2. Generate a titration curve (using a graph paper) for the reaction between 0.0195 M C6H₂O2 (quinone) with 25.0 mL 0.050 M Ti³+. The solution is 1.0 M in HCIO4 throughout the titration. Given, EºTio²+/T₁³+ = + 0.1 V and Eº C6H40₂/C₂H₁(OH)₂ = +0.699 V. Show ALL the calculation works of the following stages : (a) 6 points for pre-equivalence including 0.00 mL

An aqueous solution is prepared in which 0.00113 mol Ni(NO3)2 and 0.484  mol NH3 are dissolved in a total volume of 1.00 L.Kf for Ni(NH3)62+ is equal to 5.5 × 108. Write the expression for the dissociation constant needed to determine the concentration of Ni(H2O)62+ ions at equilibrium in the solution formed in Part 1.   Kd​ = [Ni2+][NH3​]^6/[Ni(NH3​)^2+ 6​]​   Calculate the concentration of Ni(H2O)62+ ions at equilibrium in the solution formed.

Titration of 25.0 mL of a 0.0500 M Zn2+ solution with 0.0550 M EDTA in a solution buffered at pH 8. Assume that the temperature is 25 oC and that the formation constant for Zn2+ is 3.13 x 1016 at this temperature. What is the pZn of the solution after 30 mL of titrant have been added?

What is the solubility of calcium fluoride in a buffer solution containing 0.650 M HCHO2 (formic acid) and 0.20 M NaCHO2? K„(CaF2) = 3.4 x 10-11 K.(HF) = 6.8 × 10¬4 K.(HCHO2) =1.7 × 10¬4 (Hint: Consider the equation CaF2 (8) + 2H30*(ag) = Ca²+ (aq) + 2HF(aq) + 2H20(1) and solve the equilibrium problem.) Solubility = |м

The % w/w I– in a 0.6712-g sample was determined by Volhard titration. After adding 50.00 mL of 0.05619 M AgNO3 and allowing the precipitate to form, the remaining silver was back titrated with 0.05322 M KSCN, requiring 35.14 mL to reach end point. AgNO3 + I– → AgI (s) + NO3– Ag+ + SCN– → AgSCN                                                                                                                                                     How many millimoles of AgNO3 reacted with I–?                                                                                                                                                     How many milligrams of I– (MM=126.9 g/mol) were present in the sample?                                                                                                                                                     What is the % w/w I– in the sample?

Calculate the solubility of AgCN(s) (Ksp = 2.2 x 10-12) in a solution containing 1.0 MH+. (Ka for HCN is 6.2 x 10-10.)

0.683 grams of iron mineral dissolves with acid. Fe2+ is titrated with Ce(SO4)2 solution. Calculate the percentage of 38.5 mL of 0.161 M titrant consumed in the titration (Fe: 56 g/mol). 25.4150.82C.76,23D.32,45

Calculate the equilibrium concentrations of Fe3+, Fe2+, Mn2+, MnO4 in a solution prepared by mixing 20.0 mL of 0.100 M KMnO4 and 20.0 mL of 0.100 M FeSO4 at pH 0.00. EMnO/Mn²+ = 1.51 V, Efe+/Fe²+ = 0.771 V MnO + 5Fe2+ + 8H+ - Mn²+ + 5Fe³++ 4H₂O 4

Calculate the molar solubility of calcium oxalate in a solution that has been buffered so that its pH is constant and equal to 4.00. Data: Kps of CaC2O4 = 1.7x10 -9 ; Ka1 of H2C2O4 = 5.60x10-2 and Ka2 = 5.42x10-5 . Use systematic treatment.

Calculate the pH at the equivalence point for the titration of 0.25 M CH3COOH with 0.25 M NaOH. (For CH3COOH, K₂= 1.8 x <10-5.) Hint: This is the pH of a salt solution. Assume the concentration of CH3COONa is 0.125 M. Kw = K₂ x Kb = 1.0 x 10-14 8.92 8.72 8.52 8.32 None of the above