Exp: Acetic Acid Titration

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Last Updated: 
2013 - 10/17

Molarity, Acid Base, Neutralization, Titrations, Volumetric Glassware

Introduction & Objectives: 

Using a standardized base solution, students will use titration to determine the mass percent and molarity of acetic acid in vinegar. Vinegar is a solution. The solvent is water and the primary solute is acetic acid. The molecular formula for acetic acid is HC2H3O2. It is also commonly abbreviated at HOAc.


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

H+(aq) + OH- (aq) → H2O(l)

In this experiment, you will use this reaction to determine accurately the concentration of an unknown acid solution. To do this, you will accurately measure, with a buret, the volume of your standard base that is required to exactly neutralize the acid present in the unknown solution. The technique of accurately measuring the volume of a solution required to react with another reagent is termed titration. A titration is a technique where a solution of known concentration is used to determine the concentration of an unknown solution. Typically, the titrant (the known solution) is added from a buret to a known quantity of the analyte (the unknown solution) until the reaction is complete. Knowing the volume of titrant added allows the determination of the concentration of the unknown. Often, an indicator is used to usually signal the end of the reaction, the endpoint. Titration is also called volumetric analysis, which is type of quantitative chemical analysis.

An indicator solution is used to determine when a base has exactly neutralized an acid 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.

In this experiment, your will be given a standardized solution of NaOH, i.e., the concentration of the standard solution has been determined ahead of time and is known quite precisely. You will use the standardized NaOH solution to determine the acetic acid concentration in an unknown vinegar solution. The equation for this reaction is shown below:

HC2H3O2(aq) + NaOH(aq) → H2O(l) + NaC2H3O2(aq)


Centigram Balance ( ± 0.01 g )
50 mL Buret
10 mL graduated pipet or 10 mL volumetric pipet
Buret Clamp
Ring Stand
Funnel for Refilling Buret
Clean 125 mL Erlenmeyer Flask
Clean 150 mL or 250 mL Beaker (2) (for vinegar and standardized base soutions)
Sheet of white paper


Standardized ~ 0.500 M sodium hydroxide solution
Phenolphthalein indicator solution
Vinegar solution ~ 45 mL


Vinegar has a pungent aroma. Dispose of your vinegar and sodium hydroxide solutions in the container labeled "acid waste" located in the fume hood.


Soap has traditionally been made by mixing sodium hydroxide with animal fat. If during this lab your hands feel soapy, you may have spilled some sodium hydroxide on your hands. The soapy sensation is because your skin is reacting with the sodium hydroxide to form soap. Immediately rinse your hands with large volumes of water until the soapy feeling goes away.

Do not stand on lab stools or chairs to read the buret. Lower the buret to your eye level instead prior to taking any measurements. Use caution when filling and moving burets, burets are fragile and produce sharp ends if broken that are a common cause of puncture wounds in the laboratory.

Safety goggles must be warn at all times in this lab.


For this experiment, two students will be assigned to each titration apparatus. Each student will be assigned about 50 mL of vinegar solution of unknown concentration.


  1. Collect about 80 mL each of the standardized sodium hydroxide solution and unknown acetic acid solution (vinegar) in each of your two beakers.
  2. Collect the titration apparatus. Throughly clean your buret and then rinse it out with about 5 mL of deionized water and then 5 ml of sodium hydroxide solution. Attach the buret clamp to the ring stand and install the buret in the buret clamp.
  3. Clean and rinse the pipet with about 5 ml of dionized water and then 5 ml of the acetic acid solutoin.

Trial Process (repeat for 3-4 trials):

