In this experiment you will use the technique of titration to determine the concentration of solutions of acids and bases. Using a buret to dispense the solution you can determine within 0.05 mL the volume of base solution needed for exact neutralization of either a given volume of acid solution or a known weight of dissolved solid acid.
The experiment will require two lab periods. During the first period you will perform some practice titrations and make up a base solution to be standardized. During the second period you will use a primary standard, potassium hydrogen phthalate (KHC8H4O4), to standardize your base solution. You will then use this standard base solution to titrate an unknown sample of acid to determine its concentration.
Be sure you understand these terms BEFORE you attempt this experiment.
(1) Molarity (M): The number of moles of solute per liter of solution or the number of millimoles of solute per milliliter of solution.
(2) Primary Standard: A substance of known purity that can be weighed directly and used to standardize other solutions. In this experiment the primary standard is potassium hydrogen phthalate.
(3) Secondary Standard: A solution whose concentration has been determined by reaction with a primary standard. It is used to find the concentration of unknown solutions. In this experiment the secondary standard is the base solution.
(4) Equivalence Point: The point at which chemically equivalent amounts of the reacting substances have been added.
(5) Endpoint: The point at which an observable change in some physical property (such as the color of an indicator) takes place. Ideally the end point and the equivalence point should be the same.
Wear safety goggles at all times in the laboratory. Acid or base solutions in the eyes cause damage so rapidly that it is difficult or impossible to reach the eyewash in time.
Check out two burets from the stockroom. Read the information on the reading, use and cleaning of the burets in the general equipment selection.
A. Practice Titrations
1. Use the 0.1 M solutions of NaOH and HCl that are set out in large bottles in the laboratory. Obtain about 150 mL of each solution in carefully marked beakers.
2. Label one buret "base" (or B) label the other buret "acid" (or A). Rinse the base buret twice with five to ten milliliter portions of the base solution and drain enough liquid through the stopcock so that no air bubble remains in the tip of the buret. Add base to the buret and adjust the liquid level in the buret to between 0.00 and 1.00 mL. Do not waste time trying to hit 0.00 mL exactly. Record the value as the "initial buret reading" of your base buret on the report sheet. Read the value to the nearest 0.01 mL.
3. Use the same procedure to rinse and fill the acid buret. Adjust the liquid level to between 0.00 and 1.00 mL and record as the "initial buret reading" of your acid buret on the report sheet.
4. Permit approximately 20 mL of the acid solution to run into an empty, clean Erlenmeyer flask and add two drops of phenolphthalein solution. Be sure to add the indicator or you won't know when to stop titrating. Also, use only two drops of the indicator. The end point will not coincide with the equivalence point if more or less indicator is used.
5. Add base to the acid solution rapidly until a pink color persists in the flask. Next, add a small amount of acid to destroy the pink color. Rinse the sides of the flask with a small amount of distilled water from your wash bottle. Then add base drop-wise with swirling until one drop causes a pink color that persists for thirty seconds or more. Read and record on your report sheet the final volume from each buret.
6. Calculate the molarity of the acid solution assuming that the base solution is 0.1000 M. Repeat the titration. If the molarity obtained in the second titration is within 1% of the molarity obtained in the first titration, a third titration is not needed. If the difference between the two is greater than 1%, another titration should be performed. Calculate the average molarity from your two or three titrations.
B. Preparation of the Base Solution
1. Make 500 mL of approximately 0.1 M NaOH as follows: using a balance accurate to 0.01 g., rapidly weigh out the appropriate amount of solid NaOH (you must calculate how much is needed) into a small preweighed beaker. (Do not weigh the solid pellets directly on the balance pan or on a piece of weighing paper! If you spill solid sodium hydroxide clean up the pellets immediately. Put the pellets in the sink and wash them down with lots of water. Do not put it into the waste chemical container.)
2. Place the solid pellets in a 600 mL beaker. Add distilled water slowly, a little at a time with stirring. SAFETY GOGGLES MUST BE WORN! Some heat is evolved during the dissolving process. Add enough water to make 500 mL of solution. It is permissible to use the marking on the side of the beaker to tell when 500 mL of solution has been made since the solution will be standardized later to get the exact concentration. Note: The markings on the beaker are only accurate to plus or minus 5% so don't use them when the volume must be accurately known. When the solution has cooled to room temperature pour it into a plastic bottle and store it in your locker until the next lab period. Label the bottle.
3. Thoroughly clean at least four Erlenmeyer flasks. Keep them in your lab drawer to air-dry for the use of next session.
C. Standardizing the Base Solution
1. Tare the first of three labeled Erlenmeyer flasks on a Mettler balance. Add approximately 0.6 to 0.8 grams of potassium hydrogen phthalate to the flask and record the exact weight. Enter this weight as the sample weight for sample 1.
2. Repeat this for the other two flasks and record the exact weights of the acid for sample 2 and sample 3.
3. To each flask add about 50 mL of distilled water (graduated cylinder accuracy is good enough). Don't worry if all of the solid will not dissolve (it will as the titration proceeds IF you swirl the flask as you titrate). Add two drops (no more and no less) of phenolphthalein to each flask.
4. Titrate each sample separately by adding base to the flask with constant swirling. Record the initial buret reading in your report form! You will not be able to back titrate so you must attempt to judge how close to the endpoint you are from the length of time the pink persists in the solution. The closer you are to the endpoint the longer the pink color will persist at the point where the base entered the solution before the swirling causes it to disappear. As you approach the endpoint rinse the sides of the flask with distilled water to bring down into the solution any unreacted acid that may cling to the walls of the container. When you are close to the end point add the base one drop at a time until one drop causes the color of the solution to go from colorless to pink.
5. Record the final volume from the base buret and calculate the molarity of the base.
6. Repeat for the other two samples of the acid. Calculate the other two molarities.
7. If there is substantial disagreement among the three values prepare more samples of the primary standard acid and do more titrations. Once you get three satisfactory values calculate the average molarity of the base and use this value in the calculations required in the next section.
D. Analysis of the Unknown Acid Solution
Obtain an unknown acid solution from the stockroom and determine its concentration in exactly the same way you did the the practice titrations. The base you use is the base you standardized in "Part C" above and the acid you use is your unknown acid solution from the stockroom. Be sure to rinse the acid buret with small portions of the unknown acid before filling it with the unknown acid. If you get two titrations that agree to within 1% you don't have to do a third titration. If the two do not agree to within 1% do a third titration.