Standard Reduction Potential and The Chemical Equation Lab Report

 

 

Many chemical reactions can be classified as Oxidation-Reduction reactions. These reactions involve either the partial transfer of electrons (displacement of electron density) or complete rtransfer of electrons from the species being oxidized to the species being reduced. If the species involved are touching each other the electrons can transfer directly. However, if the species are separated from each other, the reaction will only take place if there is a path for the electrons to travel. The species of the two that has a greater tendency to lose electrons will lose its electrons (oxidation, which occurs at the anode) and these loose electrons will remain there. The species of the two with a lower tendency to lose electrons will keep its electrons (cathode). The free electrons will travel through a wire from greater concentration (at the anode) to the lower concentration (cathode has fewer free electrons). These electrons will travel from the anode to the cathode. In order to maintain neutrality through the system, negative charge must travel back toward the anode through a salt bridge or through an ionic solution. When these electrons through travel through a wire, we create a voltaic cell (also known as a galvanic cell), or a battery. Electrons will travel from the “less positive” or negative electrode (anode) to the “more positive” electrode (cathode). We label the “more positive” cathode positive and the “less positive” anode negative.

We can measure the force that drives the electrons with a voltmeter, measuring this potential difference in volts. This force is dependent on the combination of the oxidation half-reaction and the reduction half-reaction, occurring at the electrodes. By convention, the red lead on a voltmeter is attached to the cathode (+ or “more positive”), where reduction takes place, and the black lead is attached to the anode (- or “less positive”), where oxidation takes place. The electrons will travel from the anode to the cathode (toward the + electrode). If in fact the voltmeter is connected in this manner and the electrons are traveling through the wires from the black lead to the red lead, the voltage reading will be positive. If the leads are attached in reverse, the voltmeter will read a negative voltage. However, it is not possible to have a negative voltage and this only indicates the leads are connected in reverse. Regardless, the voltage is always a positive value, no matter what the voltmeter indicates.

We can write the half-reactions that occur at each electrode. A common voltaic cell is the Daniell cell (Figure 1), first studied in 1836 by Frederic Daniell. It has two separate half cells, connected with a salt bridge which allows the electrical current circuit to be completed by enabling exchange of ions between solutions between solutions in the half cells.

DiagramDescription automatically generated

Figure 1: Daniell cell is made up of a zinc anode in a zinc sulfate solution, where oxidation occurs. The electrons travel to the cathode, a copper electrode in copper sulfate solution. The electrons left on the anode as the zinc ions enter the solution travel through a voltmeter and solution neutrality is maintained by ion transfer through the salt bridge. From https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/20%3A_Electrochemistry/20.3%3A_Voltaic_Cells (accessed 10/22/20).

 

The oxidation – reduction half-reactions describe the half-reactions in each half-cell. The oxidation half-reaction (equation 1) shows the zinc metal being converted into zinc ion. Note that this is reverse of the reduction half-reaction that is given in tables of reduction half reactions and has an opposite voltage. In the other cell, the reduction of copper(II) ion occurs (equation 2), producing copper metal which collects on the electrode.

Zn → Zn2+ + 2eE° = 0.76 V (1)

Cu2+ + 2e → Cu E° = 0.34 V (2)

The voltage of the cell is the sum of the standard reduction potential and standard oxidation potential (these potentials written as reduction half reactions can be found in Appendix D (pg A-21) in the textbook). When we do the calculation or measure the voltage of the Daniell cell under standard conditions, we get 1.1 V. So, we can report the overall reaction occurring in the cell as (Equation 3).

Zn + Cu2+ → Cu + Zn2+ E° = 1.10 V (3)

The reverse of the reaction in equation 3 would have a voltage of -1.10 V and would not be spontaneous. In other words, electrons will not flow from the copper metal to the zinc. However, if attach these to a power supply, we can force the electrons in the reverse direction. This is called an electrolytic cell.

These standard voltages are relative to each other and the numeric value or each reduction half reaction is relative to the hydrogen (H2) half reaction, which is described in Equation 4.

2H+ + 2e → H2 E° = 0 V (4)

This lab has two parts. The goal of the first part is to identify the metal strips in your lab kit. You have a piece of zinc, aluminum, and iron as well as a bare copper wire. You will create voltaic cells in both acid, which will measure the oxidation and reduction of the metals, and in a salt solution, which will create an air cell.

