Titration of Acids and Bases

 

Purpose:    To become familiar with the techniques of titration, a volumetric method of analysis; to determine the amount of acid in an unknown.

 

            50 mL Erlenmeyer flask                        pint bottle with rubber stopper

            50 mL buret                                         phenolphthalein solution

            buret clamp                                          potassium acid phthalate

            ring stand                                                (primary standard)

            10 M NaOH                                        250 mL Erlenmeyer flasks (3)

            balance                                     weighing bottle

            wash bottle                                           unknown acid

            600 mL beaker

 

            One of the most common and familiar reactions in chemistry is the reaction of an acid with a base. This is termed neutralization, and the essential feature of this process in aqueous solution is the combination of hydronium ions with hydroxide ions to form water.

 

H3O+(aq) + OH-(aq) → 2H2O(l)

 

             In this experiment you will use this reaction to determine accurately the concentration of a sodium hydroxide solution that you have prepared. The process of determining the concentration of a solution is called standardization. Next you will measure the amount of acid in an unknown. To do this, you will accurately measure, with a buret, the volume of your standard base that is required to exactly the acid present in the unknown. The technique of accurately measuring the volume of a solution required to react with another reagent is termed titration.

            An indicator solution is used to determine when an acid has exactly neutralized a base, or vice versa. A suitable indicator changes colors when equivalent amounts of acid and base are present. The color change is termed the end point of the titration. Indicators change colors at different pH values. Phenolphthalein, for example, changes color from colorless to pink at a pH of about 9; in slightly more acidic solutions it is colorless, whereas, in more alkaline solutions it is pink. ( If you have not done Experiment 17, read its discussion of indicators.)

            In this experiment your solution of NaOH will be standardized by titrating it against a very pure sample of potassium hydrogen phthalate, KHC8H4O4, of known weight. Potassium hydrogen phthalate ( often abbreviated as KHP) has only one acidic hydrogen. Its structure is shown below. It is a monoprotic acid with the acidic hydrogen bonded to oxygen and has a molar mass of 204.2 g.

 

 

 

 

 

 

 

 

 

 

 

The balanced equation for the neutralization of potassium hydrogen phthalate is given in Equation [1]:

            KHC8H4O4 ( aq) + NaOH(aq)à H2O(l) + KnaC8H4O4(aq)

 

In titration of the base NaOH against KHP, an equal number of moles of KHP and NaOH are present at the equivalence point. In other words, at the equivalence point

 

                                    Moles NaOH = Moles KHP

 

The point at which stoichiometrically equivalent quantities are brought together is known as the equivalence point of the titration.

            It should be noted that the equivalence point in a titration is a theoretical point. It can be estimated by observing some physical change associated with the condition of equivalence, such as the change in color of an indicator, which is termed the end point.

            The most common way of quantifying concentrations is molarity ( symbol M), which is defined as the number of moles of solute per liter  of solution, or the number of millimoles of solute per millimeter of solution:

 

                                    M=                        Moles Solute        

                                            Volume of solution in liters

                                               

                                        = 10-3 mole

                                           10-3 liter

 

                                        = mmol

                                            mL

 

From equation [3] the moles of solute ( or mmol solute) are related to the molarity and the volume of the solution as follows:

           

                        M  X Liters = moles solute and M X mL = mmol solute

 

Thus, if one measures the volume of base, NaOH, required to neutralize a known weight of KHP, it is possible to caculate the molarity of the NaOH solution.

 

 

 

 

 

 

 

 

 

 

Procedure:        Preparation of Approximately 0.100 M Sodium Hydroxide*

                                    Heat more than 500 mL of distilled water to boiling in a 600-mL flask**, and after cooling under the water tap, transfer to a 1-pt bottle fitted with a rubber stopper.+ Add 5 mL of stock solution of carbonate-free sodium hydroxide ( approximately 10 M) and shake vigorously for at least 1 min.

*both bases and acids in any concentration should be considered harmful and protective eye ware and aprons should be worn.

** anything larger would also work

                                    Preparation of a Buret for Use  Clean a 50-mL buret with soap solution and a buret brush and thoroughly rinse with tap water. Then rinse with at least five 10-mL portions of distilled water. The water must run freely from the buret without leaving any drops adhering to the sides. Make sure that the buret does not leak and that the stopcock turns freely.

 

             Reading a Buret        All liquids, when placed in a buret, form a curved meniscus at their upper surfaces. In the case of water or water solutions, this meniscus is concave, and the most accurate buret readings are obtained by observing the position of the lowest point on the meniscus on the graduated scale.

            To avoid parallax errors when taking readings, the eye must on a level with the meniscus. Wrap a strip of paper around the buret and hold the top edges of the strip are evenly together. Adjust the strip so that the front and back edges are in line with the lowest part of the meniscus and take the reading by estimating to the nearest tenth of a marked division (0.01mL).

