Materials Included In Kit

Additional Materials Required

Prelab Preparation

Prepare 100 mL each of six different “unknown” hard water samples for students to analyze in the guided-inquiry activity. See the following calculations for the amount of 2 M CaCl₂ stock solution that should be diluted to prepare 100 mL of each sample.

1. Using the 2 M calcium chloride stock solution provided in the kit, measure out V1 using a graduated cylinder or volumetric pipet.
2. Add to a 100-mL volumetric flask or graduated cylinder and fill to equal 100 mL with deionized or distilled water.
3. Mix well and store each solution in a labeled beaker or bottle.

Safety Precautions

Sodium carbonate is irritating to body tissues. Calcium chloride is moderately toxic by ingestion and the anhydrous generates a great deal of heat when dissolved in water. Avoid contact of all chemicals with eyes and skin. Antacid tablets used in the lab are considered laboratory chemicals and may not be removed from the lab. Do not taste or ingest any materials in the chemistry lab. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The leftover solid calcium chloride and sodium carbonate may be packaged for landfill disposal according to Flinn Suggested Disposal Method #26a. The leftover calcium chloride or sodium carbonate solutions may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Quantitative transfer techniques are essential for accurate results in this investigation. Students should practice pouring liquids, rinsing beakers or flasks and filtering mixtures. The type of glassware used will also affect the precision of the results.
• If a vacuum filtration apparatus is not available, students may filter the precipitates by gravity filtration. For smaller amounts of precipitate, such as 0.10 M and 0.05 M samples of CaCl₂ (samples 4 and 5), it is best to filter by gravity in order to prevent loss of product.
• If some of the solid CaCO₃ precipitate goes through the filter paper, students should stop the filtration, pour the filtrate back into the beaker, and use two filter papers to filter a second time.
• Quantitative filter paper is needed for this lab. We recommend Extra-Fine Quantitative Filter Paper (11 cm) from Flinn Scientific, Catalog No. AP7628, if it fits the funnel being used.
• Experimental errors can result if the students do not dissolve the reactants completely in distilled water before mixing. Some of the solid filtered out may be the undissolved reactant(s). Tap water should not be used because the impurities in it, such as calcium and magnesium ions in hard water, can also precipitate out and interfere with the desired results. The filter paper and precipitate must be completely dried and any water of hydration must be removed by the drying process. Usually, the most accurate method to ensure complete dehydration is by heating until a constant mass is achieved.
• Students may calculate mass and mass percent of Ca²⁺ in the hard water samples collected.
• Students may choose to bring their own samples of over-the-counter calcium tablets for analysis. Results may be compared with the manufacturer’s label.
• Students can compare results obtained via EDTA titration versus gravimetric analysis for hard water determination. In the Water Softening Guided-Inquiry Kit, available from Flinn Scientific, Catalog No. AP7352, students treat water samples of unknown hardness with ion-exchange resin and analyze by EDTA titration. The guided-inquiry format allows students to identify the minerals that are removed through ion exchange and understand the practical benefits of water softening.
• If there is an aquatics center on campus have the chemistry students perform the water hardness test on pool water. The recommended water hardness level for swimming pools and spas is 200–300 ppm. Compare the results obtained in the lab with those obtained poolside using typical “dip and read” test papers or methods. Also, introduce the routine tests used to monitor pool chemistry and water quality: chlorine, pH, alkalinity, total dissolved solids, cyanuric acid, copper and iron.
• The three principal ions that contribute to water hardness are calcium, magnesium and iron. While there are no national, state or local standards for calcium and magnesium in drinking water, the EPA recommends a limit of 0.3 mg/L for the concentration of iron in drinking water. This is a recommended value and is not legally enforceable.
• Set up a collaborative classroom research project. Using the Internet, look up typical surface water, groundwater, or municipal water hardness levels in major cities or geographic areas for every state. Plot the data on a large map of the United States using a color-coded system for the USGS water hardness categories (see the Background section). Analyze the resulting national map in terms of geographic features and properties. Although about 85% of the country has hard water, there are wide variations depending on geography. New England and the Pacific Northwest generally have soft water, while very hard water (≥1000 ppm) is common in many areas of the Southwest such as Texas, New Mexico, Arizona and Utah.

Further Extensions

Opportunities for Undergraduate Research
The amount of calcium in an antacid tablet may also be determined using gravimetric analysis. Note that antacid tablets contain binders and other inert ingredients or additives. Design a procedure to dissolve and separate the calcium carbonate from the binders in the tablet and analyze the amount of calcium.

