fix the laratory Report

Determination of Purity Using Titration

INTRODUCTION

A common problem in chemistry is determining whether or not a substance is impure and to what extent it is impure. Chemicals used in experiments or for commercial purposes often need to be pure to ensure safety and efficiency during reactions, but determining whether or not a given substance is pure can be difficult. If the chemical is an acid or a base, titration with a standardized solution can be used to assess its purity.

In a generic titration, “one solution of known concentration is used to determine the concentration of another solution through a monitored reaction.”1 An acid-base titration works on the principle that the acidic solution will combine the basic solution to neutralize, altering the pH of the solution.

H3O+(aq) + OH- (aq) 2H2O(l) (1)

Eq. 1 shows that when equal moles of the H3O+ion and the OH- ion are present, the solution is completely neutralized. 2 When this dynamic equilibrium occurs in a titration, it is called the equivalence point, and at that point a given reaction has a characteristic pH. Thus, indicators, substances that can change color depending on the pH, are valuable tools for monitoring the progress of titration reactions. A few drops of indicator are added to one of the solutions and the other solution is added dropwise. When the moles of acid and base are equivalent, the equivalence point is reached and adding one more drop of the titrant will cause the indicator to change color, signaling the end point of the reaction. The end point occurs just after the equivalence point, but the extra titrant required to reach the end point is negligible in calculations.

The titration method is useful in determining purity only if a standardized solution, one with a known molarity, is available for the titration. If a known volume of standardized solution is used in a titration, then the moles of both acid and base can be determined. From the moles, the mass of the pure substance can be determined and compared to the mass of the impure substance to find the percent purity.

In this experiment, a sodium hydroxide, NaOH, solution was standardized by titration with pure hyrdrochloric acid, HCl.

HCl (aq) + NaOH (aq) H20 (l) + NaCl (aq) (2)

This reaction was monitored using phenolphthalein indicator, which changes from clear to pink near a pH of 8, corresponding to the pH at the reaction’s equivalence point. Once the molarity of the standard NaOH solution was known, the solution was used to titrate an impure sample of industrial grade muriatic acid, HCl, again using phenolphthalein indicator, and the purity of the muriatic acid sample was determined.

EXPERIMENTAL

25 mL of pure 1.000 + 0.003 M HCl was measured using a 25.00 mL graduated pipet and added to a 250 mL flask along with 2 drops of phenolphthalein indicator. Next, an NaOH solution of approximately .1 M was created by diluting 75 mL of 1 M NaOH to 750 mL with DI water. The NaOH solution was then placed in a 50 mL buret and added dropwise to the HCl solution until the indicator in the solution turned pink. The above procedure was repeated three times to standardize the NaOH solution, at which point three samples of approximately 35 mL impure muriatic acid were prepared and titrated again using the NaOH solution and phenolphthalein.

DATA

The results obtained from the above procedure may be found in Data Tables 1, 2, and 3. In Table 1, the moles of HCl were obtained from the measured volumes and then equated to moles of NaOH. The molarity of the NaOH solution was calculated by dividing the moles of NaOH by the volume of liters of NaOH delivered during titration.

Moles HCl = Moles NaOH=Molarity x Liters HCl (3)

Molarity, NaOH = Moles Solute/ Liter Solution (4)

Table 1: Standardization of NaOH Solution

In Table 2, the volume of the impure muriatic acid was measured using a volumetric pipet, but the number of moles of HCl in the samples was determined by equating the moles of NaOH used to titrate to the moles of HCl. The percent purity of the sample was the mass of the HCl in the sample divided by the total sample mass. Moles of pure HCl in the standard HCl and moles of muriatic acid were converted to grams present and the % purity calculated according to Eq. 5.

% Purity = Mass of Pure Substance / Mass Standard x 100% (5)

Table 3 summarizes the results of Tables 1 and 2: the mean and standard deviation for NaOH Molarity and HCl purity.

Table 2: Percent Purity of Impure Muriatic Sample

Table 3: Statistcal Analysis of NaOH Molarity and KHP Percent Purity Results

RESULTS AND DISCUSSION

The molarity of the NaOH solution was determined to be .095 ± .001 M and the percent purity of the muriatic acid sample was determined to be 31.8 ± .2 %. In determining the molarity of NaOH solution, the first trial was not used because it was visibly over-titrated. Trial 3 used a higher volume of muriatic acid than trial 1, but needed less titrant, indicated, in addition to the visible evidence, that trial 1 was over-titrated. Trials 2, 3, and 4 appear consistent: the trial with the highest volume required the most NaOH and the trial with the lowest volume required the least NaOH.

In the second part of the experiment, only three trials were performed because none of the solutions appeared over-titrated. However, trials 2 and 3 are not consistent: trial 2 used more HCl than trial 3, but needed less NaOH. Trial 2 may have been slightly under-titrated or trial 3 may have been slightly over-titrated. All of the trials in the second part of the experiment required less NaOH than the trials from the first part, even though a greater volume of HCl was used. This was a clear indication that the samples in the second phase were highly impure.

The most probable cause of error in this experiment was over-titration. It is difficult to add extremely small drops toward the end of titration; often the last drop or splash added causes the solution to change from clear to dark pink rather than a pale pink. If over-titration occurred in only the first phase of the experiment, it would deflate the molarity of the NaOH and subsequently the percent purity of the muriatic acid. If over-titration occurred in only the second phase of the experiment, it would inflate the percent purity of the HCl.

Another problem was losing drops of NaOH solution: drops of solution stuck to the sides of the flask, occasionally slipped out of the loose-fitting stopcock, and often a single drop lingered on the buret’s tip at the end of titration. Losing drops of NaOH would make the volume of NaOH used too high, having the same effect on the results as over-titration. In order to reduce these effects, the sides of the flask were rinsed with DI water, the stopcock was held tightly on the buret, and the buret was turned off quickly at the end point instead of slowly to avoid drop formation.

Problems may also have occurred if any HCl was lost during measurement with the pipet. The volume on graduated pipets can be difficult to read because the meniscus level of measured liquid is often difficult to steady. Overestimating the amount of hydrochoric acid in the flask only in the first part of the experiment would inflate the molarity of the NaOH solution and subsequently the percent purity of the HCl. If HCl were lost only in the second part of the experiment, the calculated percent purity of the HCl would be too low.

CONCLUSION

The molarity of the NaOH solution was determined to be .095 + .001 M and the percent purity of the muriatic sample was determined to be 31.8 ± .2 %. Though the true molarity and percent purity were unknown, the standard deviation in these numbers was fairly low, indicating precise, if not accurate, results. The accuracy of this experiment could be improved by running more trials with both the pure and the impure muriatic acid. This experiment is significant, because being able to determine accurately and precisely the molarity of a solution is an important step in many experiments, and determining the percent purity of a sample is necessary in both scientific and commercial sectors.

REFERENCES

1. Anliker, Keith, et. al. Laboratory Manual for Experimental Chemistry I, 2002. 46-52.

2. Silberberg, Martin. Chemistry: The Molecular Nature of Matter and Change, 3rd ed. 2003.143-146.