Inert electrolytes do not contribute in the overall

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inert electrolytes do not contribute in the overall reaction; they are neither oxidized nor reduced [ CITATION Ray101 \l 13321 ]. In each of the half-cell, the electrode is in equilibrium with the metal ions from the electrolyte. Considering the flow for spontaneous reactions, Zn losses its electrons easier compared to that of Cu, which means that Zn is a stronger reducing agent. The electrons would flow from the zinc electrode to the copper electrode to achieve equilibrium. Basically, the spontaneous reaction occurs as a result of the different abilities of these metals to give up their electrons. Voltaic cells are important because these are usually formed into batteries which are, up to date, one of the most convenient sources of electric energy. The electrochemical cell potential is dependent upon the temperature and composition of the reaction mixture. This is relevant on the solute concentrations and the partial pressures exhibited by gases. This can all be related mathematically and is shown as:
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∆G = G o + RT ln Q wherein ∆G is the difference in free energy of a certain reaction under conditions that are nonstandard, G o is the difference in free energy of a reaction occurring under standard-state conditions, and Q is the reaction quotient. The Nernst equation can be derived from the presented equation above. It is shown mathematically as: E = E o RT nF ln Q and because of a specific connection between the voltage of a cell and its pH, the Nernst equation can be written in terms of base-10 logarithms, log Q , instead of ln Q . It will be in the form: E = E o 0.0592 V n log Q The Nernst equation is useful since it can give values of cell potentials that are occurring under nonstandard-state conditions. For instance, the galvanic cell 2 + ¿ + H 2 ( g ) + ¿ →Zn ( aq ) ¿ Zn ( s ) + 2 H ( aq ) ¿ The cell potential at 25˚C when the concentration of hydrogen ion is 1.0 M, zinc ion is 0.0010 M, and the pressure of hydrogen is 0.10 atm can be calculated using the Nernst equation. E = E o 0.0592 V n log Q 2 + ¿ Zn ¿ ( P H 2 ) ¿ + ¿ H ¿ ¿ ¿ 2 ¿ ¿ ¿ ¿ E = E o 0.0592 V n log ¿ E = 0.76 V 0.0592 V 2 log ( 0.0010 )( 0.10 ) ( 1.0 ) 2
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E = 0.88 V at 25˚C The standard reduction potential, denoted as E˚, is the voltage related to an electrode’s reduction reaction at an instance when all existing solutes are 1 M and all gases are at 1 atm. In other words, it is the measure of the tendency of a specific chemical species to be reduced. The standard reduction potential values of said chemical species can be determined experimentally. In such cases, a hydrogen electrode is used because under standard-state conditions, its potential value (E˚) is exactly zero. Meaning, its reduction and oxidation potentials are equal to zero. The hydrogen electrode is called the standard hydrogen electrode (SHE). With the use of a voltaic cell wherein the other side is a SHE, it is possible to determine unknown reduced species [ CITATION Sta19 \l 13321 ]. Standard reduction potentials are useful in the determination of the standard cell potential. By convention, E˚ of a cell is given by E cell o = E cathode o E anode o wherein both the E cathode o and the E anode o are the values of standard reduction potentials of the electrodes [ CITATION Jes19 \l 13321 ].
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