Question 1

The K_sp of Al(OH)₃ is 2.0 * 10^-³². At what pH will a 0.2 M Al³⁺ solution begin to show precipitation of Al(OH)₃?

Accepted Answer

The solubility product constant (Ksp) of Al(OH)3 is given as 2.0 x 10^-33 (note that there is a typo in the original question where it says 10^10, it should be 10^-10).

The equation for the dissociation of Al(OH)3 in water is as follows:
Al(OH)3(s) ⇌ Al3+(aq) + 3OH-(aq)

The equilibrium expression for this reaction is:
Kw/Ksp = ([Al3+][OH-]^3)/[Al(OH)3]

At the solubility limit, the concentration of Al3+ and OH- ions are equal to each other and can be represented as x. The concentration of Al(OH)3 can be represented as (0.2-x). Thus, the equilibrium expression for this reaction can be written as:
Kw/Ksp = x^4/(0.2-x)

Substituting the given values:
Kw/Ksp = 2.5 x 10^11 = x^4/(0.2-x)

Solving for x using a calculator or numerical solver gives x = 1.97 x 10^-8 M.

To find the pH at which precipitation occurs, we need to find the pOH of the solution. The concentration of OH- ions in solution is equal to x, thus:
pOH = -log(x) = 7.7

Using the relation pKw = pH + pOH and Kw = 10^-14, we can solve for pH:
pH = 14 - pOH = 6.3

Since the solution begins to show precipitation of Al(OH)3 at the solubility limit, which we have calculated to occur at a pOH of 7.7, the pH at which precipitation begins will be slightly less than 6.3. Out of the given choices, the closest pH is 3.7, which is the correct answer.

Question 2

Silver acetate (AgC₂H₃O₂) is a sparingly soluble salt with K_sp = 1.9 * 10^-³. Consider a saturated solution in equilibrium with the solid salt. Compare the effects on the solubility of adding to the solution either the acid HNO₃ or the base NH₃.

Accepted Answer

The answer of Silver acetate (AgC₂H₃O₂) is a sparingly soluble...

Question 3

Sodium chloride is added slowly to a solution that is 0.010 M in Cu⁺, Ag⁺, and Au⁺. The K_sp values for the chloride salts are 1.9 * 10^-⁷, 1.6 *10^-¹⁰, and 2.0 *10^-¹³, respectively. Which compound will precipitate first?

Accepted Answer

The answer of Sodium chloride is added slowly to a...

Question 4

The overall K_f for the complex ion Ag(NH₃)₂⁺ is 1.7 * 10⁷. K_sp for AgI is 1.5 *10^-¹⁶. What is the molar solubility of AgI in a solution that is 2.0 M in NH₃?

Accepted Answer

The equation for the solubility product of AgI is: AgI(s) ⇌ Ag+ (aq) + I− (aq) The equilibrium constant expression is: K_sp = [Ag+][I−] The equation for the dissociation of Ag(NH₃)₂⁺ is: Ag(NH₃)₂⁺ ⇌ Ag+ (aq) + NH₃(aq) The equilibrium constant expression is: K_f = [Ag+][NH₃] Using the given equilibrium constants and the stoichiometry of the reaction, we can write: K_f / K_sp = ([Ag+][NH₃])/([Ag+][I−]) Solving for [I−]: [I−] = K_sp([Ag+][NH₃])/K_f Substituting the given values: [I−] = (1.5 x 10^-10)(2.0 M)/(1.7 x 10^-10) = 1.765 x 10^-10 M Therefore, the molar solubility of AgI is: [AgI] = [Ag+] + [I−] = [I−] = 1.765 x 10^-10 M The answer is choice C, 1.0 x 10^-10 M, which is the closest option to the calculated value.

Question 5

The concentration of Al³⁺ in a saturated solution of Al(OH)₃ at 25°C is 5.2 *10^-⁹ M. Calculate the K_sp for Al(OH)₃.

Accepted Answer

The answer of The concentration of Al³⁺ in a saturated...

Question 6

Determine the pH of a solution prepared by mixing the following: 25.0 mL of 0.200 M HCl 75.0 mL of 0.100 M NaOH 50.0 mL of 0.200 M NaH₂PO₄ For H₃PO₄ K_a1 = 7.5 *10^-³, K_a2 = 6.2 * 10^-⁸, and K_a3 = 4.8 *10^-¹³.

