Materials Included In Kit

Additional Materials Required

Safety Precautions

Ammonia solution is slightly toxic by ingestion and inhalation; both liquid and vapor are irritating, especially to the eyes. Cobalt(II) sulfate solution is irritating to eyes, skin and respiratory tract. It is slightly toxic. Copper(II) nitrate solution is moderately toxic by ingestion. Copper(II) sulfate solution is slightly toxic by ingestion. Ethylenediamine solution is slightly toxic by inhalation and skin absorption; irritant to skin and eyes. Sodium hydroxide solution is a corrosive liquid and is dangerous to the eyes; skin burns are possible. Sodium nitrite solution is moderately toxic by ingestion. Avoid contact of all chemicals with skin and eyes. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. 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. Provide each lab group with a 250-mL beaker to dispose of the waste solutions of copper and a 50-mL beaker to dispose of the cobalt waste solutions. The cobalt waste solutions may be disposed of according to Flinn Suggested Disposal Method #2. The copper waste solutions may be disposed of according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Enough materials are provided in this kit for 30 students working in pairs or for 15 groups of students. Both parts of this laboratory activity may reasonably be completed in one 50-minute class period.
• Less than 1 mL of solution may be used per test tube in the experiment as long as equal volumes are combined.
• Dropper bottles can be used for the solutions. To avoid backups in using the reagents, multiple lab stations for each part of the experiment can be set up, with the groups split between the two parts.
• Students can premark the test tubes for solution levels. Have the students add 1 mL of deionized water to the test tube, marking the level on the test tube and another 1 mL, mark this level and finally add one more mL and mark the 3 mL volume level.

Teacher Tips

• The lab can be extended to determining the spectrochemical series of a list of ligands for the copper(II) ion. By determining the color absorbed by the ligand–copper complexes and the corresponding wavelengths, the ligands can be ranked according to the relative energies absorbed. This ranking is called the spectrochemical series. Possible ligands to use, in addition to the ones in the kit, are iodide ion, bromide ion, thiocyanate ion, acetate ion and ethylenediamine.

Answers to Prelab Questions

1. Cobalt has a coordination number of six. If cobalt nitrate, Co(NO₃)₂, is added to a saturated solution of sodium chloride, NaCl, what will be the formula of the complex ion that is formed?

   CoCl₆^4–(aq)

2. The oxalate ion, C₂O₄^2–, is bidentate. If an excess of sodium oxalate, Na₂C₂O₄, is added to a solution of zinc nitrate, Zn(NO₃)₂, what will be the formula of the complex ion that is formed? The coordination number of zinc is four.

   Since the oxalate ion is bidentate, each ion occupies two coordination sites. Therefore, the complex ion formula is Zn(C₂O₄)₂^2–(aq).

3. Refer to the color wheel above. If the complex ion Cr(NH₃)₆³⁺ is yellow in solution, what color and wavelengths of light are absorbed by the complex?

   The complement color for yellow is deep blue and therefore the wavelength of light being absorbed by the compound is 420 to 450 nm.

4. Copper(II) nitrate, Cu(NO₃)₂, forms a light blue precipitate when added to a solution of sodium phosphate, Na₃PO₄. When copper(II) nitrate is added to a solution of sodium nitrite, NaNO₂, a green solution results. If equal concentrations of sodium phosphate and sodium nitrite are added to a solution of copper(II) nitrate, a light blue precipitate forms. Write the formula for each complex ion or precipitate formed and determine which complex ion or precipitate has the higher relative stability. The coordination number of copper is four.

   For the nitrite ion, NO₂^–, the complex ion formed is Cu(NO₂)₄^2–(aq). 
   Because a precipitate is formed with the phosphate ion, PO₄^3–(aq), the formula of the compound is Cu₃(PO₄)₂.
   Since a blue precipitate forms when equal amounts of the two ligands are added to the copper nitrate solution, the copper(II) phosphate has the higher relative stability.

Sample Data

Part 1

Part 2

Answers to Questions

1. Determine the formula of the complex ion or ionic compounds formed during Part 1 in test tubes 2, 3, 5 and 6 in Part 1. Record these values in Part 1 of the data table.
2. Using the color wheel in the Background section, determine the color and wavelength range of the visible light absorbed by the complex ion or precipitate in each test tube 1 through 6. Record these values in Part 1 of the data table.
3. Determine the formula of each complex ion or ionic solid formed in Part 2. The coordination number of the copper is four. Write a balanced chemical equation for the formation of each complex ion or ionic compound and the expression for its stability constant (K_f).
   {10604_Answers_Equation_5}
4. From the results, rank the stability constants for the five complex ions or ionic compounds in Part 2 from highest to lowest. For example, if adding hydroxide ion to the copper–ammonia complex ion solution causes a precipitate and/or color change, then the stability constant for the copper–hydroxide complex ion is greater than that for the copper–ammonia complex ion.

