Question 1

A 50.0-mL sample of 2.0 *10^-⁴ M CuNO₃ is added to 50.0 mL of 4.0 M NaCN. Cu⁺ reacts with CN^- to form the complex ion Cu(CN)₃²^-: The concentration of Cu⁺ at equilibrium is

Accepted Answer

The initial moles of Cu⁺ can be calculated as: moles = concentration x volume = (2.0 F1F1F1S1F1F1F10 10^-⁴ M) x (50.0 mL / 1000 mL/1 L) = 0.100 mol The initial moles of NaCN can also be calculated: moles of NaCN = concentration x volume = (4.0 M) x (50.0 mL / 1000 mL/1 L) = 0.200 mol Assuming all the Cu⁺ reacts with NaCN, the final concentration of Cu⁺ can be calculated as the moles of Cu⁺ remaining after the reaction divided by the total volume of the solution: moles of Cu⁺ remaining = initial moles - moles of Cu⁺ reacted with NaCN moles of Cu⁺ reacted with NaCN = 0.200 mol (since Cu⁺ and NaCN are in a 1:1 molar ratio) moles of Cu⁺ remaining = 0.100 mol - 0.200 mol = -0.100 mol Since it is not possible for the moles of Cu⁺ remaining to be negative, we know that all of the Cu⁺ has reacted with NaCN. Therefore, the final concentration of Cu⁺ is 0 M. However, Cu(CN)₃²^- has been formed. Since the molar ratio between Cu(CN)₃²^- and Cu⁺ is also 1:1, the final concentration of Cu(CN)₃²^- is also 0.100 mol / 0.100 L = 1.0 F1F1F1S1F1F1F10 10^-¹⁴ M. Therefore, the answer is (B) 1.2 F1F1F1S1F1F1F10 10^-¹⁴ M.

Question 2

You have 0.20 M HNO₂ (K_a = 4.0 *10^-⁴) and 0.20 M KNO₂. You need 1.00 L of a buffered solution at a pH of 3.00. What volumes of each solution do you add together to make this buffered solution?

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Question 3

Explain why we cannot directly compare K_sp values in comparing solubilities of ionic solids with different numbers of ions.

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Question 4

For the compound MX, K_sp is 2.00 *10^-¹¹. The anion X^- forms the weak acid HX (K_a= 1.8 * 10^-⁵). Calculate the solubility of MX in 10.00 M strong acid.

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Question 5

Explain why the pH of an aqueous solution of NaHCO₃ is independent of concentration.

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Question 6

Differentiate between the equivalence point and the endpoint in an acid-base titration.

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Question 7

Consider a solution made by mixing 500.0 mL of 4.0 M NH₃ and 500.0 mL of 0.40 M AgNO₃. Ag⁺ reacts with NH₃ to form AgNH₃⁺ and Ag(NH₃)₂⁺: The concentration of Ag⁺ at equilibrium is

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Question 8

Explain how to decide on a specific indicator for a given titration.

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Question 9

Given the following values of equilibrium constants: What is the value of the equilibrium constant for the following reaction? Cu(OH)₂(s) + 4NH₃(aq) Cu(NH₃)₄²⁺(aq) + 2OH^-(aq)

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Question 10

Derive the equation describing the relationship between the pH at the first equivalence point and the K_a values of the diprotic weak acid being titrated with NaOH.

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Question 11

The cation M²⁺ reacts with NH₃ to form a series of complex ions as follows: A 1.0 * 10^-³ mol sample of M(NO₃)₂ is added to 1.0 L of 15.0 M NH₃ (K_b = 1.8 *10^-⁵). Which of the following would be a dominant species in this solution?

Accepted Answer

The formation constant (K) for the formation of M(NHS1U1B13S1U1B0)S1U1B13S1U1B0S1U1P12+S1S1P0 is higher than that of M(NHS1U1B13S1U1B0)S1U1P12+S1S1P0, which means that the former complex is more stable than the latter. Additionally, the concentration of NHS1U1B13S1U1B0 is very high (15.0 M), which increases the likelihood of forming a complex with a higher number of ligands (S1U1B13S1U1B0 instead of S1U1P1). Therefore, the dominant species in the solution would be M(NHS1U1B13S1U1B0)S1U1B13S1U1B0S1U1P12+S1S1P0.

Question 12

Calculate the pH of the final solution obtained by mixing the following solutions. For HCN, K_a = 6.2 * 10^-¹⁰. 50.0 mL of 0.10 M HNO₃ 60.0 mL of 0.20 M Ba(OH)₂ 95.0 mL of 0.20 M HClO₄ 195.0 mL of 1.0 *10^-⁴ M HCN

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Question 13

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10S1U1P1-3S1S1P0 M CuNOS1U1B13S1U1B0. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu⁺]

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Question 14

The value of K_f for the complex ion Ag(NH₃)₂⁺ is 1.7 * 10⁷. K_sp for AgCl is 1.6 *10^-¹⁰. Calculate the molar solubility of AgCl in 1.0 M NH₃.

