Titration of Strong and weal acid
Procedures
Experiment 1
The procedures described in this lab assume that you have already done the Titration Tutorial and are familiar with the technique. If you have not yet done the Titration Tutorial Lab, please do it now.
Place a clean Erlenmeyer flask from the Containers shelf onto the workbench.
Titrate: Add 25 mL hydrochloric acid (HCl) of unknown concentration from the Materials shelf to the Erlenmeyer flask.
Indicator: Add 2 drops of phenolphthalein to the flask.
Place a burette from the Containers shelf onto the workbench.
Titrant: Add 50 mL of 1 M sodium hydroxide (NaOH) solution from the Materials shelf to the burette.
Place a pH Meter from the Instruments shelf onto the flask to attach it. Record the initial pH of the solution in your Lab Notes.
Move the Erlenmeyer flask on the lower half of the burette - this will connect them.
Perform a Coarse Titration:
Add large amounts of the sodium hydroxide from the burette by pressing and holding the black knob at the bottom of the burette. Check the volume remaining in the burette after each addition of NaOH.
For every 1 mL added, record the volume and pH. As the sodium hydroxide is added to the hydrochloric acid solution the pH increases.
Watch for the phenolphthalein color change. When the end point is either reached or crossed, the pink indicator color appears. Record the burette volume and pH for both the last volume where the solution was colorless and the first volume where the solution was pink.
This gives you the range within which to do the fine titration.
Add additional titrant in 1 mL increments and record the pH at each point. Add a total of 5 mL extra. This data is useful in constructing and analyzing the titration curve later.
Clear your workbench.
Repeat steps 1 through 8 to setup for the next titration. You can skip step 6; the pH meter is not mandatory for the fine titration.
Perform a Fine Titration
Quickly add the titrant by clicking and holding the burette's black knob. Add 1 mL less than was needed for a color change in your course titration. This is near, but not yet at, the titration's end point.
Begin to add the titrant one drop at a time. Continue several drops after the end point is reached to confirm .
Move the Erlenmeyer flask and burette into the recycling bin.
Repeat the fine titration once more. Record the results.
If the results from the two fine titrations do not closely agree do a third fine titration to determine which titration was done incorrectly.
Experiment 2
In this procedure you will titrate a weak acid, acetic acid (CH3COOH), to see the difference in the titration curve of a weak acid as opposed to a strong acid.
Repeat the procedure in Experiment 1, steps 1 through 10, using these materials:
Titrate: 5 mL Acetic acid of unknown concentration, diluted with 20 mL water
Indicator: 2 drops phenolphthalein
Titrant: 50 mL of 1 M sodium hydroxide
Experiment #1
1. Create a graph of pH vs. volume of added NaOH (mL) to the HCl solution of unknown concentration, to produce a titration curve.
2. Find the equivalence point on the graph. What is the pH and equivalence volume of NaOH at this point?
3. a. At the equivalence point, how many moles of 1 M NaOH were added to the flask. To get the number of moles, multiply the molarity by the number of LITERS of NaOH added.
b. At the equivalence point, the number of moles of NaOH equals the number of moles of HCl in the flask. How many moles of HCl were in the flask?
4. Using the number of moles of HCl in the flask and the fact that 25 mL of acid were added, calculate the molarity of the HCl solution. (molarity = moles / LITER).
5. An alternative way of calculating the molarity of the HCl solution is to use the fact that as a strong acid, all of the H+ and Cl- dissociate. The pH = -log10([H+]) and thus the concentration of HCl is given by [HCl] = 10^-pH. Using this relationship and the pH of the solution before any NaOH was added, calculate the molarity of the HCl solution.
6. How do the two molarities compare? What could explain any discrepancies? Which calculated molarity is more accurate?
Experiment 2
1. Create a graph of pH s. volume of added NaOH (mL) to the acetic acid solution of unknown concentration, to produce a titration curve.
2. Find the equivalence point on the graph. What is the equivalence volume of NaOH at this point?
3. a. Calculate the unknown molarity of the diluted acetic acid from the volumes of acid and base at the equivalence point and the molarity of the NaOH Ma × Va = Mb × Vb.
b. Once you find the molarity of your diluted solution use that to calculate the molarity of the original solution using the equation M1 × V1 = M2 × V2 a second time.
4. In experiment 1, you were able to calculate the concentration of the HCl solution using the initial pH. Would this same approach work with the acetic acid? Why or why not?
5. How did the titrations of the two acids compare? Did the results align with the differences between strong and weak acids?