Answers to Prelab Questions

1. Define the term gravimetric analysis. Describe the procedure used in this activity, and identify two other common examples of gravimetric analysis.

   Gravimetric analysis is used to determine the amount of a substance in a sample by measuring the mass of a solid product or precipitate obtained by reaction of the substance with a precipitating agent. In this investigation, calcium is determined by precipitation of calcium carbonate upon addition of sodium carbonate to a solution of the sample. Other common gravimetric analysis procedures include the determination of sulfate ions by precipitation of barium sulfate and the determination of silver ions by precipitation of silver chloride.

2. Write the balanced chemical equation for the reaction between calcium chloride and sodium carbonate.

   CaCl₂(aq) + Na₂CO₃(aq) → CaCO₃(s) + 2NaCl(aq)

3. Calculate the number of moles of each reactant in the Introductory Activity (see steps 1 and 2). Identify the limiting reactant in the reaction and determine the theoretical amount of CaCO₃ that should be produced.

   2.0 g CaCl₂110.98 g/mole = 0.0180 moles of CaCl₂
   2.5 g Na₂CO₃/105.9 g/mole = 0.0237 moles of Na₂CO₃
   The limiting reactant is calcium chloride. Sodium carbonate is present in approximately 31% excess relative to the amount needed based on the stoichiometric mole ratio.
   [(0.0237 moles – 0.0180 moles)/0.0180 moles] x 100% = 32%
   The molar mass of CaCO₃ is 100.1 g/mole, and the theoretical yield of calcium carbonate is 0.0180 moles x 100.09 g/mole = 1.80 g.

4. As noted in the Background section, hardness levels are calculated by assuming that all the “hard” metal ions come from dissolved calcium carbonate and are reported in mg CaCO₃/L. Calculate the equivalent water hardness in mg CaCO₃/L for a calcium chloride solution containing 0.1 M Ca²⁺ ions.

   The number of moles of calcium ions is obtained by multiplying the solution molarity by the volume. Since water hardness is expressed in terms of equivalent weight of calcium carbonate per liter of water, it is convenient to start with the moles per liter of calcium ion, or 0.1 moles.
   (0.1 moles CaCl₂/liter of solution) x 1 L = 0.1 moles Ca²⁺
   Multiply the moles of calcium by the molar mass of calcium carbonate to determine the equivalent mass of calcium carbonate. Convert the result to mg/L.
   (0.1 moles/L) x (100.09 g/mole) x (1000 mg/g) = 10,000 mg/L of CaCO₃

Sample Data

Introductory Activity

Mass of calcium carbonate obtained = 1.774 g
Percent yield = (1.774 g/1.80 g) x 100% = 98.6%

Gravimetric Analysis to Determine Water Hardness

^aPercent yield calculated based on the known molarity of CaCl₂ provided to the students.
^bWater hardness is calculated in units of mg/L of calcium carbonate based on the assumption that the water samples analyzed by the students actually represented a 100-fold concentration of the original water sample. Thus, for sample 5, which gave 0.072 g of CaCO₃ from a 20-mL solution, the calculated water hardness is:

where the factor 1/100 accounts for the fact that the sample analyzed was 100 times more concentrated than the original water sample.

Conclusion
The gravimetric analysis procedure is not accurate for “soft” water samples containing < 100 mg/L.

Sample Procedure and Data for Analyzing Antacid Tablets
Antacid tablets are provided in this inquiry kit as an optional opportunity for undergraduate research. Antacid tablets should be ground using a mortar and pestle. Collect the pulverized solid and measure the mass. In a 100-mL beaker, dissolve the solid with 20 mL of 3 M hydrochloric acid and 20 mL of deionized or distilled water. Stir the mixture for 5–6 minutes on a stir plate or using a stirring rod. Heating gently, on a low heat setting, is optional to aid in dissolving of the solid. After stirring, filter the binder and other insoluble material. To the resulting filtrate, slowly add about 60 mL of 1 M sodium carbonate solution. During this step, the filtrate will bubble and a white solid, calcium carbonate, precipitates. Isolate the precipitate by gravity or vacuum filtration. Allow to dry and mass.

Sample Data
Mass of calcium carbonate: 0.541 g
According to the manufacturer’s label, the antacid tablet kit contains 500–675 mg of calcium carbonate.

Answers to Questions

Guided-Inquiry Design and Procedure

1. The ideal precipitate in a gravimetric analysis procedure should be insoluble and have a known composition. Using reference texts, such as The Merck Index or the Handbook of Chemistry and Physics, look up the properties of calcium carbonate and discuss its advantages and possible disadvantages for gravimetric analysis of calcium.

