Experiment: TITRATION OF AN ACID WITH A BASE

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1 Experiment: TITRATION OF AN ACID WITH A BASE INTRODUCTION: The word titration is derived from a French word, "titre" which means "to bestow a title upon or to standardize". The purpose of a chemical titration is to standardize a solution, that is, to determine the concentration of a substance in a solution. The concentration of a substance in a solution is the amount of the substance dissolved in a certain volume of solution. Often, concentrations of solutions are given in terms of molarity, M, which is defined as the number of moles in a liter of solution. For example, the formula weight of HCl is 36.5 g/mole. If 73 g (2 X 36.5) are dissolved in sufficient water to make a liter of solution, the concentration of the HCl in the solution is 2 molar (2M), since 2 moles are dissolved in the liter of solution. Mathematically: Molarity moles liters mass formula weight X liters Another concept which must be understood when titrating an acid with a base is that of neutralization. For example, if HCl (acid) is titrated with NaOH (base), the reaction is: HCl + NaOH > NaCl + H2O In other words, an acidic solution and a basic solution, if mixed in the right proportions can be made neutral. (Here the products NaCl and H2O are plain salt water.) The proper proportion for this particular reaction is, from the reaction equation, one mole of HCl for every one mole of NaOH or a mole ratio of one to one. With these two concepts in mind - molarity and neutralization - the concentration of an acid or base can now be determined by titration. In the titration procedure, a known volume of an acid with a known molarity is placed in a beaker or flask. The base is slowly added from a buret until the two solutions neutralize each other. An indicator is used to tell when neutralization occurs. The volume of base used is then determined from the buret reading. The number of moles of acid used is determined from the equation written above but solved for moles. This is: moles = molarity X liters of solution Since the mole ratio for the reaction of HCl and NaOH is 1:1, the same number of moles of base must have been used in neutralization. Knowing the number of moles and the volume of base, the molarity can be calculated.

2 If the mole ratio of H + ions in the acid to the OH - ions in the base is not 1:1, then the number of moles of acid will not equal the number of moles of the base at neutralization. Part 1- Standardization of NaOH with a 1.0 M HCl solution. 1. A buret has been set up which contains the 1.0 M HCl solution. Record the initial buret reading. Place a 250 ml Erlenmeyer flask under the buret. (The flask should be clean, but it need not be dry.) Open the buret valve and withdraw approximately 10 ml of the acid. Record the final buret reading. Subtract these readings to find the exact volume of acid solution in the flask. Using this volume and the molarity of the acid solution (1.0 M), calculate the moles of HCl in the flask. 2. Add several drops of phenolphthalein (the indicator) to the flask. Fill a 150 ml beaker with distilled water. Add approximately 50 ml of this distilled water to the erlenmeyer flask. Note: This dilution does NOT change the number of moles of HCl in the flask. 3. Pour into a clean DRY beaker approximately 150 ml of the NaOH solution into a 250 ml beaker. 4. Using a clean buret, rinse it with about 5 ml of the NaOH solution making sure all the inner surface comes in contact with the rinse solution. Discard the rinse solution. Repeat this process two more times. 5. AFTER MAKING SURE THE BURET VALVE is closed, pour your NaOH solution into the buret until it is above the top calibration marks. Open the valve and lower the meniscus of the solution carefully until it is at (or near) the zero point. Record the initial reading of the buret. 6. Place the Erlenmeyer flask containing the acid solution under the buret and lower it until it is below the rim of the flask. Put a white background under the flask so that color changes are more easily seen. 7. While swirling the acid solution with one hand, open the valve. The approach of the endpoint will be signaled by the appearance of a pink color. As the endpoint gets closer, the color persists longer. 8. Just prior to the endpoint, only one drop at a time should be added. A distilled water rinse bottle should be used to wash down the sides of the flask to ensure all the solution is thoroughly mixed. 9. The titration is complete when the first barely visible but permanent pink color appears. Record the final reading of the buret. Subtract the buret readings to find the volume of NaOH used. 10. Using the balanced equation for the reaction, calculate the moles of NaOH used. (Hint: the moles of HCl in the Erlenmeyer flask determines the moles of NaOH needed for neutralization). NaOH(aq) + HCl(aq) > H2O + NaCl(aq)

