ABSTRACT/ SUMMARY
The purposes of this experiment are to determine the percent by mass of acetic acid in vinegar by titration with standardized sodium hydroxide (NaOH) solution. Vinegar is a dilute solution of acetic acid. Since vinegar is an acid, it can be titrated with a base. In this experiment, the titration comprised of two stages. The first stage is to standardize NaOH solution with KHP. 0.6 M NaOH was used in this standardization. The second stage is the titration between NaOH and commercial vinegar sample. The purpose of titration is to determine the equivalence point of the reaction. At the equivalence point is occurs when an equal number of moles of acid and base are mixed, the pH of the solution will change drastically, as the solution is being flooded with excess base (NaOH). The percent of acetic acid in vinegar and the percentage error for the experiment are 5.6983 % and 42.457 % error. The conclusion is the experimental percentage of CH3COOH in the vinegar sample is higher than the true value which is 4.00%.
1
INTRODUCTION
According to Stanton et. al., 2009, concentration of solution is the amount of solute (species dissolved) in a given amount solvent (dissolving agent). A concentrated solution contains a relatively great quantity of solute in a given amount of solvent. In another hand, dilute solution contains relatively little solute in a given amount of solvent. Chemist uses specific terms to express the concentration of solution. Two of these terms are molarity and percent by mass:
Molarity is the number of moles of solutes per liter of solution.
Molarity (M) =
(Equation 1-1)
Percent by mass is the mass in grams of solute per 100 grams of solution.
Percent solute =
× 100%
(Equation 1-2)
Vinegar or French for sour wine is formed by aerobic bacteria oxidizing grain alcohol to acetic acid and water (Scharf and Malerich, 2010). The smell of vinegar attributed to acetic acid (an organic acid). The molecular formula for acetic acid is CH3COOH. According to the U.S. Food and Drug Administration (FDA), for a manufacturer of vinegar to use the word “vinegar” in its labeling of the product, it must contain a minimum of 4.0 grams of acetic acid per 100 mL of solution, which corresponds to a concentration of 0.67 M. The rest of the solution is primarily water, with trace amounts of alcohol, phosphoric acid, sugar, and glycerol present. Since vinegar is an acid, it can be titrated with a base to determine the molarity and percent by mass of acetic acid in a vinegar solution. Titration is the process of adding a known amount of a solution of known concentration to a known amount of solution of unknown concentration. The more accurately the concentration of the solution of known concentration is known, the more accurately the concentration of the unknown solution can be determined. The concentration of unknown solution will be determined by knowing the volume of known solution needed to complete the titration. The purpose of titration is to determine the equivalence point of the reaction. The equivalence point of a titration is the point at which the acid has completely reacted with or been neutralized by the base. At this point, the pH changes rapidly with a small addition of acid or base (Chang, 2007).
2
OBJECTIVES
The purpose of this experiment was to determine the molarity of a solution and percent by mass of acetic acid in vinegar by titration with standardized sodium hydroxide solution.
