Materials Included In Kit

Additional Materials Required

Safety Precautions

Silver nitrate solution is toxic by ingestion and irritating to body tissue. It also stains skin and clothing. Lead nitrate solution is a possible carcinogen. It is also moderately toxic by ingestion and inhalation, and is irritating to eyes, skin and mucous membranes. Zinc nitrate solution is slightly toxic by ingestion and corrosive to body tissue/severe tissue irritant. Copper(II) nitrate solution is slightly toxic by ingestion; it is irritating to skin, eyes and mucous membranes. Iron(III) nitrate solution is corrosive to body tissue. Magnesium nitrate solution is a body tissue irritant. Wear chemical splash goggles and chemical-resistant gloves and apron. Remind students to wash hands thoroughly with soap and water before leaving the laboratory. Please review 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 lead nitrate solution may be disposed of according to Flinn Suggested Disposal Method #27f. The silver nitrate solution may be disposed of according to Flinn Suggested Disposal Method #11. The remaining solutions may be disposed of according to Flinn Suggested Disposal Method #26b.

Lab Hints

• When using the voltmeter, be sure that connections are tight and that the metal electrodes are shiny. The maximum voltage will be less than 3 volts. Students may need help in setting the ideal voltage range on the meter. Use the smallest range that gives a reading on the scale. Be sure to use the connections for DC voltage. If the voltmeter gives a negative voltage, reverse the connections. When a positive voltage is obtained, the electrode connected to the positive terminal is the cathode and is undergoing reduction, while oxidation is occurring at the electrode connected to the negative (or common) terminal, the anode.
• A strip of filter paper soaked in potassium nitrate solution is used for the salt bridge. Use a fresh strip of paper for each measurement.
• In Part 3, solutions are diluted by counting the drops of solution and distilled water used. Caution students to be consistent in their technique when measuring drops and to hold the Beral-type pipet at the same angle throughout. Remind students to remove any air bubbles from the pipet before counting drops.
• Students may need help in understanding the calculations using the Nernst equation. Remember that solid substances do not appear in equilibrium expressions, and only the ion concentrations are in the calculations for the values of the equilibrium constant, K, or the reaction quotient, Q.
• An alternative procedure for Part 3 is the determination of the formation constant of Cu(NH₃)₄²⁺. The formation equation is:

  Cu²⁺(aq) + 4NH₃(aq) → Cu(NH₃)₄²⁺(aq)

  The formation constant, K_f, is equal to:
  {13803_Hints_Equation_4}
  Use a solution of 10 mL of 6.0 M NH₄OH combined with one drop of 1 M Cu(NO₃)₂. Pour this solution in a well and measure the voltage versus the zinc electrode. A representation of the cell is:

  Zn(s) | Zn²⁺(1.0 M) || Cu²⁺(unknown) | Cu(s)

  Because of the excess ammonia, the equilibrium concentration of ammonia is still 6.0 M, and all the Cu²⁺ in solution is present as the Cu(NH₃)₄²⁺ complex. The Cu(NH₃)₄²⁺ concentration can be calculated from the initial Cu(NO₃)₂ concentration (1.0 M), the initial volume of 1 drop (approximately 0.025 mL), and the final volume (10 mL). The uncomplexed Cu²⁺ concentration is calculated from the measured cell potential using the Nernst equation. The accepted value of K_f is 2 x 10¹³.

Teacher Tips

• Metals: Cut 12 2-cm pieces of metal foil or ribbon for each metal—copper, lead, magnesium, silver and zinc. An iron nail is used for the iron electrode.
• Voltmeter: The most convenient type of voltmeter to use is a digital voltmeter that automatically sets the voltage range, but an analog meter may also be used. Review the specific operation of the voltmeters before the students begin the lab procedure.
• Filter paper: Cut the filter paper into strips approximately 4 cm long x 0.6 cm wide. Soak for a few seconds in the potassium nitrate solution.
• 24-Well Plate: The 24-well reaction plates are most convenient for preparing the small cells. However, small beakers may also be used.

