Redox Reactions in Electrochemical Cell
Redox reactions take place in the electrochemical cell converted chemical energy into electrical energy by electron transfer electrolysis reaction. Therefore, the electrochemical or simply chemical cell is the device in which energy is produced due to the indirect redox reaction of the polar molecules and converted electricity for mankind. Galvanic and Voltaic cell reactions are examples of redox cell according to the name of the scientist, Luigi Galvani, and Alessandro Volta. The electrode potentials of redox reaction can define as the tendency of an element to lose or gain electron particles when contact with its own solution.
When electrode losses electron particles called oxidation potential and gain electrons called reduction potential of the redox reaction and these two chemical reactions reverse to each other. Therefore, the standard electrode potentials of the redox electrode calculated from the chemical equilibrium constant of these chemical reactions.
Redox Reactions Electrode Potential
When oxidation (loss of electrons) reactions take place on the redox electrode, the chemical potential of the electrode is called oxidation potential but reduction (gain of electrons) takes place on an electrode is called reduction potential. If the oxidation potential of an electrode = x volt, then its reduction potential = -x volt. For example, two half cells chemical reactions of ferric ion and ceric ion in the presence of dilute sulfuric acid show two electrode potentials in balancing chemical equations of the redox reaction. Therefore, Ce+4 + Fe+2 ⇄ Ce+3 + Fe+3 and half cell reactions with electrode potentials, Ce+4 + e ⇄ Ce+3 (E0 = 1.44 volt) and Fe+2 ⇄ Fe+3 + e (E0 = -077 volt).
Electrode Potential in Chemical Cells
When the metal rod suspended in a solution of its ions having 1M concentration and 298K temperature, the electrode potential of the half cell defines as the standard electrode potential of the redox reaction. If gas molecules are involved, the standard pressure (1 atm) should be used for this calculation. Standard electrode potentials of oxidation-reduction or redox reactions are represented by two types, E0ox and E0red.
Standard redox electrode potentials can be derived from chemical equilibrium constant and thermodynamics free energy changes reactions by Nernst equation in the standard state of chemistry or physics. Therefore, ΔG0 = – RT lnk = -2.303 RT logk, where ΔG0 = standard free energy. Again ΔG0 = -nFE0, where, E0 = standard electrode potential. Hence, E0 = (RT/nF) lnk = (0.059/n) lnk. Normal hydrogen electrode reactions represented as, H2 gas, 1atm (pt) | H+ (a = 1), in which pure hydrogen atom in gas phases at unit pressure is kept in contact with hydrogen ion of unit activity uses as a standard redox electrode. The potential of the normal hydrogen electrode is taken zero at all temperatures and uses as the reference for calculating electrode potential.
Electrochemical Daniell cell
Electrochemical equations of Daniell cell examples provide the quantitative basis of the redox reactions. Chemicals are arranged in this cell to produce electric potential energy out of spontaneous oxidation-reduction (redox) reactions. Daniell cell is an example of this type of electrochemical cell. Therefore, the electron charge transferred from the zinc atom to the copper atom in this cell reaction consists of two-electrode cathode and anode. Therefore, at the Zinc electrode, Zn (s) → Zn+2 (aq) +2e and at the copper electrode, Cu+2 (aq) + 2e → Cu (s). Electrode, where oxidation (loss of electrons) occurs, is called the anode. Conversely, the electrode where reduction (addition of electrons) occurs called the cathode in learning chemistry.
Examples of Redox Reactions in Chemistry
Ceric sulfate is a very good oxidizing agent used in the redox titration reactions process. The ionic half cell reaction of ceric sulfate, Ce+4 → Ce+3 + e. Ce+4. Therefore, during the redox reactions, exists as an anionic complex in the sulfuric acid solution. Hence the formal redox reaction potential of ceric ion in sulfuric acid solution = 1.42 V.
Crystalline potassium permanganate widely used oxidant in the laboratory with high redox reaction potential. Permanganate dissolve in benzene solution uses as a potential oxidant for organic hydrocarbon and also oxidizes hydrogen peroxide. Acidic permanganate solution very useful to define redox titration because no special indicator is needed to note the endpoint of the chemical reaction. Therefore, the redox reactions of permanganate in acid solution or low pH solution represented as, 2MnO4– + 8H+ + 5e ⇄ Mn+2 + 4H2O, where, E = +1.51 V.
Electrolysis Reaction in Chemistry
An electrolytic cell is a device where a non-spontaneous chemical reaction (redox chemistry) takes place by an external source of electricity. The electrolysis process occurring in the electrolytic cells. The very useful application of electrolysis is found in commercial electroplating and purification of the periodic table elements. For copper plating material molecule obtained from our environment first cleaned and dipped into the suspension of aqueous copper sulfate solution. Therefore, the electrolysis or redox chemical reaction takes place at the two electrodes. Cu+2 + 2e → Cu (cathode) and Cu → Cu+2 + 2e (anode).
Redox Reactions in Fuel cells
Fuel cell reactions are readily electrochemical redox cell reactions in which the chemical energy available from the specific heat or fuel like methane, hydrogen, etc converted into electrical energy by chemical reactions. At the cathodes of the cell, oxygen is dissolved to form OH– ions.
Such cells provide power at comparatively low cost and low maintenance. The efficiency depends particularly on the surface of the electrodes, which serve as the chemical catalyst for the attainment of equilibrium. The electric polarization due to the adoption and poisoning of the redox electrode reduces the cell performance considerably. Therefore, the density of the electrolyte in redox fuel cell reaction and the temperature has a great influence on the e.m.f of the fuel cell.