The second part of the lab is the electrolysis or water. Direct current from a battery are connected to metal electrodes that are separated submerged in an ionic solution. The current is caried through the wires and ionic solution to complete the circuit. Reactions will occur at the electrodes and the liquid in the solution is electrolyzed. In an electrolytic cell the electrons are forced to travel in the direction of the battery. Reduction will occur at the electrode connected to the negative pole of the battery. The species reduced will be a metallic cation, H+ ion, or water depending on which requires the least energy. Our electrolysis will occur in an acidic solution, so the H+ ion will be reduced, as described in Equation 4.

Prelab Questions:

  1. What is the standard reduction potential and the chemical equation of the half-reaction for the reduction of aluminum ion?
  2. What is the standard reduction potential and the chemical equation of the half-reaction for the reduction of iron ion?
  3. What is the net chemical equation of the reaction in a voltaic cell for the reaction of aluminum in Al(NO3)3 solution and iron in Fe(NO3)2? Combine the half-reactions in questions 1 and 2 to produce a positive voltage.
  4. What is the voltage of the reaction described in #3?
  5. Describe an air cell.
  6. What is the reduction half-reaction and its standard voltage that occurs in an air cell?

Safety: Wear safety goggles.

Chemical Disposal: All chemicals in this experiment can be disposed of in the sink. The metals can be returned with your kit.

Materials:

Vinegar (5% acetic acid)

Sodium Chloride

250 mL beaker

Zinc Metal Strip

Iron Metal Strip

Aluminum Metal Strip

Copper Wire

Multimeter

Insulated copper wires

Test tube

Procedure:

Part 1

  1. Attach the red lead to the Voltage plug (center) on the multimeter. Attach the black lead to the COM plug (right) on the multimeter. Turn the multimeter to read 20 Volts. See Figure 2 for proper set up.

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Figure 2: The dial on the multimeter should be set at 20V when in use. Be sure to turn it off when you are not using it. The red lead should be attached to the center, V plug, and the black lead should be attached on the right, COM plug. The alligator clip leads should be attached to the metal and the leads attached to the meter. Do not let the electrodes touch each other when you are making the reading.

Acetic Acid Cell

  1. Add 150 mL of vinegar to your 250 mL beaker.
  2. Label each of the metal strips (#1, #2, and #3) and describe your observations in the data table.
  3. Using different combinations of the four metals (the three metal strips and the bare copper wire), set up voltaic cells and record the voltages. Put both electrodes into the solution and read the voltmeter (Figure 2). Be sure to note which electrode the red lead is attached to. The reading may jump around a lot. You may want to clean the electrode with a piece of steel wool or sandpaper or crumbled up printer paper. Do not let the electrodes touch each other when you are making your reading.
  4. Rinse off your metals and clean your beaker.

NaCl cell

  1. Add 150 mL of water to your 250 mL beaker
  2. Dissolve 2 scoopfuls of sodium chloride in the water
  3. Using different combinations of the four metals (the three metal strips and the bare copper wire), set up voltaic cells and record the voltages. Be sure to note which electrode the red lead is attached to.
  4. Rinse off your metals and clean your beaker

Part 2 – Electrolytic Cell

  1. Dissolve about 2 g of NaCl in 75 mL vinegar in your 250 mL beaker
  2. Fill a test tube completely with vinegar. Cover the top with your finger (if your finger is large enough) and invert the tube in the beaker. Do not uncover the tube until it is below the surface of the vinegar in the beaker. If your finger is not large enough, take a piece of paper towel folded over in a couple of layers, wet it with water, and cover the tube with the paper, holding it tight. Invert the tube into the beaker and remove the paper after it is below the level of the liquid. If the liquid comes out of the tube, refill it with the solution in your beaker and try again.
  3. With a scissors, carefully remove about 1 cm of insulation from each end of both insulated copper wires. Twist the stranded wire together after the insulation is removed. See the introduction video.
  4. Insert one end of the copper wire bent into the test tube. The part with the insulation removed must be completely in the test tube (see Figure 3).

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Figure 3 The electrolytic cell. The red lead is on the negative terminal of the battery (labeled on the battery and the knurled terminal) and attached to the electrode that is inserted into the test tube. The black lead is on the positive terminal (labeled on the battery and the flat terminal) and attached to the electrode that is in the beaker.