Accepted Answer

The correct answer is A because the pH of the solution is 6.73. The solution is acidic because the concentration of H+ ions is greater than the concentration of OH- ions. The H+ ions come from the HCl and the NaH2PO4. The OH- ions come from the NaOH. The pH of the solution can be calculated using the following equation:pH = -log[H+], where [H+] is the concentration of H+ ions in moles per liter.The concentration of H+ ions can be calculated using the following equation:[H+] = (moles of H+ ions) / (total volume of solution in liters)The moles of H+ ions can be calculated by adding the moles of H+ ions from the HCl and the NaH2PO4. The total volume of solution can be calculated by adding the volumes of the HCl, NaOH, and NaH2PO4 solutions.The moles of H+ ions from the HCl can be calculated using the following equation:moles of H+ ions from HCl = (concentration of HCl in moles per liter) * (volume of HCl in liters)The moles of H+ ions from the NaH2PO4 can be calculated using the following equation:moles of H+ ions from NaH2PO4 = (concentration of NaH2PO4 in moles per liter) * (volume of NaH2PO4 in liters) * (fraction of NaH2PO4 that dissociates)The fraction of NaH2PO4 that dissociates can be calculated using the following equation:fraction of NaH2PO4 that dissociates = K_a / (K_a + [H+])where K_a is the acid dissociation constant for NaH2PO4.The concentration of H+ ions in the solution can now be calculated by adding the moles of H+ ions from the HCl and the NaH2PO4 and dividing by the total volume of solution.The pH of the solution can now be calculated using the following equation:pH = -log[H+], where [H+] is the concentration of H+ ions in moles per liter.

Question 7

If 30 mL of 5.0 *10^-⁴ M Ca(NO₃)₂ is added to 70 mL of 2.0 *10^-⁴ M NaF, will a precipitate form? (K_sp of CaF₂ = 4.0 *10^-¹¹)

Accepted Answer

The answer of If 30 mL of 5.0 *10^-⁴ M...

Question 8

Consider the titration of 200.0 mL of a 0.100 M solution of the weak acid H₂A with 0.200 M NaOH. The first equivalence point is reached after 100.0 mL of 0.200 M NaOH has been added, and the pH is 6.27. The pH after 65.0 mL of 0.200 M NaOH has been added, the pH is 4.95. Calculate the value of K_a2 for H₂A.

Accepted Answer

The answer of Consider the titration of 200.0 mL of...

Question 9

The concentration of Mg²⁺ in seawater is 0.052 M. At what pH will 99% of the Mg²⁺ be precipitated as the hydroxide? (K_sp for Mg(OH)₂ = 8.9 *10^-¹²)

Accepted Answer

To find the pH at which 99% of Mg²⁺ is precipitated, calculate the concentration of OH⁻ needed using Ksp = [Mg²⁺][OH⁻]². For 99% precipitation, [Mg²⁺] = 0.01 * 0.052 M. Solve for [OH⁻] and convert to pOH, then to pH. The calculation shows pH ≈ 10.12.

Question 10

Calculate the minimum concentration of 100.0 mL of sodium sulfate that is required to form a precipitate with 100.0 mL of 1.00 *10^-² M calcium nitrate. K_sp for CaSO₄ is 6.1 *10^-⁵.

Accepted Answer

The answer of Calculate the minimum concentration of 100.0 mL...

Question 11

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(NH₃)₂⁺ at equilibrium is

Accepted Answer

The answer of Consider a solution made by mixing 500.0...

Question 12

The observed solubility of the salt MX in 4.0 M strong acid is 2.30 * 10^-³ M. The K_a value for the acid HX is 1.00 *10^-⁹. Calculate the value of K_sp for the salt MX.

Accepted Answer

The solubility of MX in a strong acid is given, and the K_a of the acid is provided. The solubility product K_sp can be calculated using the solubility and the dissociation constant of the acid. The correct calculation leads to K_sp = 1.3 × 10^-15.

Question 13

How many moles of Ca(NO₃)₂ must be added to 1.0 L of a 0.12 M HF solution to begin precipitation of CaF₂(s)? For CaF₂, K_sp = 4.0 *10^-¹¹.

Accepted Answer

The answer of How many moles of Ca(NO₃)₂ must be...

Question 14

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^-: What is the concentration of CN^- at equilibrium?

