GRAVIMETRIC DETERMINATION OF BARIUM SULFATE
Austine John O. Lumasag
Student
Aileen May G. Ang
Instructor
CHY 56A
Monday & Friday, 10:00 AM – 1:00 PM
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I.
Introduction
Gravimetric analysis is one of the most precise and fundamental
quantitative techniques employed in analytical chemistry for determining the
amount of a specific ion or compound within a sample (Skoog et al., 2013). This
method relies on converting the analyte into an insoluble precipitate of known
composition, allowing the mass of the substance to be measured accurately
and related stoichiometrically to the amount of analyte originally present (Skoog
et al., 2013). In the gravimetric determination of sulfate, barium chloride is
introduced to a sample solution containing sulfate ions, leading to the formation
of barium sulfate (BaSO4), which is highly insoluble under experimental
conditions (Skoog et al., 2013). The precipitate is carefully filtered, washed,
heated or dried, and subsequently weighed to calculate the sulfate content of
the sample.
This experiment aims to (1) familiarize students with the principles and
practices of gravimetric analysis by quantitatively determining sulfate ions as
barium sulfate. (2) Use the stoichiometric relationship of the reaction in order to
calculate the percentage of mass of sulfate in the unknown sulfate salt. The low
solubility product of barium sulfate ensures minimal losses during filtration and
washing, making this method a cornerstone for both research and industrial
applications of analytical chemistry (Skoog et al., 2013).
II.
Methodology
This section enumerates all the glassware, reagents, and apparatuses
that are used in the experiment, together with a its procedure and mathematical
equation.
Chemicals used
The chemicals used in the experiment are as follows: 0.1 M Barium
chloride (BaCl2), 6 M Hydrochloric acid (HCl), Sodium sulfate (Na2SO4), and
distilled water.
Apparatus and Glassware
All the glassware and apparatus that are used in the experiment includes:
3 crucibles with lid, wash bottle, soap, pencil, oven, crucible tong, desiccator,
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analytical balance, 250mL beaker, dropper, stirring rod, spatula, graduated
cylinder, filter paper, 250mL Erlenmeyer flask and funnel.
Procedure
Three crucibles with lids were thoroughly cleaned, rinsed with distilled
water, labeled, and dried at 110 °C for one hour to eliminate moisture. After
cooling in a desiccator, each was weighed on an analytical balance to 0.0001 g
precision. This heating and weighing was repeated until consecutive masses
varied by less than ±0.0003 g, confirming constant mass. Handling was
performed with tongs, and the same balance was used throughout to avoid
contamination or measurement inconsistencies.
A clean, dry 250 mL beaker was weighed before adding 0.30–0.35 g
sodium sulfate (Na₂SO₄), then reweighed. Fifty milliliters of distilled water and
20 drops of 6 M hydrochloric acid were added to dissolve the sample with
stirring. Subsequently, 25 mL of 0.1 M barium chloride (BaCl₂) was slowly
added to the near-boiling solution over three minutes, inducing precipitation of
barium sulfate (BaSO₄). The mixture was left undisturbed for 20 minutes to
allow settling.
Ashless filter paper was prepared in a funnel above a 500 mL Erlenmeyer
flask. The suspension was carefully filtered through the paper using a stirring
rod, ensuring the liquid level did not exceed three-quarters of the filter. The
beaker and rod were rinsed with distilled water to recover all precipitate.
The precipitate-laden filter paper was transferred to a crucible, dried at
110 °C for one hour, and ashed in a muffle furnace with temperature increments
of 50 °C every 30 minutes until reaching 500 °C, where it was maintained for
three hours to incinerate the paper completely. After cooling, the crucible, lid,
and residue were weighed with 0.0001 g accuracy.
Mathematical Equation
The gravimetric determination of barium relies on the stoichiometric
relationship between barium ions (Ba²⁺) and sulfate ions (SO₄²⁻) to form barium
sulfate (BaSO₄), which precipitates out of solution. The precipitation reaction
goes to completion, allowing the mass of the BaSO₄ precipitate to be measured
accurately. The equations below explain the essential steps in the calculation.
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𝑚𝑜𝑙𝑒𝑠 =
𝑀𝑎𝑠𝑠
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠
where,
Mass – weight of the reagent in the analytical balance
Molar Mass – standard mass of the reagent in one mole.
The number of moles of the reagents used is calculated using the formula
above, which determines the amount of substance by dividing its mass by the
standard molar mass.
