Gravimetric Determination of chloride in a soluble sample

Abstract
The chloride present in an unknown soluble sample was precipitated into silver chloride through precipitation gravimetry. The colloidal silver chloride originally formed was converted to a crystalline solid by controlling certain parameters of the experiment such as temperature, pH of the solution, and concentration of AgNO3. Once the solid was large enough, it could be washed, filtered, and weighed. The percentage of chloride present was found to be 24.73695223 (±0.00000006) %.

Introduction
The purpose of this lab was to determine the percentage of chloride in an unknown soluble salt sample through a gravimetric method known as precipitation gravimetry. The objective was to accurately and precisely convert the unknown sample containing the chloride analyte to another known solid, silver chloride. Once the mass of the known solid is measured, the percentage of chloride analyte can be calculated using stoichiometry.
The chemical reaction occurring in this particular experiment involved the mixing of silver nitrate and the soluble salt in water containing nitric acid. Both reactants break into their ionic components allowing the silver ion and chloride ion to form silver chloride. The overall and net ionic equations for the reaction are as follows:
Overall Equation
AgNO(aq) + XCl(aq)  AgCl(s) + XNO3(aq)
Net Ionic Equation
Ag+(aq) + Cl-(aq) AgCl(s)
Precipitation gravimetry for this experiment was performed by using silver nitrate as the precipitating reagent then digesting the newly formed silver chloride sample in a solution containing ions. Through heating and the high concentration of the ions, the silver chloride coagulated to a desirable size that could be filtered and washed with nitric acid through a previously weighed crucible. Once the sample was washed free of impurities, it could be dried to a constant mass which allowed for the mass of chloride to be determined using a calculation explained later.
As it was alluded to in the previous paragraph, the silver chloride, when first precipitated by the silver nitrate, was originally colloidal in size, making the solution appear a milky. By adding the ions and increasing the kinetic energy of the solution through heating and stirring, the colloidal particles were forced together to form crystalline particles.
The effects of certain factors--- such kinetic energy (i.e. the temperature), solubility of the precipitate, concentration of the reactants, and the rate at which the reactants are mixed--- on coagulation of the colloid silver chloride can be explained by a property known as relative supersaturation. The equation for relative supersaturation is
Relative supersaturation = where Q represents the concentration of the solute at any point in time and S represents the solubility of the solid at equilibrium, a constant value. A high relative supersaturation produces colloidal solids while a low relative supersaturation produces crystalline solids. In other words, relative supersaturation is inversely related to particle size. Because of this relation, the concentration of the solute (Q) at any point of time must be kept low as low as possible for a crystalline solid to form. The concentration of the solute can be managed through temperature regulation and the other factors mentioned previously.
Certain parameters during the precipitation, washing, and drying of the solid were controlled in order to maximize the amount of chloride collected and to minimize certain systematic errors. These parameters included the addition of nitric acid as an electrolyte to the original solution and to the wash solution, the elevation of the solution temperature as the silver nitrate was added, slow addition of the precipitating reagent (silver nitrate) with constant stirring, the addition of a slight excess of silver nitrate, and minimization of the precipitate to light exposure.
Nitric acid was added to the original solution in order to keep the solution slightly acidic and prevent weak acids from forming undesirable solids. The purpose of adding nitric acid to the wash solution was to prevent the diminishing of the counter-ion layer formed by the nitrate ions surrounding the adsorbed silver ions. If the solid had been washed with pure water, then the layer would have been disturbed causing the particle to shrink in size. The temperature of the solution was elevated in order to increase the kinetic energy of the particles. A higher kinetic energy allows for the particles to overcome the repulsion force of the electric double layer (the adsorbed layer plus the counter-ion layer) and form larger particles. Stirring the solution as the sliver nitrate was added also contributed to the increase of kinetic energy while the slow addition of the silver nitrate helped control the concentration of the solute (Q). Because the concentration of the solute in the solution was kept low, the relative supersaturation remained low which helped form larger particles. The excess of silver nitrate also prevented the value of Q from becoming too large by reducing the solubility of the silver chloride. The excess silver ions formed the primary adsorbed layer surrounding the solid while the nitrate ions from the nitric acid and the silver nitrate form the counter-ion layer. The ions provide the solution with a high electrolyte concentration that works to decrease the volume of the counter-ion layer and enlarge the particles. Once the precipitate was formed, it was exposed to as little light as possible to prevent the photodecomposition of silver chloride to silver and chlorine.
The collected solid could be dried and weighed, and the total mass of the sample could be determined by subtracting initial mass of the crucible from the final mass of the crucible with the sample. From this total mass, the mass of chloride can be determined using the mass of the measured silver chloride by the following stoichiometric equation:

The percentage of chloride is then calculated by taking the mass of the chloride and dividing by the total mass of the sample and multiplying by one hundred as seen below: percent of chloride = x 100 % As the percentage for the two measured samples is determined, the standard deviation can also be calculated using the formula listed: standard deviation =
Once the standard deviation is known the relative standard deviation can be calculated by the relative standard deviation formula relative standard deviation in ppt = x 1000
A low relative standard deviation would indicate good precision in measurements.
Materials and Methods
Materials
Porosity sintered glass crucible (2)
Crucible holder with funnel
Vacuum tube
Vacuum
Oven (at 110 ºC)
Stirring rod with rubber policeman (2)
Aluminum foil
Tongs
Hot mitts
Desiccator
Analytical balance
Hot plate
5 mL test tube
500 mL filtering flask (1)
400 mL beaker (2)
600 mL beaker
250 mL beaker (2)
Deionized water
Concentrated HNO­3
6 M NH3
6 M HNO3
0.2 M AgNO3
HCl

MSDS
Chemical Data
Physical Data
Hazards
PPE
Nitric acid
Formula: HNO3
MW: 63.01 g/mol
Liquid with strong odor.
Colorless to light yellow.
Density: 1.408 g/mL
Boiling point: 121 ºC
Melting point: 41.6 ºC
Soluble in water and diethyl ether.

Corrosive, irritant, and permeator.
Hazardous in case of skin or eye contact and in case of ingestion.
Lung sensitizer. May cause burns.
Health Rating - 3
Flammability Rating - 0
Reactivity Rating - 0
Contact Rating – Extreme
Wear gloves, closed-toes shoes, and lab coat to prevent skin exposure.
Wear splash goggles to prevent eye exposure.
Use under hood to prevent inhalation.
Ammonia
Formula: NH3
MW: 17.00 g/mol
Gas or liquid with strong odor.
Colorless.
Density: 0.90 g/mL
Boiling point: -33.4 ºC
Melting point: -77.7 ºC
Soluble in water.

Corrosive to upper respiratory tract.
Causes burns to any area of contact.
Avoid eye and skin contact. Avoid ingestion (extremely corrosive).
Health Rating - 3
Flammability Rating - 1
Reactivity Rating - 2
Contact Rating – Extreme
Wear gloves, closed-toes shoes, and lab coat to prevent skin exposure.
Wear splash goggles to prevent eye exposure.
Use under hood to prevent inhalation.
Silver nitrate
Formula: AgNO3
MW: 169.87 g/mol
Liquid with no odor.
White or colorless.
Boiling point: 440 ºC
Melting point: 212 ºC
Soluble in cold water, hot water, and diethyl ether.
Slightly soluble in acetone.
Corrosive and oxidizer. May cause cyanosis and liver and kidney damage. Causes eye and skin burns. Severe irritant to respiratory tract (possible burns).
Health Rating- 2
Flammability Rating – 0
Reactivity Rating- 0
Contact Rating – Extreme
Wear gloves, closed-toes shoes, and lab coat to prevent skin exposure.
Wear splash goggles to prevent eye exposure.
Use general exhaust to prevent inhalation.
Silver chloride
Formula: AgCl
MW: 143.32 g/mol
Solid with no odor.
White or silver color.
Density: 5.56 g/mL
Boiling point: 1550 ºC
Melting point: 455 ºC
Slightly soluble in cold water.
Soluble in AgNO3 and MgNO3.
Eye and skin irritant.
Hazardous if inhaled or ingested.
Health Rating- 2
Flammability Rating – 0
Reactivity Rating- 0
Contact Rating – Extreme
Wear gloves, closed-toes shoes, and lab coat to prevent skin exposure.
Wear splash goggles to prevent eye exposure.
Use general exhaust to prevent inhalation.
Hydrochloric acid
Formula: HCl
MW: 36.46 g/mol
Liquid with pungent, irritating odor.
Colorless or light yellow.
Density: 1.18 g/mL
Boiling point: 53 ºC
Melting point: -74 ºC
Soluble in cold water, hot water, and diethyl ether.
Corrosive. Causes severe burns to all body tissues. Fatal if swallowed or inhaled. Can cause lung damage. Avoid eye and skin contact due to burn possibilities and irritation.
Health Rating- 3
Flammability Rating – 0
Reactivity Rating- 2
Contact Rating – Extreme
Wear gloves, closed-toes shoes, and lab coat to prevent skin exposure.
Wear splash goggles to prevent eye exposure.
Use under hood to prevent inhalation.

