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International Journal of Chemical Engineering and Analytical Science Vol. 2, No. 1, 2017, pp. 1-8 3 2.2. Experimental Protocol of Electrolysis A high purity sodium chloride solution (26% NaCl, 80°C) was fed to the anode compartment, to produce chlorine gas; whereas, hot water (80°C) was fed to the cathode side to produce hydrogen and sodium hydroxide. Electrolysis occurs when direct current electricity flows between anodes and cathodes, through the electrolytes (brine and water). Chlorine then generated at the anode, bubbling up through the brine, and is carried away by the collecting system. Simultaneously, the other gas generated hydrogen and is similarly collected. The hydrogen gas like chlorine, at approximately 80°C, is water-saturated when the cell exits. It was cooled at around 15°C and dehydrated using a wash bottle (500 mL) filled with the concentrated sulfuric acid (96-98%). Similarly, the chlorine gas was cooled and dehydrated. The gas production was measured by water replacement. 2.3. Theoretical Cell Voltages Electrolysis of brine is decomposed into chlorine (at the anode) and hydrogen (at the cathode) gases due to the passage of an electric current. With established reversibility and absence of cell current between the two different electrode reactions, the equilibrium cell voltage is defined as the equilibrium potential difference between the respective anode and cathode. It is described by equation (7). E=E −E (7) Equation (8) reveals the change in the Gibbs free energy ∆G of the electrochemical reaction to the equilibrium cell voltage as follows: ΔG = nF (8) Where n is the number of moles of electrons transferred in the reaction, and F the Faraday constant. In electrochemical systems there are three fundamental variables: current, potential and time [16]. The relationship between current and potential curves, gives information about the overall performance of the system. The output cell potential is the sum of all these contributions (thermodynamics, reaction kinetics, charge and mass transport) and is given by the following equation: = − + + +Σ (9) Where Eanode and Ecathode are the thermodynamic potential of anode and cathode reactions, and are the anodic and cathodic activation overpotentials, and Σ is the ohmic resistance to the charge transport through the electrolytes and the membrane. 2.4. Electrodes from Used Batteries Recycling In this investigation, the graphite electrodes were used. These electrodes are from recycling used batteries. Indeed, alkaline batteries contain manganese dioxide, graphite, steel and zinc. Thus, the used batteries Tiger Head brand (type R20 UM-1) were collected and shredded to recover the electrodes (graphite). The electrodes was cleaned in acid bath contain hydrochloric acid 0.1N. Afterwards, they were treated with an abrasive paper at their surface, in order to remove impurities. The last step was to clean the electrodes with 0.1N hydrochloric acid solution. 2.5. Calculation and Measurement The efficiency of electrochemical reactions is often expressed in terms of current efficiency. Two types of current efficiency were calculated based on Faraday’s law of electrolysis: chlorine production at the anode compartment and hydrogen production at the cathode compartment. The current efficiency (CE) was calculated as follows [1]. Current efficiency for Cl2 production at the anode compartment: CE % = !"#$%&'(% × +, × -./ × 100 (10) !) + 0×1×23 Current efficiency based on H2 production (CE6 ) at the cathode compartment: !"#$%&'(% +, -./ CE6 % = !) × + × 0×1×23 × 100 (11) Where Vproduced is gas volume produced from the cathode compartment (L/hr), Vm= 22.414 (L/mol) is the molar volume of a gas at standard temperature and pressure, T0 = 273 (K) is standard temperature, T is room temperature (K), NA = 6.022×1023 (mol-1) is the Avogadro constant, C is Coulomb, which is equal to the charge of approximately proton number of electrons (6.241 × 1018 /C), and I is the actual current applied at the electrolytic cell (ampere, C/sec). Laboratory pressure was assumed to be standard pressure (1 atm.). The voltage and current was monitored and registered by a digital multimeter (Jeulin Evolution®, R30, France). The pH was measured using a bench-top pH meter (Oakton Instruments Co., Ltd., USA) and the electrical conductivity was measured by a bench-top conductivity meter (Mectler-Toledo Co., Ltd., USA). The gas production was measured by water replacement. 3. Results and Discussion 3.1. Effect of Electrolytes Concentration on Conductivity To study the effect of concentration of electrolytes onPDF Image | Electrolysis Parameters for Chlorine and Hydrogen Production
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