Answers to Prelab Questions

The following data were measured using a nickel electrode as the reference standard:
Potential, volts
Cu²⁺(aq) + 2e^– → Cu(s) +0.62
Ni²⁺(aq) + 2e^– → Ni(s) +0.00
Fe²⁺(aq) + 2e^– → Fe(s) –0.15
Al³⁺(aq) + 3e^– → Al(s) –1.38

1. Which ion is most easily reduced?

   Cu²⁺ (most positive potential)

2. Which metal is most easily oxidized?

   Al (most positive oxidation potential, the negative of the reduction potential)

3. The copper and aluminum electrodes are connected in a battery.
   1. Which is the anode?

      Aluminum

   2. Which is oxidized?

      Aluminum

   3. What will be the battery voltage?

      2.00 volts

   4. Write a balanced net ionic equation for the reaction that takes place.

      3Cu²⁺(aq) + 2 Al(s) → 3Cu(s) + 2Al³⁺(aq)

4. A solution is prepared in which a trace or small amount of Fe²⁺ is added to a much larger amount of solution in which the [OH^–] is 1.0 x 10^–2 M. Some Fe(OH)₂ precipitates. The value of K_sp of Fe(OH)₂ = 8.0 x 10^–10.
   1. Assuming that the hydroxide ion concentration is 1.0 x 10^–2 M, calculate the concentration of Fe²⁺ ions in the solution.

      K_sp = [Fe²⁺] [OH^–]2 = 8.0 x 10^–10

      {13803_PreLabAnswers_Equation_5}
   2. A battery is prepared using the above solution with an iron wire dipping into it as one half-cell. The other half-cell is the standard nickel electrode. Write the balanced net ionic equation for the cell reaction.

      Ni²⁺(aq) + Fe(s) → Ni(s) + Fe²⁺(aq)

   3. Use the Nernst equation to calculate the potential of the above cell.
      {13803_PreLabAnswers_Equation_6}

Sample Data

Part 1. Data Table

Voltage of each half-cell versus the zinc electrode

Predicted and Measured Cell Potentials Part 2. Data Table Part 3. Data Table

Answers to Questions

Calculations

Part 1 

1. Write reduction equations for each metal ion, arranging the equations in decreasing order of measured potential in thefollowing table. Include zinc in the table, using 0.00 volts as the potential of the Zn | Zn²⁺ half-cell. Record the accepted standard potentials using the hydrogen electrode as standard and calculate the difference between the two standard values.
   
   Reduction Equations for Each Ion Arranged in Decreasing Order of Potential
   {13803_Answers_Table_5}
2. Use the electrode potentials from the above table to predict the voltages of the six half-cell combinations selected in Part 1, step 10. Record this value and which metal is the cathode and which is the anode in the second Part 1 Data Table for each combination. Compare the predicted and measured potentials.

Part 2
Write a balanced net ionic equation for the reaction occurring in the cell in Part 2. Record this equation in the Part 2 Data Table. Use the Nernst equation to calculate what the expected voltage should be. Record this value in the Part 2 Data Table. Compare this value to the measured voltage.

A balanced net ionic equation for the reaction is: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s).

The measured voltage for the cell: 0.89 V. Zinc is the anode, copper is the cathode.

Calculate the expected voltage of the cell:

The calculated voltage and the measured voltage agree.

Part 3

1. Write a balanced net ionic equation for the reaction occurring in the cell. Use the Nernst equation to calculate the concentration of the Ag⁺ ion. Record this value in the Part 3 Data Table.
2. Calculate the value of the solubility product of AgCl. Compare the calculated value to a reported value. Record this value in the Part 3 Data Table.
   A balanced net ionic equation for the electrode reaction is: Zn(s) + 2Ag⁺(aq) → Zn²⁺(aq) + 2Ag(s).
   The measured voltage for the cell: 0.84 V. Zinc is the anode, silver is the cathode.
   Calculate the concentration of silver ion in the half-cell:
   {13803_Answers_Equation_9}
   [Ag⁺] = 2.4 x 10^–10 mol/L
   Calculate the solubility product constant of AgCl:

   K_sp = [Ag⁺] [Cl^–] = (2.4 x 10^–10) x (1.0) = 2.4 x 10^–10

   The accepted value for K_sp of AgCl is 1.8 x 10^–10. The calculated value is close to the accepted value.

Post-Laboratory Review

1. What is an electrode potential?

   An electrode potential is the voltage generated when a half-cell is connected to a standard half-cell, usually the hydrogen electrode that is assigned a potential of zero volts. In this experiment the zinc electrode was assigned a zero volt potential.

2. Did the ranking of reduction equations agree with that in a published chart of E° values?

   Yes, the ranking of reduction equations agreed exactly with the ranking in a published chart of E° values.

3. How should the values found using the zinc electrode as a standard compare with those in the E° table that are based on the standard hydrogen electrode? Did they?