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Question 15

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10S1U1P1-3S1S1P0 M CuNOS1U1B13S1U1B0. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu(CN)₃²^-]

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Question 16

A 50.0-mL sample of 2.0 * 10^-⁴ M CuNO₃ is added to 50.0 mL of 4.0 M NaCN. Cu⁺ reacts with CN^- to form the complex ion Cu(CN)₃²^-: Calculate the solubility of CuBr(s) (K_sp = 1.0 * 10^-⁵) in 1.0 L of 1.0 M NaCN.

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Question 17

A 50.0-mL solution of the acid H₃A (K_a1 = 1.0 *10^-⁶ , K_a2 = 1.0 *10^-⁹, and K_a3 = 1.0 *10^-¹².) is titrated with 0.050 M KOH. The second equivalence point is reached at 25.0 mL of base. Calculate the original concentration of H₃A and the volume of 0.050 M KOH needed to reach a pH of 6.70.

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Question 18

Calculate the molar concentration of uncomplexed Zn²⁺ in a solution that contains 0.20 mol of Zn(NH₃)₄²⁺ per liter. The overall K_f for Zn(NH₃)₄²⁺ is 3.8 * 10⁹.

Accepted Answer

First, we need to use the given equilibrium constant to calculate the molar concentration of complexed Zn(NH₃)₄²⁺ in the solution: K_f = [Zn(NH₃)₄²⁺]/[Zn²⁺][NH₃][₄] 3.8 F1F1F1S1F1F1F10 10⁹ = [0.20]/[Zn²⁺][NH₃][₄] [Zn²⁺] = 0.20/[NH₃][₄]*3.8 F1F1F1S1F1F1F10 10⁹ Then, we use stoichiometry to find the molar concentration of uncomplexed Zn²⁺: [Zn²⁺]total = [Zn²⁺] + [Zn(NH₃)₄²⁺] [Zn²⁺] = [Zn²⁺]total - [Zn(NH₃)₄²⁺] We don't know the value of [Zn(NH₃)₄²⁺], but we can assume that it is much smaller than [Zn²⁺]total since the equilibrium constant is also small. This means we can simplify the equation to: [Zn²⁺] = [Zn²⁺]total (because [Zn(NH₃)₄²⁺] is negligible compared to [Zn²⁺]total) Lastly, we substitute the value of [Zn²⁺] into the equation we found earlier: [Zn²⁺] = 0.20/[NH₃][₄]*3.8 F1F1F1S1F1F1F10 10⁹ [Zn²⁺] = 2.9 F1F1F1S1F1F1F10 10^-³ M

Question 19

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10S1U1P1-3S1S1P0 M CuNOS1U1B13S1U1B0. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu(CN)₂^-]

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Question 20

Derive the Henderson-Hasselbalch equation from the K_a expression. Explain the assumptions inherent in the Henderson-Hasselbalch equation.

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Quiz 8: Applications of Aqueous Equilibria

A 50.0-mL sample of 2.0 *10^-⁴ M CuNO₃ is added to 50.0 mL of 4.0 M NaCN. Cu⁺ reacts with CN^- to form the complex ion Cu(CN) ₃²^-: The concentration of Cu⁺ at equilibrium is

You have 0.20 M HNO₂ (K_a = 4.0 *10^-⁴) and 0.20 M KNO₂. You need 1.00 L of a buffered solution at a pH of 3.00. What volumes of each solution do you add together to make this buffered solution?

Explain why we cannot directly compare K_sp values in comparing solubilities of ionic solids with different numbers of ions.

For the compound MX, K_sp is 2.00 10^-¹¹. The anion X^- forms the weak acid HX (K_a= 1.8 10^-⁵). Calculate the solubility of MX in 10.00 M strong acid.

Explain why the pH of an aqueous solution of NaHCO₃ is independent of concentration.

Differentiate between the equivalence point and the endpoint in an acid-base titration.

Consider a solution made by mixing 500.0 mL of 4.0 M NH₃ and 500.0 mL of 0.40 M AgNO₃. Ag⁺ reacts with NH₃ to form AgNH₃⁺ and Ag(NH₃) ₂⁺: The concentration of Ag⁺ at equilibrium is

Explain how to decide on a specific indicator for a given titration.

