Materials Included In Kit

Additional Materials Required

Safety Precautions

This activity requires the use of hazardous components and/or has the potential for hazardous reactions. Sodium hydroxide solution is corrosive and may cause skin burns. Iron(III) nitrate, copper(II) chloride and sodium phosphate solutions may be skin/tissue irritants. All of these chemicals are toxic by ingestion. Avoid all body tissue contact when working with these chemicals. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

All solutions may be flushed down the drain with plenty of water. The solids should be disposed of in the solid waste disposal. Wood splints and Beral-type pipets may be rinsed and reused or disposed of in the solid waste disposal.

Teacher Tips

• This kit is an introductory kit designed to aid in the teaching of writing ionic chemical formulas.

  Parts A & B—Day 1

• In part A, students will experimentally determine the formula of Iron(III) hydroxide by performing a precipitation reaction using various ratios of Iron(III) nitrate and sodium hydroxide. The precipitate formed will be the rust-colored Iron(III) hydroxide, Fe(OH)₃. Students should determine a 1 to 3 ratio of Fe³⁺ to OH^– by measuring the height of the precipitate formed in each tube. Sample data is provided in this section.
• In part B, students will experimentally determine the formula of copper(II) phosphate by performing another precipitation reaction using various ratios of copper(II) chloride and sodium phosphate. The precipitate formed will be the blue-colored copper(II) phosphate, Cu₃(PO₄)₂. Students should determine a 3 to 2 ratio of Cu²⁺ to PO₄^3– by measuring the height of the precipitate formed in each tube. Sample data is provided in this section.
• For both parts A & B, students will construct bar graphs of their data and answer questions. The questions are aimed at making students aware of the charges on the ions and how this relates to the combining ratio of the two ions to form the compound.
• Students will need two charts in order to make their predictions and answer the questions. Masters for these charts are provided in the kit and may be reproduced for student use. The Ion Formula chart includes the names and charges of some common ions. The Solubility Rules chart provides some general solubility rules for inorganic compounds.

Sample Data

Part A.

Answers to Questions

Part A 

1. Graph: part A—Iron(III) nitrate and sodium hydroxide
   {12837_Answers_Figure_1}
2. Write the name of the solid that formed in the tubes when you mixed Iron(III) nitrate and sodium hydroxide.

   The solid that formed is Iron(III) hydroxide.

3. Look at your experimental results.
   1. Which test tube had the greatest amount of precipitate?

      Test Tube ___2___

   2. Was your prediction correct about which tube would have the most precipitate?

      Student answers will vary.

   3. What is the ratio of the Iron(III) nitrate to sodium hydroxide in the tube with the greatest amount of precipitate?

      Ratio of Fe(NO₃)₃:NaOH = ___1 to 3___

4. Write the formula (using the ratio from Question 3) for the precipitate, based on your experimental results.

   The correct formula for the precipitate is Fe(OH)₃.

5. 1. Which reagent is in excess in test tube 1?

      ___NaOH___

   2. Which reagent is the limiting reagent in test tube 1? (Hint: Which ran out first?)

      ___Fe(NO₃)₃___

6. 1. Which reagent is in excess in test tube 7?

      ___Fe(NO₃)₃__

   2. Which reagent is the limiting reagent in test tube 7? (Hint: Which ran out first?)

      ___NaOH___

7. Write the formulas for both the iron(III) ion and the nitrate ion. (Hint: Refer to the Ion Formula chart.)

   Iron(III) ion = Fe³⁺, nitrate ion = NO₃^–

8. Write the formulas for both the sodium ion and the hydroxide ion. (Hint: Refer to the Ion Formula chart.)

   Sodium ion = Na⁺, hydroxide ion = OH^–

9. Based on the formulas and charges for the four ions listed in Questions 7 and 8, what would you predict is the correct chemical formula for the precipitate?

   The correct chemical formula for the precipitate is Fe(OH)₃.

10. Do your experimental results for the formula of this precipitate (from Question 4) agree with your predicted chemical formula in Question 9? Explain.

    Student answers will vary.

11. Write the complete balanced equation for the reaction between Iron(III) nitrate and sodium hydroxide. Include physical states, using (aq) for aqueous, (s) for solid, (l) for liquid and (g) for gas. Include the names of the two products.

    The complete balanced equation is Fe(NO₃)₃(aq) + 3NaOH(aq) → Fe(OH)₃(s) + 3NaNO₃(aq)
    The two products that form are solid Iron(III) hydroxide and aqueous sodium nitrate.

