Lab: Acid-Base Titration

Objective

In this lab students will determine the concentration of a sodium hydroxide (NaOH) solution using the primary standard potassium hydrogen phthalate (KHP). Once standardized, students will use the NaOH solution to determine the concentration of an acetic acid (CH₃COOH) solution, in this case white table vinegar labeled as 5% by mass.

Introduction

Part 1: Standardization of Sodium Hydroxide Solution

Very precise measurements require great care. It is possible to use a 50 mL buret to make measurements of differences in volume to the nearest 0.01 mL in a procedure called a titration. For a total volume larger than 10 mL this gives four significant figures. This degree of precision depends on the skill of the user, which takes time and practice to acquire. The main challenge is knowing when to stop adding the solution that is in the buret to the solution underneath it. This is normally determined by the change in color of an indicator at what is called the endpoint of a titration. An indicator is selected by the designer of a procedure to change color over a range of pH values that includes the value at the moment of stoichiometric equivalence. The person performing a titration aims to add the final drop that causes a permanent color change in such a way as to be sure not to have gone beyond equivalence. This is accomplished by adding the final amounts slowly, in small increments.

Solid sodium hydroxide (NaOH) attracts water from the water vapor in air, which adds to the mass measured when it is placed on a lab balance. In addition, the presence of this water allows minute amounts of carbon dioxide (CO₂) to dissolve, which then reacts with the sodium hydroxide to make sodium carbonate (Na₂CO₃):

Because the mass of solid sodium hydroxide as measured in the lab cannot be guaranteed to be 100% NaOH it is impossible to make a solution with a precise concentration. This is important because if a sodium hydroxide solution is going to be used to measure the concentration of an acid then it must have a precise concentration. To overcome this obstacle the solution must be standardized. This means that its concentration must be determined by reacting the solution with a precise amount of an acid so that both the volume of the sodium hydroxide solution and the number of moles in that volume may be known.

The usual primary standard used to standardize a sodium hydroxide solution is solid KHP (potassium hydrogen phthalate, KHC₈H₄O₄). When dried in an oven KHP can be weighed on a balance and its mass can be used to precisely calculate the number of moles. Then it is possible to react the KHP with a sodium hydroxide solution to determine the number of moles of sodium hydroxide in a precisely measured volume of solution. A buret is used to measure the volume of solution to the nearest 0.01 mL. Phenolphthalein is an acid-base indicator which changes color at just the right pH to show show when the reaction of KHP with NaOH is complete. The chemical reaction below shows the neutralization of KHP by NaOH (the KHP is shown as the anion of the salt, leaving out the potassium ion and sodium ions).

Here is an example calculation to find the concentration of an approximately 0.5 M solution of sodium hydroxide using a mass of KHP of 1.057 g. The measured volume of NaOH solution is 10.52 mL.

            1 mol KHP    1 mol NaOH
1.057 g × ------------ × ---------- = 5.176 × 10–3
            204.22 g     1 mol KHP     mol NaOH

    so the concentration is
     5.176 × 10–3 mol NaOH
    ----------------------- = 0.4920 M
           0.01052 L            

Note that since your sodium hydroxide solution should have a concentration near 0.5 M the volume of that solution needed to react with approx. 1 gram of KHP is near 10 mL. This suggests that you add up to 8.5 or even 9 mL before slowly dripping in the last of the solution to precisely determine the end point volume of the titration.

Part 2: Determining the Acetic Acid Concentration in Vinegar

Vinegar is a solution of acetic acid (CH₃COOH) in water. As a pure substance, acetic acid is extremely corrosive and has an overwhelming odor of vinegar. It is sold in stores in diluted form with a concentration of 5% by mass. This means that 5 g out of every 100 g of the mixture is acetic acid. You will use the sodium hydroxide solution you standardized in Part 1, along with a stoichiometric calculation, to measure by titration the number of moles of acetic acid per liter of vinegar solution. You will also determine the mass % of acetic acid in the vinegar by weighing the sample of vinegar, calculating the mass of acetic acid dissolved in the sample, and calculating a percent. The neutralization reaction is written below, followed by an example showing how to calculate the concentration of acetic acid and an example showing how to calculate the mass % of acetic acid.

Given a standardized sodium hydroxide solution with a concentration of 0.4920 M and a volume of 10.00 mL of vinegar, what is the concentration of the acetic acid in the vinegar if 16.56 mL of sodium hydroxide solution are required to reach the endpoint of the titration?

              1 L        0.4920 mol NaOH     1 mol CH3COOH
16.56 mL × ---------- × ---------------- × ---------------- = 8.148 × 10–3 mol CH3COOH
            1000 mL           1 L            1 mol NaOH

So the concentration of acetic acid in the vinegar is: (8.148 × 10^–3 mol CH₃COOH)/(0.0100 L) = 0.8148 M.