Accepted Answer

This is a simple equilibrium problem involving the reaction of Cu⁺ with CN^- to form the complex ion Cu(CN)₃²^-. The balanced equation is: Cu⁺ + CN^- → Cu(CN)₃²^- We are given the initial concentrations of Cu⁺ and CN^-, and we assume that the reaction goes to equilibrium. Let x be the concentration of CN^- that reacts with Cu⁺ to form the complex ion. Then the equilibrium concentrations are: [Cu⁺] = (2.0 M) - x [CN^-] = (4.0 M) - x [Cu(CN)₃²^-] = x The equilibrium constant, K, for the reaction is: K = [Cu(CN)₃²^-]/([Cu⁺][CN^-]) Substituting the equilibrium concentrations into K, we get: K = x/[(2.0 - x)(4.0 - x)] We are asked to find the concentration of CN^- at equilibrium, so we need to solve for x. Since K is given as 1.6 x 10^-8, we can write: 1.6 x 10^-8 = x/[(2.0 - x)(4.0 - x)] Expanding the denominator and simplifying, we get a quadratic equation: 1.6 x 10^-8 = x/(8 - 6x - x^2) 1.6 x 10^-8 (8 - 6x - x^2) = x 1.28 x 10^-7 - 9.6 x 10^-8 x - 1.6 x 10^-8 x^2 = x 1.6 x 10^-8 x^2 + 1.056 x 10^-7 x - 1.28 x 10^-7 = 0 Using the quadratic formula, we get: x = 6.15 x 10^-9 M or x = -8.26 x 10^-6 M Since x represents a concentration, it must be positive, so we choose the positive root: x = 6.15 x 10^-9 M This is the concentration of CN^- that reacts with Cu⁺ to form the complex ion at equilibrium. The concentration of CN^- remaining at equilibrium is: [CN^-] = (4.0 M) - x = 3.99999999385 M ≈ 4.0 M (to three significant figures) Therefore, the best choice is D) 2.0 M.

Question 15

The K_sp for Mn(OH)₂ is 2.0 * 10^-¹³. At what pH will Mn(OH)₂ begin to precipitate from a solution in which the initial concentration of Mn²⁺ is 0.10 M?

Accepted Answer

To find the pH at which Mn(OH)₂ begins to precipitate, use the Ksp expression: Ksp = [Mn²⁺][OH⁻]². Given Ksp = 2.0 × 10⁻¹³ and [Mn²⁺] = 0.10 M, solve for [OH⁻]: [OH⁻] = √(Ksp / [Mn²⁺]) = √(2.0 × 10⁻¹³ / 0.10) = √(2.0 × 10⁻¹²) = 1.41 × 10⁻⁶ M. Convert [OH⁻] to pOH: pOH = -log(1.41 × 10⁻⁶) ≈ 5.85. Then, pH = 14 - pOH = 14 - 5.85 = 8.15.

Question 16

A 200-mL solution contains 0.018 mol each of I^-, Br^-, and Cl^-. When the solution is mixed with 200 mL of 0.24 M AgNO₃, how much AgCl(s) precipitates out?

Accepted Answer

The answer of A 200-mL solution contains 0.018 mol each...

Question 17

A 0.012-mol sample of Na₂SO₄ is added to 400 mL of each of two solutions. One solution contains 1.5 * 10^-³ M BaCl₂; the other contains 1.5 *10^-³ M CaCl₂. K_sp for BaSO₄ = 1.5 * 10^-⁹ and K_sp for CaSO₄ = 6.1 *10^-⁵. Which of the following statements is true?

Accepted Answer

BaSO₄ would precipitate because its ion product exceeds its Ksp, while CaSO₄ would not precipitate as its ion product is less than its Ksp.

Question 18

Titrating 30.00 mL of a saturated calcium iodate solution requires 28.91 mL of a 0.092 M solution of Na₂S₂O₃ according to the equation IO₃^- + 6S₂O₃²^- + 6H⁺$\rightarrow$ I^- + 2S₃O₆²^- + 3H₂O Calculate K_sp for Ca(IO₃)₂.

Accepted Answer

The answer of Titrating 30.00 mL of a saturated calcium...

Question 19

Calculate the solubility of Ag₂SO₄ [K_sp = 1.2*10^-⁵] in a 2.0 M AgNO₃ solution.

Accepted Answer

The solubility of Ag2SO4 in a solution containing a common ion (Ag+) is calculated using the Ksp expression: Ksp = [Ag+]^2[SO4^2-]. In a 2.0 M AgNO3 solution, [Ag+] is approximately 2.0 M. Solving for [SO4^2-] gives a solubility of 3.0 x 10^-6 mol/L.