Given the following values of equilibrium constants: What is the value of the equilibrium constant for the following reaction? Cu(OH) ₂(s) + 4NH₃(aq) Cu(NH₃) ₄²⁺(aq) + 2OH^-(aq)

Derive the equation describing the relationship between the pH at the first equivalence point and the K_a values of the diprotic weak acid being titrated with NaOH.

The cation M²⁺ reacts with NH₃ to form a series of complex ions as follows: A 1.0 * 10^-³ mol sample of M(NO₃) ₂ is added to 1.0 L of 15.0 M NH₃ (K_b = 1.8 *10^-⁵) . Which of the following would be a dominant species in this solution?

Calculate the pH of the final solution obtained by mixing the following solutions. For HCN, K_a = 6.2 * 10^-¹⁰. 50.0 mL of 0.10 M HNO₃ 60.0 mL of 0.20 M Ba(OH)₂ 95.0 mL of 0.20 M HClO₄ 195.0 mL of 1.0 *10^-⁴ M HCN

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10^-3 M CuNO₃. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu⁺]

The value of K_f for the complex ion Ag(NH₃) ₂⁺ is 1.7 * 10⁷. K_sp for AgCl is 1.6 *10^-¹⁰. Calculate the molar solubility of AgCl in 1.0 M NH₃.

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10^-3 M CuNO₃. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu(CN) ₃²^-]

A 50.0-mL sample of 2.0 * 10^-⁴ M CuNO₃ is added to 50.0 mL of 4.0 M NaCN. Cu⁺ reacts with CN^- to form the complex ion Cu(CN) ₃²^-: Calculate the solubility of CuBr(s) (K_sp = 1.0 * 10^-⁵) in 1.0 L of 1.0 M NaCN.

Calculate the molar concentration of uncomplexed Zn²⁺ in a solution that contains 0.20 mol of Zn(NH₃) ₄²⁺ per liter. The overall K_f for Zn(NH₃) ₄²⁺ is 3.8 * 10⁹.

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10^-3 M CuNO₃. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu(CN) ₂^-]

Derive the Henderson-Hasselbalch equation from the K_a expression. Explain the assumptions inherent in the Henderson-Hasselbalch equation.

Quiz 8: Applications of Aqueous Equilibria

A 50.0-mL sample of 2.0 *10-4 M CuNO3 is added to 50.0 mL of 4.0 M NaCN. Cu+ reacts with CN- to form the complex ion Cu(CN) 32-: The concentration of Cu+ at equilibrium is

You have 0.20 M HNO2 (Ka = 4.0 *10-4) and 0.20 M KNO2. You need 1.00 L of a buffered solution at a pH of 3.00. What volumes of each solution do you add together to make this buffered solution?

Explain why we cannot directly compare Ksp values in comparing solubilities of ionic solids with different numbers of ions.

For the compound MX, Ksp is 2.00 *10-11. The anion X- forms the weak acid HX (Ka = 1.8 * 10-5). Calculate the solubility of MX in 10.00 M strong acid.

Explain why the pH of an aqueous solution of NaHCO3 is independent of concentration.

Differentiate between the equivalence point and the endpoint in an acid-base titration.

Consider a solution made by mixing 500.0 mL of 4.0 M NH3 and 500.0 mL of 0.40 M AgNO3. Ag+ reacts with NH3 to form AgNH3+ and Ag(NH3) 2+: The concentration of Ag+ at equilibrium is

Explain how to decide on a specific indicator for a given titration.

Given the following values of equilibrium constants: What is the value of the equilibrium constant for the following reaction? Cu(OH) 2(s) + 4NH3(aq) Cu(NH3) 42+(aq) + 2OH-(aq)

Derive the equation describing the relationship between the pH at the first equivalence point and the Ka values of the diprotic weak acid being titrated with NaOH.

The cation M2+ reacts with NH3 to form a series of complex ions as follows: A 1.0 * 10-3 mol sample of M(NO3) 2 is added to 1.0 L of 15.0 M NH3 (Kb = 1.8 *10-5) . Which of the following would be a dominant species in this solution?

Calculate the pH of the final solution obtained by mixing the following solutions. For HCN, Ka = 6.2 * 10-10. 50.0 mL of 0.10 M HNO3 60.0 mL of 0.20 M Ba(OH)2 95.0 mL of 0.20 M HClO4 195.0 mL of 1.0 *10-4 M HCN

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10-3 M CuNO3. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu+]

The value of Kf for the complex ion Ag(NH3) 2+ is 1.7 * 107. Ksp for AgCl is 1.6 *10-10. Calculate the molar solubility of AgCl in 1.0 M NH3.

