Acid and Base Titrations: Preparing Standardized Solutions Introduction:
This experiment focuses on titrations of acids and bases. Titrations involve the addition of a known volume of solution and are a type of volumetric analysis. Many titrations involve acid-base reactions or oxidation-reduction reactions. In this experiment, we perform one of each. We monitor the pH of the reaction using a color indicator. We also learn about the standardization of bases, such as NaOH, and acids, like HCl, which involves dilution to change the molarity. The first reaction involves titrating NaOH into KHP (potassium acid phthalate): K+[HC8H4O4]- + Na+OH- => K+Na+[HC8H4O4]- + H2O
The procedure for the second titration involved combining hydrochloric acid (HCl) and sodium hydroxide (NaOH) according to the following equation: HCl(aq) + NaOH(aq) => NaCl(aq) + H2O(l).
Procedure:
To make a 0.1 M solution, calculate the volume of 3 M NaOH needed for 500 mL. Mix the solution in a 500 mL flask and cover it with a stopper temporarily. Take two 250 mL flasks and label them as “Sample 1” and “Sample 2.” Place a plastic weigh boat on a scale and zero it. Carefully add 0.4 to 0.5 grams of potassium acid phthalate to the weigh boat and record the final mass with four decimal places. Transfer this to the flask labeled “Sample 1.” Repeat the process for the other flask labeled “Sample 2.” Keep track of which flask has which mass of KHP. Clean the buret by rinsing it with 3-4 mL of the NaOH solution. Fill the buret up to the zero level mark. Open the stopcock fully to fill the buret tip and eliminate air bubbles, with a waste beaker underneath. Adjust the liquid level in the buret to be between 0.00-2.00 mL and record this initial reading in your notebook, estimating all readings to two decimal places. Add about 50 mL of distilled water to the flask labeled “Sample 1” and gently swirl until the KHP is completely dissolved. Then add 3 or 4 drops of phenolphthalein indicator solution and begin titration, swirling the flask gently as the NaOH solution is added slowly.
Approach the end point by adding very small amounts of the solution until it turns pink briefly. Continue adding small amounts until the solution stays light pink for over ten seconds. Record the final volume and use it to calculate the volume of base required to reach the end point. Follow the same procedure to titrate sample two, but it may proceed faster since we know approximately how much volume is needed. Remember to record the initial volume, go slow towards the end point, and record the final volume. After calculating the molarity and finding the average, store the NaOH in a bottle for future experiments. Clean the 250 flasks and 500 mL flask for the next titration.
To calculate the volume of 3 M HCl needed for a 500 mL solution with a concentration of 0.1 M, measure the HCl volume in a flask and then add enough water to reach 500 mL. Secure the stopper and gently swirl the mixture.
Switch the buret clamp so that one side has an empty buret and the other side has a NaOH buret. Rinse the empty buret with a few mL of the HCl solution and fill it with approximately 40-45 mL of HCl. Note down this initial volume.
In the cleaned flask labeled “Sample 1”, carefully add around 35 mL of HCl and record the final volume. Calculate the exact amount of HCl added to the flask. Also, add 3 or 4 drops of phenolphthalein solution to this flask.
Switch back to the NaOH buret and start titration following similar steps as in KHP titration. Remember to record both initial and final volumes of HCl used during titration.
Repeat this process for another flask called “Sample 2.” Use both results to calculate acid solution molarity.
Finally, store this HCl solution in another empty bottle for future lab experiments.
The molecular weight of KHP is 204.22 g/mol.
The experiment involves the 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
The molecular weight of potassium hydrogen phthalate (KHP) is provided as 204.22 grams per mole.
During the experiment, the goal is to calibrate or determine the concentration of sodium hydroxide (NaOH) solution accurately.
Two samples are labelled as Sample 1 and Sample 2.
For Sample 1, the mass of KHP used is measured as approximately 0 .4628 grams and for Sample 2 it is approximately 0 .4926 grams.
By performing calculations based on these masses, it was determined that there are approximately
2 .266 millimoles (mmol) in Sample 1 and about
2 .412 mmol in Sample two .
To measure accurately, a buret was utilized to measure out specific amounts:
For sample one: The final buret reading after dispensing NaOH was recorded at approximately
24 .3 milliliters (mL), while the initial buret reading before dispensing NaOH was roughly 0 .7 mL.
Therefore, the volume of NaOH used for sample one is approximately
23 .6 mL. The molarity of NaOH solution in this case was determined to be approximately
0 .096 M.
For sample two: The final buret reading after dispensing NaOH was recorded as approximately
30 .2 mL, while the initial buret reading before dispensing NaOH was about
5 .65 mL.
Thus, the approximate volume of NaOH used for sample two is around
24 .55 mL and the molarity of NaOH solution for sample two is approximately
0 .098 M.
To obtain an overall average value, both samples were taken into consideration and it was determined that the average molarity of NaOH solution from these measurements is around 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:
The standardization of NaOH solution is important because it is used for both the titration of the unknown substance and the standardization of HCl. To determine the molarity of NaOH, I reacted it with KHP, a primary standard. To ensure accuracy, I made two precise measurements. Molarity refers to the concentration of a solution, with a 1 M solution containing one mole per liter. To simplify calculations, I converted the values to millimoles and milliliters: Molarity = (moles / liters) = (10-3 x moles / 10-3 x liters) = (mmol / mL). During titration, the mmol of KHP matched the mmol of NaOH added. By the end of the first titration, I had enough information to calculate the molarity of NaOH using this equation: M base = (g KHP x 103) / (mL base)(mw KHP).
Conclusion:
The acid base titrations yielded valuable insights into the interactions of the acid, base, and water. It was observed that when equivalent concentrations of a strong acid and a strong base react, the pH of the solution becomes neutral, indicating the end point. The molarity of both the bases and acids can be determined by knowing certain information about the acid-base system. In summary, this experiment provided important observations and achievements that highlight significant qualities of acid-base equilibrium. These qualities have practical implications for real-world systems. The ability to predict values such as pH, molecular weight, and molarity, which were utilized in this experiment, offers significant power and understanding in comprehending the chemical system.