3 11. Using the moles of NaOH and the volume of the NaOH solution required in the titration, calculate the molarity of the NaOH solution. 12. Wash and rinse out your Erlenmeyer flask, top-off the buret with more NaOH solution and repeat the titration process two more times. Part II: Analysis of vinegar (MICROSCALE TITRATION) Vinegar is a dilute solution of acetic acid in water. Using titration techniques the concentration of the acetic acid in the vinegar can be determined. The experiment will be completed using microscale titration techniques. There is great interest within the scientific community in microscale techniques. The techniques significantly reduce the cost of the chemicals used by colleges and industry, greatly reduce chemical waste disposal costs, and reduce the exposure of students, faculty, and research scientists to potentially harmful chemicals. While the materials we will use today are inexpensive and easily disposed of, the experiment provides an opportunity to try microscale techniques. 1. Your instructor has set up a station for dispensing the vinegar. Using a disposable glass pipet (equipped with a syringe or pipet pump), extract a 1.0 ml sample of vinegar and put the sample into one of the dry cells in the well plate. Repeat this procedure two more times. (At this point each of three cells will contain 1.0 ml of vinegar.) Add 1 drop of phenolphthalein indicator to each of the three wells. Place a white paper towel under your well plate so that color changes are more easily seen. 2. Take one polyethylene pipet. Grasp the bulb end in one hand and wrap the long skinny end around the index finger of your other hand. Pull the two ends in opposite directions. When you pull hard enough, a region of the long tube will narrow slightly. Cut the long end off the pipet in this narrowed region. (If you pull too hard, you may break the pipet. If this happens, try again with a new polyethylene pipet.) You now have a "mini-buret" which dispenses very small drops. Fill the pipet with the standardized NaOH solution. Practice the technique of extracting drops from this micropipet. The empty well plates can be used to hold your micropipet between trials. 3. Rather than counting drops or measuring volume changes, we will measure the difference in the weight of the micropipet containing the NaOH solution before and after the titration process. Find the initial weight of the micropipet containing the NaOH solution. 4. Using your micropipet add the NaOH solution dropwise to one of the wells. Continue to add the NaOH solution until the permanent pink color is observed. During the titration, move the well plate (on the paper towel) in a gentle circular pattern to mix the solutions.

4 5. Measure the mass of the micropipet containing the remaining NaOH solution after the first titration is complete. Subtract to find the mass of the NaOH solution used in the first titration. Assuming the density of the NaOH is equal to the density of water (1.0 g/ml), calculate the volume of the NaOH solution required to neutralize the acetic acid in vinegar. Using the molarity of the NaOH solution (determined in Part I), calculate the moles of NaOH necessary to neutralize vinegar solution. 6. Using the balanced equation for the reaction, calculate the moles of acetic acid in the vinegar sample. (Hint: the moles of NaOH used can be used to calculate the moles of acetic acid.) HC2H3O2(aq) + NaOH(aq) > H2O + NaC2H3O2(aq) Acetic Acid 7. From the volume of vinegar used (how much did you put in the well?) and moles of acetic acid in the sample, calculate the Molarity of the acetic acid in vinegar. 8. Refill the micropipet with the NaOH solution. Starting from Step 3, repeat the experiment by titrating the vinegar in the remaining two wells.

5 Experiment - Titration of an Acid with a Base Report Sheet Name Chem 121 Lab Partner Data: STANDARDIZATION OF NaOH TRIAL 1 TRIAL 2 TRIAL 3 Notes: The HCl is 1.0 M and volumes must be in liters when you complete your calculations. HCl BURET READINGS: Initial buret reading Final buret reading Volume of HCl Moles of HCl NaOH BURET READINGS: Initial buret reading Final buret reading Volume of NaOH Write a balanced equation for the reaction of HCl with NaOH What is the mole ratio of the HCl to NaOH in your balanced equation? CALCULATIONS: Moles of NaOH Molarity of NaOH solution FIND THE AVERAGE VALUE FOR THE MOLARITY OF THE NaOH SOLUTION. M

6 The NaOH solution is STANDARDIZED. That is, we now know its molarity and this NaOH solution can be used for the analysis of vinegar. Data: ANALYSIS OF VINEGAR The Molarity of the NaOH solution (as determined in Part I) is: The Density of the NaOH solution is 1.0 g/ml. TRIAL 1 TRIAL 2 TRIAL 3 Initial mass of micropipet with NaOH solution Mass of NaOH micropipet after titration Mass of NaOH solution used in the titration Volume of NaOH solution used in the titration (Hint: use Density) Moles of NaOH used in the titration (Hint: use Molarity) Write a balance equation for the reaction of the acetic acid solution and the sodium hydroxide solution. What is the ratio of the NaOH to HC2H3O2 in your balanced equation? Moles of HC2H3O2 in vinegar sample Volume of vinegar in well Molarity of vinegar FIND THE AVERAGE VALUE FOR THE MOLARITY OF THE ACETIC ACID IN THE VINEGAR. M

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