To calculate the mass % and compare your result with the label on the bottle it is necessary also to know the mass of the vinegar solution. Usually, this is very nearly precisely 10.00 g since the acetic acid adds little mass at only 5% and the density of the water it is dissolved in is 1.00 g/mL. For the purposes of this example, the sample was found to have a mass of 10.023 g. The first step is to find the number of grams of acetic acid in the sample. Then the mass % may be calculated:

                    60.035 g
8.148 × 10–3 mol × ------------ = 0.4891 g CH3COOH
                     1 mol

    so the mass % is 
     0.4891 g
    ---------- × 100% = 4.880% 
     10.023 g

Since there is a lot to keep track of, here is a template for a data table you may use for this lab. It may be simplest to cut it out and paste it into your lab notebook.

|(value 1) – (value 2)|
------------------------ × 100%
    average value

|(std. value) – (exp. value)|
------------------------ × 100%
    std. value

Pre-lab Questions

Prior to taking part in this lab, complete the following questions.

1. The titration of 1.106 g of KHP required 9.85 mL of NaOH solution. What is the concentration of NaOH in the solution?
2. It required 15.39 mL of 0.5035 M NaOH solution to titrate a sample of 10.00 mL of white vinegar. What is the molar concentration of the vinegar?
3. If the sample of vinegar had a mass of 10.02 g then what is the mass percent of acetic acid in the vinegar?
4. When dissolving the KHP, as you see must be done according to the procedure, does it matter if you use a little more water than the procedure calls for? Why or why not?
5. How would it affect your experimental determination of the concentration of NaOH solution if you think you have 1.3 g of KHP but you really only had 1.1 g?
6. How would it affect the outcome of your experimental measurement of the concentration of acetic acid in the vinegar if due to faulty equipment your measurement of the concentration of NaOH is higher than the true measurement?
7. Hypothetically, consider the possibility you only had 9.10 mL of the vinegar rather than the 10.00 mL the pipet is designed to deliver. This could come about if the glassware is faulty or if some vinegar solution were accidentally left in the pipet. How would this quantitatively change your determination of the concentration of acetic acid in the vinegar if you thought you had the full 10.00 mL?

Procedure

Materials

Safety

The following list does not cover all possible hazards, just the ones that can be anticipated. Move slowly and carefully in the lab: haste and impatience have caused more than one accident.

• If you choose not to wear safety glasses you are choosing to sit out the lab.
• Sodium hydroxide is a strong base and can cause severe damage to skin and eyes. Avoid contact and wash hands after handling it. If in eyes, rinse with cold water as soon as possible for up to 15 minutes.
• Potassium hydrogen phthalate is a acid. Its dust is an eye and respiratory tract irritant. It is also toxic by ingestion. Avoid contact and wash hands after handling it. If exposed to dust, remove victim to fresh air. If in eyes, rinse with cold water for up to 15 minutes.
• Phenolphthalein is a 1% solution in denatured ethanol.
• Denatured alcohol is toxic by ingestion and is an eye irritant. If in eyes, rinse with cold water as soon as possible for up to 15 minutes.
• Denatured alcohol is highly flammable. Keep flames and sources of ignition far from acohol.
• White vinegar is an eye irritant. If in eyes, rinse with cold water for up to 15 minutes.
• Be careful using electrical equipment near water and chemical solutions. Avoid spilling water where it may short out a piece of electrical equipment, which can cause fires.

Part I: Standardization of NaOH Solution

In this procedure your main goal is to use a primary standard (potassium hydrogen phthalate, KHP) to quantitatively determine the concentration of the sodium hydroxide solution. The primary standard can be weighed in order to determine a precise number of moles. By titrating this acid against your sodium hydroxide solution, and by doing a simple stoichiometric calculation, you can calculate the concentration of NaOH in your base solution.