Question 20

A 50.00-mL sample of 0.100 M Ca(NO₃)₂ is mixed with 50.00 mL of 0.200 M NaF. When the system has come to equilibrium, which of the following sets of conditions will hold? The K_sp for CaF₂ is 4.0 * 10^-¹¹.

Accepted Answer

The answer of A 50.00-mL sample of 0.100 M Ca(NO₃)₂...

Quiz 8: Applications of Aqueous Equilibria

The K_sp of Al(OH) ₃ is 2.0 * 10^-³². At what pH will a 0.2 M Al³⁺ solution begin to show precipitation of Al(OH) ₃?

Silver acetate (AgC₂H₃O₂) is a sparingly soluble salt with K_sp = 1.9 * 10^-³. Consider a saturated solution in equilibrium with the solid salt. Compare the effects on the solubility of adding to the solution either the acid HNO₃ or the base NH₃.

Sodium chloride is added slowly to a solution that is 0.010 M in Cu⁺, Ag⁺, and Au⁺. The K_sp values for the chloride salts are 1.9 * 10^-⁷, 1.6 10^-¹⁰, and 2.0 10^-¹³, respectively. Which compound will precipitate first?

The overall K_f for the complex ion Ag(NH₃) ₂⁺ is 1.7 * 10⁷. K_sp for AgI is 1.5 *10^-¹⁶. What is the molar solubility of AgI in a solution that is 2.0 M in NH₃?

The concentration of Al³⁺ in a saturated solution of Al(OH) ₃ at 25°C is 5.2 *10^-⁹ M. Calculate the K_sp for Al(OH) ₃.

Determine the pH of a solution prepared by mixing the following: 25.0 mL of 0.200 M HCl 75.0 mL of 0.100 M NaOH 50.0 mL of 0.200 M NaH₂PO₄ For H₃PO₄ K_a1 = 7.5 10^-³, K_a2 = 6.2 10^-⁸, and K_a3 = 4.8 *10^-¹³.

If 30 mL of 5.0 10^-⁴ M Ca(NO₃) ₂ is added to 70 mL of 2.0 10^-⁴ M NaF, will a precipitate form? (K_sp of CaF₂ = 4.0 *10^-¹¹)

The concentration of Mg²⁺ in seawater is 0.052 M. At what pH will 99% of the Mg²⁺ be precipitated as the hydroxide? (K_sp for Mg(OH) ₂ = 8.9 *10^-¹²)

Calculate the minimum concentration of 100.0 mL of sodium sulfate that is required to form a precipitate with 100.0 mL of 1.00 10^-² M calcium nitrate. K_sp for CaSO₄ is 6.1 10^-⁵.

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(NH₃) ₂⁺ at equilibrium is

The observed solubility of the salt MX in 4.0 M strong acid is 2.30 * 10^-³ M. The K_a value for the acid HX is 1.00 *10^-⁹. Calculate the value of K_sp for the salt MX.

How many moles of Ca(NO₃) ₂ must be added to 1.0 L of a 0.12 M HF solution to begin precipitation of CaF₂(s) ? For CaF₂, K_sp = 4.0 *10^-¹¹.

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^-: What is the concentration of CN^- at equilibrium?

The K_sp for Mn(OH) ₂ is 2.0 * 10^-¹³. At what pH will Mn(OH) ₂ begin to precipitate from a solution in which the initial concentration of Mn²⁺ is 0.10 M?

A 200-mL solution contains 0.018 mol each of I^-, Br^-, and Cl^-. When the solution is mixed with 200 mL of 0.24 M AgNO₃, how much AgCl(s) precipitates out?

A 0.012-mol sample of Na₂SO₄ is added to 400 mL of each of two solutions. One solution contains 1.5 * 10^-³ M BaCl₂; the other contains 1.5 10^-³ M CaCl₂. K_sp for BaSO₄ = 1.5 10^-⁹ and K_sp for CaSO₄ = 6.1 *10^-⁵. Which of the following statements is true?

Titrating 30.00 mL of a saturated calcium iodate solution requires 28.91 mL of a 0.092 M solution of Na₂S₂O₃ according to the equation IO₃^- + 6S₂O₃²^- + 6H⁺ $\rightarrow$ I^- + 2S₃O₆²^- + 3H₂O Calculate K_sp for Ca(IO₃) ₂.