  1. Fill or refill the buret with the standardized sodium hydroxide solution to wihtin 3 mL of it's capacity. Record the initial buret read to two decimal places (± 0.01 mL).
  2. Clean and rinse the 125 mL Erlenmyer flask with dionized water. Dry the outside of the Erlenymyer flask and weigh it.
  3. Weigh and record the flask weight to within ± 0.01 g. It is not necessary for the inside of the flask to be dry. You must weight the flask again with every trial as droplets of dionized water will change the weight.
  4. Pipet roughly 10.0 mL of the acetic acid solution into the Erlenmyer flask from your beaker of acetic acid solution. Record the volume of acetic acid to within 0.01 mL.
  5. Re-weigh the flask with acetic acid solution. Record the combined weight of the flask and acetic acid solution. Subtract the weight of the empty flask from this number to find the weight of the acetic acid solution.
  6. Add 1 drop of phenolphthalein indicator to the acetic acid solution.
  7. Add roughly 20 mL of dionized water to the aceit acid solution.
  8. Place a sheet of white paper underneath the Erlenmyer flask (so you can more easily spot the persistent pink tink of solution that will indicate the titration is complete).
  9. Titrate the vinegar solution to the endpoint using the standardized sodium hydroxide solution. Record the buret reading at the endpoint, to two decimal places (± 0.01 mL).

After the first titration, dump the contents of the beaker down the sink, rinse the beaker well with tap water, shake it dry, and carefully dry off the outside of the flask with a paper towel. The flask does not have to be absolutely dry, but it should be dry on the outside. You will need to weigh it before each trial, because it will contain a slightly different amount of water each time. Weigh the flask again, and continue with the next trial. You should conduct at least three trials on the vinegar solution. Your objective is to generate at least three results that do not differ by more than 1.5%. Use the formula below to calculate the % difference between runs:

% difference = [ Molarityhigh - Molaritylow x 100 ] / Molarityaverage


You will need to complete the following calculations to complete your data table. In your calculations section show your work for each calculation below for one example trial.

C1: Calculate the moles of NaOH used in each trial. To do this multiply the volume of sodium hydroxide solution you added with your buret by the molarity of that solution. Be careful to convert your volume from mL to L before multiplying.

C2: Calculate the moles of AcOH in the flask, for each trial. Using the balanced equation above, determine the mole ratio between NaOH and AcOH. Use the mole ratio to determine the moles of AcOH neutralized by the sodium hydroxide you added.

C3: Calculate the molarity of the AcOH, for each trial. You determined the moles of AcOH in the acetic acid sample you added to your Erlenmeyer flask. You carefully measured the volume. Divide the number of moles by the volume (in Liters) to determine the molarity.

C4: Calculate the average molarity and the percent difference for your trials.

C5: Calculate the mass of acetic acid in the sample of acetic acid solution (vinegar), for each trial. Use the moles of acetic acid calculated in C2 and the molar mass of acetic acid to calculate the mass of acetic acid in your sample.

C6: Calculate the mass percent of acetic acid in the acetic acid sample, for each trial. Divide the mass of acetic acid you calculated in C5 by the total mass of the acetic acid solution sample. The total mass of the solution is just the mass of the flask with vinegar less the mass of the empty flask.

C7: Calculate the average of the mass percents and also the percent different of mass percents for your trials.

Presenting your Conclusions (how to prepare your report): 

A cover sheet for this experiment will be available from your instructor. Follow the instructions on the cover sheet with regard to preparing each section of your report (usually DATA, CALCULATIONS, QUESTIONS, and CONCLUSIONS).

In your DATA section, organize all your measurements and observations into a data table and include it in this section of your report. If your instructor provided a special data table for this lab you may use that, otherwise use either Word or Excel to make your data table. In this experiment, your data is includes the weights and volumes described in the procedures section.

In your CALCULATIONS section, complete all the calculations described above. For each calculation C1 - C7, show your work using data from one trial, to demonstrate how each calculation

Be sure to include the questions listed in this lab with answers in the QUESTIONS section of your report.

Your most significant conclusions are your average percent mass and average molariy. In your CONCLUSIONS section, list your conclusions for this experiment.


Q1: What are equivalence points and end points, and how do they differ?

Q2: Why can’t you add water to the vinegar solution at any time before it is transferred to the titration flask?

Q3: If 25.00 mL of a sulfuric acid solution (H2SO4 – two titratable protons) is titrated with sodium hydroxide, and if it requires 35.88 mL of 0.1127 M NaOH to reach the equivalence point, what is the molarity of the sulfuric acid solution?

Q4: A 2.353 g sample of vinegar was titrated with 0.08751 M NaOH and it requires 22.31 mL of NaOH to reach the endpoint. Calculate the mass percent of acetic acid in the vinegar sample.


Brown & LeMay, Chemistry the Central Science: Chapter 4