12. Using your results from Table 6, construct a bar graph using the test tube graph provided. To do this, plot the height of the precipitate (in mm) versus the test tube number by shading in the tubes to the appropriate level. Above each tube, label the ratio of the two reactants using the whole-number ratios that you determined in Table 6.
    {12837_Answers_Figure_2_Copper(ll) chloride and sodium phosphate graph}
13. Write the name of the solid that formed in the tubes in part B when you mixed copper(II) chloride and sodium phosphate.

    The solid that formed is copper(II) phosphate.

14. Look at your experimental results.
    1. Which test tube had the greatest amount of precipitate?

       Test Tube ___5___

    2. Was your prediction correct about which tube would have the most precipitate?

       Student answers will vary.

    3. What is the ratio of the copper(II) chloride to sodium phosphate in the tube with the greatest amount of precipitate?

       Ratio of CuCl₂:Na₃PO₄ = ___3:2___

15. Write the formula (using the ratio from Question 14) for the precipitate, based on your experimental results.

    The correct formula for the precipitate is Cu₃(PO₄)₂.

16. Write the formulas for both the copper(II) ion and the chloride ion. (Hint: Refer to the Ion Formula chart.)

    Copper(II) ion = Cu²⁺, chloride ion = Cl^–

17. Write the formulas for both the sodium ion and the phosphate ion. (Hint: Refer to the Ion Formula chart.)

    Sodium ion = Na⁺, Phosphate ion = PO₄^3–

18. Based on the formulas and charges for the four ions and charges listed in Questions 16 and 17, what would you predict is the correct chemical formula for the precipitate?

    The correct chemical formula for the precipitate is Cu₃(PO₄)₂.

19. Do your experimental results for the formula of this precipitate (from Question 15) agree with your predicted chemical formula in Question 18? Explain.

    Student answers will vary.

20. Write the complete balanced equation for the reaction between copper(II) chloride and sodium phosphate. Include physical states, using (aq) for aqueous, (s) for solid, (l) for liquid and (g) for gas. Include the names of the two products.

    3CuCl₂(aq) = 2Na₃PO₄(aq) → Cu₃(PO₄)₂(s) = 6NaCl(aq)
    The two products that form are solid copper(II) phosphate and aqueous sodium chloride.

21. What do you think determines how much precipitate will be made in each tube?

    Student answers may vary.
    The amount of precipitate depends on having the correct proportion or ratio of the two components according to the coefficients in the balanced equation. For example, in part A each Iron(III) ion combines with three hydroxide ions until one of the reactants is used up. Having an excess amount of one of the reactants may possibly cause some of the precipitate to dissolve, leaving less solid than may be expected in the tube.

22. What if you had not tried combining the correct ratio of reactants?

    Student answers may vary.
    If the correct combining ratio of reactants was not tried, the experimenter may be misled. For this reason, graphing the data is important. Students should look for a maximum point on the bell-like curve. If two tubes contain the same amount of precipitate and there is no maximum, students should be alerted that there may be a point between the two ratios which needs to be tried.

23. Discuss any possible sources of error which may have occurred in this experiment. Discuss ways in which the experiment may be modified or improved.

    Possible errors may be inconsistent drop size, inconsistent number of drops, failure to allow precipitate to settle before measuring the height in the tube, or inaccurate measurement of height of the precipitate. The experiment may be improved by performing multiple trials for each set of reactants, carefully counting drops and measuring precipitate heights, holding the pipet vertically for consistent drop size, and allowing the precipitate to settle fully before measuring the height.

Teacher Handouts

12837_Teacher1.pdf

References

Herron, J. D.; Sarquis, J. L.; Schrader, C. L.; Frank, D. V.; Sarquis, M.; Kukla, D. A. Chemistry; D. C. Heath: Boston, MA, 1996; Chapter 19.

LeMay, H. E.; Beall, H.; Robblee, K. M.; Brower, D. C. Chemistry: Connections to Our Changing World; Prentice Hall: Upper Saddle River, NJ, 1996; Chapter 7.

Wilbraham, A. C.; Staley, D. D.; Simpson, C. J.; Matta, M. S. Chemistry; Addison-Wesley: Menlo-Park, CA, 1987, pages 64-65.

Introduction

Water, salt, vitamin C, rust... these are just a few of the multitudes of chemical compounds that have been identified. Each compound has its own unique properties and each is represented by its own chemical formula. Investigate how chemical formulas can be experimentally determined.