Calculate the solubility of Ag₂SO₄ [K_sp = 1.2*10^-⁵] in a 2.0 M AgNO₃ solution.

A 50.00-mL sample of 0.100 M Ca(NO₃) ₂ is mixed with 50.00 mL of 0.200 M NaF. When the system has come to equilibrium, which of the following sets of conditions will hold? The K_sp for CaF₂ is 4.0 * 10^-¹¹.

Quiz 8: Applications of Aqueous Equilibria

The Ksp of Al(OH) 3 is 2.0 * 10-32. At what pH will a 0.2 M Al3+ solution begin to show precipitation of Al(OH) 3?

Silver acetate (AgC2H3O2) is a sparingly soluble salt with Ksp = 1.9 * 10-3. Consider a saturated solution in equilibrium with the solid salt. Compare the effects on the solubility of adding to the solution either the acid HNO3 or the base NH3.

Sodium chloride is added slowly to a solution that is 0.010 M in Cu+, Ag+, and Au+. The Ksp values for the chloride salts are 1.9 * 10-7, 1.6 *10-10, and 2.0 *10-13, respectively. Which compound will precipitate first?

The overall Kf for the complex ion Ag(NH3) 2+ is 1.7 * 107. Ksp for AgI is 1.5 *10-16. What is the molar solubility of AgI in a solution that is 2.0 M in NH3?

The concentration of Al3+ in a saturated solution of Al(OH) 3 at 25°C is 5.2 *10-9 M. Calculate the Ksp for Al(OH) 3.

Determine the pH of a solution prepared by mixing the following: 25.0 mL of 0.200 M HCl 75.0 mL of 0.100 M NaOH 50.0 mL of 0.200 M NaH2PO4 For H3PO4 Ka1 = 7.5 *10-3, Ka2 = 6.2 * 10-8, and Ka3 = 4.8 *10-13.

If 30 mL of 5.0 *10-4 M Ca(NO3) 2 is added to 70 mL of 2.0 *10-4 M NaF, will a precipitate form? (Ksp of CaF2 = 4.0 *10-11)

The concentration of Mg2+ in seawater is 0.052 M. At what pH will 99% of the Mg2+ be precipitated as the hydroxide? (Ksp for Mg(OH) 2 = 8.9 *10-12)

Calculate the minimum concentration of 100.0 mL of sodium sulfate that is required to form a precipitate with 100.0 mL of 1.00 *10-2 M calcium nitrate. Ksp for CaSO4 is 6.1 *10-5.

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(NH3) 2+ at equilibrium is

The observed solubility of the salt MX in 4.0 M strong acid is 2.30 * 10-3 M. The Ka value for the acid HX is 1.00 *10-9. Calculate the value of Ksp for the salt MX.

How many moles of Ca(NO3) 2 must be added to 1.0 L of a 0.12 M HF solution to begin precipitation of CaF2(s) ? For CaF2, Ksp = 4.0 *10-11.

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-: What is the concentration of CN- at equilibrium?

The Ksp for Mn(OH) 2 is 2.0 * 10-13. At what pH will Mn(OH) 2 begin to precipitate from a solution in which the initial concentration of Mn2+ is 0.10 M?

A 200-mL solution contains 0.018 mol each of I-, Br-, and Cl-. When the solution is mixed with 200 mL of 0.24 M AgNO3, how much AgCl(s) precipitates out?

A 0.012-mol sample of Na2SO4 is added to 400 mL of each of two solutions. One solution contains 1.5 * 10-3 M BaCl2; the other contains 1.5 *10-3 M CaCl2. Ksp for BaSO4 = 1.5 * 10-9 and Ksp for CaSO4 = 6.1 *10-5. Which of the following statements is true?

Titrating 30.00 mL of a saturated calcium iodate solution requires 28.91 mL of a 0.092 M solution of Na2S2O3 according to the equation IO3- + 6S2O32- + 6H+ →\rightarrow→ I- + 2S3O62- + 3H2O Calculate Ksp for Ca(IO3) 2.

Calculate the solubility of Ag2SO4 [Ksp = 1.2*10-5] in a 2.0 M AgNO3 solution.

A 50.00-mL sample of 0.100 M Ca(NO3) 2 is mixed with 50.00 mL of 0.200 M NaF. When the system has come to equilibrium, which of the following sets of conditions will hold? The Ksp for CaF2 is 4.0 * 10-11.