   The first row data indicates that only the hydroxide ligand has a larger stability constant than the ammonia ligand. Since the other three failed to cause a precipitation or color change with the copper–ammonia complex, each has a lower stability constant than ammonia.
   Cu(OH)₂(s) > Cu(NH₃)₄²⁺(aq) > CuCl₄^2–(aq), Cu(C₂O₄)(s), and Cu(NO₂)₄^2–(aq)
   
   The second row data indicates that four species have a greater stability constant than the chloride ion. The hydroxide ion, OH^–, caused a light blue precipitation. The oxalate ion, C₂O₄^2–, formed two species. The light blue solid copper(II) oxalate, CuC₂O₄, precipitated from solution, then dissolved to form a turquoise solution of the complex ion Cu(C₂O₄)₂^2–. The nitrite ion, NO₂^–, formed a green solution. The order of stability is:
   Cu(OH)₂(s), CuC₂O₄(s), Cu(C₂O₄)₂^2–(aq), Cu(NO₂)₄^2–(aq) > CuCl₄^2–(aq)
   and
   Cu(C₂O₄)₂^2–(aq) > CuC₂O₄(s)
   
   The third row data indicates that the hydroxide ion, OH^–, has a larger stability constant than either the oxalate ion, C₂O₄^2–, or the nitrite ion, NO₂^–.
   Cu(OH)₂(s) > CuC₂O₄(s), Cu(C₂O₄)₂^2–(aq), Cu(NO₂)₄^2–(aq)
   
   The fourth row indicates that the oxalate ion, C₂O₄^2–, has a larger stability constant than the nitrite ion.
   CuC₂O₄(s), Cu(C₂O₄)₂^2–(aq) > Cu(NO₂)₄^2–(aq)
   
   The overall series of the relative stability constant is:
   Cu(OH)₂(s) > Cu(NH₃)₄²⁺(aq) > Cu(C₂O₄)₂^2–(aq) > CuC₂O₄(s) > Cu(NO₂)₄^2–(aq) > CuCl₄^2–(aq)

References

Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers in Chemistry; University of Wisconsin: Madison, WI; 1983; Vol. 1, pp 260–263.

Slowinski, E. J., Wosley, W. C., Masterton, W. L. Chemical Demonstrations in the Laboratory; Saunders College: Fort Worth, TX; 1990; p 223.

Introduction

Complex ions are interesting substances that contain transition metals and are usually highly colored when in solution. In this lab, the structure, stability and reactions of a series of complex ions will be investigated.

Concepts

• Complex ion
• Ligand
• Coordination number
• Color and energy absorption
• Relative stability
• Transitions metals

Background

When transition metal cations are in solution, they do not exist as free ions, but are bonded to neutral molecules or ions. The molecule or ion, called a ligand, has one or more lone pair of electrons that “coordinate with” and form a bond to the metal ion.

The number of ligand atoms bonded to the metal ion is called the coordination number of the metal. This value is usually, but not always, 2, 4 or 6, and depends on both the particular metal ion and the nature of the ligand.

In forming a bond with a metal ion, the ligand acts as a Lewis base (electron donor), while the metal ion acts as a Lewis acid (electron acceptor).

The bonding in complex ions, also called coordination compounds, involves the overlap of the metal d-orbitals with the ligand lone pair orbitals. If two atoms in a ligand each donate a lone pair of electrons to form separate single bonds with the metal ion, then the ligand is said to be a bidentate ligand. If the ligand donates more than two lone pairs of electrons, it is called a polydentate ligand. As many as six atoms in an individual ligand may be bonded with the metal ion. Ethylene diamine, C₂H₈N₂, is an example of a bidentate ligand. It has two nitrogen atoms with lone pairs of electrons capable of binding to a metal ion.

