Experiment Seven - Electrochemical Cells

Electrochemical Cells Lab Lance A. Schell Park Hill South High School Block One Advanced Placement (AP) Chemistry Thursday, April 30, 2009 Abstract In this experiment, the purpose was to prepare several different half-cells using zinc as the standard electrode, connect them to find the voltages they generated, and construct a table of relative electrode potentials. This was done by predicting the electrode potentials as well as actually measuring the various electrode potentials using a voltmeter and a 24-well plate. In the results of the experiment, the electrodes followed the general rule of thumb in that the more negative an electrode potential was, the more it favored the reverse reaction and was thus the cathode, while the more positive an electrode potential was, the more it favored the forward reaction and was thus the anode. In the first part of the experiment, zinc was the anode in all five cases (silver, copper, iron, magnesium, and lead), and resulted in four positive potentials and one negative potential – 1.32 V, 0.94 V, 0.43 V, -0.65 V, and 0.40, listed with respect to order of cells performed (indicated above). In the second part of the experiment where the cells were chosen from a pool of varying combinations, the same general rule of thumb held true. The only difference that was noticed between the two sections was the differing potential of the zinc and iron cell, which was 0.43 V in the first and 1.18 V in the second, and the differing potential of the zinc and lead cell, which was 0.40 V in the first and 0.36 V in the second. The zinc and lead cell remains close enough for a negligible difference, while the differing potentials of the two measures of the zinc and iron cell could be due to the time at which the readings were read (in reference to the time the salt bridge had been connecting the two electrodes before the voltmeter was connected, as the voltage was read immediately upon connection).

Introduction An electrochemical cell comes from the ambiguous conjugation of the words electrochemical, meaning part of a science that deals with the relation of electricity to chemical changes and with the interconversion of chemical and electrical energy1, and cell, meaning a receptacle containing electrodes and an electrolyte either for generating electricity by chemical action or for use in electrolysis.2 The cell itself is usually always in the same solution, but not always the same container. A cell such as a Duracell® battery exists in one cell with a separator to separate the anode and cathode ends, while a cell such as a camping battery exists in the same solution in two different containers with a salt bridge connecting both containers of solution. Aside from the use and application in batteries, the electrochemical cells and the concept behind them allow scientists to conduct tests to prepare conditions and constants for lab experiments, such as testing various work surfaces for static electric charges [which in certain experiments can be detrimental to the test subject(s)], evidenced in a document published with the World Intellectual Property Organization (WIPO).3 The first coinage of the term “electrochemical cell” or “voltaic cell” was around 1771, when Luigi Galvani observed a twitching of a frog as the frog was dissected, leading him to the conclusion that “animal electricity” remained in the frog, even after death.4 While scientifically incorrect, the concept has brought about the concept of electrochemical cells through the years. It was not for twenty more years (around 1791) before a scientist by the name of Alessandro Volta made the conclusion that the spasms were a result of two pieces of metal connecting via the bloodstream of the frog (a cell utilizing blood as the solution).5 There are many types and variants of electrochemical cells in the world today, and usually refer to the standard electrode potential (Eº) which is where hydrogen is used as the anode and compared against (the hydrogen

standard potential in that case is 0.00 V). Other standard electrodes may be used, with varying results, such as platinum or zinc. Zinc was first used around the time John Frederick Daniell developed a voltaic cell in 1836, where zinc was used as the anode and copper as the cathode.6 The hence referred to “Daniell Cell” held the two electrodes in the same container and solution, separated by a porous barrier that was small enough to allow transmittance of SO42- ions to flow from one side to the other, but blocked the passage of the Zn2+ and Cu2+ ions.6 Without the movement of the sulfate ion, there would be no external electric current generated. Electrochemical cells have since played a role and continue to play a big role in an abundance of various industries in the world today. The SI unit for measuring electrode potential remains the same (the volt) which is equal to one joule per coulomb. The concept to remember is that the oxidation is the driving force which produces a finite amount of energy. After the cell has been allowed to run, the potential constantly decreases and when the potential eventually reaches zero, the cell itself is in equilibrium. Without a salt bridge in a cell with electrodes in two separate containers, the potential will also show zero, meaning there is no potential as there is no flow of ions (the salt bridge maintains this flow as well as the charge of each side of the cell). Experimental The experiment was taken from a handout published by Flinn Scientific.7 Modifications to the procedure included not performing parts three through five (only parts one and two), and the use of potassium nitrate in the salt bridge instead of 6M ammonia. Data and Calculations To reiterate the concept of electrochemical cells, the standard potential (generally when dealing with hydrogen) may be found using the equation:𝐸°𝐶𝑒𝑙𝑙 = 𝐸𝐶𝑎𝑡 ℎ𝑜𝑑𝑒 + 𝐸𝐴𝑛𝑜𝑑𝑒 . This

