GRAVIMETRIC DETERMINATION OF SULFATE
Purpose
You will be given a solid powder, which has been dried to constant mass. The sample is water soluble.
The purpose of this experiment is to determine the concentration of sulfate in your solid sample using
gravimetric analysis.
Background
Gravimetric Analysis
The goal of most quantitative chemical analysis measurements is to estimate the relative abundance of
an analyte in a chemical sample. For solid and liquid samples, a very common expression of analyte
content is based on the mass fraction. This fraction is commonly expressed as a percentage (w/w %),
as parts per million (ppm) or as parts per billion (ppb), depending on the concentration level of the
analyte. Classical chemical analysis methods are excellent for the determination of analyte
concentrations in the range of 1−100 w/w %. In order to estimate the analyte mass fraction of any solid
sample, we typically need two measurements, one to estimate the sample mass and one to estimate the
analyte mass.
In gravimetric analysis, measurements of mass are used to determine the analyte concentration in a
sample. In the most common form of gravimetric analysis, the analyte is separated from the rest of the
sample in the form of a solid precipitate. In this experiment, the analysis of sulfate is performed using
barium chloride as a reagent to precipitate sulfate from the dissolved sample. The mass of sulfate in the
sample can be calculated by simple stoichiometry from the mass of the weighing form of the sulfate
precipitate.
In order to ensure maximum filterability and minimum contamination, the sulfate is precipitated under
conditions that maximize the particle size. These conditions include: (a) avoiding excessively high
concentrations of precipitating reagent; (b) slow addition of reagent to a hot solution, with vigorous
stirring; (c) avoiding a great excess of reagent. The digestion of the precipitant after the reaction
minimizes errors due to coprecipitation.
For accurate measurements, we must be sure that the composition of the weighing form is known
exactly. You must be very careful in your mass measurements throughout this experiment − for
example, do not handle your weighing bottle with your bare hands, to avoid leaving fingerprints that
would increase the mass of the bottle. Of course, the recovery of the sulfate precipitate must be
quantitative, and the final weighing form must be pure.
GRAVIMETRIC ANALYSIS: Procedure
Chemicals and materials:
6 M HCl
0.05 M BaCl2 (note: you will be provided with 0.5 M BaCl2)
Whatman #42 ashless filter paper
Place the sample in a beaker and cover it with a watch glass to protect your sample from dust in the air.
Dry the sample for at least one hour at 110 °C. Allow the sample to cool in a desiccator for at least 30
minutes before weighing. Weigh three samples by difference (your instructor will show you how to do
this) into 400 mL beakers. The samples should be approximately 0.4 to 0.5 g in size, and their mass
must be known accurately (to the nearest 0.1 mg). Be sure to label the beakers; you do not want to mix
up your samples.
While the sample is drying, clean three porcelain crucibles and covers and identify each numerically
according to the instructor’s direction. Heat the crucibles without covers to constant weight at the
highest temperature of the burner. To heat the crucibles to constant weight, place the crucibles in
triangles, bring the tip of the flame close to the bottom of the crucible, and heat for 15 minutes (see
Harris Fig 29-1). Allow the crucibles to cool in air until only moderate radiated heat is felt when the
back of the hand is brought near the crucible (about one minute or so), then transfer the crucibles to the
desiccator. After the crucibles reach room temperature in the desiccator (about 30 minutes), weigh the
crucibles accurately to the nearest 0.1 mg. Repeat the heating, cooling and weighing until successive
weights are within ±0.2 mg.
Dissolve each sample in approximately 200 mL of deionized water and add 4 mL of 6 M HCl to each.
Heat nearly to boiling.
You will need to add an excess of precipitating reagent (BaCl2) to the samples in order to precipitate
the sulfate. To calculate how much reagent to use, use the following calculation: assume that the
samples are pure sodium sulfate (they aren’t) and calculate the volume of 0.05 M BaCl2 needed to
precipitate the largest of your samples. Add to the calculated amount a 10% excess of BaCl2 solution.
This is the volume of reagent that will be added to each of your samples. You will be provided with 0.5
M BaCl2; you must make your precipitating reagent from this solution.
Heat three portions of reagent BaCl2 solution nearly to boiling. Slowly pour the hot BaCl2 solution into
the hot sample solution with vigorous stirring during the addition. Use a separate stirring rod for each
sample and leave it in the solution throughout the determination. Allow the precipitate to settle and test
for complete precipitation by adding a few drops of BaCl2 solution and observing the result.
After the precipitation is complete, cover each beaker with a watch glass and digest just below the
boiling point for at least one hour. Longer periods are desirable, if time permits. If digestion is
completed near the end of the laboratory period, so that there is insufficient time for filtration, set the
samples in the cupboard where they will not be disturbed until the next laboratory period. You will
want to heat the solution nearly to boiling before filtering at the beginning of the next laboratory
period.
Carefully decant the supernatant liquid through Whatman #42 ashless filter paper (obtained from your
instructor). Wash the precipitate three times by stirring with warm distilled water and decanting the
water after the precipitate has settled. Empty the filtrate from the beaker below the funnel unless the
filtrate is cloudy. If the filtrate is cloudy, refilter the filtrate through another piece of ashless filter
paper.
Using a stirring rod and wash bottle, transfer the filtrate to the filter paper. Police the walls of the
beaker to remove the last traces of the precipitate. Wash the precipitate in the filter three times with 3-5
mL portions of warm distilled water each time.
Fold the filter paper and place in the previously weighed porcelain crucible (see Harris fig 29-2). Place
the crucible in a triangle and heat very gently to dry the paper. Next char the paper with gentle heat.
The heat should be regulated so that the paper does not burst into flame. A clean crucible cover should
be located nearby, ready for use if necessary. If the paper does burst into flame, the burner must be
quickly removed and the cover used to snuff out the flame.
After the paper has burned away, you will want to burn away deposited carbon particles by applying
increased heat to the crucible. Do not point your flame directly into the crucible − you will probably
blow your precipitate across the room. You should always position your flame so that the precipitate
has access to atmospheric oxygen; otherwise you might reduce your precipitate from BaSO4 to BaS.
When the carbon is gone from the crucible, place the crucible upright in the triangle and ignite for at
least fifteen minutes at the highest temperature of the burner. You should note that the hottest portion
of the flame is just above the top of the inner blue cone. After ignition, allow the crucible to cool until
the red glow disappears and then place it in the desiccator until it cools to room temperature. Weigh
each crucible to the nearest 0.1 mg. Finally, bring each crucible to constant weight, as directed
previously.
GRAVIMETRIC ANALYSIS: DATA SHEET
Directions: fill in this sheet and turn it in with your abstract. Also turn in a sheet showing your
calculations (i.e., how you transform this raw data into your final results).
Unknown #:
Name:
sample A
sample B
sample C
before:
g
g
g
after:
g
g
g
crucible mass
(heated to constant mass)
g
g
g
mass: crucible + ppt
(heated to constant mass)
g
g
g
dry sample mass
(mass by
difference)
estimated analyte conc
(as a confidence interval)