The copper(II) ion has a coordination number of four. In solution, the ion exists as a complex ion with four water ligands, Cu(H₂O)₄²⁺. The water molecules are loosely bonded and are thus easily displaced by other ligands (:L) that will form more stable complex ions with the copper(II) ion (Equation 2). In general, any ligand can be displaced in turn by a ligand that forms a more stable complex with the metal ion. The relative stability of different ligand complexes of a metal in solution can be determined by comparing the formation constant values for the different complexes. For the reaction of a hydrated copper ion with a specific ligand: The formation constant (K_f) is defined as
The higher the formation constant, the more stable the complex ion. With some ionic ligands (X^–) the metal may precipitate as an ionic solid rather than form a complex ion (Equation 4). If equal concentrations of different ligands are added to a solution of hydrated copper ions, the complex ion or precipitate with the higher stability will form.

The color of a complex ion results from the absorption of light by the metal. The particular wavelength of light absorbed depends on the ligand. The actual color of the ion in solution is the complement of the absorbed color (white light with the absorbed color subtracted). Referring to the following diagram when a complex ion absorbs one color, it appears as the opposite, or complementary, color: Thus, a green copper(II) complex ion results from the absorption of red light having wavelengths in the range 630–720 nm.

Experiment Overview

The purpose of this experiment is to investigate the properties of complex ions. In Part 1, the absorbed energy and chemical formula of four complex ions will be determined. In Part 2, the relative stability of five different ligand combinations with the copper(II) ion will be studied.

Materials

Prelab Questions

1. Cobalt has a coordination number of six. If cobalt nitrate, Co(NO₃)₂, is added to a saturated solution of sodium chloride, NaCl, what will be the formula of the complex ion that is formed?
2. The oxalate ion, C₂O₄^2–, is bidentate. If an excess of sodium oxalate, Na₂C₂O₄, is added to a solution of zinc nitrate, Zn(NO₃)₂, what will be the formula of the complex ion that is formed? The coordination number of zinc is four.
3. Refer to the color wheel above. If the complex ion Cr(NH₃)₆³⁺ is yellow in solution, what color and wavelengths of light are absorbed by the complex?
4. Copper(II) nitrate, Cu(NO₃)₂, forms a light blue precipitate when added to a solution of sodium phosphate, Na₃PO₄. When copper(II) nitrate is added to a solution of sodium nitrite, NaNO₂, a green solution results. If equal concentrations of sodium phosphate and sodium nitrite are added to a solution of copper(II) nitrate, a light blue precipitate forms. Write the formula for each complex ion or precipitate formed and determine which complex ion or precipitate has the higher relative stability. The coordination number of copper is four.

Safety Precautions

Ammonia solution is slightly toxic by ingestion and inhalation; both liquid and vapor are irritating, especially to the eyes. Cobalt(II) sulfate solution is irritating to eyes, skin and respiratory tract. It is slightly toxic. Copper(II) nitrate solution is moderately toxic by ingestion. Copper(II) sulfate solution is slightly toxic by ingestion. Ethylenediamine solution is slightly toxic by inhalation and skin absorption; irritant to skin and eyes. Sodium hydroxide solution is a corrosive liquid and is dangerous to the eyes; skin burns are possible. Sodium nitrite solution is moderately toxic by ingestion. Avoid contact of all chemicals with skin and eyes. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Note: Only take small amounts of reagents. Never replace reagents and pour back into bottle. A mL is about 10 mm in the test tubes.

Part 1 

1. Obtain six test tubes, label them 1 to 6, and place them in a test tube rack.
2. Using a pipet or medicine dropper, add approximately 1 mL of 1 M copper(II) sulfate solution to the first three test tubes. Record the initial color of the copper(II) sulfate solution in Part 1 of the data table.
3. Using a clean pipet or medicine dropper, add several drops of 1 M sodium nitrite solution to the second test tube and swirl the test tube. Continue adding drops and swirling until a permanent color change occurs. Record the color of the resulting solution in Part 1 of the data table.
4. Repeat step 3, adding 0.1 M ethylenediamine solution to test tube 3.
5. Add approximately 1 mL of 0.1 M cobalt(II) sulfate solution to test tubes 4, 5, and 6. Record the initial color of the cobalt(II) sulfate solution in Part 1 of the data table.
6. Repeat steps 3 and 4 adding 1 mL of 1 M potassium oxalate solution to test tube 5 and 1 mL of 1 M ethylenediamine solution to test tube 6. Record the color change in the data table.
7. Dispose of solutions as directed by the instructor and thoroughly rinse pipets and test tubes with distilled water from a wash bottle.

