH ( K _a = 1.8×10^{- 4}) and HCOO^- was submitted for chemical analysis. The results were: [HCOOH] = 0.050 M, [HCOO^-] = 0.15 M. Calculate the pH.

Accepted Answer

The correct answer is A because the pH of the solution can be calculated using the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]), where pKa is the negative logarithm of the acid dissociation constant, [A-] is the concentration of the conjugate base, and [HA] is the concentration of the acid. In this case, pKa = -log(1.8×10^-4) = 3.74, [A-] = 0.15 M, and [HA] = 0.050 M. Substituting these values into the equation gives pH = 3.74 + log(0.15/0.050) = 4.22.

Question 5

Consider the buffer pair, NH₄Cl/NH₃. What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted to 12.00 and the p K _a for NH₄⁺ equals 9.25?

Accepted Answer

At a pH of 12, the solution is very basic and the concentration of H+ ions is very low. Therefore, the buffer pair is in its deprotonated form, with NH₄⁺ and NH₃ both present. The p K_a of NH₄⁺ is 9.25, which means that at pH 9.25, the concentration of NH₄⁺ and NH₃ will be equal. At higher pH values, the concentration of NH₃ will be greater than that of NH₄⁺, so choice B is correct.

Question 6

Which pair(s) of substances would make a suitable buffer pair ? I. HCl and NaCl II. HF and NaF III. NH₄Cl and NH₃

Accepted Answer

A suitable buffer pair consists of a weak acid and its conjugate base or a weak base and its conjugate acid. HF is a weak acid and NaF is its conjugate base, making them a suitable buffer pair. The substances in option III represent a weak base and its conjugate acid, also making them a suitable buffer pair. HCl is a strong acid and NaCl is its salt, which does not make a suitable buffer pair because strong acids fully dissociate in water and do not form buffers.

Question 7

A buffer solution maintains a constant pH level to a certain extent when a strong base such as NaOH is added. Which of the following buffer solutions has the greatest buffering capacity to consume added NaOH?

Accepted Answer

The answer of A buffer solution maintains a constant pH...

Question 8

What is the pH of a solution that is 0.100 M methylamine (CH₃NH₂) and 0.200 M methyl ammonium chloride (CH₃NH₃Cl)? The K _b of methylamine is 3.70×10^{- 4}.

Accepted Answer

The solution is a buffer system consisting of a weak base (methylamine) and its conjugate acid (methyl ammonium chloride). The pH can be calculated using the Henderson-Hasselbalch equation: pH = pKa + log([base]/[acid]). First, calculate pKa from pKb: pKa = 14 - pKb = 14 - (-log(3.70×10^-4)) = 10.43. Then, pH = 10.43 + log(0.100/0.200) = 10.43 - 0.301 = 10.13. However, due to rounding and approximation, the closest answer is 10.27.

Question 9

Which of the following would not make a suitable buffer pair?

Accepted Answer

HNO3 is a strong acid and NaNO3 is its salt. A buffer solution is made from a weak acid and its salt or a weak base and its salt. So, HNO3/NaNO3 is not a suitable buffer pair.

Question 10

Consider the two buffer solutions prepared below: Solution A: [HA] = 0.15 M and [A^-] = 0.25 M Solution B: [HA] = 1.50 M and [A^-] = 2.50 M Which statement below is true regarding the pH and buffering capacity of these two solutions?

Accepted Answer

The pH of a buffer solution depends on the ratio of the concentration of the weak acid ([HA]) and its conjugate base ([A-]), not on their absolute concentrations. Since the ratio of [HA] to [A-] is the same in both solutions, their pH will be identical. However, Solution B has higher concentrations of both components, providing a larger reservoir of acid and base to neutralize added H+ or OH-, thus offering a larger buffering capacity.

Question 11

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions does the pH just equal the p K _a of the weak acid?

Accepted Answer

When [HX] = [NaX], the concentration of the weak acid and its conjugate base are equal, resulting in the buffer solution being at its optimal buffering capacity, and the pH being equal to the p K _a of the weak acid (assuming ideal conditions with no complications such as any common ion effects).

Question 12

Calculate the pH of a solution that is 0.78 M NH₃ and 0.140 M NH₄NO₃. K _b = 1.8×10^{- 5} for NH₃.