THEORY
In this experiment, the equivalence point is reach when the added quantity of base is the exact amount necessary for stoichiometric reaction with acid (Stanton et. al., 2009). For example, the stoichiometric amount of 1 mole of the strong base, sodium hydroxide (NaOH), is necessary to neutralize 1 mole of the weak acid, acetic acid (CH3CO2H), as shown in Equation 1-3: NaOH(aq) + CH3CO2H(aq) NaCH3CO2(aq) + H2O(l)
(Equation 1-3)
At the equivalence point of titration, the pH of the solution will change drastically, as the solution is being flooded with excess strong base or strong acid (depending on what is being added). pH in an aqueous solution is related to its hydrogen ion concentration. Symbolically, the hydrogen ion concentration is written as [H3O+]. The molar concentration of hydrogen ions in aqueous solution is frequently very small. For convenience, therefore the pH of a solution is defined as the negative logarithm of the hydrogen ion concentration (Brown et. al., 2009). pH = -log10 [H3O+]
(Equation 1-4)
pH scale is used to express the acidity or basicity of solution. Solutions with pH < 7 are acidic, pH = 7 are neutral and pH > 7 are basic. For example a solution having [H3O+] concentration of 4.23 × 10-3 would have a pH of 2.374 and is acidic. There are a couple of methods that can be used to monitor the pH change in a solution. In this experiment, pH electrodes will be used to constantly measure and record the pH of the solution. The titration is initiated by using a pH electrode to measure the pH of acid solution (pH within 3-5). As sodium hydroxide, NaOH, is incrementally added to the acid solution, some of the hydrogen ions will be neutralized. The pH of the solution will gradually increase as the hydrogen ion concentration decreases. When sufficient NaOH is added to completely neutralize the acid (most of the H3O+ ions are removed from the solution), the next drop of NaOH added will cause sudden change in pH (increase sharply). This shows that the equivalence point of titration is reached. The volume of based needed to completely neutralized the acid is 3
obtained from the equivalence point of titration. From the graph of pH versus the volume of base added plotted, the equivalence point is the middle of the vertical part of the curve (Chang, 2007).
Titration of vinegar sample with a standardized sodium hydroxide solution will be performed in this experiment. To standardize the sodium hydroxide solution, a primary standard acid solution is initially prepared. In general, primary standard solutions are produce by dissolving a weighed quantity of pure acid or base in a known volume of solution.
Sodium hydroxide (NaOH) is one of the bases commonly uses in laboratory. However, it is difficult to obtain solid sodium hydroxide in a pure form because it has tendency to absorb moisture from air, its solution react with carbon dioxide and often appears wet. For these reasons, a solution of sodium hydroxide must be standardized before it can be used in accurate analytical work. The sodium hydroxide solution can be standardized by titrating it against an acid solution of accurately known concentration. The acid chosen for this experiment is a monoprotic acid called potassium hydrogen phthalate (KHP), for which the molecular formula is KHC8H4O4 (Chang, 2007). The equation for the reaction of potassium hydrogen phthalate with sodium hydroxide is: KHC8H4O4(aq) + NaOH(aq) KNaC8H4O4 + H2O(l)
(Equation 1-5)
Once the sodium hydroxide solution has been standardized, it will then used to titrate 10.0 mL aliquots of vinegar. The equation for the reaction between vinegar and sodium hydroxide is: CH3COOH(aq) + NaOH(aq) NaCH3COO(aq) + H2O(l)
(Equation 1-6)
By knowing the standardized NaOH concentration and using Equation 1-6, the molarity and mass percentage of acetic acid in the vinegar solution can be determined.
4
APPARATUS/ MATERIALS
pH meter
Vinegar
Magnetic stirrer
Weighing balance
Stir bar
Burette and burette clamp
Beaker (100, 250 mL)
Measuring cylinder
Conical flask
Retort stand
10 mL volumetric pipette
0.6 M NaOH solution
1.5 g potassium hydrogen phthalate (KHP)
Diagram of titration
.
5
PROCEDURES
Standardization of sodium hydroxide solution
1. 250 mL of approximately 0.6 M sodium hydroxide solution was prepared from NaOH solid. The solution was initially prepared in a beaker to make it easier to transfer into the burette for titration. 2. A beaker was placed on the balance and tare. 1.5 grams of KHP was added to the beaker. The mass of KHP to the nearest 0.001 g was recorded and 30 mL of distilled water was added to the beaker. The solution was stirred until the KHP had dissolved completely. 3. This solution was titrated with NaOH and the pH with 1 mL additions of NaOH solution was recorded. 4. Steps 1 to 3 were repeated and two more solutions for NaOH standardization were prepared. 5. The graph of pH versus NaOH was plotted. From the plotted graph, the volume of NaOH required to neutralize the KHP solution in each titration was determined. 6. The molarity of sodium hydroxide for titrations 1, 2 and 3 was calculated. 7. The average molarity of the sodium hydroxide solution was calculated. The result of sodium hydroxide concentration was used in part B of the experiment.