Method
Week One
1. Place two porosity sintered glass crucibles in a 600 mL beaker then put under the hood. Clean both using about 5 mL of concentrated HNO3. Leave for 5 minutes.
2. Set up a vacuum filtration using the 500 mL filtering flask crucible holder with funnel, vacuum, and vacuum tube.
3. Pour HNO3 into the waste container under the hood then rinse with deionized water at least 3 times.
4. Take the crucibles back to the hood and add about 5 mL of 6 M NH3. Leave for 5 minutes.
5. Pour out the 6 M NH3 into the waste container then wash crucibles with deionized for 6-8 times.
6. Label then place the crucibles in separate 250 mL beakers. Dry the crucibles for 45 minutes in the oven set at 110 ºC.
7. Let the crucibles cool to room temperature then weigh on the analytical balance. Record the measurement then place in the oven for an additional 20 minutes.
8. Cool to room temperature and reweigh. If crucibles’ mass are within ±0.0004 g of each other, then the crucibles are dry. If not, repeat drying, cooling, and reweighing until the measurements lie within the desired range. Record the measurements.
9. Weigh 2 clean 400 mL beakers and record. Then add approximately 0.15 g of unknown sample to each beaker with each sample only varying as much as 0.01 g from 0.15 g. Record the new weight. (Weighing the sample is done by difference.)
10. Label the beakers as beaker 1 and beaker 2.
11. Add approximately 100 mL of deionized water and 2-3 mL of 6 M HNO3 to each sample.
12. Heat each sample until warm (the hot plate temperature should be about 230 ºC).
13. Calculate the volume of 0.2 M AgNO3 needed to add to the solutions by using NaCl as the stoichiometric converter.
14. Once the correct volume is calculated and poured into the solutions, add an additional 3-5 mL of AgNO­3. Note: make sure to stir continuously when adding the AgNO3.
15. Continue stirring and heating until a solid is formed and the solution appears clear above the solid.
16. Once clear, place about 2-3 drops of AgNO3 into each beaker. If more AgCl coagulates, add an additional 3 mL of AgNO3 then let the solution clear and test again with 2-3 drops of AgNO3.
17. If no precipitate forms, let the solutions cool then completely cover each beaker with aluminum foil and store for 1 week.
Week Two
1. Set up vacuum filtration once again using the 500 mL filtering flask, crucible holder with funnel, vacuum tube, and vacuum.
2. Place crucible #1 in the funnel then decant the supernatant liquid from beaker #1 using the rubber policeman to pour. Do not let the liquid fill over 2/3 of the crucible.
3. Once the liquid is removed, pour as much precipitate as possible into the crucible.
4. Remove any remaining precipitate from the beaker using a solution containing of 2-5 mL of HNO3 per liter of deionized water. Try to get as much precipitate as possible into each wash using the rubber policeman.
5. Test the washes by adding about 5 mL of wash solution to the crucible then collecting the washes into a 5 mL test tube.
6. Add about 10 drops of HCl to the solution in the test tube. If the solution remains clear, the precipitate in the crucible is free of Ag+ ions.
7. Place the crucible with the precipitate in a 250 mL beaker then dry in the oven set at 110 ºC for 1 hour. Let the crucible cool to room temperature in desiccator then weigh. Record this measurement.
8. Place the crucible in the oven once again for 20 minutes, cool in desiccator, and weigh. Repeat heating, cooling, and weighing until the crucible masses lay within ±0.0004 g of each other. Record the mass measurements.
9. Repeat steps 2-8 for crucible #2 and beaker #2.
10. Remove the precipitates from the crucibles using a rubber police man to gently dislodge the solid. Pour into the solid waste.
11. Add 6 M NH3 to each crucible and let stand for 5 minutes. Pour the NH3 into the liquid aqueous waste then wash each crucible using vacuum filtration and deionized water unto the crucibles are free of NH3.
12. Calculate the percent Cl- in the samples using stoichiometry and calculate the standard deviation.
Data and Results
Table 1- Beakers with Unknown Sample
Beaker #
Initial Mass (g)
Mass with Unknown Sample (g)
Unknown Sample Mass (g)
Volume of AgNO3
1
17.3480
17.1858
0.1619
13.87 mL
2
17.1856
17.0366
0.1490
12.74 mL

