Gravimetric Analysis of a Chloride Salt

Name: Yuseung Ohn

Date: November 10,, 2014

Group: Tuesday AM Group E

Purpose:

This experiment was conducted in order to derive the chloride content of an unknown soluble

salt through the process of a precipitate-based gravimetric analysis.

Theory:

Some chemicals and solutions can quickly disassociate into their respective ions and components

when dissolving in other solutions. On the other hand, some solvents are insoluble in other

solutions which means that they remain solid when mixed. With this, researchers and chemists

can selectively precipitate ions from stoichiometric analysis. This applies to this lab because the

silver and chloride ions produce a silver chloride precipitate. Silver chloride has a very low Ksp

value and its reverse reaction takes place at a very low rate (Jonte 1958). The yield is considered

to be almost zero.

Ag+(aq) + Cl-

(aq) => AgCl(s)

AgCl(s) ⇌   Ag+(aq) + Cl-

(aq) Ksp = [Ag+(aq)]·[Cl-

(aq)] = 1.6 x 10-10

The chloride ions in the solution try to form a product with an equal amount of silver ions in the

solution to produce the silver chloride. With the knowledge of the Ksp value, the maximum

solubility of the silver chloride can be used to find the maximum amount of chloride. From the

following equation, the chloride ions can be represented by x:

Ksp = [Ag+(aq)]∙[Cl-

(aq)] = 1.6 x 10-10

[x]∙[x]=1.6 x 10-10

x =1.3 x 10-5

The reaction allows the precipitate to form small, crystalline structures but with heating and

stirring, they will collide in the presence of nitric acid. With this new solution, the colloid

coagulates into even larger and more lattice pieces which can then be filtered, such as through a

crucible.

The silver chloride crystals actually decomposes back into its ions if it is in an area with a

presence of light. This allows the mass of the final precipitate and affect the other calculations

after the filtering processes. The colour change from white to light-purple was credited to the

silver build-up from the light. The outer layers of the precipitate slowly turn into a solid silver

coating, which changes the final mass of the solid crystals.

AgCl(s) => Ag(s) + (1/2)Cl2(g)

3Cl2(g) + 3H20(l) + 5Ag+(aq) => 5AgCl(s) + CIO3

-(aq) + 6H+

(aq)

To calculate the chloride content of the unknown soluble salt,

1) Find moles of chloride: mass of AgCl / molar mass of AgCl

2) Convert moles of chloride to mass of chloride: moles of Cl- * molar mass of Cl-

3) Calculate percent chloride of original sample: (mass of Cl- / mass of sample) * 100%

The speed of the reaction can actually allow many other unwanted ions in to the solution, such as

nitrate and hydroxide (Ajmal 1993). The reaction must take place in an acid because this

prevents the anions of weak acids from co-precipitating. In addition to this, the then high amount

of silver ions may then co-precipitate with the unwanted added ions which will increase the mass

entirely. The ions that were entrapped through co-precipitation can be removed with 100mL of

water and the 15mL of acetone. Some silver chloride may be rinsed along with the unwanted

ions so to account for this,

Loss in moles = (Ksp)(.1L)

= 1.6 x 10 -11 moles

Converting to grams:

(moles of AgCl)(molar mass of AgCl) = mass lost

(1.6 x 10-11mol)(143.32 g/mol) = 2.3 x 10-9 g lost

Procedure:

At the beginning of the lab, a mysterious soluble salt was given and its sample number along

with its actual chloride content was recorded in the lab notebook. A sample of 0.1246g of the salt

was weighed through the method of “weighing out by difference”. Next, 100mL of distilled

water was poured in along with 1mL of dilute (6M) nitric acid. The beaker with the solution was

placed on a hotplate which was set to around 4-5, and the necessary 24.304mL and extra 5mL of

a 0.1M silver nitrate was slowly poured into the heated solution. The beaker was gently stirred as

it heated near boiling. After around 25 minutes of waiting, a few drops of silver nitrate were

placed into the beaker to check for the completion of the reaction. This made sure that no more

new silver chloride ions formed in the solution. The beaker was then moved from the hotplate to

a dark, cool area, such as the inside of a cabinet at the lab bench. As it cooled, a sintered glass

crucible was weighed on the analytical balance and was 32.6394g. The crucible was grabbed

with tissues and was not touched by fingers to avoid contamination. After approximately 10

minutes of waiting for the beaker to cool, the vacuum filtration apparatus was set with the help of

the TA. The empty crucible was attached to the apparatus and the solution was decanted through

the glass filter while leaving the precipitate in the beaker. To transfer all the precipitate, 5mL of