   Since zinc was assigned a zero voltage in this experiment, the voltages measured should differ from those published in a table in which the hydrogen electrode was the standard voltage by the value of the zinc half-cell potential in the standard voltage table, 0.76 volts. Most differed by about 0.6 volts, except for the magnesium, which had a difference of 1.75 volts.

4. What factors can cause a difference between experimental and reported values?

   The published values are obtained in systems with a minimal resistance. The systems used in this experiment probably have a larger resistance causing a lower voltage. If connections are not tight, lower voltages will also result.

5. What does a negative value for a standard potential indicate?

   A negative value means that the reduction reaction occurs less readily than that of the standard zinc ions.

6. How did the change in concentration of the copper ions in Part 2 affect the cell potential? Is this change in agreement (qualitatively) with that which would be predicted by Le Chatelier’s Principle? Did the calculated and measured values agree?

   The voltage for the zinc versus copper cell was 0.98 volts when solutions were one molar. The equation for the reaction is: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). A decrease in Cu²⁺ concentration should drive the reaction to the left, according to Le Chatelier. This would have the effect of lowering the voltage. In the cell in which the copper ion concentration is 0.0010 M, the measured voltage was lowered to 0.89 volts as predicted by Le Chatelier’s Principle. The calculated and measured values did agree.

7. Explain how the AgCl solubility product was determined.

   To determine the solubility product of silver chloride, a 1.0 M solution of chloride ion was obtained. A trace of silver ion was added to this solution. Almost all of the silver ions precipitated. An electrochemical cell was constructed with the silver ions and silver metal acting as one half-cell and the zinc half-cell as the other. The voltage was measured and the Nernst equation used to calculate the silver ion concentration. The chloride ion concentration was 1.0 M since such a small amount precipitated. These values were substituted into the solubility product expression to calculate the value of K_sp.

Introduction

Oxidation–reduction reactions form a major class of chemical reactions. From the reactions of oxygen with sugars, fats, and proteins that provide energy for life to the corrosion of metals, many important reactions involve the processes of oxidation and reduction. In this three-part lab, these reactions are studied by constructing various electrochemical cells and measuring the voltage generated. From these measurements, a reduction series is generated, the concentration of copper ions in solution determined, and the K_sp of silver chloride calculated.

Concepts

• Half-cell reaction
• Standard reduction potential
• Nernst equation
• Spontaneous reaction

Background

An electrochemical cell results when oxidation and reduction reactions occur, and the resulting electron transfer between the two processes occurs through an external wire. The oxidation and reduction reactions are physically separated from each other and are called half-cell reactions. A half-cell can be prepared from almost any metal in contact with a solution of its ions. Since each element has its own electron configuration, each element develops a different electrical potential, and different combinations of oxidation and reduction half-cells result in different voltages for the completed electrochemical cell.

The standard reduction potential is the voltage that a half-cell, under standard conditions (1 M, 1 atm, 25 °C), develops when combined with the standard hydrogen electrode, which is arbitrarily assigned a potential of zero volts. A chart of reduction half-cell reactions, arranged in order of decreasing standard reduction potential, shows the relative ease of reduction of each substance listed. The more positive the reduction potential, the easier the reduction. A spontaneous cell (a battery) can be constructed if two half-cells are connected internally using a salt bridge, and externally using a metallic connector. In an electrochemical cell, the reaction listed in the standard reduction potential chart with the more positive voltage occurs as a reduction, and the reaction listed with the less positive voltage reverses and occurs as an oxidation reaction. The cell voltage can be found by adding the voltages listed in the table, with the value of the voltage for the oxidation reaction becoming the negative of its reduction reaction voltage.

As an example, consider a cell made up of copper and aluminum half-cells.

Cu²⁺(aq) + 2e^– → Cu(s)    E° = 0.34 V
Al³⁺(aq) + 3e^– → Al(s)      E° = –1.66 V

The copper reaction has the more positive potential and remains a reduction reaction. The aluminum reaction with the less positive (more negative) potential is reversed and becomes an oxidation reaction. Its potential is now an oxidation potential:

Al(s) → Al³⁺(aq) + 3e^–    E° = +1.66 V

The reduction potential and the oxidation potential are added to find the cell voltage:

3Cu²⁺(aq) + 2Al(s) → 3Cu(s) + 2Al³⁺(aq)
E°_cell = E°_reduction + E°_oxidation
E°cell = 0.34 V + 1.66 V = 2.00 V

A positive value for E°_cell indicates that the oxidation–reduction reaction, as written, is spontaneous.