 

A.     Standardization of sodium Hydroxide Solution

Prepare about 400-450 mL of CO2 – free water by boiling for about 5 minutes. Weigh from a weighing bottle (your lab instructor will show you how to use a weighing bottle if you don’t already know) triplicate samples of between 0.4 and 0.6g each of pure potassium acid phthalate into three separate 250 mL Erlenmeyer flasks; accurately weigh to four significant figures. Do not weigh the flasks. Record the weights and label the three flasks in order to distinguish among them. Add to each sample about 100 mL of distilled water that has been freed from carbon dioxide by boiling, and warm gently with swirling until salt is dissolved. Add to each flask two drops of phenolphthalein indicator solution.

Rinse the previously cleaned buret with at least four 5 mL portions of the approximately 0.100 M sodium hydroxide solution that you have prepared. Discard each portion. Do not return any of the washings to the bottle. Completely fill the buret with the solution and remove the air from the tip by running out some of the liquid into an empty beaker. Make sure that the lower part of the meniscus is at the zero mark or slightly lower. Allow the buret to stand for at least 30 seconds before reading the exact location of the meniscus. Remove any hanging drop from the buret tip by touching it o the side of the beaker used for the washings. Record the initial buret reading.

Slowly add the sodium hydroxide solution to one of your flasks of potassium hydrogen phthalate solution while gently swirling the contents of the flask. As the sodium hydroxide solution is added, a pink color appears where the drops of the base come in contact with the solution. This coloration disappears with swirling. As the end point is approached, the color disappears more slowly, at which the sodium hydroxide should be added drop by drop. It is most important that the flask be swirled constantly throughout the entire titration. The end point is reached when one drop of the sodium hydroxide turns the whole flask from colorless to pink. The solution should remain pink when it is swirled. Allow the titrated solution to stand for at least one minute so the buret will drain properly. Remove any hanging drop from the buret tip by touching it to the side of the flask and wash down the sides of the flask with a stream of water from the wash bottle. Record the buret reading. Repeat this procedure with the other two samples.

From the data you obtain in the three titrations, calculate the molarity of the sodium hydroxide solution to four significant figures.

The three determinations should agree within 1.0 percent. If they do not, the standardization should be repeated until agreement is reached. The average of the three acceptable determinations is taken as the molarity of the sodium hydroxide. Calculate the standard deviation of you results.

 

B.     Analysis of an Unknown Acid

Measure out a given sample of a weak acid (record this volume) and add two drops of phenolphthalein indicator solution. Titrate with your standard sodium hydroxide solution to the faintest visible shade of pink (not red) as described above in the standardization procedure. As you titrate measure the pH of the system using a pH meter at given interval (every ½ ml recommended). Use this data in answering the graph questions in the analysis.  Calculate the concentration of the unknown acid. For good results, repeat 3 times and the three determinations should agree within 1.0 percent. Your answers should have four significant figures. Compute the standard deviation of your results.

      Test your results by computing the average deviation from the mean. If one result is noticeably different from the others, perform an additional titration. If any result is more than two standard deviations away from the mean, discard it and titrate another sample.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Analysis Questions:

 

  1. Define standardization and state how you would go about doing it.
  2. Define titration.
  3. Why would you want to repeat these measurements 3 different times?
  4. Define molarity.
  5. Why do you weigh by difference?
  6. What are equivalence points and end points and how do they differ?
  7. What is parallax and why should you avoid it?
  8. Why is it necessary to rid the distilled water of CO2?
  9. What is the concentration of the measured KHP system you produced?
  10. What is the calculated (standardized NaOH) concentration?
  11. Construct a graph of pH (y axis) vs. amount of Titrant (volume NaOH) used.
  12. From this graph, mark and estimate the starting pH point and its pH value, the half titrated point and value, the equivalence point and value (inflection point) and the ending point and its pH value.
  13. Where is the inflection point located? Does this make sense?
  14. Calculate (Show equilibrium equations / calculations and the Henderson Hasselbalch equation calculations) for the pH staring point, the half equivalence, equivalence, and ending pH point.
  15. How do the calculated values compare to the estimated points?
  16. What is the molarity of solution that contains 1.89g of H2C2O4. 2H2O in 100 mL of solution?
  17. If 50.0 mL of NaOH solution is required to react completely with 1.24 g KHP, what is the molarity of the NaOH solution?
  18. In the titration of an impure sample of KHP, it was found that 29.4 mL of 0.100 M NaOH was required to react completely with 0.745 g of sample. What is the percentage of KHP in this sample?

19. Given that 2, 4, 6, 8 and 10 grams of NaOH are added to 100 ml of 2.0 M HC2H3O2 (Ka=1.8 X 10-5). Detail (show all equations and calculations) the pH behaviors at each of the above conditions.