Textbook Question
Use the given molar solubilities in pure water to calculate Ksp for each compound. b. PbF2; molar solubility = 5.63⨉10-3 M c. MgF2; molar solubility = 2.65⨉10-4 M
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Ch.17 - Aqueous Ionic Equilibrium
Chapter 17, Problem 7
A 15.0-mL sample of 0.100 M Ba(OH)2 is titrated with 0.125 M HCl. Calculate the pH for at least five different points throughout the titration curve and sketch the curve. Indicate the volume at the equivalence point on your graph.
Verified step by step guidance1
<insert step 1: Determine the initial pH of the Ba(OH)_2 solution before any HCl is added. Calculate the concentration of OH^- ions using the formula [OH^-] = 2 * [Ba(OH)_2] because each Ba(OH)_2 dissociates into two OH^- ions. Then, find the pOH and convert it to pH using pH + pOH = 14.>
<insert step 2: Calculate the pH after adding a small volume of HCl, such as 5.0 mL. Determine the moles of HCl added and the moles of OH^- initially present. Subtract the moles of HCl from the moles of OH^- to find the remaining moles of OH^-. Calculate the new concentration of OH^- and find the pH.>
<insert step 3: Determine the pH at the equivalence point. At this point, all the OH^- ions have reacted with HCl, forming water and BaCl_2. Calculate the volume of HCl needed to reach the equivalence point using the stoichiometry of the reaction. Since the solution is neutral at the equivalence point, the pH is 7.>
<insert step 4: Calculate the pH after adding a volume of HCl beyond the equivalence point, such as 20.0 mL. Determine the excess moles of HCl and calculate the concentration of H^+ ions in the solution. Use this concentration to find the pH.>
<insert step 5: Sketch the titration curve using the calculated pH values at different points. Label the initial pH, the pH at the equivalence point, and the pH after adding excess HCl. Indicate the volume of HCl at the equivalence point on the graph.>
Related Practice
Textbook Question
Use the given molar solubilities in pure water to calculate Ksp for each compound. b. Ag2SO3; molar solubility = 1.55⨉10-5 M c. Pd(SCN)2; molar solubility = 2.22⨉10-8 M
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Textbook Question
Refer to the Ksp values in Table 17.2 to calculate the molar solubility of each compound in pure water. b. Mg(OH)2 c. CaF2
Open Question
A 20.0-mL sample of a 0.125 M diprotic acid (H2A) solution is titrated with 0.1019 M KOH. The acid ionization constants for the acid are Ka1 = 5.2 * 10^-5 and Ka2 = 3.4 * 10^-10. At what added volume of base does each equivalence point occur?
Open Question
After a 120.0-mL sample of a solution that is 2.8 * 10^-3 M in AgNO3 is mixed with a 225.0-mL sample of a solution that is 0.10 M in NaCN and the solution reaches equilibrium, what concentration of Ag+(aq) remains?
Open Question
A 0.5224-g sample of an unknown monoprotic acid was titrated with 0.0998 M NaOH. The equivalence point of the titration occurred at 23.82 mL. Determine the molar mass of the unknown acid.