The reactivity of the metal species can be determined from the activity series of the metals which enlists metals are organized according to their reactivity (oxidation number) in single displacement reactions, from higher to lower order (Runo and peters 1993). It defines the ability of a metal to lose electrons to another element below it in the series. The donation of electrons from an electronegative atom/molecule to a less electronegative species is the basis of redox (reduction-oxidation) reactions which are one of the fundamental chemical reaction types. This electron transfer will affect proton: electron ratio of the reactants and will thus shift their respective oxidation number. Donor ion is now oxidized with higher oxidation number while the acceptor atom is in reduced state lower oxidation number.
Redox reactions form the basis of electrochemistry science. One application of the electrochemical reactions is that if the oxidizing and reducing species are connected with a salt bridge, two electrodes, and a wire, an electrochemical cell is created (Winter and Brodd 2004). A salt bridge is introduced in these cells to make electrical contact feasible between the electrodes, forming an internal circuit. The two electrodes (or half-cells) are an anode, where oxidation reaction proceeds, and cathode, the site of the reduction reaction (Bockris, Reddy, and Gamboa-Aldeco 1998). The wire serves as an electron circuit for electrons to travel from reduced species to oxidized one, creating an electrical current. The electrons will travel spontaneously unless the specific redox reaction is not thermodynamically spontaneous, producing the current. This type of chemical reactions that generate current spontaneously is typical to galvanic/voltaic cells (a type of electrochemical cells) (Schimdt-Rohr 2018). In galvanic cells, electrodes are separated to stall the redox reactions to occur spontaneously. The electrical neutrality in each half-cell is maintained by the internal circuit which causes in-flow and outflow of ions in the salt bridge. A schematic representation of a galvanic cell is depicted in figure 1. Galvanic cells are employed in batteries and other applications.
Figure 1. A scheme of a typical Galvanic cell.
In a way of example, the following equation represents a redox reaction:
Cu (s) + 2 Ag+(aq) à Cu2+(aq) + 2 Ag(s)
Equation (1)
wherein the copper (Cu) is getting oxidized in the reaction:
Cu (s) à Cu2+(aq) + 2 e-
Equation (2)
causing reduction of silver (Ag) ions:
e- + Ag+2(aq) à Ag(s)
Equation (3)
Since the final state of these metal ion is caused by each other, therefore, copper ion is a reducing agent and silver ion is an oxidizing agent. The redox reaction between these two species, if not separated by any means, is thermodynamically spontaneous.
The purpose of this report to analyze redox reactions taking place at half-cells in a series of galvanic cells created in the laboratory. Henceforth, this report proposes following hypotheses:
H: The metals selected for the electrodes of galvanic cells react simultaneously when mixed directly.
H0: Redox reaction is not spontaneous between the metal species constituting electrodes of a galvanic cell if mixed.
Method
Solutions of cuprous nitrate (Cu(NO3)2), ferrous sulphate (FeSO4), magnesium nitrate (Mg(NO3)2), and zinc nitrate (Zn(NO3)2) were used for the following experiments.
Part A: Any two of these solutions were mixed to observe the chemical process, if any, takes place between them. The visual observations were recorded.
Part B: In order to construct different galvanic cells, some half-cells were created. The combinations were decided as per the result of redox reactions carried out in part A. Those combinations were selected to construct the galvanic cell wherein the redox reaction took place spontaneously. The choice of ions for cathode and anode were made by their reactivity. Highly active species was the anode, and the lesser was the cathode. A solution of potassium nitrate (KNO3) was introduced in the salt bridge. Voltage differences between the electrodes per cell were recorded with the help of a voltage probe, in all possible ways. The cell potential for each combination was calculated as per Nernst’s equation (Stock and Orna 1989):
E = E0 – (RT/veF)ln Q
Equation (4)
where E stands for observed cell potential, E0 is theoretical/standard cell potential, R refers to the universal gas constant, T for temperature, ve is the count of electrons transferred, F is Faraday’s constant, and Q stands for the reaction quotient.
