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9 Nov 2019

Silver acetate, AgC2H3O2, is a sparingly soluble salt with Ksp =1.9x 10^-3. Consider a saturated solution in equilibrium with thesolid salt. Compare the effects on the solubility of adding to thesolution either the acid HNO3 or the base NH3.

Either substance would increase the solubility.

NH3 would decrease the solubility, butHNO3 would increase it.

Either substance would decrease the solubility.

NH3 would increase the solubility, butHNO3 would decrease it.

NH3 would increase the solubility, butHNO3 would have virtually no effect

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Related questions

1) Consider the following equation: Rate = k [CH3I]2 [C15H11N]

i) If the concentration of C15H11N is doubled, does the rate constant:

A) Increase

B) Decrease

C) Remain unchanged

ii) If the concentration of of CH3I is decreased by a factor of 3, the time it takes for the reaction to finish will:

A) Increase by a factor of 3

B) Increase by a factor of 9

C) Decrease by a factor of 9

D) Remain Unchanged

iii) If the temperature increases how will the activation energy change:

A) It will increase

B) It will decrease

C) It will remain unchanged

2) During a titration of a weak acid with a strong base you add an additional small amount of base after reaching the equivalence point.

i) How does the pH of the solution change?

A) Increase

B) Decrease

C) Remain unchanged

ii) How does the Ka of the acid change?

A) Increase

B) Decrease

C) Remain unchanged

3) You combine a solution of silver nitrate (AgNO3) and a solution of sodium carbonate (Na2CO3) to form solid silver carbonate (Ag2CO3) and aqueous sodium nitrate (NaNO3), the net ionic equation is listed below. You observe the white solid, Ag2CO3 in the test tube. You add a strong acid (HNO3) to the mixture and it reacts with the carbonate ions as listed in the equations below. What would you expect to observe when adding the strong acid to the test tube?

2 Ag+(aq) + CO32-(aq) <--> Ag2CO3(s)

2 H+(aq) + CO32-(aq) <--> H2CO3

H2CO3 -> H2O(l) + CO2(g)

A) An increase in the amount of white solid

B) A color change of the solid

C) A decrease in the amount of white solid

D) No change in the amount of white solid

4) Ksp of BaF2 at room temperature is 1.84 x 10-7. It is also known that the reaction of dissolution of BaF2 is endothermic. Imagine that you have a saturated solution of BaF2 at equilibrium, that has some undissolved BaF2 precipitated on the bottom of the test tube.

4.1) If the temperature is increased how would that affect:

i) The solubility of BaF2

A) Increase

B) Decrease

C) Remain Unchanged

ii) The solubility constant of BaF2

A) Increase

B) Decrease

C) Remain unchanged

4.2 An additional amount of BaF2 was added to the solution how would that affect:

i) The solubility of BaF2

A) Increase

B) Decrease

C) Remain unchanged

ii) The solubility constant of BaF2

A) Increase

B) Decrease

C) Remain unchanged

4.3 If NaF is added to the solution how will it affect:

i) The solubility of BaF2

A) Increase

B) Decrease

C) Remain Unchanged

ii) The solubility constant of BaF2

A) Increase

B) Decrease

C) Remain unchanged

Part A - Calculate the molar solubility in water

Mg(OH)2 is a sparingly soluble salt with a solubility product, Ksp, of 5.61×10?11. It is used to control the pH and provide nutrients in the biological (microbial) treatment of municipal wastewater streams.

Based on the given value of the Ksp, what is the molar solubility of Mg(OH)2 in pure H2O?

Express your answer with the appropriate units.

molar solubility =

2.41×10?4 M

Calculate the molar solubility in NaOH

Based on the given value of the Ksp, what is the molar solubility of Mg(OH)2 in 0.120 M NaOH?

Express your answer with the appropriate units.

molar solubility =

3.90×10?9 M

Part C: Calculate how many times more soluble Mg(OH)2 is in pure water

Based on the given value of the Ksp, calculate the ratio of solubility of Mg(OH)2 dissolved in pure H2O to Mg(OH)2 dissolved in a 0.120 M NaOH solution.

A buffer is a mixed solution of a weak acid or base, combined with its conjugate. Note that this can be understood essentially as a common-ion problem: The conjugate is a common ion added to an equilibrium system of a weak acid or base. The addition of the conjugate shifts the equilibrium of the system to relieve the stress of the added concentration of the common ion. In a solution consisting of a weak acid or base, the equilibrium shift also results in a pH shift of the system.

It is the presence of the common ion in the system that results in buffering behavior, because both added H+ or OH? ions can be neutralized.

Part D

What is the pH change of a 0.240 M solution of citric acid (pKa=4.77) if citrate is added to a concentration of 0.175 M with no change in volume?

QUESTION 1

Equal molar amounts of nitric acid (a strong acid) and methylamine (a weak base) are combined in an aqueous solution.

The pH of the resulting solution is:

greater than 7

close to 7

less than 7

QUESTION 2

A 1.00 L aqueous solution contains 0.10 mol acetic acid and 0.10 mol sodium acetate. The initial pH is 4.74. Calculate the pH after the addition of 0.049 mol of OH-.

QUESTION 3

Using the Henderson-Hasselbalch equation, calculate the pH of a buffer solution that contains 0.96 M NH4Cl and 0.33 M NH3.

QUESTION 4

A 10.0 mL volume of 0.49 M HNO3 is combined with a 15.0 mL volume of 0.58 M NaOH. Calculate the pH.

QUESTION 5

A diprotic acid is being titrated with a strong base. A volume of 0.0500 L of 0.0750 M diprotic acid (H2A+) is combined with 0.0150 L of 0.250 M NaOH.

Ka1 for the diprotic acid is 3.6 x 10-3.

Ka2 for the diprotic acid is 4.7 x 10-7.

Calculate the pH of the solution.

QUESTION 6

A saturated solution of Ca3(PO4)2 has a concentration of Ca2+ = 5.65 x 10-7 M and a concentration of PO43- = 9.80 x 10-8 M. Calculate Ksp for Ca3(PO4)2.

Answer in exponential notation. Examples:

Scientific Notation Exponential Notation

1.00 x 10-10 1.00e-10

5.00 x 105 5.00e5

QUESTION 7

A solution is prepared by combining 0.125 L of 0.020 M Pb(NO3)2 and 0.125 L of 0.050 M PbBr2. Does a precipitate of PbBr2 form? Ksp of PbBr2= 6.6 x 10-6

Yes, a precipitate forms.

No, a precipitate does not form.

QUESTION 8

Calculate the molar solubility of MgF2 in 0.37 M NaF at 25°C. Ksp = 7.4 x 10-11.

Answer in exponential format.

QUESTION 9

Determine the concentration of Ag+ in a solution prepared by adding 0.10 mol AgNO3 to 1.0 L of 3.9 M NH3.

Ag+(aq) + NH3(aq) [left right double arrow] AgNH3+(aq) K1 = 2.1 x 103

AgNH3+(aq) + NH3(aq) [left right double arrow] Ag(NH3)2+(aq) K2 = 8.1 x 103

Ag+(aq) + 2 NH3(aq) [left right double arrow] Ag(NH3)2+(aq) Kf = 1.7 x 107

QUESTION 10

Determine the molar solubility of MgCO3 at 25°C in pure water.

Ksp = 6.8 x 10-6

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