   Calcium carbonate is considered virtually insoluble in water and is thus an ideal candidate for the gravimetric determination of calcium ions. (The Handbook of Chemistry and Physics lists the solubility as 0.0014 g in 100 mL of cold water and 0.0019 g in 100 mL of hot water.) It has a stable composition, but also exists as a potential hexahydrate. This could complicate the percent yield determination if the solid is not thoroughly dried to remove waters of hydration. The anhydrous solid is easily obtained by heating to approximately 100 °C. Another potential disadvantage of calcium carbonate is that heating at elevated temperatures (> 800 °C) may cause decomposition to give calcium oxide and carbon dioxide. As long as the solid is dried in an oven, and not in a burner flame, to drive off residual water, decomposition of the product is not likely to affect the results.

2. Based on solubility rules, what ions in water might interfere with the analysis of calcium ions by precipitation of calcium carbonate?

   Other carbonate salts that are insoluble in water and might interfere with gravimetric determination of calcium include the alkaline earth metal cations Mg²⁺, Sr²⁺ and Ba²⁺ as well as Fe³⁺ and other transition metal cations, such as Cu²⁺ and Ag⁺. Of these cations, only Mg²⁺ and Fe³⁺ are likely to be present in water.

3. Precipitate particles in gravimetric analysis must be large enough to be collected by filtration—smaller particles may pass through or clog the filter. Discuss how the following techniques will help prevent product loss and ensure product purity in a gravimetric procedure.
   • Add the precipitant slowly with vigorous mixing.
   • “Digest” the precipitate by allowing it to stand in contact with the solution and/or heating the mixture for 10−15 minutes.
   • Rinse the precipitate with a small amount of water after filtration.

   The precipitating agent (sodium carbonate in this example) should be added slowly to ensure even mixing and avoid local areas of supersaturation, which would lead to small particles or even a colloidal solution. Digesting the precipitate allows time for the particles to coagulate into larger particles. Heating and then cooling the mixture essentially promotes recrystallization of the solid, which leads to larger crystal or particle growth and also helps remove impurities that might be adsorbed on the surface. Finally, rinsing or washing the precipitate with water also helps remove impurities.

4. Calculate the number of moles of Ca²⁺ ion in 20 mL of each solution and the theoretical amount of CaCO₃ that can be obtained by reacting 20 mL of each solution with excess sodium carbonate. Enter the results in the table.

   Sample Calculation (sample 1):
   0.400 moles/L x 0.020 L = 0.008 moles Ca²⁺
   0.008 moles Ca²⁺ x 100.09 g/mole CaCO₃ = 0.801 g (theoretical yield)
   See Sample Data for the results of all calculations.

5. Excess sodium carbonate solution (precipitant) is recommended to ensure that all of the calcium ions in solution are converted to product. For each sample, determine the volume of 0.5 M sodium carbonate solution providing the stoichiometric number of moles of Na₂CO₃ to react completely with 20 mL of the CaCl₂ solution. Multiply the result by 1.2 to provide a 20% excess, and enter the results in the table.

   Sample Calculation (sample 1):

   {13837_Answers_Equation_3}
   See Sample Data for the results of all calculations.
6. Calculate the theoretical water hardness in mg CaCO₃/L for each water sample. Recall that each sample has been concentrated by a factor of 100 to provide the solution shown in column 2. The calculation for sample 1 is shown below as a guide.

   Sample Calculation (sample 1): (0.801 g CaCO₃/0.020 L) x (1000 mg/g) x (1/100) = 400 mg CaCO₃/L
   Note that the factor 1/100 accounts for the concentration of the original water sample to the final analyzed volume of 20 mL.
   See the following data table for the results of all calculations.

   {13837_Answers_Table_5}

Post-Laboratory Review

1. Calculate the number of moles of (a) copper(II) chloride and (b) aluminum that reacted.
   {13837_Answers_Equation_4}
2. What is the mole ratio of copper(II) chloride to aluminum metal? Express this to the nearest whole number ratio.
   {13837_Answers_Equation_5}
   The nearest whole number ratio of CuCl₂ to Al is 3:2.
3. What happened to the aluminum metal that was consumed in this reaction? Write the formula of the most probable aluminum-containing product.

   The aluminum metal dissolved and was converted to aluminum (Al³⁺) ions upon reaction with Cu²⁺ ions. The most probable aluminum-containing product is AlCl₃.