1. Obtain two samples of KHP. Each should have a mass between 1.000 and 1.200 g. The most important thing is to record exactly how much you have. Transfer these samples quantitatively into labeled 125-mL flasks for trial 1 and trial 2. (A quantitative transfer means that every last speck of solid material ends up in the flask. You can do this by rinsing with water from a wash bottle).
2. Adding no more than about 30 mL total, pour distilled water into the flasks containing your KHP samples. Use a magnetic stirrer to dissolve or simply swirl the water around until all of the solid has dissolved. Be aware that any liquid that splashes up onto the sides of the flask will contain some of your KHP.
3. Add 1 - 2 drops of phenolphthalein indicator and stir. This indicator will remain colorless until the pH reaches a value of about 9. At that point, the end point of the titration, it turns bright pink. The volume of base added at the endpoint is taken to contain the stoichiometric equivalent amount of NaOH for the amount of KHP.
4. Put a magnetic stir-bar into your flask, if you have not already done so. Place the flask on a piece of white paper on top of a magnetic stirrer. Set it stirring slowly. If it goes too fast it splashes solution up onto the sides of the flask and it may be this will prevent some of the KHP from reacting.
5. Use a 100-mL beaker to obtain about 55 mL of the sodium hydroxide (NaOH) solution. The labeled concentration is “approximately 0.5 M”.
6. Make sure the stopcock of the buret is closed. Use a small funnel to fill your buret with about 5 mL of NaOH solution. Take out the funnel and swirl the solution around inside the buret to coat the inside. Discard this solution by pouring it out the top. This is called “rinsing in” and it ensures that nothing inside the buret will change the concentration of the solution you put in it.
7. Clamp the buret and use the funnel again to fill it up. Do not go over the zero mark at the top. When you finish filling the buret, remove the funnel. This prevents small drops of liquid from descending later, altering the volume you are measuring.
8. Use a 50-mL beaker to collect a small amount of NaOH solution by opening the stopcock of the buret. This action fills the tip. Make sure there are no air bubbles.
9. Read the volume of NaOH solution in your buret to the nearest 0.01 mL by estimating between the marks, which show every 0.1 mL. Note that the scale is upside-down and zero is at the top.
10. See the illustration above and make sure everything is set up exactly as you see it there.
11. Assume that the concentration of NaOH is 0.5 M. Carry out calculations to determine the approximate volume you will need to react with the exact mass of KHP in your flask. Write this volume down.
12. Add the volume you just calculated to your initial buret reading. This is the mark on your buret where you will likely stop your titration, having reached the end point.
13. Begin adding NaOH solution to the flask of KHP solution. Add it carefully but not too slowly until you get within 2 to 3 mL of the predicted final volume. Then add NaOH solution more slowly. The phenolphthalein will change color as you add basic solution and the color will change back more and more slowly the closer you get to the end point.
14. Ideally you will add the NaOH solution one drop at a time when you get very close to the end point. At the end point the color should change to pink for the whole volume of solution and stay pink.
15. When you have reached the end point read the volume on the buret to the nearest 0.01 mL. Enter this reading into your data table.
16. Refill the buret with a little more NaOH solution and repeat the titration with the second sample of KHP solution. Use the same steps as above.
17. Enter all data into your data table and perform the necessary calculations to determine the concentration of the NaOH solution. If the two trials do not at all agree with one another (they are more than about 5% different) then perform another KHP trial, if time allows. Or, if you know that one of your trials is probably no good, then take note of that and repeat the titration if you think it is needed.

Part II: Determining Acetic Acid Concentration

After completing Part I, above, you will have determined the concentration of the NaOH solution that has been made available to you. Now it can be used to determine the concentration an acid. By measuring a very precise volume of vinegar (an acetic acid solution) and titrating it with your NaOH solution you can determine the number of moles of acetic acid in that volume of vinegar. From this you can calculate the molarity. With a simple mass measurement you can also calculate the mass percent of acetic acid in the vinegar, which you can compare to the value printed on the label.

1. Obtain two clean 125-mL Erlenmeyer flasks. Label them Trial 1 and Trial 2.
2. Measure the mass of each flask.
3. Use a 100-mL beaker to obtain about 30 mL of white table vinegar from the bottle provided.
4. Use a 10-mL volumetric pipet to measure precisely 10.00 mL of the vinegar into each flask. To use a volumetric pipet first rinse it twice with a few mL the solution to be measured, discarding the rinses. Do this by drawing liquid up using a suction bulb. Be careful not to get any liquid into the bulb! Next, draw enough solution into the pipet to go up over the calibration mark. Slowly let the solution drain out until the bottom of the meniscus in the pipet just touches the calibration mark. Being careful not to allow any solution to cling to the exterior of the pipet, empty it into a waiting flask. Do not blow out the last drop! Rest the end of the pipet against the glass and allow it to drain until it stops. The liquid that remains inside the pipet is meant to stay there as part of the design of the instrument. If this seems a bit intimidating, feel free to practice with plain water a few times before you try to measure the vinegar solution.
5. Weigh the flasks again so that it will be possible to know the precise mass of the vinegar sample. This is important for determining the mass percent of acetic acid in the vinegar solution.
6. Add 2 - 3 drops of phenolphthalein solution. In a happy coincidence, acetic acid and KHP have titration equivalence points at similar pH values and as a result the same indicator can be used to determine the end point of the titration.
7. The molar concentration of acetic acid is near 0.8 M. Use this value to calculate the approximate volume of NaOH solution you will need.
8. Use the calculation to help you to perform a titration of the vinegar exactly as you did when titrating the KHP. Record data as appropriate.
9. Before leaving the lab, perform all calculations necessary to determine the molar concentration and the percent mass concentration of acetic acid in the vinegar. If the difference between your two measurements of the molar concentration is greater than 5% you may wish to consider doing a third titration, if you have time.

Grading

Address the following items as part of your formal lab report. Include them in the Analysis portion of the report in paragraph form rather than question and answer format.

1. Why was it necessary to measure the concentration of the sodium hydroxide solution (using KHP) prior to using that solution to measure the concentration of the vinegar?
2. What was the concentration of NaOH in your solution?
3. Calculate and comment on the percent difference between your standardization trials for NaOH. What specific circumstance in your work may have led to the size of the difference between your trials?
4. What was the molar concentration of acetic acid in the vinegar according to your analysis? What was the percent acetic acid by mass?
5. Calculate and comment on the percent error between your result for the mass % of acetic acid in vinegar. What specific circumstance in your work may have led to the size of the difference between your result and the label on the bottle? (Acetic acid is volatile and evaporates over time so it is possible that the true mass % is less than 5%).