  1. Put the other copper wire into the beaker.
  2. Attach the alligator clip leads to the terminals on the 9V battery. Attach the red wire to the negative terminal and the black wire to the positive terminal. Note: The positive terminal is identified on the side of the batter (flat button terminal, not the curled one).
  3. Attach the other end of the red lead wire to the copper wire that leads into the test tube.
  4. Attach the other end of the black lead wire to the copper wire that is loose in the beaker. Observe the reaction for several minutes. If hooked Cu lead is not bubbling in the test tube, be sure the black and red leads are attached correctly. You may need to switch the black and red battery terminals.
  5. Allow the reaction to run for about 45 minutes and make observations periodically. Look closely at changes in the test tube, changes to the wires, and changes to the color of the solution in the beaker (compare it to water in the 150 mL beaker. You may need to use a white paper background to see the color more clearly. Record your observations.
  6. After at least 45 minutes remove the lead wires from the copper wires. If gas collects in the test tube, without lifting the tube out of the solution in the beaker, remove the copper wire from tube and the beaker. Record your observations
  7. Leaving the test tube upside down, remove the test tube from the liquid. Do not turn it over.
  8. Light a match and quickly turn the tube over and insert the lit match into the mouth of the tube. Record your observations.

Part 1:

Data And Observations:

Metal Strip Observations

#1

 

#2

 

#3

 

Acetic Acid Cell

Lead Color

Electrode

Lead color

Electrode

Voltage

 

Copper

 

#1

 
 

Copper

 

#2

 
 

Copper

 

#3

 
 

#1

 

#2

 
 

#1

 

#3

 
 

#2

 

#3

 

NaCl Cell

Lead Color

Electrode

Lead color

Electrode

Voltage

 

Copper

 

#1

 
 

Copper

 

#2

 
 

Copper

 

#3

 
 

#1

 

#2

 
 

#1

 

#3

 
 

#2

 

#3

 

Part 2 observations:

Post-Lab Questions

Part 1

  1. Identify the Anode and Cathode in each cell in the acetic acid solution.

Anode or Cathode

Electrode

Anode or Cathode

Electrode

 

Copper

 

#1

 

Copper

 

#2

 

Copper

 

#3

 

#1

 

#2

 

#1

 

#3

 

#2

 

#3

  1. Identify the Anode and Cathode in each cell in the NaCl solution.

Anode or Cathode

Electrode

Anode or Cathode

Electrode

 

Copper

 

#1

 

Copper

 

#2

 

Copper

 

#3

 

#1

 

#2

 

#1

 

#3

 

#2

 

#3

  1. Write the chemical equations of the half-reactions as well as the net reaction for each of the following voltaic cells under standard conditions (they will have positive voltages) and give the standard voltage. Remember, one will have to be an oxidation and one will have to be a reduction.
    1. Cu and Al
    2. Cu and Zn
    3. Cu and Fe
    4. Al and Zn
    5. Zn and Fe
  2. What are the identities of your metals? Justify each identity with your data. Be sure to correlate your voltages with expected voltages from all of your data from both the acetic acid cells and the NaCl cells. You need to devise a logic to justify your identifications by comparing your results.

#1

#2

#3

  1. Why do some of the combinations have the same voltages in the two solutions (identify them) and some have the different voltages (identify them)? Consider the possibility of an air cell. Be specific about the reaction taking place at the anode and cathode in each set.

Part 2

  1. Identify the gas produced in the test tube.
  2. Write the chemical equation of the half-reaction that occurred at the cathode in the test tube.
  3. What ion caused the color in the solution in the beaker?
  4. Write the chemical equation of the half-reaction that occurred at the anode in the beaker.
  5. If you collected 28.92 mL of gas in the test tube at a temperature of 22.5°C and a pressure of 747 mmHg, how many moles of the gas are produced? Show all your work.
  6. If the anode was an unknown metal, rather than the copper wire we used, and the anode was oxidized so that the mass of the anode decreased by 0.0635g during the electrolysis. We do know that the metal forms a 2+ ion and has the half-reaction of its oxidation can be described by the chemical equation below.

M(s) → M2+(aq) + 2e

Write the net chemical equation of the reaction, combining the chemical equations of the two half-reactions, and do stoichiometry. What is the molar mass of the unknown metal?

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