4. Write a balanced chemical equation for the single replacement reaction of copper(II) chloride with aluminum.

   3CuCl₂(aq) + 2Al(s) → 3Cu(s) + 2AlCl₃(aq)

References

Harris, D.C. Exploring Chemical Analysis, 3rd ed.; W. H. Freeman and Company: New York, 2005.

Introduction

In certain regions, the presence of hard water poses significant problems in water supply systems. Various water softening techniques are used to remove the cations responsible for water hardness. This investigation involves the application of gravimetric analysis to test samples for the amount of water hardness and calcium ions.

Concepts

• Water hardness
• Double replacement reactions
• Gravimetric analysis
• Stoichiometry

Background

Water from natural sources may contain a number of dissolved substances. The amount and nature of these dissolved substances varies depending on the geography of the area and the journey the water has taken. As water travels through the ground or over the surface of the land, it can dissolve naturally occurring minerals. As minerals dissolve in the water, the compounds separate into their respective cations and anions. Common cations in water include Na⁺, Ca²⁺, Mg²⁺ and Fe³⁺ while the principal anions in water are Cl^–, HCO₃^–, NO₃^– and SO₄^2–. The main ions contributing to water hardness are Ca²⁺, Mg²⁺ and, to a lesser extent, Fe³⁺. Their presence makes it difficult for soaps to lather and also causes a “scum” to form. Equation 1 (where R is a long hydrocarbon chain) shows the precipitation reaction between alkyl sulfate anions in a typical soap with calcium ions in hard water. The main problem due to water hardness in industrial pipes or boilers is the buildup of solid CaCO₃, which precipitates out and causes thick deposits to form in pipes and other appliances.

There are many different ways to “soften” water. One of the most common ways to remove ions is by ion exchange. The ion exchange process uses a resin to replace some of the ions that cause hardness with ions that do not. Hardness is commonly measured in units of grains per gallon or milligrams per liter (also known as parts per million), and is classified by the U.S. Department of the Interior and the Water Quality Association as follows: Although several ions contribute to water hardness, the units of mg/L or ppm are defined in terms of the equivalent mass (milligrams) of CaCO₃ that would be present per liter of water. In this investigation, gravimetric analysis will be used to precipitate and isolate solid CaCO₃ from water samples and determine water hardness. Many municipal water treatment plants use soda ash (sodium carbonate, Na₂CO₃) and lime (calcium hydroxide, Ca(OH)₂) to chemically remove calcium and magnesium ions, respectively, from hard water.

Experiment Overview

The purpose of this inquiry lab is to investigate the suitability of gravimetric analysis for determining the amount of water hardness in the form of calcium carbonate, CaCO₃, in various water samples. Six samples, representing a wide range of potential water hardness, from 50 ppm to 500 ppm, will be analyzed by various student groups as part of a cooperative class investigation to determine the accuracy and sensitivity of gravimetric analysis for water hardness testing. Note that all water samples have been concentrated by a factor of 100 for the purpose of quantitative analysis. The lab begins with an introductory activity to develop skill in the calculations and techniques of gravimetric analysis, in particular, quantitative transfer and vacuum or gravity filtration. The precipitation reaction involves preparing and combining solutions of Na₂CO₃ and CaCl₂. The balanced chemical equation for this reaction predicts the amount of precipitate that will be formed. Careful isolation, drying and weighing of the precipitate will confirm the calculations and the percent yield. The procedure provides a model for guided-inquiry design of the cooperative class investigation described above. Antacid tablets are also provided as an opportunity for further inquiry—the use of gravimetric analysis to determine the amount of calcium in an over-the-counter medication.

Materials

Prelab Questions

1. Define the term gravimetric analysis. Describe the procedure used in this activity, and identify two other common examples of gravimetric analysis.
2. Write the balanced chemical equation for the reaction between calcium chloride and sodium carbonate.
3. Calculate the number of moles of each reactant in the Introductory Activity (see steps 1 and 2). Identify the limiting reactant in the reaction and determine the theoretical amount of CaCO₃ that should be produced.
4. As noted in the Background section, hardness levels are calculated by assuming that all the “hard” metal ions come from dissolved calcium carbonate and are reported in mg CaCO₃/L. Calculate the equivalent water hardness in mg CaCO₃/L for a calcium chloride solution containing 0.1 M Ca²⁺ ions.

Safety Precautions

Sodium carbonate is irritating to body tissues. Anhydrous calcium chloride is moderately toxic by ingestion and generates a great deal of heat when dissolved in water. Avoid contact of all chemicals with eyes and skin. Antacid tablets used in the lab are considered laboratory chemicals and may not be removed from the lab. Do not taste or ingest any materials in the chemistry lab. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines.

