The Concentration of Acetic Acid in Vinegar
To determine the concentration of acetic acid in vinegar
Primary standards, use of volumetric glassware, acid-base titrations
This experiment is in two parts: Part A involves standardization of an unknown sodium hydroxide solution and part of your pre-lab assignment is to calculate the concentrations of solutions needed for the experiment to work. In Part B, you will use the sodium hydroxide to determine the strength of commercial vinegar. It is impossible to prepare solutions of most solutes in accurately-known concentrations. For example, sodium hydroxide tends to absorb both water and carbon dioxide from the air. The former affects the mass of material, although it does not change it chemically; the latter reacts with NaOH, forming, in turn, bicarbonate and carbonate ions. For that reason, sodium hydroxide is prepared in approximately the concentration desired, then standardized against a primary standard. A primary standard is a substance that is obtainable in high purity and whose mass can be measured with high precision. The primary standard usually chosen for strong base solutions is potassium hydrogen phthalate, or KHP, for short. Figure 4-1 shows the structure of KHP. In this short-hand form, the comers of the hexagon are understood to be occupied by carbon atoms. The four carbon atoms that do not have carboxyl groups attached each have one hydrogen atom to complete their valence configuration. You will be provided with a sodium hydroxide solution with a molar concentration of approximately 0.1 M You will standardize the solution against potassium acid phthalate primary standard to get a more precise value for its concentration, in units of moles of NaOH per gram of sodium hydroxide solution. If possible, all masses should be determined using an analytical balance so that the concentrations will have three significant-figure precision.
In Part B you will use the standardized sodium hydroxide solution to determine the concentration of acetic acid in ordinary household vinegar. Since the concentration of acetic acid is in the range 0.8-1.0 M, it will be necessary to dilute the vinegar tenfold; you will use volumetric glassware for this purpose. The diluted vinegar will be titrated with the standard NaOH to establish the concentration of the diluted acid, again in moles of acetic acid per gram of diluted vinegar. Using your dilution data, you will calculate the density of the diluted vinegar, and then use the density to determine the mass percent of acetic acid in both the diluted and the full-strength vinegar solutions
Read the entire experiment before you begin. Prepare a data table for Part A that includes spaces for recording the mass of KHP used, the initial and final masses of the NaOH pipet, and the mass of NaOH used. Set your table up for three trials, but leave space in case one or more additional trials is needed. (See the instructions following step 4.)
1. Pure acetic acid is a liquid, with density 1.0492 g/mL at 25°C. An aqueous solution that is 0.1000 molar (0.1000 M) in acetic acid contains 6.0053 g of solute per liter of solution. Calculate: (a) the respective volumes of acetic acid and water needed to make a total volume of 1.000 L, and (b) the' concentration of the solution in mol HC2H302/gram of solution. At 25°C, the density of pure water is 0.9969 g/mL. (Click here for help)
2. What mass of 1.00 x 10-1 M NaOH solution will be needed to titrate 25.0 mg of KHP? Assume a density of 1.04 g/mL for the NaOH solution. (Click here for help)
3. Using the information from question 2, convert the concentration given for the NaOH to moles of solute per gram of solution.
4. Commercial vinegar is approximately 5 % acetic acid by mass. Assuming a density of 1.006 g/mL, calculate the molar concentration of acetic acid in vinegar. If 1.00 g of this solution is to be neutralized by 0.10 M NaOH, what volume of base will be required? A microtip pipet holds about 3.5 grams of solution; will one pipet-full of base be enough to titrate 1.00 g of full-strength vinegar? Show calculations to defend your answer. (Click here for help)
1. Chemical splash-protective eyewear must be worn at all times in the laboratory.
2. Sodium hydroxide is caustic and particularly dangerous to the eyes. In case of spills, flood the affected area with water for 5 minutes.
balance, analytical or milligram
10- or 25-mL Erlenmeyer flasks
pipets, microtip (2) for NaOH and vinegar
volumetric flask, 10- or 25-mL
volumetric or Mohr pipet, 1.0 mL or 2.5 mL
sodium hydroxide, NaOH(aq),
potassium hydrogen phthalate (KHP), solid
distilled water, in wash bottle
phenolphthalein, 0.5 (alcohol)
commercial "white" vinegar
Note: The potassium hydrogen phthalate primary standard must
be dried in a 110°C oven for at
least two hours before use. The dried solid acid is transferred to a desiccator to cool for a
minimum of one hour, or until it is needed.
Part A: Standardization of Sodium Hydroxide by Potassium Hydrogen Phthalate
1. Using an analytical balance if available, place between 20 and 25 mg (± 0.1 mg) of primary standard KHP in each of three 10-mL Erlenmeyer flasks. Add about 1-1.5 mL of distilled water to each flask and swirl to dissolve the solid acid, followed by one drop of 0.5 phenolphthalein indicator solution.
