Gravimetric Analysis of a Chloride Salt

Name: Andrew Yaksic

Purpose: To illustrate typical techniques used in gravimetric analysis by quantitatively determining the amount of chloride in an unknown.

Equipment: Analytical balance, 6 250-mL beakers, Bunsen burner, 3 funnels, funnel support, plastic wash bottle, 10-mL graduated cylinder, 100-mL graduated cylinder, ring stand, iron ring, wire gauze, 3 stirring rods, 3 rubber policemen, 3 pieces of #1 Whatman filter paper, 3 watch glasses, weighing paper

Materials: unknown chloride sample (sample #1), 0.5 M AgNO3, 6 M HNO3, acetone, distilled water

Introduction: Quantitative analysis is that aspect of analytical chemistry concerned with determining how much of one or more constituents is present in a particular sample of material. Information such as percentage composition is essential in analytical chemistry for determining chemical formulas for compounds. Two common methods for quantitative analysis are gravimetric analysis and volumetric analysis.

Gravimetric analysis derives its name from the process of isolating the desired constituent in weighable form. Volumetric analysis derives its name from the process of measuring the volume of a reagent. Gravimetric analysis, in short, involves changing one compound containing the constituent into another compound containing that constituent and measuring the percent chloride in the new compound to determine the percent chloride in the previous compound.

In this experiment, silver chloride will be produced from an unknown chloride compound. The percent chloride will then be determined based on the amount of silver chloride recovered from the precipitation reaction. In order to recover the precipitate, the following steps must be made (references to this figure will be made in the Procedure):

This experiment also illustrates the concept of stoichiometry. Stoichiometry is the determination of the proportions in which chemical elements combine and the mass relations in any chemical reaction. This experiment involves the proportions of chloride in two chloride compounds.

Procedure:

1. About .4 g of the unknown sample was weighed on weighing paper. The sample was transferred to a clean 250-mL beaker. The beaker was labeled #1.
2. Approximately 150 mL of distilled water and 1 mL of 6 M HNO3 were added to the beaker.
3. Steps 1 and 2 were repeated twice, with subsequent beakers labeled #2 and #3.
4. Using a different glass rod for each solution, the contents of the beakers were stirred until the sample had dissolved in each. The stirring rods were left in the beakers after stirring.
5. While stirring each solution, approximately 20 mL of 0.5 M AgNO3 solution was added. A watch glass was placed over each beaker. Each solution was warmed gently with a Bunsen burner and was kept warm for 10 minutes. Solutions were not allowed to come to boil.
6. Three filter papers were weighed. The masses were recorded. The paper was folded as illustrated in Figure 8.1 in the Introduction. The filter paper was placed into a funnel. Each filter paper was wet with distilled water to hold it in place in the funnel.
7. The precipitate and the warm water from each beaker were transferred into each funnel. A rubber policeman and a wash bottle were used to obtain the last traces of precipitate from each beaker. The level of solution in each filter funnel was kept below the top of the filter paper.
8. After all of the precipitate had been transferred from beaker to funnel, approximately 5 mL of distilled water was poured through the filter three times. 5 mL of acetone was poured through the filter.
9. The filter paper was removed from each funnel and placed on a numbered watch glass. It was stored away from light overnight.
10. After the precipitate on the filter paper had been thoroughly dried overnight, the mass of each filter paper was taken.

Observations:

Results:

Mass of AgCl. (Mass of filter paper + AgCl minus mass of filter paper.) =
Trial 1. 1.8288 ± .0001 g – 1.1090 ± .0001 g = .7198 ± .0002 g.
Trial 2. 1.7845 ± .0001 g – 1.0599 ± .0001 g = .7246 ± .0002 g.
Trial 3. 1.8024 ± .0001 g – 1.0956 ± .0001 g = .7068 ± .0002 g.