allows us to calculate an estimated potential based off of a standard reduction potentials table (included in tables). This equation is used to find the estimated potentials of each of the cells in the experiment (zinc and iron cell shown below). 𝑬°𝑪𝒆𝒍𝒍 = 𝑬𝑪𝒂𝒕𝒉𝒐𝒅𝒆 + 𝑬𝑨𝒏𝒐𝒅𝒆 𝑬°𝑪𝒆𝒍𝒍 = 𝑬𝑰𝒓𝒐𝒏 + 𝑬𝒁𝒊𝒏𝒄 𝑬°𝑪𝒆𝒍𝒍 = 𝟎. 𝟕𝟕 𝑽 − 𝟎. 𝟕𝟔 𝑽 𝑬°𝑪𝒆𝒍𝒍 = + 𝟎. 𝟎𝟏 𝑽 The other main calculation this experiment involves subtracting the standard zinc potential from the standard potential for the respective cell, per the equation below. Again, an example using the zinc and iron cell is shown. 𝐸° − 𝐸𝑍𝑛 𝟎. 𝟎𝟏 𝑽 − 𝟎. 𝟒𝟑 𝑽 − 𝟎. 𝟒𝟐 𝑽 Conclusion In conclusion, several different half-cells were prepared and connected together, and a table of relative electrode potentials was constructed based on the information. The general rule of thumb held true in this experiment in that the more negative the potential was, the more it favored the reverse reaction and was thus the cathode. The more positive the potential was, the more the electrode favored the forward reaction and was thus the anode. The only error reported was the differing measurements of the zinc and iron cell in parts one and two. There was no percent error applicable, and thus the experiment was completed.

Graphs and Figures

Relative Electrode Potentials (EZn) 1.5 1.32 1 0.94 0.5 0.43

0.4

0 Zn Vs. Ag

Zn Vs. Cu

Zn Vs. Fe

Zn Vs. Mg

-0.5

Zn Vs. Pb

-0.65

-1

Relative Electrode Potentials (Choice Cells) 2.50

2.00 1.94 1.50

1.00

1.60 1.18

0.50 0.51 0.36

0.25 0.00 Fe Vs. Zn

Cu Vs. Mg

* Zinc is anode in all cases.

Fe Vs. Cu

Pb Vs. Cu

Ag Vs. Mg

Pb Vs. Zn

Tables Standard Reduction Potentials In Aqueous Solution At 25 °C Half-Reaction Equation Standard Electrode Potential [Eº (V)]* Ag+ + e-  Ag(s) 0.80 3+ 2+ Fe + e  Fe 0.77 Cu2+ + 2e-  Cu(s) 0.34 Pb2+ + 2e-  Pg(s) - 0.13 Zn2+ + 2e-  Zn(s) - 0.76 Mg2+ + 2e-  Mg(s) - 2.37

Cell Zn Vs. Ag Zn Vs. Cu Zn Vs. Fe Zn Vs. Mg Zn Vs. Pb

Relative Electrode Potentials Voltage (V) Anode + 1.32 Zinc (Zn) + 0.94 Zinc (Zn) + 0.43 Zinc (Zn) - 0.65 Zinc (Zn) + 0.40 Zinc (Zn)

Reduction Equation Ag+ + e-  Ag(s) Cu2+ + 2e-  Cu(s) Fe3+ + e-  Fe2+ Mg2+ + 2e-  Mg(s) Pb2+ + 2e-  Pg(s)

Reduction Equations Electrode Potentials (EZn) Accepted Potentials* 1.32 V 0.80 V 0.94 V 0.34 V 0.43 V 0.77 V - 0.65 V - 2.37 V 0.40 V - 0.13 V

Anode Iron (Fe) Copper (Cu) Iron (Fe) Lead (Pb) Lead (Pb) Silver (Ag)

Cathode Silver (Ag) Copper (Cu) Iron (Fe) Magnesium (Mg) Lead (Pb)

Predicted and Measured Differentials Cathode Equation Predicted* 2+ 3+ Zinc (Zn) Fe(s) + Zn  Fe + Zn(s) 0.01 V Magnesium (Mg) Cu(s) + Mg2+  Cu2+ + Mg(s) - 2.03 V Copper (Cu) Fe(s) + Cu2+  Fe3+ + Cu(s) 1.11 V Copper (Cu) Pb(s) + Cu2+  Pb2+ + Cu(s) 0.21 V 2+ 2+ Zinc (Zn) Pb(s) + Zn  Pb + Zn(s) - 0.89 V Magnesium (Mg) Ag(s) + Mg2+  Ag+ + Mg(s) -1.57 V

* Note: Predicted values are based off of standard hydrogen electrode values

𝑬𝒁𝒏 − 𝑬° 0.52 V 0.60 V - 0.34 V 1.72 V 0.53 V

Measured 1.18 V 1.60 V 0.25 V - 0.51 V - 0.36 V - 1.94 V

References - Merriam-Webster. http://www.merriam-webster.com/dictionary/electrochemical - Merriam-Webster. http://www.merriam-webster.com/dictionary/cell - World Intellectual Property Organization (WIPO). http://www.wipo.int/pctdb/en/wo.jsp?wo=2002025249 - California Energy Commission. http://www.energyquest.ca.gov/scientists/galvani.html - Hebrew University of Jerusalem. http://chem.ch.huji.ac.il/history/volta.htm - Georgia State University (GSU) Department of Physics. http://hyperphysics.phyastr.gsu.edu/hbase/chemical/electrochem.html - Flinn Scientific, Inc. 1996. Handout “Electrochemical Cells”.