Accepted Answer

The answer of Calculate the pH of a solution that...

Question 13

Calculate the pH of a buffer prepared by dissolving 0.10 mol of NH₄Cl in 1.00 L of 0.15 M NH₃. K _b = 1.8×10^{- 5} for NH₃.

Accepted Answer

The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation: pH = pKa + log([base]/[acid]). First, calculate pKa from pKb: pKa = 14 - pKb = 14 - (-log(1.8×10^-5)) = 9.25. Then, use the concentrations: [NH3] = 0.15 M, [NH4+] = 0.10 M. pH = 9.25 + log(0.15/0.10) = 9.43.

Question 14

What is the pH of an aqueous solution that is 0.55 M HNO₂ and 0.75 M KNO₂? K _a(HNO₂) = 7.1×10^{- 4}

Accepted Answer

To find the pH of the solution, we need to consider the dissociation of HNO₂ in water. The balanced equation for the dissociation is: HNO₂ + H2O ↔ H3O+ + NO₂- The equilibrium constant expression for this reaction is: Ka = [H3O+][NO₂-]/[HNO₂] We can use the given Ka value to find the concentration of hydronium ions in the solution: Ka = 7.1×10^-5 [HNO₂] = 0.55 M [KNO3] = 0.75 M Since HNO₂ and KNO3 are both strong electrolytes, we can assume that they dissociate completely in water. This means that the concentration of NO₂- ions in the solution is equal to the concentration of KNO3, which is 0.75 M. We can use this information to set up an ICE table for the dissociation: HNO₂ + H2O ↔ H3O+ + NO₂- I 0.55 M 0.75 M C -x x x E 0.55 - x x 0.75 + x Substituting these values into the Ka expression and solving for x gives: 7.1×10^-5 = x^2/(0.55 - x) x = 0.014 Therefore, [H3O+] = 0.014 M, and the pH of the solution is: pH = -log[H3O+] = 3.28 The best choice is B, which is the answer obtained after solving the problem.

Question 15

What is the pH of an aqueous solution that contains 0.085 M HNO₂ and 0.10 M potassium nitrite (KNO₂)? K _a(HNO₂) = 4.5×10^{- 4}

Accepted Answer

The answer of What is the pH of an aqueous...

Question 16

Consider the buffer pair, HF/F^-. What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted tO₄.00 and the p K _a for HF equals 3.17?

Accepted Answer

The pH is higher than the pKa, indicating that the solution contains more of the conjugate base (F-) than the acid (HF).

Question 17

A solution is made by dissolving 0.100 mole NH₃ ( K _b = 1.8×10^{- 5}) and 0.200 mole of NH₄Cl in water and diluting to 1.00 L. What is the pH of the resulting solution?

Accepted Answer

The correct answer is A because the solution is basic due to the presence of NH3 and the Kb value of NH3 is greater than Kw.

Question 18

What is the pH of a solution prepared by mixing 0.250 mol of hydrazoic acid, HN₃, and 0.500 mol of sodium azide, NaN₃, to make 1.00 liter of solution? ( K _a = 1.9×10^{- 5})

Accepted Answer

The answer of What is the pH of a solution...

Question 19

Which of the following mixtures would not be suitable for preparing a buffer solution?

Accepted Answer

The answer of Which of the following mixtures would not...

Question 20

A buffer solution maintains a constant pH level to a certain extent when a strong acid such as HCl is added. Which of the following buffer solutions has the greatest buffering capacity to consume added HCl?

Accepted Answer

The greatest buffering capacity is achieved with the highest total concentration of the acid and its conjugate base. Option C has the highest total concentration (1.0 M acid + 1.5 M base = 2.5 M), providing the greatest capacity to neutralize added HCl.

Quiz 16: Reactions Between Acids and Bases

Which one of these combinations would give a buffer that would be most effective in the pH range 3.0 tO₄.0?

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions is the pH greater than the p K _a of the weak acid?

What is the pH of a buffer solution that consists of 0.83 M HBrO and 0.53 M KBrO? K _a(HBrO) = 2.3×10^{- 9}

A solution of HCOOH ( K _a = 1.8×10^{- 4}) and HCOO^- was submitted for chemical analysis. The results were: [HCOOH] = 0.050 M, [HCOO^-] = 0.15 M. Calculate the pH.