A cell representation such as the following: Zn(s) | Zn²⁺(1.0 M) || Cu²⁺(0.0010 M) | Cu(s) means that a cell is constructed of zinc metal dipping into a 1.0 M solution of Zn²⁺. The symbol “|” refers to a phase boundary. The symbol “||” indicates a salt bridge between the zinc ion solution and the copper ion solution. The second half-cell is copper metal dipping into a 0.0010 M solution of copper ions. The anode is on the left (where oxidation occurs) and the cathode is on the right (where reduction occurs).

In this laboratory a “standard” table of electrode potentials is constructed. A value of 0.00 volts is assigned to the electrode made from zinc metal in a 1.0 M solution of zinc ions. The voltage values should correlate with those found in published tables, differing only by the value of E° for the standard zinc electrode. Published standard values are measured in solutions that have very small electrical resistance. The resistance of the experimental cell will probably cause a lowering of measured values from the ideal values.

The table of standard potentials assumes that all ion concentrations are 1.0 M, gas pressures are 1 atm, and temperature is 25°C. Calculations of potentials under nonstandard conditions can be made using the Nernst equation: where E = the measured cell potential, E° = the standard cell potential, R is the gas constant (8.314 J/mol•K), T is the temperature (K), n = the number of moles of electrons transferred as shown by the oxidation–reduction equation, and F is the Faraday constant (9.65 x 10⁴ C/mol). Q is the reaction quotient: the actual concentrations of products and reactants substituted into the equilibrium constant expression.

Using base 10 or common logarithms the expression can be written: Substituting for the constants 2.303, R and F, and using a temperature of 25 °C (298 K) the expression can be simplified to: A measurement of the cell potential, E, under nonstandard conditions, can be used to calculate the value of Q, which can then be used to determine unknown concentrations of ions actually present in a solution.

Experiment Overview

The purpose of Part 1 of this laboratory is to construct a table listing the reduction potentials of a series of metal ions, in order of ease of reduction. The series of microscale half-cells is constructed by placing a piece of metal into a 1.0 M solution of its ions for each metal in the series. The metals chosen are copper, iron, lead, magnesium, silver and zinc. The half-cells are connected by a salt bridge constructed of a strip of filter paper soaked in a solution of potassium nitrate. The zinc half-cell is chosen as the reference standard, and all potentials are measured with respect to the zinc electrode.

In Part 2, the Nernst equation is applied to the voltage measurement of a cell with nonstandard copper ion concentration. A solution of 0.0010 M Cu²⁺ is prepared, and the voltage of the cell: Zn(s) | Zn²⁺(1.0 M) || Cu²⁺(0.0010 M) | Cu(s) is measured. The measured voltage is compared to that calculated from the Nernst equation. In the final section, the solubility product constant of silver chloride, AgCl, is determined from the Nernst equation and the voltage of a cell in which the zinc half-cell is connected to a solution containing a trace of Ag⁺ ions in a 1.0 M solution of sodium chloride, NaCl.

Materials

Prelab Questions

The following data were measured using a nickel electrode as the reference standard:
Potential, volts
Cu²⁺(aq) + 2e^– → Cu(s) +0.62
Ni²⁺(aq) + 2e^– → Ni(s) +0.00
Fe²⁺(aq) + 2e^– → Fe(s) –0.15
Al³⁺(aq) + 3e^– → Al(s) –1.38

1. Which ion is most easily reduced?
2. Which metal is most easily oxidized?
3. The copper and aluminum electrodes are connected to form a battery.
   1. Which is the anode?
   2. Which is oxidized?
   3. What will be the battery voltage?
   4. Write a balanced net ionic equation for the reaction that takes place.
4. A solution is prepared in which a trace or small amount of Fe²⁺ is added to a much larger amount of solution in which the [OH^–] is 1.0 x 10^–2 M. Some Fe(OH)₂ precipitates. The value of K_sp for Fe(OH)₂ = 8.0 x 10^–10.
   1. Assuming that the hydroxide ion concentration is 1.0 x 10^–2 M, calculate the concentration of Fe²⁺ ions in the solution.
   2. A battery is prepared using the above solution with an iron wire dipping into it as one half-cell. The other half-cell is the standard nickel electrode. Write the balanced net ionic equation for the cell reaction.
   3. Use the Nernst equation to calculate the potential of the above cell.