2. Fill a microtip pipet with the NaOH that is to be standardized. Wipe the tip of the pipet dry, and then determine the mass of pipet and contents (± 0.1 mg).
3. Titrate the first KHP sample by adding a few drops at a time of the NaOH and swirling after each addition. Continue in this fashion until you achieve a faint pink color that persists throughout the solution for at least 30 seconds. Record the final mass of your NaOH pipet.
4. Repeat Steps 2 and 3 with each of your other two KHP samples, refilling and reweighing the NaOH pipet between trials as necessary.
Precision check: Calculate the mass of NaOH solution used per gram of KHP for each of the three trials. Determine the mean value for the ratio, and then determine the relative deviation. If that relative deviation exceeds 5 of the mean, do one or more additional trials until you have three that show less than 5% deviation.
Part B: Determination of the Mass Percent of Acetic Acid Commercial Vinegar
5. Dilution of commercial vinegar: Weigh a clean, dry 50 mL volumetric flask, then
use a volumetric pipet to transfer exactly 5 mL of a commercial vinegar product to the flask. Weigh
the flask and contents once again, and then use distilled water to bring the volume of solution in the flask up to exactly 10.00 mL; then weigh the flask and contents to the nearest 0.1 mg (if possible). Your teacher will show you how to ensure that the diluted vinegar solution is thoroughly mixed. Weigh the flask and contents once again.
6. You will now use the standardized solution of sodium hydroxide to titrate three (at least) samples of the diluted vinegar. Following the procedure of Part A, prepare three fresh sample vessels, each containing 1.0-1.5 gram of the diluted vinegar solution in place of the potassium acid phthalate. As before, add phenolphthalein to all three and titrate each with the sodium hydroxide, which you standardized in Part A.
Precision check: Calculate the mass of NaOH solution used per gram of diluted vinegar for each of the three trials. Determine the mean value for the ratio, and then determine the relative deviation. If that relative deviation exceeds 5 of the mean, do one or more additional trials until you have three that show less than 5 deviations.
The contents of the titration vessels can be safely rinsed down the drain with water, as can leftover diluted vinegar. Unused NaOH must first be neutralized. If you have approximately equal volumes of NaOH and dilute vinegar, it is safe to combine them and rinse down the drain. For excess NaOH, add a few drops of phenolphthalein, and then add vinegar until the pink color appears.
Analysis and Conclusions (click her for help on 1-3)
1. Use your data from Part A to determine the concentration of the sodium hydroxide solution in mol NaOH/g solution. Show your work for one trial and report the average concentration and average deviation.
2. U sing the average value for the concentration of NaOH, determine the number of moles of acetic acid in each sample of the diluted vinegar that you titrated in Part B. Convert these to moles of acetic acid per gram of dilute vinegar by dividing the number of moles of acetic acid by the mass or the respective samples. As before, calculate individual values for each trial, then report an average value and average deviation.
3. Based on your average value for the number of moles of acetic acid per gram of diluted vinegar and using the data from the dilution, determine:
a. The density of the diluted vinegar
b. The mass percent of acetic acid in the original (undiluted) vinegar solution
4. The Introduction describes two successive side-reactions that can change the make-up of sodium hydroxide. Write balanced equations showing: (a) the reaction between carbon dioxide and hydroxide ions to form bicarbonate ions, and (b) the further reaction between bicarbonate and excess hydroxide to form carbonate ions. Does either of the two processes represent an acid-base reaction according to the Bronsted-Lowry system? Discuss.
5. In preparing a sodium hydroxide solution that is to be used for quantitative work, the usual technique is to boil the water first to drive off dissolved CO2• Based on the equilibria you wrote in response to Question 4, what effect (if any) would the presence of carbonates or bicarbonates have on the ability of the solution to neutralize hydronium ions? Explain.
6. Also mentioned in the Introduction is the fact that sodium hydroxide aggressively absorbs water vapor from the air. What effect would this have on the concentration of a solution prepared by dissolving 40.00 grams of solid sodium hydroxide in water to make 1.000 L of solution? Be very specific in your response.
7. At room temperature, the molar solubility of carbon dioxide in water is about 0.033 M. would you expect the solubility of carbon dioxide in a solution containing sodium hydroxide to be greater or less than 0.033 mol/L? Explain using the equations for the reaction(s) involved.
8. Use appropriate equations to explain why a white crust will sometimes form when sodium hydroxide that has been exposed to air is dissolved in "hard" tap water. (Hint: "hard" water contains calcium and/or magnesium ions. Consider the solubility rules.)
9. Write balanced net-ionic equations for: (a) the reaction
between hydroxide ions and hydrogen phthalate ions, HC4Hg04 -, with
the first H being the acidic hydrogen; and (b) the reaction
between (molecular) acetic acid and hydroxide ions. (See Figure 4-1.)