Molarity of acetic acid and mass percent in vinegar
1. A 10 mL volumetric pipette was used to transfer 10.0 mL of vinegar to a clean, dry 100 mL beaker. Sufficient water, 75 to 100 mL was added to cover the pH electrode tip during the titration. 2. 1 mL of NaOH was added to the vinegar solution and the pH was recorded. 3. The above steps were repeated twice more. 4. The graph of pH versus NaOH volume added was plotted. The volume of NaOH required to neutralize the vinegar in each titration was determined from the plotted graph and the data was recorded. 5. The molarity of acetic acid in vinegar for titrations 1, 2 and 3 were calculated. 6. The average molarity of acetic acid for each titration was calculated. 7. The percent by mass of acetic acid in vinegar for titrations 1, 2 and 3 were calculated. 8. The percent by mass of acetic acid in vinegar was calculated.
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RESULTS
Standardization of sodium hydroxide solution
Mass of KHP (g)
TITRATION 1
TITRATION 2
TITRATION 2
1.529
1.5047
1.5057
12.27
11.82
12.29
0.6102 M
0.6234 M
0.6000 M
Volume of NaOH to neutralize the KHP solution (mL) Molarity of NaOH Average molarity of
0.6112 M
NaOH for each titration
Molarity of acetic acid and mass percent in vinegar
TITRATION 1
TITRATION 2
TITRATION 3
15.54
15.48
15.55
0.9498 M
0.9461 M
0.9504 M
Volume of NaOH to neutarlize the vinegar solution (mL) Molarity of acetic acid in vinegar Average molarity of acetic acid for each
0.9488 M
titration % by mass of acetic acid in vinegar
5.7045 %
5.6823 %
5.7081 %
Average percent by mass of acetic acid in
5.6983 %
vinegar % error
42.457 % error
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Standardization of sodium hydroxide solution
KHP TITRATED WITH NaOH (TITRATION 1) 14 12 10
12.27 mL
pH
8 6 4 2 0 0
2
4
6
8
10
12
14
16
VOLUME OF NaOH (mL)
KHP TITRATED WITH NaOH (TITRATION 2) 14 12 10 8 pH
11.82 mL
6 4 2 0 0
2
4
6
8
10
12
14
16
VOLUME OF NaOH (mL)
8
KHP TITRATED WITH NaOH (TITRATION 3) 16 14 12
pH
10 8
12.29 mL
6 4 2 0 0
2
4
6
8
10
12
14
16
VOLUME OF NaOH (mL)
Molarity of acetic acid and mass percent in vinegar
VINEGAR TITRATED WIYH NaOH (TITRATION1) 14 12
pH
10 8
15.54 mL
6 4 2 0 0
2
4
6
8
10
12
14
16
18
VOLUME OF NaOH (mL)
9
VINEGAR TITRATED WITH NaOH (TITRATION 2) 14 12 10
pH
8
15.48 mL
6 4 2 0 0
2
4
6
8
10
12
14
16
18
VOLUME OF NaOH (mL)
VINEGAR TITRATED WITH NaOH (TITRATION 3) 14 12 10
pH
8
15.55 mL
6 4 2 0 0
2
4
6
8
10
12
14
16
18
VOLUME OF NaOH (mL)
10
CALCULATION TO PREPARE 0.6 M NaOH
× 250 mL ×
×
= 6 g NaOH solid
SAMPLE OF CALCULATION FOR STANDARDIZING BASED WITH KHP
Titration 1
Calculate the moles of KHP used in the titration.
1.529 g KHC8H4O4 ×
= 0.007488 mol KHC8H4O4
Calculate the moles of NaOH required neutralizing the moles of KHP.
0.007488 mol KHC8H4O4 ×
= 0.007488 mol NaOH
Calculate the molarity of NaOH solution.
