Determination of Concentration Using Titration
Introduction:
In chemistry, it is important to know the concentration of all solutions that we work with so that they may be safely and accurately used in chemical reactions. By determining the concentration of these solutions, it is also determined whether the solution is pure or impure. In order to determine the concentration and purity of a solution (specifically acids and bases), titration can be used. During titration, an unknown solution is placed in a beaker beneath a burette filled with a solution of known concentration. The known solution, called the titrant, is slowly added to the unknown until the reaction reaches neutralization, which is normally seen through a change in color due to an added indicator.
Pure solutions are defined as “substances that are made of only one type of atom or molecule” (Crawford “Pure Substance in Chemistry”). If a solution is impure, it means that the solution is a mixture of two types of atoms or molecules. An example of this is vinegar. Vinegar is composed of acetic acid (CH3COOH) and water. CH3COOH alone is a pure substance, but when it is mixed with water, it becomes an impure substance because it is no longer the only type of atom or molecule in the solution. Acid-base titrations (also known as neutralization reactions) are the most common type of titrations. These titrations occur between an unknown base/acid and a known base/acid. During these titrations, the final goal is to reach the equivalence point of the reaction. The equivalence point is the point where the moles of known substance are equal to the moles of the unknown substance in the solution, or more simply, where the reaction is neutralized. However, because we cannot physically see the equivalent point of a reaction, it is considered nearly impossible to reach this point.
In order to effectively determine the concentration/purity of a substance then, we use pH indicators (a compound that can be added to the solution which causes the solution to change color based upon acidity) to determine the end of the neutralization reaction, known as the end point. The pH indicators will cause the solution to change color when the endpoint is reached and the reaction has been neutralized. Although the endpoint is reached after the equivalence
point, the two points are so close together that the extra amount of known solution (titrant) is insignificant in the calculations.
When there is a known amount of a known standardized solution (a solution with known concentration), then we can calculate the moles and mass of both the known and unknown solutions after titration. Using these moles/masses, we can also calculate the mass % composition (concentration) of the unknown solution, as well as the percent error of the calculations (how inaccurate the calculations from the titrations were in comparison to the theoretical concentration of the solution).
The purpose of this lab is to determine the concentration of Acetic acid from an impure solution of acetic acid and water (Vinegar). In order to determine the concentration of pure acetic acid from the impure solution, NaOH (strong base) will be used to titrate acetic acid in a known amount of vinegar. The following formula will be used in the experiment:
NaOH(aq) + CH3COOH(l) NaCH3COOH + H2O
In this reaction, the concentration of CH3COOH is determined by titrating it with a known amount of standardized NaOH. From the results of this titration, we can calculate the moles/mass of both NaOH and CH3COOH, and the concentration of CH3COOH.
Experimental:
A 50 mL buret with a stopcock was filled with a .300 M NaOH solution and placed in a buret clamp. 10 mL of 5% stock vinegar was then measured using a graduated cylinder and added to a 250 mL Erlenmeyer flask along with 3 drops of phenolphthalein indicator and 10mL of deionized water. This flask was placed beneath the 50mL buret along with a sheet of white paper beneath the flask. The NaOH solution was then added to the vinegar in the flask 1 mL at a time while swirling the solution in the flask. This continued until the phenolphthalein indicator in the solution caused the solution to turn pink and remain that color for 30 seconds, indicating a change from an acidic to a slightly basic solution. The final volumes of NaOH and CH3COOH were recorded and used as a guide for the following trials. After the initial trial, this process was repeated three more times to determine an average experimental value of acetic acid in vinegar.
DATA
The results obtained in the experimental section above can be found below in the tables. In table 1, The moles of NaOH and Vinegar were found from the measured volume of NaOH which was then converted into moles of NaOH and Vinegar. The mass of vinegar was measured using a top-load balance by taking the difference between the mass of an empty beaker and the mass of a beaker container vinegar. The mass of acetic acid in vinegar was found by converting the moles of NaOH found in data table 1 to moles and then grams of Acetic Acid. The mass percent of acetic acid was obtained by dividing the mass of the acetic acid found by the total mass of vinegar, using the equation:
Mass % = (mass of solute/mass of solution) x 100%
Table 1 – Mass Percent of Acetic Acid in Vinegar
Trial 1
Trial 2
Trial3
Molarity of NaOH Solution (M)
0.3
0.3
0.3
Mass of Vinegar (g)
9.92
9.97
9.75
Volume of 0.3 M NaOH delivered (L)
.02307
.02161
.02232
Moles NaOH and Vinegar (mol)
.006921
.006483
.006696
Mass % of Acetic Acid in Vinegar (%)
4.19
3.90
4.12
In table 2, the mean, standard deviation, and percent error were calculated for the mass percent of acetic acid in vinegar.
Table 2 – Statistical Analysis of Mass percent of Acetic Acid in Vinegar
Mass Percent of Acetic Acid in Vinegar
Mean
4.1%
Standard Deviation
+ 0.2%
Mean +/- Std. Dev.
4.1 + 0.2 %
Percent error
82.0%
Results and Discussion:
The mass percent of acetic acid in vinegar was determined to be 4.1%. The molarity of each NaOH solution was 0.3M. Trial 1 used the largest volume of NaOH to standardize the vinegar solution. All three trials were fairly consistent and only had a standard deviation of + 0.2%. Three trials were performed to standardize the NaOH solution so that accurate as well as precise data could be obtained. The mass of vinegar in trial 3 was slightly less than that in trials 1 and 2, which caused it to have a lower mass percent of acetic acid in vinegar. The difference in masses between trial 3 and trials 1 and 2 caused the average mass percent of acetic acid in vinegar to be slightly lower than it would have been as well.
There were possible human errors in this lab, as well as procedural errors. A common procedural error was that a small drop of titrant could have been released from the buret but stuck on the tip of the buret so that they did not enter the vinegar solution. This could cause the end-point volume of NaOH higher than it should be, which would make the result of mass % of acetic acid in vinegar higher than it should have been.
Another source of error could have been if a drop of NaOH remained on the wall of the flask during titration and never reached the reaction mixture. This would cause a higher end-point volume of NaOH as well, which would make the result of mass % of acetic acid in vinegar inaccurately high. In order to minimize the errors in the experiment due to these two errors, it was ensured that the sides of the flask were rinsed with deionized water to wash all NaOH into the solution, and the stopcock was kept securely closed so that no extra drops could escape.
Human error could have caused inaccurate measurements in this lab as well. The volume on the buret could have been read inaccurately, which would result in inaccurate data. To avoid this, both partners read the volume on the buret and decided on an answer together.
Conclusion
The average mass percent of acetic acid in vinegar was determined to be 4.1 + 0.2 %. The percent error of the mass % of acetic acid in vinegar was calculated to be 18%. The true value of the mass percent of acetic acid in vinegar was 5.0%, so the results determined in lab were very precise, albeit not perfectly accurate. The accuracy of this experiment could have been improved by running more trials of the experiment and titrating with more caution to prevent over-titration of the acid. This experiment is extremely important, as it shows the process by which the unknown concentration of a solution can be determined. This process is used in both the scientific field to determine concentrations of solutions, as well as the medical field to accurately determine the dosages for patients.
Citations:
1.) Crawford, Nathan. “Pure Substance in Chemistry: Definition, Properties & Examples.” Study.com, Study.com, study.com/academy/lesson/pure-substance-in-chemistry-definition-properties-examples.html. Accessed 19 Sept. 2017.