Concepts

• Chemical formulas
• Ionic compounds
• Polyatomic ions
• Precipitate
• Solubility

Background

Atoms of different elements combine chemically with one another to form chemical compounds. A given compound always has the same relative number and kinds of atoms. In other words, a specific compound will form when certain atoms combine together in a fixed whole-number ratio. The law of definite composition states that the proportion of the elements in a given compound is fixed. This is one of the postulates stated in John Dalton’s atomic theory of matter which was first presented in 1808.

The chemical formula of a compound indicates the kinds of atoms and the relative number of each atom that is combined together to form the compound. Subscripts, the numbers written after the chemical element symbols, are used to indicate the number of atoms of each element in the compound. Water, the most abundant chemical on Earth, has the chemical formula H₂O with a whole-number hydrogen:oxygen ratio of 2 to 1. The hydrogen:oxygen ratio in water can never be 1 to 1, 3 to 1, or even 2½ to 1. It must be 2 to 1. The ratio and kinds of atoms in a compound defines that compound, which is thus considered a pure substance. All elements and compounds are pure substances. The only way to know the formula of a compound is to experimentally determine the ratio of the elements in the compound. The formula of the compound is then very useful in predicting further chemical reaction of that compound.

In many compounds, such as water or carbon dioxide, the atoms are bound together by bonds called covalent bonds. Covalent bonds form by the sharing of electrons between two atoms, most often between two nonmetal atoms, resulting in the formation of a molecule. A molecule is an electrically neutral group of atoms that acts as a unit. Compounds composed of molecules are molecular compounds.

Not all compounds are composed of molecules. Some compounds such as sodium chloride, which is common table salt, are composed of ions. Ions are atoms or groups of atoms that have a positive or a negative charge. Ions may join together because of an oppositely charged or electrostatic attraction between a positive cation and a negative anion resulting in an ionic bond. Compounds held together by ionic bonds are ionic compounds. Ionic compounds are generally composed of an extended network of positively charged metal ions and negatively charged nonmetal ions. Although composed of ions, ionic compounds are electrically neutral. That is, the total positive charge is equal to the total negative charge. The chemical formula indicates the smallest whole-number ratio of each element in the network of ions and is called a formula unit. For example, magnesium chloride has the chemical formula MgCl₂. The magnesium cation (Mg²⁺) and the chloride anion (Cl^–) combine in a 1:2 ratio to form MgCl₂. The overall charge on the ionic compound must be zero, thus the compound needs two Cl^– ions for each Mg²⁺ ion.

Some ions consist of a group of covalently bonded atoms that tend to stay together as if they were a single ion. Such ions are called polyatomic ions. An example is the nitrate ion, NO₃^–. This polyatomic anion contains one nitrogen atom and three oxygen atoms and has an overall charge of –1. If a calcium cation, Ca²⁺, is combined with a nitrate anion, NO₃^–, there must be two nitrate ions for every calcium ion in order to balance the positive and negative charges and achieve an overall charge of zero. Thus the formula for calcium nitrate is written as Ca(NO₃)₂, showing that for every calcium ion there are two nitrate ions. Parentheses must be used around the polyatomic ion to show that the subscript pertains to the polyatomic ion as a whole. Notice that subscripts of one are commonly omitted when writing chemical formulas as they are assumed.

Experiment Overview

In this lab, the formulas for various ionic compounds will be verified. In each part of the lab, solutions of two ionic compounds will be combined in various ratios in test tubes. Some ionic compounds dissolve in water and thus are soluble, existing as separate ions in solution. Other ionic compounds do not dissolve in water and are called insoluble, forming a solid material called a precipitate. Using a table of solubility rules which lists whether or not a compound is soluble or insoluble in water, a prediction will be made as to which ions will combine together to form a precipitate. By observing which reactant ratio produces the most solid product, the chemical formula for the ionic compound will be written.