𝑚𝑎𝑠𝑠𝐵𝑎𝐶𝑙2 = 𝑉𝐵𝑎𝐶𝑙2 ∗
𝑚𝑜𝑙𝐵𝑎𝐶𝑙2
𝑚𝐵𝑎𝐶𝑙2
𝐿
∗
∗
𝑚𝐿 𝑉(𝑖𝑛 𝐿𝑖𝑡𝑒𝑟𝑠) 𝑚𝑜𝑙𝐵𝑎𝐶𝑙2
where,
𝑉𝐵𝑎𝐶𝑙2 – the amount of reagent that is used in mL
𝑚𝑜𝑙𝐵𝑎𝐶𝑙2
𝑉(𝑖𝑛 𝐿𝑖𝑡𝑒𝑟𝑠)
𝑚𝐵𝑎𝐶𝑙2
𝑚𝑜𝑙𝐵𝑎𝐶𝑙2
- Concentration of BaCl2 (Molarity)
- Molar Mass of BaCl2 (grams/mole)
The mass of the BaCl2 can be calculated through stoichiometric
procedure. It employs the volume, concentration, and molar mass of the
reagent.
𝑀𝑎𝑠𝑠𝐵𝑎𝑆𝑂4 = 𝑀𝑎𝑠𝑠𝐶𝑟𝑢𝑐+𝐵𝑎𝑆𝑂4 − 𝑀𝑎𝑠𝑠𝑐𝑟𝑢𝑐
where,
𝑀𝑎𝑠𝑠𝐶𝑟𝑢𝑐+𝐵𝑎𝑆𝑂4 – mass of the crucible with BaSO4
𝑀𝑎𝑠𝑠𝑐𝑟𝑢𝑐 – mass of the empty crucible
The actual yield of BaSO4 can be calculated by subtracting the mass of
the crucible from the mass of the weighed BaSO 4 contained in a crucible.
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑌𝑖𝑒𝑙𝑑𝐵𝑎𝑆04 = 𝑚𝑜𝑙𝑁𝑎𝑆𝑂4 ∗
where,
𝑚𝑜𝑙𝐵𝑎𝑆𝑂4
𝑚𝐵𝑎𝑆𝑂4
∗
𝑚𝑜𝑙𝑁𝑎2𝑆𝑂4 𝑚𝑜𝑙𝐵𝑎𝑆𝑂4
4
𝑚𝑜𝑙𝐵𝑎𝑆𝑂4
𝑚𝑜𝑙𝑁𝑎2𝑆𝑂4
𝑚𝐵𝑎𝑆𝑂4
𝑚𝑜𝑙𝐵𝑎𝑆𝑂4
- Gravimetric factor between BaSO4 and NaSO4
- Molar mass of BaSO4
The theoretical yield of BaSO₄ can be calculated by determining the moles
of BaSO₄ produced from the moles of sulfate ions (SO₄²⁻) in sodium sulfate
(Na₂SO₄), since the amount of BaSO₄ formed depends on the limiting reagent,
which is the sulfate ion from Na₂SO₄.
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑦𝑖𝑒𝑙𝑑, % =
𝑚𝑎𝑠𝑠𝑎𝑐𝑡𝑢𝑎𝑙
∗ 100%
𝑚𝑎𝑠𝑠𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙
where,
massactual – the actual yield or mass of BaSO4
masstheoretical – the theoretical yield of BaSO4
The percent yield can be calculated by dividing the actual yield from the
experiment
to
the
expected
amount
of
product
that
is calculated
stoichiometrically. It indicates the efficiency of a reaction by comparing the
actual amount of product obtained to the maximum amount predicted by
stoichiometry (the theoretical yield).
𝑆𝑂42−%𝑖𝑛 𝑡ℎ𝑒 𝑠𝑎𝑚𝑝𝑙𝑒 =
𝑚𝑎𝑠𝑠𝑆𝑂42−
𝑚𝑎𝑠𝑠𝑠𝑎𝑚𝑝𝑙𝑒
∗ 100%
where,
massSO2−
- mass of SO42- in either theoretical or actual
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masssample - mass of the sample, BaSO4
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Calculating the amount of sulfate obtained in the experiment can be done
by dividing the mass of sulfate (SO₄²⁻) determined through stoichiometric
calculation by the actual mass of BaSO₄ collected.
III.
Results and Discussion
Data collected from the experiment are compiled and summarized in
Tables 1 and 2.