Safety Precautions

Silver nitrate solution is toxic by ingestion and irritating to body tissue. It also stains skin and clothing. Lead nitrate solution is a possible carcinogen. It is also moderately toxic by ingestion and inhalation; irritating to eyes, skin and mucous membranes. Zinc nitrate solution is slightly toxic by ingestion; it is corrosive to body tissue/severe tissue irritant. Copper(II) nitrate solution is slightly toxic by ingestion and irritating to skin, eyes and mucous membranes. Iron(III) nitrate solution is corrosive to body tissue. Magnesium nitrate solution is a body tissue irritant. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part 1. Determine Reduction Potentials

1. Prepare a test cell to measure the voltage of the copper and zinc half-cells. Using a graduated Beral-type pipet, transfer approximately 2 mL of 1.0 M Zn(NO₃)₂ solution to one of the center wells of a 24-well plate. With a new pipet, place approximately 2 mL of 1.0 M Cu(NO₃)₂ in an adjacent well.
2. Polish small strips of zinc and copper metal with sandpaper or steel wool and place each metal in the appropriate well containing the solution of its ions.
3. Take a small strip of filter paper that has been soaked in KNO₃ solution, and drape it across the wells so that one end dips in the solution in each well. This acts as the salt bridge. Use a fresh strip of paper for each measurement in the procedure.
4. Use a voltmeter to measure the potential difference between the two half-cells. Connect the negative terminal of the voltmeter to the zinc electrode. Use the most sensitive scale that is practical. Make note as to which electrode is the anode and which is the cathode. When the voltmeter reads a positive voltage, the electrode connected to the positive terminal is the cathode and is undergoing reduction, while oxidation is occurring at the electrode connected to the negative (or common) terminal, the anode.
5. Prepare half-cells in other wells of the 24-well plates. Make a diagram of the order of the solutions in the wells. The other four solutions are 1.0 M Fe(NO₃)₃, 1.0 M Pb(NO₃)₂, 1.0 M Mg(NO₃)₂ and 1.0 M AgNO₃. An example is shown.
   {13803_Procedure_Figure_1}
6. Use a new pipet to transfer approximately 2 mL of each 1.0 M solution to their designated wells.
7. Polish the metals with sandpaper or steel wool so that they are shiny, and place them in the wells that contain the ion of the same metal.
8. Use fresh strips of filter paper soaked in 1.0 M potassium nitrate as salt bridges. The electrodes to be tested are:

   Ag | Ag⁺
   Cu | Cu²⁺
   Fe | Fe³⁺
   Mg | Mg²⁺
   Pb | Pb²⁺
   Zn | Zn²⁺

9. Designate the zinc electrode as the standard electrode. Measure the potential difference between the zinc electrode and each of the other electrodes. Note which terminal is the anode and which is the cathode in each case. Record the data in the Part 1 Data Table.
10. Measure the potential difference between at least six combinations of the various electrodes. Record your data, including the equation for each cell reaction, in the second Part 1 Data Table. An example of six combinations set up in the 24-well plate is shown.
    {13803_Procedure_Figure_2}

Part 2. Change the Ion Concentration
Prepare a 0.0010 M Cu(NO₃)₂ solution as follows (steps 1–3).

1. Dilute the 1.0 M Cu(NO₃)₂ to 0.0010 M. Begin by counting 18 drops of distilled water into a small test tube and then adding 2 drops of the 1.0 M Cu(NO₃)₂ solution.
2. Mix well by pouring back and forth from one test tube to another. This solution is now 0.10 M.
3. Repeat this dilution process two more times to give a final concentration of 0.0010 M.
4. Pour some of this 0.0010 M Cu(NO₃)₂ solution into one of the wells in the well plate. Add a piece of polished copper foil and measure the voltage against the standard zinc electrode. A representation of the cell is: Zn(s) | Zn²⁺(1.0 M) || Cu²⁺(0.0010 M) | Cu(s)
5. Record the data in the Part 2 Data Table.

Part 3. Solubility Product of AgCl

1. Pour 10 mL of 1.0 M NaCl solution into a 50-mL beaker.
2. Add one drop of 1.0 M AgNO₃ to the NaCl solution and stir well. Almost all of the silver ions will combine with chloride ions to precipitate AgCl. Since there is such a large excess of Cl^– ions, it can be assumed that the concentration of Cl^– is still 1.0 M. The concentration of Ag⁺, which is limited by the K_sp of AgCl, will be very small.
3. Pour some of the solution into one of the wells in the well plate and add a silver metal electrode. Measure the potential difference between this half-cell and the zinc half-cell. A representation of the cell is: Zn(s) | Zn²⁺(1.0 M) || Ag⁺(unknown M) | Ag(s)
4. Record the voltage in the Part 3 Data Table.

Student Worksheet PDF

13803_Student1.pdf