0.01M nitric acid was added to the beaker and was gently swirled in order to decant the washings

through the filter again. The previous step was repeated by washing the beaker with nitric acid,

but the solids were also transferred to the filter. All the remaining precipitate in the beaker was

dislodged with the rubber policeman to maximize all precipitate. Next the precipitate in the glass

crucible was washed with 0.01M nitric acid again by pouring 10mL over the crucible. The

filtered solution was emptied into a 250mL beaker and poured in a test tube. The collected nitric

acid was tested for turbidity by getting a TA to check it with washings of hydrochloric acid.

After passing the turbidity test, the crucible of precipitate was washed with 15mL of acetone.

The crucible was then placed in the Fisher Isotemp Oven Junior Model at 118˚C for about 25

minutes. After being heated, the crucible was placed in the desiccator to cool. After 5 minutes of

cooling, the crucible and dried sample was weighed on the analytical balance to be 32.8770g.

Each lab partner was responsible for performing their own trials to compare results with one

another.

Observations:

Table 1: Qualitative descriptions of sample soluble salt #352 of its precipitate and different

tests

Sample #352 (60.20%)

Sample Description

Times washed with 0.01M nitric acid

Test for completion of precipitation

Precipitate description

Sample 1 White powder 3 Passed with one test with HCl

Light purple, powder texture

Sample 2 White Powder 5 Passed with two test with HCL

Light purple, powder texture

Data:

Table 2: Quantitative analysis for both samples 1 and 2

Sample 1 2Actual Chloride Content 60.20%Sample of Salt in Beaker 0.1246±0.001g 0.1420±0.001gRequired Amount of 0.1M Silver Nitrate With added excess 5mL

24.3±0.803% mL 27.0±0.704% mL

Empty Crucible Mass 32.6394±0.001g 32.6275±0.001gOven Model and Temperature

Fisher Isotemp Junior Model at 118˚C

Drying Time 30 minutesCooling Time 5 minutesCrucible and Precipitate Mass

32.8770±0.001g 32.8520±0.001g

Precipitate Mass 0.2376±0.002g 0.2245±0.002g

Table 3: Further calculated quantitative values regarding precision and accuracy

Sample 1 2Chloride ion Mass 0.05877g 0.842% 0.05553g 0.891%Salt Chloride Content by Weight

47.17% 1.645% 39.11% 1.595%

Average Chloride Content by Weight

43.14% 3.240%

Relative Error -0.28%

Relative Spread 186.83ppt

Calculations:

Calculating amount of silver nitrate necessary:

Ag+(aq) + Cl-

(aq) => AgCl(s)

The ratio of Ag to Cl is a 1 to 1 molar ratio and this was derived from basic stoichiometric

analysis. Assuming the chloride content by weight is 55% in the sample,

(sample mass)(55% chloride by mass)/(molar mass of Cl-) = moles Cl-

Sample 1: (0.1246±0.001g)(0.55)/35.453g/mol = 0.00193±0.803% mol

Sample 2: (0.1420±0.001g)(0.55)/35.453g/mol = 0.0022±0.704% mol

Because of the 1:1 ratio, moles of Cl- will equal the moles of Ag+. Next, convert to mL…

(moles of Ag+)/(mol/L of AgNO3)(1000mL/L) = mL of AgNO3

Sample 1: ((0.00193mol 0.803%)/(0.10 moles/litre))(1000 ml/Litre) = 19.3mL 0.803%