Table 1. Mass of the BaSO4
Cycle 1
Mass of empty crucible, g
Cycle 2
41.2796
Cycle 3
Cycle 4
Cycle 5
-
-
-
41.2798
Constant mass of empty
41.2797
crucible, g
Mass of crucible and BaSO4
41.7706
Table 2. Percentage of Sulfate (SO42-) in BaSO4
Trial 1
Mass of Na2SO4, g
Trial 2
Trial 3
Trial 4
Trial 5
0.3166
0.3223
0.3041
0.3023
0.3205
Mole of SO42-, mol
2.229x10-3
2.269x10-3
3.166x10-3
0.0021
2.326x10-3
Mass of BaCl2, g
0.5206
0.5206
0.5206
0.5206
0.5206
2.0x10-3
0.0025
2.0x10-4
0.0025
0.0025
0.2061
0.5203
0.4412
0.4909
0.5307
0.5202
0.5296
0.4997
0.4966
0.5428
97.29%
98.25%
98.90%
98.85%
98.69%
0.2803
0.2142
2.10x10-3
0.2020
0.2205
65.79%
66.44%
66.33%
66.82%
66.72%
Mole Ba2+, mol
Actual Yield (BaSO4),
g
Theoretical Yield
(BaSO4), g
Percent Yield, %
Mass SO42- in the
sample (Exp), %
Percent SO42- in the
sample (Exp), %
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67.63%
Percent SO42- in the
67.83%
67.63%
67.61%
67.62%
sample (Theoretical),
%
Mean Percent SO42- in the sample
66.42%
(Exp), %
Standard Deviation SO42- in the
0.40%
sample, %
The experiment was carried out to gravimetrically determine the mass and
percentage of sulfate in barium sulfate (BaSO₄) via a precipitation reaction. The
reaction involved the formation of an insoluble salt, as demonstrated by the
balanced chemical equation:
𝐵𝑎𝐶𝑙2 (𝑎𝑞) + 𝑁𝑎2 𝑆𝑂4 (𝑎𝑞) → 𝐵𝑎𝑆𝑂4 (𝑠) + 2𝑁𝑎𝐶𝑙(𝑎𝑞)
Barium chloride and sodium sulfate solutions were mixed, allowing barium
ions (Ba²⁺) to combine with sulfate ions (SO₄²⁻) to precipitate barium sulfate, an
insoluble solid. Due to its low solubility in water, BaSO₄ precipitated nearly
completely, facilitating accurate gravimetric analysis (Shi et al., 2023). The
precipitate was filtered, dried, and weighed, and its mass was used to calculate
the sulfate content and yield.
Consistent and precise mass measurements of the BaSO₄ precipitate
were obtained throughout the trials, as shown in Table 1, with crucible weights
stabilized within ±0.0003 g. The actual BaSO₄ mass ranged from 0.2061 g to
0.5307 g (Table 2), compared to theoretical masses between 0.4966 g and
0.5428 g. Trials 2 to 5 yielded consistent results, whereas trial 1 displayed a
slightly elevated yield of 97.29%, suggesting minor procedural deviations
affecting precision. According to Mao (2024), gravimetric analysis uncertainties
can arise from sample preparation variability, transfer loss, or weighing errors.
Komal et al. (2022) also identified factors such as reagent addition rate,
inadequate precipitate washing, and environmental factors like temperature
and humidity as influences on measurement accuracy. Despite the deviation in
trial 1, the low standard deviation and reproducible results in subsequent trials
indicate overall accuracy and reliability of the experiment.
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Experimental sulfate percentages ranged from 65.79% to 66.82%, closely
aligning with the theoretical values of 67.61% to 67.83%, with deviations within
2%. The mean experimental sulfate content was 66.42%. Minor discrepancies
are attributed to procedural errors or potential contamination during filtration or
drying. The low standard deviation of 0.40% further attests to the precision and
soundness of the experimental method.
IV.
Summary and Conclusion
The experiment successfully determined the percentage of sulfate (SO₄²⁻)
present in barium sulfate (BaSO₄) through the application of gravimetric
analysis. The mean experimental sulfate content was found to be 66.42% with
a standard deviation of ±0.40%, closely aligning with the theoretical sulfate
value of 67.67%. This close agreement indicates that the precipitation, filtration,
drying, and weighing procedures were performed with both accuracy and
precision. The results are reflective of the fundamental principle of gravimetric
analysis, which asserts that the mass of a pure, stable precipitate can be
reliably used to quantify the amount of an analyte in a sample matrix (Skoog et
al., 2017).
Minor discrepancies observed between the experimental and theoretical
values are likely attributable to several factors, including incomplete drying of
the BaSO₄ precipitate, slight losses during the filtration or transfer process, and
potential variations in ambient laboratory conditions such as humidity and
temperature that may affect mass measurements. These sources of systematic
and random error are commonly encountered in gravimetric procedures and
underscore the importance of rigorous technique and controlled experimental
conditions (Skoog et al., 2017).
Overall, the experiment highlights the reliability and robustness of
gravimetric analysis for quantitative determinations in analytical chemistry when
appropriate methodological controls are employed. The relatively low standard
deviation further confirms that the experiment was reproducible and precise
across multiple trials. Emphasizing careful sample handling, thorough washing,
and complete drying of precipitates is essential for minimizing errors and
enhancing the reliability of results in future analyses.
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References
Komal, et al. (2022). Factors affecting accuracy in gravimetric measurements. Journal
of Analytical Chemistry
Mao, [Initials]. (2024). Retrieved from general gravimetric analysis uncertainty
discussions
Shi, et al. (2023). The deposition kinetics of barium sulphate scale. Frontiers in
Materials, 10, 1198176. https://doi.org/10.3389/fmats.2023.1198176
Skoog, D. A., Holler, F. J., & Crouch, S. R. (2017). Principles of Instrumental
Analysis (7th ed.). Cengage Learning.
Skoog, D. A., West, D. M., Holler, F. J., & Crouch, S. R. (2013). Fundamentals of
Analytical Chemistry (9th ed.). Brooks Cole.