12.27 mL NaOH ×
M1 =
= 0.01227 L NaOH
=
= 0.6102 M NaOH
* The same calculation was used for titrations 2 and 3 to find the molarity of NaOH solution (Appendix 1).
Calculate average molarity of sodium hydroxide for each titration.
Mav =
=
= 0.6112 M NaOH
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SAMPLE
OF
CALCULATIONS
FOR
DETERMINING
THE
ACETIC
ACID
CONCENTRATION IN VINEGAR BY TITRATION WITH STANDARD BASE
Titration 1
Calculate the moles of NaOH that reacted.
15.54 mL NaOH ×
= 0.01554 L NaOH
0.01554 L NaOH ×
= 0.009498 mol NaOH
Calculate the moles of CH3COOH neutralized by the moles of NaOH.
0.009498 mol NaOH ×
= 0.009498 mol CH3COOH
Calculate the molarity of the CH3COOH solution.
10 mL CH3COOH × M1 =
= 0.010 L CH3COOH solution =
= 0.9498 M CH3COOH
Calculate the mass of acetic acid in the solution.
10 mL CH3COOH ×
= 0.010 L CH3COOH solution
0.010 L CH3COOH ×
×
= 0.57045 g
CH3COOH
Calculate the mass of acetic acid solution.
10 mL CH3COOH solution ×
= 10.0 g CH3COOH solution
Calculate the percent by mass of acetic acid in the solution.
Percent mass CH3COOH =
Percent mass CH3COOH =
× 100%
× 100% = 5.7045 % CH3COOH 12
* The same calculation was used for titrations 2 and 3 to find the molarity of NaOH solution (Appendix 1).
Calculate average molarity of acetic acid for each titration.
Mav =
=
= 0.9488 M CH3COOH
Calculate average molarity of acetic acid for each titration.
Percent mass CH3COOH av =
=
= 5.6983
% CH3COOH
SAMPLE ERROR CALCULATIOS
% error =
=
× 100% = 42.457 % error
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DISCUSSION
The objectives of this experiment are to determine the molarity of a solution and the percent by mass of acetic acid in vinegar by titration with a standardization sodium hydroxide (NaOH) solution. For the first experiment, the purposes are to standardize the sodium hydroxide solution and to determine the molarity of sodium hydroxide. Sodium hydroxide was needed to standardize because it can easily contaminated by the absorption of carbon dioxide or water from the air as it absorbs moisture from the air and usually appears wet. Thus if a solution of sodium hydroxide is prepared by weighing the sodium hydroxide, the concentration for the solution may be not the approximately the desired concentration. On another hand, potassium hydrogen phthalate (KHP) has a lesser tendency to absorb water from the air and when dried will remain dry for a reasonable period of time. Potassium hydrogen phthalate is a white, soluble solid that is commercially available in highly pure form. Potassium hydrogen phthalate is a primary standard. Besides that, KHP also must be dissolved completely and carefully at the beginning of the titration to get the accurate concentration of NaOH. This means that more accurately the concentration of the solution of known concentration is known (KHP), the more accurately the concentration of the unknown solution can be determined (NaOH).
pH is a method of expressing the acidity or basicity of a solution. Solution with pH<7 are acidic, pH=7 are neutral and pH>7 are basic. In this experiment pH was used to determine the equivalence point of the solutions. Equivalence point occurs when chemically equivalent amounts of acid and base are present. At this point, the pH will rapidly increase with a small addition of NaOH. This means that equivalence point in this experiment is the point at which the volume of NaOH required to neutralize the acid solution.