Table 2- crucibles with silver chloride sample
Crucible #
Initial Mass (g)
Crucible Mass with AgCl Mass (g)
AgCl Mass (g)
1
11.7438
12.1098
0.3661

11.7434*
12.1094

-
12.1095*

2
11.0299
11.3667
0.3367

11.0293
11.3663*

11.0296*
-

*Data used to calculate the mass of AgCl
Table 3- Mass and Percent of chloride
Sample #
AgCl Mass (g)
Mass of Chloride (g)
Percent of Chloride
Relative Standard Deviation (ppt)
1
0.3659
0.09056
24.74 %
6 x 10-6
2
0.3371
0.08329
24.74 %

Calculations
See attached paper under the heading “Calculations.”
Discussion
One source of error that may have occurred is surface adsorption, a type of coprecipitation. Since surface adsorption causes the silver nitrate to coprecipitate with the silver chloride, the measured mass of the silver chloride would be higher than the actual mass thereby causing the percentage of chloride to be higher than the actual. This type of error is systematic and can be fixed by avoiding large excess of AgNO3.
Another possible error source of error that may have occurred involves incomplete drying to the sample due to the temperature that the oven was set at. If the temperature of the oven fell below the desired 110 ºC, then the nitric acid trapped in the solid may have not volatized completely causing a positive error in measurements. This type of error is systematic and can be minimized by constant monitoring of the oven temperature.
A third error that may have occurred involves vibrations while weighing on the analytical balance. This error would have been random and would have resulted in data that deviates in either the positive or the negative direction. Such error can be minimized by zeroing the balance each time before weighing the beakers and crucibles.
One other error that may have occurred is peptization, a systematic error. Peptization is the converting of the coagulated solid back to its colloidal form during washing which would cause a negative error since some of the solid would not be captured in the filter. This error is minimized by using nitric acid in the wash solution to prevent the disruption of the counter-ion layer.
One final error that possible occurred involved the photodecomposition of silver chloride to elemental silver and chlorine. This error, because it is systematic, would cause the measurement to deviate in the negative direction since the chlorine is a gas and escapes the solid. Therefore, the percentage of chloride calculated would be less. A way to minimize this error is to keep the solution covered as much as possible before weighing. conclusion This experiment yielded sample measurements that contained the same percentage of chloride, indicating the procedure was carried out with a fair amount of precision. Also, the relative standard deviation calculated yielded very small number further proving that the measurement of chloride was precise. The percentage of chloride in the unknown sample was found to be 24.73695223 (±0.00000006) %.
References
Skoog, Douglas A., Donald M. West, and F. James. Holler. Fundamentals of Analytical Chemistry, 8th ed. Fort Worth: Saunders College Pub., 1996. Print.

References: Skoog, Douglas A., Donald M. West, and F. James. Holler. Fundamentals of Analytical Chemistry, 8th ed. Fort Worth: Saunders College Pub., 1996. Print.