Acid Base Titrations

Acid and Base Titrations: Preparing Standardized Solutions
Introduction:
This experiment focuses on titrations of acids and bases. A titration depends on addition of a known volume of solution and is a type of volumetric analysis. Many titrations involve either acid-base reactions or oxidation-reduction reactions. In this experiment we do one of each. We monitor the pH of the reaction with the use of a color indicator. We also learn about the standardization of bases (NaOH) and acids (HCl) which is basically making a dilution to change the molarity. The first reaction consists of titrating sodium hydroxide (NaOH) into potassium acid phthalate (KHP or K[HC8H4O4]):
K+[HC8H4O4]- + Na+OH- => K+Na+[HC8H4O4]- + H2O
The second titration we did was hydrochloric acid (HCl) with sodium hydroxide (NaOH):
HCl(aq) + NaOH(aq) => NaCl(aq) + H2O(l)
Procedure:
You need to calculate the volume of 3 M NaOH needed to make 500 mL of a 0.1 M solution. Mix the solution in a 500 mL flask and place the stopper on it for time being. Take two 250 mL flasks and label each one as “Sample 1” and “Sample 2.” Then take a plastic weigh boat, place it on a scare and tare (zero) it. Carefully add 0.4 to 0.5 grams of potassium acid phthalate to the weigh boat and record the final mass to 4 decimal places. Place this into the flask labeled “Sample 1.” Repeat the same thing for the other flask, “Sample 2.” Be sure to keep track of which flask has which mass of KHP. Rinse and clean the buret with 3-4 mL of the NaOH solution. Then fill the buret up to the zero level mark. With a waste beaker underneath, open the stopcock completely to fill the buret tip and remove air bubbles. Adjust the liquid level in the buret so that it’s between 0.00-2.00 mL. Record this initial reading in your notebook. Estimate all buret readings to 2 decimal places. Add about 50 mL of distilled water to the flask labeled “Sample 1” and gently swirl until the KHP has completely dissolved. Then add 3 or 4 drops of phenolphthalein indicator solution and begin your titration. The NaOH solution should run into the flask little at a time as the flask is gently swirled to mix the solution. As you approach the end point, the solution will turn pink very briefly. Add very small amounts until the solution stays very light pink for more than ten seconds. Record the final volume then and calculate the volume of base necessary to reach the end point of the reaction. Titrate sample two in the same exact way. (This titration should proceed a little faster since we now know the approximate volume needed to reach the endpoint). Just remember to record the initial volume, take it slow towards the end point, and record final volume. After you calculated the molarity and taken the average, take your NaOH and place it in a bottle to save for the next experiment. Clean out the 250 flasks and the 500 mL flask for use in the next titration. Calculate the volume of 3 M HCl needed to make 500 mL of a 0.1 M solution. Measure the volume of the HCl into a flask and then add enough water to make 500 mL. Remember to place the stopper on and swirl gently. Switch the buret clamp around so you now have the empy buret while the NaOH buret is on the other side (We will go back to that one). Rinse the empty buret with a few mL of the HCl solution. And fill it with about 40-45 mL of HCl. Record the initial volume. Take the now cleaned-out flask labeled “Sample 1” and carefully add about 35 mL of HCl into the flask. Record the final volume and calculate the exact volume of HCl added to the flask. Place 3 or 4 drops of phenolphthalein solution into the flask as well. Switch back over to the NaOH buret and begin the titration which follows the same concept of the KHP titration. Remember to record initial and final volumes of the HCl. Repeat the same process for the second flask, “Sample 2.” Calculate the molarity of the acid solution for the two results. Afterwards, take the second empty bottle and save the HCl solution for he next lab experiment as well.
Data and Results:
Molecular Weight of KHP: 204.22 g/mol
I. Standardization of the NaOH solution
SAMPLE
1
2
Wt. KHP
0.4628 g
0.4926 g mmol of KHP
2.266 mmol
2.412 mmol
Final buret reading
24.3 mL
30.2 mL
Initial buret reading
0.7 mL
5.65 mL
Volume of NaOH used
23.6 mL
24.55 mL
Molarity of NaOH
0.096 M
0.098 M
Average Molarity
0.097 M

II. Standardization of the HCl solution
SAMPLE
1
2
Final Buret reading of HCl
43.75 mL
39.8 mL
Initial Buret reading of HCl
8.6 mL
4.5 mL
Volume of HCl added to flask
35.15 mL
35.3 mL
Final Buret reading of NaOH
31.9 mL
42.1 mL
Initial Buret reading of NaOH
1.4 mL
8.2 mL
Volume NaOH used in titration
30.5 mL
33.9 mL
Calculated Molarity of HCl
0.0842 M
0.0932 M
Average
0.0887 M

Analysis and Discussion: Since NaOH solution is used for the titration of the unknown and also in the standardization of HCl, its standardization is important as well. In order to measure the molarity of the NaOH, I reacted it with the KHP with is a primary standard. I made two very precise measurements to make the molarity a much more accurate number. With molarity, a 1 M solution has one mole per liter of solution. It’s easier to work with smaller numbers so I converted them to millimoles and milliliters:
Molarity = (moles / liters) = (10-3 x moles / 10-3 x liters) = (mmol / mL)
In the titration, the mmol of KHP equaled the mmol of NaOH added to the solution. At the end of the first titration, I had the information needed to calculate the molarity of the NaOH solution using this equation:
M base = (g KHP x 103) / (mL base)(mw KHP)
Conclusion:
The acid base titrations provided key insights into the behavior of both the acid and the base with each other and with water. It was observed that when a strong acid and a strong base react with each other in equivalent concentrations, the pH of the solution is neutral, reaching an end point. It was also shown that the molarity of the bases and the acids could be determined if a few facts were known about the acid-base system. Finally, using all the information and data we obtained from the experiment, we put it all together. These observations and accomplishments demonstrate important qualities of the acid-base equilibrium. These qualities have real effects on real systems in the world. The ability to predict pH, molecular weight, molarity, or any of the other values that were used in this experiment gives great power and understanding to understanding the chemical system.