Consider the buffer pair, NH₄Cl/NH₃. What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted to 12.00 and the p K _a for NH₄⁺ equals 9.25?

Which pair(s) of substances would make a suitable buffer pair ? I. HCl and NaCl II. HF and NaF III. NH₄Cl and NH₃

A buffer solution maintains a constant pH level to a certain extent when a strong base such as NaOH is added. Which of the following buffer solutions has the greatest buffering capacity to consume added NaOH?

What is the pH of a solution that is 0.100 M methylamine (CH₃NH₂) and 0.200 M methyl ammonium chloride (CH₃NH₃Cl) ? The K _b of methylamine is 3.70×10^{- 4}.

Which of the following would not make a suitable buffer pair?

Consider the two buffer solutions prepared below: Solution A: [HA] = 0.15 M and [A^-] = 0.25 M Solution B: [HA] = 1.50 M and [A^-] = 2.50 M Which statement below is true regarding the pH and buffering capacity of these two solutions?

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions does the pH just equal the p K _a of the weak acid?

Calculate the pH of a solution that is 0.78 M NH₃ and 0.140 M NH₄NO₃. K _b = 1.8×10^{- 5} for NH₃.

Calculate the pH of a buffer prepared by dissolving 0.10 mol of NH₄Cl in 1.00 L of 0.15 M NH₃. K _b = 1.8×10^{- 5} for NH₃.

What is the pH of an aqueous solution that is 0.55 M HNO₂ and 0.75 M KNO₂? K _a(HNO₂) = 7.1×10^{- 4}

What is the pH of an aqueous solution that contains 0.085 M HNO₂ and 0.10 M potassium nitrite (KNO₂) ? K _a(HNO₂) = 4.5×10^{- 4}

Consider the buffer pair, HF/F^-. What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted tO₄.00 and the p K _a for HF equals 3.17?

A solution is made by dissolving 0.100 mole NH₃ ( K _b = 1.8×10^{- 5}) and 0.200 mole of NH₄Cl in water and diluting to 1.00 L. What is the pH of the resulting solution?

What is the pH of a solution prepared by mixing 0.250 mol of hydrazoic acid, HN₃, and 0.500 mol of sodium azide, NaN₃, to make 1.00 liter of solution? ( K _a = 1.9×10^{- 5})

Which of the following mixtures would not be suitable for preparing a buffer solution?

A buffer solution maintains a constant pH level to a certain extent when a strong acid such as HCl is added. Which of the following buffer solutions has the greatest buffering capacity to consume added HCl?

Quiz 16: Reactions Between Acids and Bases

Which one of these combinations would give a buffer that would be most effective in the pH range 3.0 tO4.0?

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions is the pH greater than the p K a of the weak acid?

What is the pH of a buffer solution that consists of 0.83 M HBrO and 0.53 M KBrO? K a(HBrO) = 2.3×10 - 9

A solution of HCOOH ( K a = 1.8×10 - 4) and HCOO - was submitted for chemical analysis. The results were: [HCOOH] = 0.050 M, [HCOO - ] = 0.15 M. Calculate the pH.

Consider the buffer pair, NH4Cl/NH3. What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted to 12.00 and the p K a for NH4+ equals 9.25?

Which pair(s) of substances would make a suitable buffer pair ? I. HCl and NaCl II. HF and NaF III. NH4Cl and NH3

A buffer solution maintains a constant pH level to a certain extent when a strong base such as NaOH is added. Which of the following buffer solutions has the greatest buffering capacity to consume added NaOH?

What is the pH of a solution that is 0.100 M methylamine (CH3NH2) and 0.200 M methyl ammonium chloride (CH3NH3Cl) ? The K b of methylamine is 3.70×10 - 4.

Which of the following would not make a suitable buffer pair?

Consider the two buffer solutions prepared below: Solution A: [HA] = 0.15 M and [A - ] = 0.25 M Solution B: [HA] = 1.50 M and [A - ] = 2.50 M Which statement below is true regarding the pH and buffering capacity of these two solutions?

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions does the pH just equal the p K a of the weak acid?

