EXPERIMENT 10:
TITRATION AND STANDARDIZATION
PURPOSE
To determine the molarity of a NaOH solution by titrating it with a standard HCl solution.
To determine the molarity of acetic acid in vinegar using the NaOH as a standard solution.
To determine the percentage of acetic acid in vinegar and compare different brands.
BACKGROUND
ACID-BASE REACTIONS: NEUTRALIZATION
A frequently used definition of the terms acid and base is attributed to Brønsted and Lowry: an acid is any substance that donates protons (H+ ions) while a base is any substance that accepts protons. When an acid reacts with a base, a neutralization reaction occurs. In this experiment, the base used is a strong base, NaOH, so the products of the neutralizations will be water and a salt.
TITRATIONS
It is sometimes necessary to experimentally determine the concentration of an acid solution or a base solution. A procedure for making this kind of determination is called an acid-base titration. In this procedure, a solution of known concentration, called a standard solution, is used to neutralize a precisely measured volume of the solution of unknown concentration. If the solution of unknown concentration is acidic, a standard base solution is added to the acid solution until it is neutralized. If the unknown is a base, a standard acid solution is added.
When carrying out an acid-base titration, you must be able to recognize when to stop adding the standard solution: that is when neutralization has been achieved. For this purpose, a few drops of an acid-base indicator are added before the titration begins. A sudden change in color of the indicator signals that neutralization has occurred. This is called the endpoint of the titration. At the equivalence point the number of hydrogen ions from the acid is equal to the number of hydroxide ions from the base. In a good titration the endpoint and the equivalence point should be the same. When the endpoint is reached, the volumes of the two solutions and the known concentration of the standard can be used to calculate the concentration of the unknown solution.
The indicator you will use in this reaction is phenolphthalein. Although
phenolphthalein is almost insoluble in water, it is very soluble in alcohol. The odor from the phenolphthalein solution is from the alcohol. Phenolphthalein is also commonly used as a cathartic in over-the-counter medication. It is the active ingredient in commercial laxative Ex-Lax.
Up to this point in your laboratory work, most of your quantitative experiments have surrounded mass relationships and mass measurements. This is known as gravimetric analysis. Titration requires the use of volume relationships and volume measurements, a technique known as volumetric analysis.
MATERIALS
250 mL Erlenmeyer flask 1.00 M HCl standard solution
ring stand unknown NaOH solutions
buret unknown vinegar solutions
buret clamp wash bottle with distilled water 250 mL beaker phenolphthalein indicator
PROCEDURE
Part A - Standardization of NaOH solution
1. Rinse a buret thoroughly, first with tap water, then with distilled water.
2. Obtain about 50 mL of one unknown NaOH solution in a clean beaker. Be sure to record which unknown you are using. Caution: Handle this solution with care. It can cause burns! Pour about 5.0 mL of the base solution into your buret and rinse the inside surface of the buret thoroughly. Allow the base to run out of the buret tip.
Pour the remainder of the base solution into your buret. Purge any bubbles from the buret tip by allowing some of the base to run out of the tip.
3. Your instructor has set up a buret for the standard acid solution. Place a 250 mL Erlenmeyer flask under the acid buret and obtain approximately 10.0 mL of the acid.
Be sure to record the concentration of the HCl solution and the initial and final volumes of the buret. You might find that holding a white sheet of paper behind the buret makes it easier to read.
4. Add approximately 10 mL of distilled water to the flask, then add 3 drops of phenolphthalein indicator to your flask and swirl until the contents are mixed.
5. Place the flask on a sheet of white paper under the buret containing the base solution. To avoid splashing, be sure the tip of the buret is below the mouth of the flask. Record the initial volume of base in the buret.
6. Begin titrating by allowing the base solution to flow into the flask until a light pink color appears in the flask. Stop the flow of base and swirl until the pink color disappears. Wash down the inside surface of the flask frequently with a little distilled water using the wash bottle. Continue this process of adding base and swirling until the pink color persists for about 10 to 15 seconds before disappearing. At this point, add the base one drop at a time, swirling constantly until the pink color persists for about 30 seconds. Record this volume. Add one more drop of base: if the solution turns a dark pink color that does not disappear, then the recorded volume is your final volume. If a light pink color appears but disappears in about 30 seconds, then record this new volume and add one more drop of base. Keep repeating until the color persists.
7. Wash and rinse the Erlenmeyer flask thoroughly. Repeat steps 3 through 6 as Trial 2. Copy the data for both trials into your notebook in a data summary table as shown below. Calculate the total amounts of acid and base used for each trial. Both the acid totals should be approximately 10.0 mL. If the base totals are very different you will need to do a third trial and discard the worst value. Consult your instructor.
BASE SOLUTION # Trial 1 Trial 2
HCl NaOH HCl NaOH
Initial volume
Final volume
Total volume
Part B - Analysis of vinegar solutions
8. Wash and rinse the Erlenmeyer flask thoroughly. Pipette 10.0 mL of one of the vinegar (acetic acid) solutions into the flask and add 10.0 mL of distilled water.
Record which of the acetic solutions you are standardizing.
9.Add 3 drops of phenolphthalein indicator to the flask and swirl. Titrate with the now standardized base as before (outline in step 6). Record appropriate data in a data summary table as shown below. Repeat step 9 for Trial 2. Calculate total volumes used. Again, if base totals vary too much, do another trial.
Vinegar solution # Trial 1 Trial 2
acid NaOH acid NaOH
Initial volume Final volume Total volume
Name: ________________________________ Date ___________
PRE-LAB QUESTIONS
1. At several points in the procedure you added distilled water. Why may this be done without disturbing the titration process?
2. Write the balanced equation for the reaction between hydrochloric acid and sodium hydroxide.
3. Write the balanced chemical equation for the reaction between vinegar (acetic acid) and sodium hydroxide.
4. Define:
a) titration
b) standard solution
c) endpoint
d) equivalence point
e) volumetric analysis
Name: ________________________________ Date __________
TITRATION AND STANDARDIZATION
Part A- Standardization of NaOH solutionBase solution # Trial 1 Trial 2
HCl(M) NaOH HCl(M) NaOH
Initial volume
Final volume
Total volume
CALCULATIONS
1. Write the balanced chemical equation for the reaction.
2. Calculate the molarity of the base solution for each trial.
3. Calculate the average molarity of the NaOH solution.
RESULTS
NaOH solution # Trial 1 Trial 2
Molarity Average Molarity
Part B - Analysis of vinegar solution
Vinegar solution # Trial 1 Trial 2
acid NaOH acid NaOH
Initial volume Final volume Total volume
CALCULATIONS
1. Write the balanced chemical equation for the reaction.
2. Calculate the molarity of the acetic acid solution for each trial.
3. Calculate the average molarity of the acetic acid solution.
RESULTS
Vinegar solution # Trial 1 Trial 2
Molarity Average Molarity
POST-LAB QUESTION
1. In a few sentences, summarize your results giving the possible sources of error.
2. Many household products such as cleaning agents, lawn care products and personal care items are sold with the active ingredient listed in mass % (weight %).
Explain why mass % is more commonly used than molarity.