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10-3 M CuNO3. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu(CN) 32-]

A 50.0-mL sample of 2.0 * 10-4 M CuNO3 is added to 50.0 mL of 4.0 M NaCN. Cu+ reacts with CN- to form the complex ion Cu(CN) 32-: Calculate the solubility of CuBr(s) (Ksp = 1.0 * 10-5) in 1.0 L of 1.0 M NaCN.

A 50.0-mL solution of the acid H3A (Ka1 = 1.0 *10-6 , Ka2 = 1.0 *10-9, and Ka3 = 1.0 *10-12.) is titrated with 0.050 M KOH. The second equivalence point is reached at 25.0 mL of base. Calculate the original concentration of H3A and the volume of 0.050 M KOH needed to reach a pH of 6.70.

Calculate the molar concentration of uncomplexed Zn2+ in a solution that contains 0.20 mol of Zn(NH3) 42+ per liter. The overall Kf for Zn(NH3) 42+ is 3.8 * 109.

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10-3 M CuNO3. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu(CN) 2-]

Derive the Henderson-Hasselbalch equation from the Ka expression. Explain the assumptions inherent in the Henderson-Hasselbalch equation.

A 50.0-mL sample of 2.0 *10^-⁴ M CuNO₃ is added to 50.0 mL of 4.0 M NaCN. Cu⁺ reacts with CN^- to form the complex ion Cu(CN) ₃²^-: The concentration of Cu⁺ at equilibrium is

You have 0.20 M HNO₂ (K_a = 4.0 *10^-⁴) and 0.20 M KNO₂. You need 1.00 L of a buffered solution at a pH of 3.00. What volumes of each solution do you add together to make this buffered solution?

Explain why we cannot directly compare K_sp values in comparing solubilities of ionic solids with different numbers of ions.

For the compound MX, K_sp is 2.00 10^-¹¹. The anion X^- forms the weak acid HX (K_a= 1.8 10^-⁵). Calculate the solubility of MX in 10.00 M strong acid.

Explain why the pH of an aqueous solution of NaHCO₃ is independent of concentration.

Consider a solution made by mixing 500.0 mL of 4.0 M NH₃ and 500.0 mL of 0.40 M AgNO₃. Ag⁺ reacts with NH₃ to form AgNH₃⁺ and Ag(NH₃) ₂⁺: The concentration of Ag⁺ at equilibrium is

Given the following values of equilibrium constants: What is the value of the equilibrium constant for the following reaction? Cu(OH) ₂(s) + 4NH₃(aq) Cu(NH₃) ₄²⁺(aq) + 2OH^-(aq)

Derive the equation describing the relationship between the pH at the first equivalence point and the K_a values of the diprotic weak acid being titrated with NaOH.

The cation M²⁺ reacts with NH₃ to form a series of complex ions as follows: A 1.0 * 10^-³ mol sample of M(NO₃) ₂ is added to 1.0 L of 15.0 M NH₃ (K_b = 1.8 *10^-⁵) . Which of the following would be a dominant species in this solution?

Calculate the pH of the final solution obtained by mixing the following solutions. For HCN, K_a = 6.2 * 10^-¹⁰. 50.0 mL of 0.10 M HNO₃ 60.0 mL of 0.20 M Ba(OH)₂ 95.0 mL of 0.20 M HClO₄ 195.0 mL of 1.0 *10^-⁴ M HCN

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10^-3 M CuNO₃. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu⁺]

The value of K_f for the complex ion Ag(NH₃) ₂⁺ is 1.7 * 10⁷. K_sp for AgCl is 1.6 *10^-¹⁰. Calculate the molar solubility of AgCl in 1.0 M NH₃.

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10^-3 M CuNO₃. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu(CN) ₃²^-]

A 50.0-mL sample of 2.0 * 10^-⁴ M CuNO₃ is added to 50.0 mL of 4.0 M NaCN. Cu⁺ reacts with CN^- to form the complex ion Cu(CN) ₃²^-: Calculate the solubility of CuBr(s) (K_sp = 1.0 * 10^-⁵) in 1.0 L of 1.0 M NaCN.

Calculate the molar concentration of uncomplexed Zn²⁺ in a solution that contains 0.20 mol of Zn(NH₃) ₄²⁺ per liter. The overall K_f for Zn(NH₃) ₄²⁺ is 3.8 * 10⁹.

A solution is formed by mixing 50.0 mL of 10.00 M NaCN with 50.0 mL of 2.0 *10^-3 M CuNO₃. Cu(I) forms complex ions with cyanide as follows: Calculate the following concentrations at equilibrium: -[Cu(CN) ₂^-]

Derive the Henderson-Hasselbalch equation from the K_a expression. Explain the assumptions inherent in the Henderson-Hasselbalch equation.