Procedure

Introductory Activity

Precipitation Reaction and Vacuum Filtration

1. Weigh 2.5 g of sodium carbonate and place in a clean, dry 150-mL beaker. Record the precise mass and dissolve the solid in 20 mL of deionized or distilled water.
2. Weigh 2.0 g of calcium chloride and place in a clean, dry 150-mL beaker. Record the precise mass and dissolve the solid in 20 mL of deionized or distilled water.
3. Combine the two solutions by slowly adding the sodium carbonate solution to the calcium chloride solution. Record all observations about the reaction.
4. Separately weigh a piece of filter paper and a watch glass and record their masses.
5. Set up a vacuum filtration apparatus as shown in Figure 1. The second filter flask is used to prevent back-up of water from the aspirator to the filter flask when the vacuum is released. {13837_Procedure_Figure_1}
6. Isolate the precipitate by vacuum filtration. Careful transfer techniques are essential for accurate results! Precipitate may also be collected by gravity filtration.
7. Place the watch glass and filter paper containing precipitate in a lab oven to dry at 100 °C for 10–20 minutes. Monitor and carefully break up the solid with a spatula to ensure complete drying.
8. Calculate the percent yield of calcium carbonate.

Guided-Inquiry Design and Procedure

Accuracy and Sensitivity of Gravimetric Analysis to Determine Water Hardness
Form a working group with other students and discuss the following questions.

1. The ideal precipitate in a gravimetric analysis procedure should be insoluble and have a known composition. Using reference texts, such as The Merck Index or the Handbook of Chemistry and Physics, look up the properties of calcium carbonate and discuss its advantages and possible disadvantages for gravimetric analysis of calcium.
2. Based on solubility rules, what ions in water might interfere with the analysis of calcium ions by precipitation of calcium carbonate?
3. Precipitate particles in gravimetric analysis must be large enough to be collected by filtration—smaller particles may pass through or clog the filter. Discuss how the following techniques will help prevent product loss and ensure product purity in a gravimetric procedure.
   • Add the precipitant slowly with vigorous mixing.
   • “Digest” the precipitate by allowing it to stand in contact with the solution and/or heating the mixture for 10−15 minutes.
   • Rinse the precipitate with a small amount of water after filtration.

Six water samples containing known concentrations of calcium chloride are available for analysis as part of a cooperative class activity. Each student group should analyze two different samples. The recommended sample volume for the precipitation reaction is 20 mL. Complete the table with the results of the calculations from Questions 4−6.

4. Calculate the number of moles of Ca²⁺ ion in 20 mL of each solution and the theoretical amount of CaCO₃ that can be obtained by reacting 20 mL of each solution with excess sodium carbonate. Enter the results in the table.
5. Excess sodium carbonate solution (precipitant) is recommended to ensure that all of the calcium ions in solution are converted to product. For each sample, determine the volume of 0.5 M sodium carbonate solution that provides the stoichiometric number of moles of Na₂CO₃ needed to react completely with the CaCl₂ solution. Multiply the result by 1.2 to provide a 20% excess, and enter the results in the table.
6. Calculate the theoretical water hardness in mg CaCO₃/L for each water sample. Recall that each sample has been concentrated by a factor of 100 to provide the solution shown in column 2. The calculation for sample 1 is shown as a guide.
   Sample Calculation (sample 1): (0.801 g CaCO₃/0.020 L) x (1000 mg/g) x (1/100) = 400 mg CaCO₃/L
   Note that the factor 1/100 accounts for the concentration of the original water sample to the final analyzed volume of 20 mL.
7. Write a detailed, step-by-step procedure for analyzing the concentration of calcium in the water samples. Include the reagents needed, the glassware and equipment that will be used, and the appropriate measurements and observations that must be made.
8. Review the hazards of the chemicals used in the procedure and write appropriate safety precautions that must be followed during the experiment.
9. Carry out the procedure and record the results in an appropriate data table.
10. Repeat the analysis as needed to check for reproducibility.

Analyze the Results
Calculate the percent yield of calcium carbonate and determine the experimental water hardness in mg/L for each sample. Classify the water hardness of each sample according to the criteria established by the U.S. Department of the Interior and the Water Quality Association (see the Background section). Compile the class data for all the samples that were analyzed and compare the accuracy and sensitivity of the gravimetric analysis procedure over the range of possible water hardness from 50 to 500 mg/L.

Student Worksheet PDF

13837_Student1.pdf