The K_sp of Al(OH) ₃ is 2.0 * 10^-³². At what pH will a 0.2 M Al³⁺ solution begin to show precipitation of Al(OH) ₃?

Silver acetate (AgC₂H₃O₂) is a sparingly soluble salt with K_sp = 1.9 * 10^-³. Consider a saturated solution in equilibrium with the solid salt. Compare the effects on the solubility of adding to the solution either the acid HNO₃ or the base NH₃.

Sodium chloride is added slowly to a solution that is 0.010 M in Cu⁺, Ag⁺, and Au⁺. The K_sp values for the chloride salts are 1.9 * 10^-⁷, 1.6 10^-¹⁰, and 2.0 10^-¹³, respectively. Which compound will precipitate first?

The overall K_f for the complex ion Ag(NH₃) ₂⁺ is 1.7 * 10⁷. K_sp for AgI is 1.5 *10^-¹⁶. What is the molar solubility of AgI in a solution that is 2.0 M in NH₃?

The concentration of Al³⁺ in a saturated solution of Al(OH) ₃ at 25°C is 5.2 *10^-⁹ M. Calculate the K_sp for Al(OH) ₃.

Determine the pH of a solution prepared by mixing the following: 25.0 mL of 0.200 M HCl 75.0 mL of 0.100 M NaOH 50.0 mL of 0.200 M NaH₂PO₄ For H₃PO₄ K_a1 = 7.5 10^-³, K_a2 = 6.2 10^-⁸, and K_a3 = 4.8 *10^-¹³.

If 30 mL of 5.0 10^-⁴ M Ca(NO₃) ₂ is added to 70 mL of 2.0 10^-⁴ M NaF, will a precipitate form? (K_sp of CaF₂ = 4.0 *10^-¹¹)

The concentration of Mg²⁺ in seawater is 0.052 M. At what pH will 99% of the Mg²⁺ be precipitated as the hydroxide? (K_sp for Mg(OH) ₂ = 8.9 *10^-¹²)

Calculate the minimum concentration of 100.0 mL of sodium sulfate that is required to form a precipitate with 100.0 mL of 1.00 10^-² M calcium nitrate. K_sp for CaSO₄ is 6.1 10^-⁵.

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(NH₃) ₂⁺ at equilibrium is

The observed solubility of the salt MX in 4.0 M strong acid is 2.30 * 10^-³ M. The K_a value for the acid HX is 1.00 *10^-⁹. Calculate the value of K_sp for the salt MX.

How many moles of Ca(NO₃) ₂ must be added to 1.0 L of a 0.12 M HF solution to begin precipitation of CaF₂(s) ? For CaF₂, K_sp = 4.0 *10^-¹¹.

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^-: What is the concentration of CN^- at equilibrium?

The K_sp for Mn(OH) ₂ is 2.0 * 10^-¹³. At what pH will Mn(OH) ₂ begin to precipitate from a solution in which the initial concentration of Mn²⁺ is 0.10 M?

A 200-mL solution contains 0.018 mol each of I^-, Br^-, and Cl^-. When the solution is mixed with 200 mL of 0.24 M AgNO₃, how much AgCl(s) precipitates out?

A 0.012-mol sample of Na₂SO₄ is added to 400 mL of each of two solutions. One solution contains 1.5 * 10^-³ M BaCl₂; the other contains 1.5 10^-³ M CaCl₂. K_sp for BaSO₄ = 1.5 10^-⁹ and K_sp for CaSO₄ = 6.1 *10^-⁵. Which of the following statements is true?

Titrating 30.00 mL of a saturated calcium iodate solution requires 28.91 mL of a 0.092 M solution of Na₂S₂O₃ according to the equation IO₃^- + 6S₂O₃²^- + 6H⁺ $\rightarrow$ I^- + 2S₃O₆²^- + 3H₂O Calculate K_sp for Ca(IO₃) ₂.

Calculate the solubility of Ag₂SO₄ [K_sp = 1.2*10^-⁵] in a 2.0 M AgNO₃ solution.

A 50.00-mL sample of 0.100 M Ca(NO₃) ₂ is mixed with 50.00 mL of 0.200 M NaF. When the system has come to equilibrium, which of the following sets of conditions will hold? The K_sp for CaF₂ is 4.0 * 10^-¹¹.