Mass of Cl in original sample.
Mass of Cl in AgCl. (Molecular weight of Cl divided by molecular weight of AgCl times mass of AgCl sample.) =

Trial 1. (.24766)(.7198 ± .0002 g)
(.24766)(.7198 ± .03% g)
.17827 ± .03% g
.17827 ± .00005 g

Trial 2. (.24766)(.7246 ± .0002 g)
(.24766)(.7246 ± .03% g)
.17945 ± .03% g
.17945 ± .00005 g

Trial 3. (.24766)(.7068 ± .0002 g)
(.24766)(.7068 ± .03% g)
.17505 ± .03% g
.17505 ± .00005 g

Mass of Cl in original sample. (Mass of Cl in AgCl.) =
Trial 1. .17827 ± .00005 g
Trial 2. .17945 ± .00005 g
Trial 3. .17505 ± .00005 g

Percent chloride in original sample. (Mass of Cl in original sample divided by mass of sample multiplied by 100.)
Trial 1. (.17827 ± .00005 g) / (.3718 ± .0001 g) x 100 = (.17827 ± .03% g) / (.3718 ± .03% g) x 100 =
47.95 ± .06%
47.95 ± .03 %
Trial 2. (.17945 ± .00005 g) / (.3673 ± .0001 g) x 100 = (.17945 ± .03% g) / (.3673 ± .03% g) x 100 =
48.86 ± .06%
48.86 ± .03 %
Trial 3. (.17505 ± .00005 g) / (.3766 ± .0001 g) x 100 = (.17505 ± .03% g) / (.3766 ± .03% g) x 100 =
46.48 ± .06%
46.48 ± .03 %

Average percent chloride. (Sum of percent chloride trials divided by 3.) = 47.76%

Standard deviation. (Square root of the sum of the squares of the deviations from the mean divided by number of observations minus one.) =
sqrt((.192 + 1.12 + 1.282) / 2) =
1.44 --> 1 = standard deviation.

Relative standard deviation. (Standard deviation divided by average percent chloride.) =
1/47.76 =
.02.

None of the results differ from the mean by more than two standard deviations. Reported percent chloride is 48 ± 1 percent chloride.

Discussion: There were several sources of error in this experiment that led to a standard deviation of 1. If any contaminants got into any of the reagents involved in this experiment, then other reactions would cause other precipitates to form, causing inaccurate mass readings for the supposed silver chloride compound. Some possible points of ingress for contaminants are during handling of compounds (sweat on hands could contaminate reagents); if tap water was accidentally used (substances in the water would react with silver nitrate); or other contaminants in the original sample. Other sources of error include bad technique. If any substances were dropped or otherwise lost during handling, the mass reported would be too low. Decomposition from light would also produce low results.

The theory associated with this experiment is the atomic theory. The chemical reaction takes place because of the ionic charges on the elements and because of the activity series. The activity series dictates which metathesis reactions will take place based on the electronegativities of elements. Relative attractions for electrons, and thus ability to form ionic bonds, dictates not only which compounds can exist, but which reactions will occur between ionic compounds.

Some ramifications for this experiment include increased laboratory experience. Laboratory techniques such as filtration and use of rubber policemen were used for the first time in the laboratory. On a broader scale, the ability to test for unknowns is useful for both research scientists and practical, commercial scientists. Separation and stoichiometric analysis of compounds is the root of all analytical chemistry, and the ability to test compounds for unknown substances is an extremely important skill. This skill is used widely by scientists researching unknown compounds in riverbeds or searching for the identity of a newly-discovered adhesive. Many practical applications of the procedures and techniques practiced in this experiment exist.

Questions:

1. a. The mean is 32.51%. The standard deviation is .26%. The relative standard deviation is .008.
b. No results can be discarded. Each result was within .52% of the mean.

2. The percentage of barium in the compound is 24.39%.

3. The percentage of sodium in table salt is 39.32%.

4. There are 786 milligrams of sodium in 2.00 g of NaCl.

5. The impure sample contains 91.55% NaCl.

6. Sources of error can be found in the Discussion section.

Conclusion: This experiment produced extremely accurate results. The results were reasonably precise. Laboratory techniques were learned and practiced. Mathematical formulas were used correctly to produce proper results.