AgNO3 + 5mL excess = 24.3mL AgNO3 0.803%

Sample 2: ((0.0022mol 0.704%)/(0.10 moles/Litre))(1000 ml/Litre)= 22.0mL 0.704%

AgNO3 + 5ml excess = 27.0mL AgNO3 0.704%

Mass of Chloride Ions:

Again, with the 1:1 ratio…

(precipitate of mass)/(molar mass of AgCl)(molar mass of Cl-) = mass of Cl-

Sample 1: ((0.2376 0.002 g)/(143.32 g/mol))(35.453g/mol) = 0.05877g 0.842%

Sample 2: ((0.2245 0.002 g)/(143.32 g/mol))(35.453g/mol) = 0.05553g 0.891%

Percent Weight of the Sample Salt of Chloride ions:

(mass of Cl-)/(mass of unknown salt sample) * 100 = Cl- % by Weight of sample

Sample 1: ((0.05877g 0.842%)/(0.1246g 0.803%)) x 100% = 47.17% 1.645%

Sample 2: ((0.05553g 0.891%)/(0.1420g 0.704%)) x 100% = 39.11% 1.595%

Average of Trials:

(Value 1 + value 2)/2 = average value

(47.17% 1.645% + 39.11% 1.595%)/2 = 43.14% 3.240%

Relative Error:

%error = (average value - accepted value)/(accepted value) * 100

= (43.14-60.2)/(60.2)*100

=-0.28%

Relative Spread:

Spread = (high value – low value)/(average value) * 1000 parts per thousand

(47.17%-39.11%)/43.14 * 1000ppt = 186.83ppt

Discussion:

The results were determined and was noticed to be significantly lower than the accepted value. A

possible reason for this occurrence may have happened due to the excessive amounts of the

silver nitrate tests. Due to impatience, the silver nitrate test happened many times which added

much more silver content to the AgCl precipitate. In essence, this had increased the silver

content and most likely decreased the chloride content. In order to prevent this, when the

solution was on the hotplate, waiting patiently and diligently for the solution to turn transparent

seemed to be highly recommended. Another factor that lowered our results was due to the

excessive stirring during the heating of the solution. The rapid stirring during the reaction may

have allowed other ions to become included into the collisions, which would increase the

external ion content but lower the chloride content. This could have been avoided by lowering

the rate of stirring and leaving the solution alone during the heating session. In addition to this,

another factor may have been the inaccurate mass of the precipitate. When using the rubber

policeman to dislodge the solid content from the beaker, not all of the precipitate may have been

dislodged into the crucible. Because of this, the mass of the precipitate was lower and effectively

lowered the results from the calculations of the chloride content. Finally, another potential reason

for lower values may have been credited to the photodecomposition of the silver nitrate. When

the solution was in the bright light in open air, this could have increased the concentration and

mass of the silver. To avoid this, after the reaction was finished, the solution should have been

moved quickly to a dark, cool area.

Conclusion:

The sample salt #352 had a chloride content of 60.20%. However, through two trials, the chloride content of the salts was averaged to be 43.14% 3.240%. The calculated values were found to not be accurate in correspondence with the accepted value as the relative error was around -0.28% and a relative spread value of 186.83 ppt.

Bibliography:

Ajmal, Mohammad, Afsar M. Sulaiman, and Akhtar Hussain Khan. 1993. "Surface Entrapment

of Toxic Metals from Electroplating Waste and Their Possible Recovery." Water, Air, &

Soil Pollution 68.3:485-92. Print.

Burk, R.C., M. Azad, X. Sun, and P.A. Wolff. 2014. Introductory Chemistry Laboratory Manual.

Print.

Jonte, J. Haworth, and Don S. Martin. 1958. "The Solubility of Silver Chloride and the

Formation of Complexes in Chloride Solution." Journal of the American Chemical Society

74.8:2052-054. Print.

Skoog, Douglas A., and Donald M. West. 1996. Fundamentals of Analytical Chemistry. 7th ed.

Fort Worth: Saunders College Pub, Chapters 5-8.