In the first experiment, 6 g of NaOH was diluted with 250 mL of distilled water in order to get 0.6 M NaOH. For this experiment the titration of KHP with NaOH was repeated for 3 times to get the average. At first titration, the pH was slowly increase from 1 to 12 mL NaOH added which are from pH 4.06 to 6.77. After 0.5 mL of NaOH was added, the pH sharply increases to 12.16. The same goes to the titration 2 and 3 at which the pH sudden change after 12 mL of NaOH was added to the KHP solution. This means that the pH of the solution had changed from acid to base. The equivalence points are 12.27 mL for the first titration, 11.82 mL for the second titration and 12.29 mL for the third titration. The molarity of NaOH for first titration is 0.6102 M, second titration, 0.6234 M and lastly third titration, 0.6000 M. The average molarity of NaOH for these three titrations is 0.6112 M. 14
In second experiment, 10 mL of vinegar was used as acid. The titration of vinegar with NaOH solution was also repeated for three times in this experiment. For the first titration, the pH of the solution was slowly increased when NaOH was added from 1 to 15 mL which are from pH 3.09 to 6.06. After 1 mL of NaOH was added, the pH drastically increases to 11.89. The same goes to the titration 2 and 3 at which the pH sudden change after 15 mL of NaOH was added to the vinegar solution. This means that the pH of the solution had changed from acid to base. The equivalence points for each titration is 15.54mL for the first titration, 15.48mL for the second titration and 15.55mL for the third titration. The molarity of acetic acid in vinegar for first titration is 0.9498 M, second titration, 0.9461 M and lastly third titration, 0.9504 M. The average molarity and percent by mass of acetic acid in vinegar for these three titrations is 0.9488 M and 5.6983 %. The percentage error for the experiment is 42.457 %.
CONCLUSION
The purpose of this experiment was to determine the molarity of a solution and the percent by mass of acetic acid, CH3COOH, in a commercial vinegar sample by titration with a standardized sodium hydroxide solution. The percentage of CH3COOH was found to be 5.6983 %. The exact value was 4.00 % from the label on the commercial vinegar sample. The deviation of the experimental value from the true value was calculated to be 1.6983. Using the true value and experimental value, a percent error was calculated. The percent error was found to be 42.457 %. This shown that the objective was achieved because for a manufacturer of vinegar to use the word “vinegar” in its labeling of the product, it must contain a minimum of 4.00 % of CH3COOH. Based on this experiment, we can conclude that the experimental percentage of CH3COOH in the vinegar sample is higher than the true value. One possible source of error to account for this different could be the titration past the true equivalence point of the reaction. Addition of more NaOH than desired volume to reach the end point would give a larger value for the number of moles of NaOH used. An error in the molarity of NaOH used would result in a larger value for the number of moles of CH3COOH in the sample. A larger number for the moles of CH3COOH would result in a larger mass of CH3COOH in the sample. The larger mass of acetic acid would result in a higher percentage of acid in the sample.
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RECOMMENDATIONS 1. Student should be given more than one weak acid. Thus, they will able to compare and study the efficiency of each weak acid.
2. The wet burette must be rinsed with NaOH before filling to make sure that the concentration of the NaOH was not contaminated with other particles.
3. Potassium hydrogen phthalate solutions must be carefully prepared to get the accurate concentration of NaOH.
4. The tip of the burette must be filled with NaOH before initial volume was recorded in order to get the accurate volume that required neutralizing the acid solution.
5. An indicator such as phenolphthalein must be used to monitor the pH change. An indicator changes color when the pH reaches a certain point. So, we will be aware with the pH changes and when it reaches the equivalence point.
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REFERENCES Brown, T. L.,LeMay, H. E.,Bursten, B. E.and Murphy, C. J. (2009). "CHEMISTRY: the central science," 11th edition/Ed. Prentice Hall. 673 - 675.
Chang, R. (2007). "Chemistry," 9th edition/Ed. McGraw-Hill. 150 - 709.
Stanton;, B.,Zhu;, L.and Atwood, C. H. (2009). "EXPERIMENTS IN GENERAL CHEMISTRY: measureNet," 2nd edition/Ed. Cengage Learning. 155 - 157.
Scharf, W.and Malerich, C. (2010). Determination of Acetic Acid Content of Vinegar. Baruch College.
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