CHEM 1B Chapter Notes - Chapter 9: Buffer Solution, Titration, Equivalence Point

The kinetics of the reaction below were studied. The volumes of 0.040 M KBrO3, 0.040 M KI,

0.00050 M Na2S2O3, 0.040 M KCl, H2O and a buffer solution employed in each trial are shown

in the table below. One drop of 1% aqueous starch solution was also added in each trial. The

buffer solution was a pH = 4.74 solution containing 0.50 M acetic acid and 0.50 M sodium acetate; Ka = 1.8 x 10−5 for acetic acid. Each trial was performed at 25◦C. For each trial, the reaction mixture was initially colorless after combining all solutions; appearance of a blue solution color in the reaction mixture signaled that all of the thiosulfate ion initially present had reacted and marked the end of the timing period. The initial reaction rate for each trial was calculated from the change in molar concentration of thiosulfate ion from each trial; the equation for calculating Initial Rate is shown below, where Δt is the elapsed time between combining all solutions and appearance of the blue solution color. Each solution volume was measured to the nearest 0.1 mL. (ie 5 mL shown in the table is 5.0 mL.)

BrO3- (aq) + 6I−(aq) + 6H+(aq) → 3H2O(l) + Br−(aq) + 3I2(aq)

Initial rate = - 13∆S2O32-∆t

Volume (mL) Used

a)Which pair of trials could be used to determine the order for I−? Justify your choice.

b)How would you determine the order for H+? Be as specific as possible.

c)What would be the value for the initial rate if 5 mL KBrO3 solution, 5 mL KI solution, 5 mL Na2S2O3 solution, 10 mL KCl solution, 15 mL pH = 4.74 buffer solution, and 1 drop of 1% starch solution were used?

d)Why do Trials 1 and 4 have the same initial rate?

e)What time was required for the color change to occur in Trial 1?

f)What was the purpose of adding 5 mL of 0.040 M KCl in Trial 2?

g)How would you experimentally verify that KCl does not cause the color change from colorless to blue? Be as specific as possible.

I am unable to solve these questions due to a time constraint. I received help on part 1 of the unknown acid. Need help with the rest! I really appreciate it!

Lab: ACID-BASE TITRATION

Objective: To use titration to determine in Part 1, the concentration (molarity) of an unknown acid solution and in Part 2, the purity of a sample of KHP acid. Background: Titration is a volumetric technique used to determine the concentrations of solutions, molar masses of solids, and purity of samples. A titration inolves the addition of a titrant (a solution of known concentration) to an analyte ( a solution of unknown concentration), or vice versa, using a piece of glassware called a burette. The titration is carried out until it reaches an equivalence point (the exact point at which the reaction between the two solutions is complete). A chemical indicator is often used to aid in the identification of the equivalence point. An indicator changes color upon reaching the equivalence point of endpoint. Since we will be carrying out an acid-base titration, the indicator must change color upon reaching the endpoint at a specific pH. An example of such an indicator is phenolphthalein, which changes from colorless to pink near the equivalence point when pH approaches 7. It should be noted that the indicator selected is dependent upon a given titration, so that the observed color change is close to the ideal equivalence point. Once the titration is completed, we can use the volumes measured for each solution, as well as the concentration of the titrant, to determine the concentration, molar mass, or purity of the unknown analyte. Procedure Note: During lab, record all measurements and data in a clearly labeled table. Part 1: 1. Obtain ~80mL of ~0.1M of base (sodium hydroxide, NaOH) into a labeled beaker. Make sure to report the actual concentration from the bottle. 2. Obtain ~25mL of unknown acid into a second labeled beaker. Record whether your acid is monoprotic, diprotic, or triprotic. 3. Using a funnel, or carefully pouring, fill the burette with sodium hydroxide NaOH. Make sure a labeled waste beaker is under the buretter to catch any drippings. Rinse your burette by filling it with ~5-10mL of NaOH. Turn the knob to let ~2 mL out the tip into the waste beaker. Make sure your waste beaker is under the burette to catch any drips. Next, turn the Burette upside down and pour the remaining NaOH into your waste beaker while twisting the burette to coat the inside with NaOH. 4. Fill your burette with sodium hydroxide. Turn the knob to fill the tip of the burette with NaOH. Record the initial volume to the correct number of significant figures. 5. Using a volumentric pipette, transfer 10.00 mL of your unknown acid into the Erlenmeyer flask. 6. Add 3 drops of pH indicator and magnetic stirrer to the Erlenmeyer. 7. Under the burette, place the Erlenmeyer on a hot plate and turn on the stir knob. Heat should be off. 8. Titrate the unknown acid until the solution turns pink and remains pink. 9. Record the final volume on the burette to the correct number of significant figures. 10. Repeat (4) to (7). Pre-plan the whole lab and who will do which task. MY RESULTS: PART 1 The actual concentration of NAOH is 0.5939, 80 mL Initial Final 24.9 mL 44.5mL 26.1 mL 45.57mL PART 2 Impure KHP (SAMPLE B ) Only Paper: 0.280 g Impure KHP Mass: 0.526 g Initial volume: 22.30 mL + water +26.10 mL Final Volume: 26.80 mL QUESTIONS: 1. Calculate the concentration (molarity) of your unknown acid. 2. Calculate how many grams in your impure KHP sample was indeed KHP. 3. Calculate the percent (by mass) KHP in the sample. Part 2: 1. What would happen to the final result( state what the final result is ) if the tip of the burette was not filled when reading the initial volume? 2. What would happen to the final answer if you used Ca(OH)2 instead of NaOH and DID NOT know it? 3. What would happen to the final result if the Erlenmeyer flask was contaminated with other types of acid? 4. Recalculate the concentration of your unknown acid if it was given as a diprotic acid. 5. Draw a molecular level picture (showing Na+, H+, OH-, H2O, etc...) of what is in your beaker before, at half point, at equivalent point, and at end point of the titration. Part 3: 1. What would have happened to the final result if the impurity in the sample was also acidic? 2. What would have happend to your final result if you used ~25mL of water to dissolve the solid acid instead of 20mL? Summary data: Create a clearly labeled table of all your final data results for each run, and athe average of multiple runs. Part 1: Unknown #, Vol base, M base, Vol acid, M acid. Part 2: Sample #, % KHP (purity)