Calculate the pH of a solution that is 0.78 M NH3 and 0.140 M NH4NO3. K b = 1.8×10 - 5 for NH3.

Calculate the pH of a buffer prepared by dissolving 0.10 mol of NH4Cl in 1.00 L of 0.15 M NH3. K b = 1.8×10 - 5 for NH3.

What is the pH of an aqueous solution that is 0.55 M HNO2 and 0.75 M KNO2? K a(HNO2) = 7.1×10 - 4

What is the pH of an aqueous solution that contains 0.085 M HNO2 and 0.10 M potassium nitrite (KNO2) ? K a(HNO2) = 4.5×10 - 4

Consider the buffer pair, HF/F - . What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted tO4.00 and the p K a for HF equals 3.17?

A solution is made by dissolving 0.100 mole NH3 ( K b = 1.8×10 - 5) and 0.200 mole of NH4Cl in water and diluting to 1.00 L. What is the pH of the resulting solution?

What is the pH of a solution prepared by mixing 0.250 mol of hydrazoic acid, HN3, and 0.500 mol of sodium azide, NaN3, to make 1.00 liter of solution? ( K a = 1.9×10 - 5)

Which of the following mixtures would not be suitable for preparing a buffer solution?

A buffer solution maintains a constant pH level to a certain extent when a strong acid such as HCl is added. Which of the following buffer solutions has the greatest buffering capacity to consume added HCl?

Which one of these combinations would give a buffer that would be most effective in the pH range 3.0 tO₄.0?

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions is the pH greater than the p K _a of the weak acid?

What is the pH of a buffer solution that consists of 0.83 M HBrO and 0.53 M KBrO? K _a(HBrO) = 2.3×10^{- 9}

A solution of HCOOH ( K _a = 1.8×10^{- 4}) and HCOO^- was submitted for chemical analysis. The results were: [HCOOH] = 0.050 M, [HCOO^-] = 0.15 M. Calculate the pH.

Consider the buffer pair, NH₄Cl/NH₃. What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted to 12.00 and the p K _a for NH₄⁺ equals 9.25?

Which pair(s) of substances would make a suitable buffer pair ? I. HCl and NaCl II. HF and NaF III. NH₄Cl and NH₃

What is the pH of a solution that is 0.100 M methylamine (CH₃NH₂) and 0.200 M methyl ammonium chloride (CH₃NH₃Cl) ? The K _b of methylamine is 3.70×10^{- 4}.

Consider the two buffer solutions prepared below: Solution A: [HA] = 0.15 M and [A^-] = 0.25 M Solution B: [HA] = 1.50 M and [A^-] = 2.50 M Which statement below is true regarding the pH and buffering capacity of these two solutions?

Consider a buffer solution containing a weak acid, HX, and a salt of the weak acid's conjugate base, NaX. Under what conditions does the pH just equal the p K _a of the weak acid?

Calculate the pH of a solution that is 0.78 M NH₃ and 0.140 M NH₄NO₃. K _b = 1.8×10^{- 5} for NH₃.

Calculate the pH of a buffer prepared by dissolving 0.10 mol of NH₄Cl in 1.00 L of 0.15 M NH₃. K _b = 1.8×10^{- 5} for NH₃.

What is the pH of an aqueous solution that is 0.55 M HNO₂ and 0.75 M KNO₂? K _a(HNO₂) = 7.1×10^{- 4}

What is the pH of an aqueous solution that contains 0.085 M HNO₂ and 0.10 M potassium nitrite (KNO₂) ? K _a(HNO₂) = 4.5×10^{- 4}

Consider the buffer pair, HF/F^-. What can be stated about the relative proportions of this buffer pair if the pH of an aqueous solution of this pair is adjusted tO₄.00 and the p K _a for HF equals 3.17?

A solution is made by dissolving 0.100 mole NH₃ ( K _b = 1.8×10^{- 5}) and 0.200 mole of NH₄Cl in water and diluting to 1.00 L. What is the pH of the resulting solution?

What is the pH of a solution prepared by mixing 0.250 mol of hydrazoic acid, HN₃, and 0.500 mol of sodium azide, NaN₃, to make 1.00 liter of solution? ( K _a = 1.9×10^{- 5})