Results
Table 1. Initial observations of metals and solutions and observations after mixing.
The appearance of metal à
Cu
Fe
Mg
Zn
Appearance of solution
Reddish brown
Light grey
Grey
Silver, shiny
Cu(NO3)2
Blue color
No reaction
The reaction, color change
The reaction, deposits on the metal surface
The reaction, color change
FeSO4
Colorless
No reaction
No reaction, bubbles appear
The reaction, deposits on the metal surface
Reaction, displacement
Mg(NO3)2
Colorless
No reaction
No reaction
No reaction
No reaction
Zn(NO3)2
Colorless
No reaction
No reaction
The reaction, deposits on the metal surface
No reaction
Table 2. The redox reaction that took place during the mixing of the solutions.
Metal
Metal ion
Reaction (from equation 1)
Fe(s)
Cu2+(aq)
Fe(s) + Cu2+(aq) à Fe2+(aq) + Cu(s)
Mg(s)
Cu2+(aq)
Mg(s) + Cu2+(aq) à Mg2+(aq) + Cu(s)
Zn(s)
Cu2+(aq)
Zn(s) + Cu2+(aq) à Zn2+(aq) + Cu(s)
Mg(s)
Fe2+(aq)
Mg(s) + Fe2+(aq) à Mg2+(aq) + Fe(s)
Zn(s)
Fe2+(aq)
Zn(s) + Fe2+(aq) à Zn2+(aq) + Fe(s)
Mg(s)
Zn2+(aq)
Mg(s) + Zn2+(aq) à Mg2+(aq) + Zn(s)
Based on the above results, the following combinations were selected to construct different galvanic cells:
1. Zn(s)/ Zn2+(aq) and Cu2+(aq)/Cu(s)
2. Zn(s)/ Zn2+(aq) and Pb2+(aq)/Pb(s)
3. Cu(s)/ Zn2+(aq) and Cu2+(aq)/Pb(s)
Table 3. Potential differences across the cell in all the combinations.
Cell combination
Theoretical cell voltage (in volts, V)
Observed cell voltage (in volts, V)
Zn(s)/ Zn2+(aq)
//Cu2+(aq)/Cu(s)
1.10
1.03
Zn(s)/ Zn2+(aq) //Pb2+(aq)/Pb(s)
0.63
0.54
Cu(s)/ Zn2+(aq)
//Cu2+(aq)/Pb(s)
0.47
0.45
Discussion
The results of this lab were in accordance with the prediction. When the solutions of two metals having a difference in their reactivity, oxidation-reduction of the reactant species took place spontaneously (Table 1). This validates the hypothesis of this report. Similar findings were reported by Schmidt-Rohr in their study (2018). Further, the data showed that magnesium was the most reactive species among the group and copper was the least. Three combinations were selected. However, the observed total cell potential was less than the theoretical value (calculated according to the Nernst equation) might be because of internal resistance within the cell or circuit, whose source is unknown. Out of all the above combinations, Zn(s)/ Zn2+(aq)
and Cu2+(aq)/Cu(s) was the best combination to construct a galvanic cell.
References
Bockris, J.O., Reddy, A.K.N. and Gamboa-Aldeco, M., 1998. Modern Electrochemistry 2A: Fundamentals of Electrodics, vol. 2A of Modern Electrochemistry. New York: KluwerAcademic.
Runo, J.R. and Peters, D.G., 1993. Climbing a potential ladder to understanding concepts in electrochemistry. Journal of chemical education, 70(9), p.708.
Schmidt-Rohr, K., 2018. How Batteries Store and Release Energy: Explaining Basic Electrochemistry. Journal of Chemical Education, 95(10), pp.1801-1810.
Stock, J. T., and Orna, M. V. 1989. Electrochemistry, past and present. Washington, DC, American Chemical Society.
Winter, M. and Brodd, R.J., 2004. What are batteries, fuel cells, and supercapacitors?.