Materials

Safety Precautions

Sodium hydroxide solution is corrosive and may cause skin burns. Iron(III) nitrate, copper(II) chloride and sodium phosphate solutions are skin and tissue irritants. All of these chemicals are slightly toxic by ingestion. Avoid all body tissue contact when working with these chemicals. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part A. Iron(III) Nitrate and Sodium Hydroxide

1. Label seven small test tubes 1–7 with a marking pen and place them in a test tube rack or in a 24-well plate.
2. Use the Solubility Rules chart to predict the name of the precipitate that will form when Iron(III) nitrate and sodium hydroxide are mixed. (Note: First determine which possible compounds may form when the two reactants are mixed and then use the table of solubility rules to determine which of the possibilities will precipitate in water.) Fill in your prediction by writing the name of the product in Table 1. Explain the basis for your prediction in Table 1.
3. Use the Ion Formula chart to predict which ratio of Iron(III) nitrate to sodium hydroxide will produce the most precipitate. Fill in your prediction by writing the ratio in Table 1. Explain the basis for your prediction in Table 1.
4. Fill a Beral-type pipet with 0.1 M Iron(III) nitrate solution. Observe the physical properties of the 0.1 M Iron(III) nitrate solution. Record your observations in Table 2. Carefully add the number of drops indicated in Table 3 to each test tube. (Note: Exact volumes are crucial, so it is important to count accurately and to hold the pipet vertically so the drop size is consistent and repeatable.)
5. Fill a different Beral-type pipet with 0.1 M sodium hydroxide solution. Observe the physical properties of the 0.1 M sodium hydroxide solution. Record your observations in Table 2. Again carefully add the number of drops indicated in Table 3 to each test tube. Note: Exact volumes are crucial, so it is important to count accurately and to hold the pipet vertically so the drop size is consistent and repeatable.
6. Use a wood splint to stir each of the mixtures in the tubes.
7. Let the tubes stand undisturbed for about 10 minutes to allow the precipitates to settle to the bottom of the tubes. During this time, determine the lowest whole-number ratio of drops of Iron(III) nitrate to sodium hydroxide. Fill in this ratio in Table 3. Go on to part B while waiting for the precipitates to settle.
8. After the precipitates have settled to the bottom of the tubes, observe the physical properties of the new substances formed Record all observations of both the solid and solution in Table 2.
9. Use a metric ruler to measure the height of the precipitate in millimeters in each test tube. Read from the top of the solid material to the bottom center of the test tube. It is important that you measure accurately so that your results will be consistent. Record each height in mm in Table 3.
10. Clean out the test tubes by pouring the solutions down the drain and dumping the solids in the solid waste disposal. Rinse out the tubes with plenty of water; use a test tube brush if necessary. Wash out the test tubes with soap and water for the next lab group.

11. Again label seven small test tubes 1–7 with a marking pen and place them in a test tube rack or in a 24-well plate.
12. As in part A, use the Solubility Rules chart to predict the name of the precipitate that will form when copper(II) chloride and sodium phosphate are mixed. Fill in your prediction by writing the name of the product in Table 4. Also write the basis for your prediction in Table 4.
13. Use the Ion Formula chart to predict which ratio of copper(II) chloride and sodium phosphate will produce the most precipitate. Fill in your prediction by writing the ratio in Table 4. Also write the basis for your prediction in Table 4.
14. Fill a Beral-type pipet with 0.1 M copper(II) chloride solution. Observe the physical properties of the 0.1 M copper(II) chloride solution. Record your observations in Table 5. Carefully add the number of drops indicated in Table 6 to each test tube. Note: Exact volumes are crucial, so it is important to count accurately and to hold the pipet vertically so the drop size is consistent and repeatable.
15. Fill a different Beral-type pipet with 0.1 M sodium phosphate solution. Observe the physical properties of the 0.1 M sodium phosphate solution. Record your observations in Table 5. Again carefully add the number of drops indicated in Table 6 to each test tube. Note: Exact volumes are crucial, so it is important to count accurately and to hold the pipet vertically so the drop size is consistent and repeatable.
16. Use a wood splint to stir each of the mixtures in the tubes.
17. Let the tubes stand undisturbed for about 10 minutes to allow the precipitates to settle to the bottom of the tubes. During this time, determine the lowest whole-number ratio of drops of copper(II) chloride to sodium phosphate. Fill in this ratio in Table 6. Go back to part A, step 8, while you wait for the precipitates to settle.
18. After the precipitates have settled to the bottom of the tubes, observe the physical properties of the new substances formed. Record all observations of both the solid and solution in Table 5.
19. Use a metric ruler to measure the height of the precipitate in millimeters in each test tube. Read from the top of the solid material to the bottom center of the test tube. It is important that you measure accurately so that your results will be consistent. Record each height in mm in Table 6.
20. Clean out the test tubes by pouring the solutions down the drain and dumping the solids in the solid waste disposal. Rinse out the tubes with plenty of water; use a test tube brush if necessary. Wash out the test tubes with soap and water for the next lab group.

Student Worksheet PDF

12837_Student1.pdf