Part 2 Stability of Initial Copper–Ammonia Complex

1. Obtain five clean test tubes (1 to 5) and place them in a test tube rack.
2. Using a clean pipet or medicine dropper, add approximately 1 mL of 1 M copper(II) nitrate solution to each test tube 1 to 5.
3. With a clean pipet or medicine dropper, add 1 M NH₃ drop by drop to test tube 1. Continue adding NH₃ and swirling the test tube until the precipitate which forms dissolves. In the Part 2 data table, record the color of the final solution in the upper left box with both the column and row labeled NH₃.
4. Repeat step 3 to prepare the copper–ammonia complex ion in each of the remaining test tubes. All solutions should be the same color.
5. Add 1 mL of 1 M sodium chloride solution to test tube 2 and swirl the test tube. If a precipitate forms, write ppt in the first row box, labeled Cl^–, of the data table. If a new (permanent) color change is observed, write the color of the complex ion in the first row box labeled Cl^– of the data table.
6. Repeat step 5 for the remaining test tubes, adding 1 mL of 1 M sodium hydroxide to test tube 3, 1 mL of 1 M potassium oxalate to test tube 4, and 1 mL of sodium nitrite to test tube 5. Remember to record either ppt or the color of the complex ion in the appropriate box in the data table.
7. Dispose of the solutions in the test tubes as directed by the instructor and thoroughly rinse test tubes with distilled water from a wash bottle.

Part 3. Stability of Copper Chloride Complex

1. Place four of the rinsed test tubes in the test tube rack. The procedure will now be repeated using the chloride ion, Cl^–(aq), as the original ion added to the copper(II) nitrate solution.
2. Add 1 mL of 1 M copper(II) nitrate solution to test tubes 1 through 4.
3. Add 1 mL of 1 M sodium chloride to each test tube and swirl to mix. Record any precipitation as ppt and any color changes in test tube 1 in the second row, second column box, labeled Cl^–, in the data table.
4. Add 1 mL of 1 M sodium hydroxide solution to test tube 2 and swirl. Record any precipitation as ppt and record any color changes in the second row box, labeled OH^–, of the data table.
5. Add 1 mL of 1 M potassium oxalate solution to test tube 3 and swirl. Record any precipitation as ppt and record any color changes in the second row box, labeled C₂O₄^2–, of the data table.
6. Add 1 mL of 1 M sodium nitrite solution to test tube 4 and swirl. Record any precipitation as ppt and record any color changes in the second row box, labeled NO₂^–, of the data table.
7. Dispose of solutions as directed the the instructor and thoroughly rinse the test tubes with distilled water from a wash bottle.

Part 4. Stability of Copper Hydroxide Complex

1. Place three test tubes in the test tube rack. The original ion added to the copper(II) nitrate solution is now OH^–.
2. Add 1 mL of 1 M copper(II) nitrate solution to three test tubes 1 through 3.
3. Add 1 mL of 1 M sodium hydroxide to each test tube and swirl to mix. Record any precipitation as ppt and record any color change in the third row box, labeled OH^–, in the data table.
4. Add 1 mL of 1 M potassium oxalate to test tube 2 and swirl the test tube to mix. Record any precipitation as ppt and record any color change in the third row box, labeled C₂O₄^2–, of the data table.
5. Add 1 mL of 1 M sodium nitrite solution to test tube 3 and swirl to mix. Record any precipitation as ppt and record any color change in the third row box, labeled NO₂^–, of the data table.

Part 5. Stability of Copper Oxalate and Copper Nitrite Complexes

1. Place a clean test tube in the test tube rack. Add 1 mL of 1 M copper(II) nitrate solution to the test tube.
2. Add 1 mL of 1 M potassium oxalate to the test tube and swirl to mix. Record any precipitation as ppt and record any color change in the fourth row box, labeled C₂O₄^2–, of the data table.
3. Add 1 mL of 1 M sodium nitrite to the test tube and swirl to mix. Record any precipitation as ppt and record any color change in the fourth row box, labeled NO₂^–, of the data table.
4. Place a clean test tube in the test tube rack. Add 1 mL of the copper(II) nitrate solution. Add 1 mL of 1 M sodium nitrite solution to the test tube and swirl. Record any precipitate as ppt and record any color change in the fifth row box, labeled NO₂^–, of the data table.
5. Dispose of the test tube solutions as directed by the instructor.

Student Worksheet PDF

10604_Student1.pdf