Electrolysis electrode half-reactions

Problems Involving Stoichiometry of Electrolysis Problems based on Faraday s law often ask you to calculate current, mass of material, or time. As we said, the electrode half-reaction provides the key to solving these problans because it is related to the mass for a certain qirantity of charge. 

Water electrolysis is a method to generate hydrogen. The water electrolysis process can be treated as a superposition of concurrent or sequential electrochemical reaction occurring in the vicinity of electrodes (half reactions) whose overall effect is to split water molecule under the influence of a direct electric current and separate gaseous products (hydrogen and oxygen) [80]. 

EXAMPLE 19-12 Predicting Electrode Half-Reactions and Overall Reactions in Electrolysis... 

At the left electrode in Figure 19-5 the halfreaction occurring is Cl- —>- Cl4.g) + e, and at the right electrode the half-reaction is Na+ + er — - Naff). Which electrode is the anode and which is the cathode With these half-reactions, balance the net reaction occurring in the electrolysis cell. 

In case (a), the galvanic cell under non-faradaic conditions, one obtains an emf of 0.34 - (-0.76) = 1.10 V across the Cu electrode ( + pole) and the Zn electrode (- pole). In case (b), the galvanic cell with internal electrolysis, the electrical current flows in the same direction as in case (a) and the electrical energy thus delivered results from the chemical conversion represented by the following half-reactions and total reaction, repsectively ... 

We must narrow the options. There will be only one cathode reaction and only one anode reaction. How do we pick the correct half-reactions If one of the half-reactions were spontaneous (positive), we would pick it for that electrode. (If more than were spontaneous, we would pick the largest positive value.) All four half-reactions in this case are nonspontaneous (negative). This is typical for electrolysis, because you are using electrical energy to force a nonspontaneous process to take place. 

The information you have just learned permits a very precise control of electrolysis. For example, suppose you modify a Daniell cell to operate as an electrolytic cell. You want to plate 0.1 mol of zinc onto the zinc electrode. The coefficients in the half-reaction for the reduction represent stoichiometric relationships. Figure 11.23 shows that two moles of electrons are needed for each mole of zinc deposited. Therefore, to deposit 0.1 mol of zinc, you need to use 0.2 mol of electrons. 

The ready availability of electricity following the invention by Alssandro Volta of his famous pile in 1800 prompted, from an early date, the study of its effects on condensed matter and, most particularly, the decomposition of water by electrolysis involving chemical reactions at the electrodes. Work developed to the point where, by the middle of the second half of the nineteenth century, well-established industrial processes for the manufacture of aluminium and chlorine gas operated by electrolysis. 

When an aqueous salt solution is electrolyzed, the electrode reactions may differ from those for electrolysis of the molten salt because water may be involved. In the electrolysis of aqueous sodium chloride, for example, the cathode half-reaction might be either the reduction of Na+ to sodium metal, as in the case of molten sodium chloride, or the reduction of water to hydrogen gas ... 

Section 2 (Fig. 1, curves A and B), usually performed at the rotating platinum electrode (anode reactions) or the dropping mercury electrode (cathode reactions), should ideally suffice to define the electroactive species and determine its half-wave potential. It may be that systems in which acid-base equilibria exist are somewhat more laborious to study due to the necessity of recording voltammetric curves over a wide pH range, but in most cases the task can be accomplished with some effort. Once the voltammetric characteristics are known, it remains to carry out preparative constant potential electrolysis (cpe) at a suitable potential in order to make sure that the electroactive species is connected with the reaction of interest. 

Write the equations describing the half-reactions and net reaction for the electrolysis of molten KF/HF mixtures. At which electrodes are the products formed What is the purpose of the HF ... 

In Davy s electrolysis of molten NaCl, sodium ions were reduced to metallic sodium at the cathode. The oxidation of chloride ions to chlorine gas occurred at the second electrode, the anode. Each half-reaction in the electrolysis of molten sodium chloride is shown. 

Magnesium is produced commercially by electrolysis fiom a molten salt using a cell similar to the one shown here, (a) What salt is used as the electrolyte (b) Which electrode is the anode, and which one is the cathode (c) Write the overall cell reaction and individual half-reactions, (d) What precautions would need to be taken with respect to the magnesium formed [Section 20.9]... 

Starting at open circuit, when the electrolysis current is gradually increased, the voltage imposed across the system also increases in turn and several simultaneous half-reactions may possibly be observed at one electrode (or both of them). Remember that the currents for each half-reaction at the same interface must be added together. This results in faradic yields that are lower than 100%. When it is possible for several reactions to occur at one of the electrodes, the main half-reaction is the one that would lead to the lowest polarisation in absolute value, for the same individual current When several reactions can be envisaged at both interfaces in the whole system, then the overall reaction which results from the two main half-reactions is the one that requires the lowest imposed voltage. 

The anodic branch at the positive electrode corresponds to a single redox half-reaction, i.e., water oxidation whatever the working point is, the anodic faradic yield along this branch is 100%. For the current to flow and therefore for electrolysis to occur, the following condition must be fulfilled -i-I. 4 V/sce- This value... 

The cathodic branch at the negative electrode corresponds to one or two redox half-reactions, depending on the potential zone the reduction of Cu " ions and the reduction of protons (or of water). Here for the current to flow and therefore for electrolysis to occur, one must have the following Ec [lTtO)<-OM l/ici- If the absolute value of the current increases, then the copper electrode potential decreases. When the potential becomes lower than -0.4 V/sce (the working point is indicated by squares in figure 2.37) two half-reactions occur simultaneously and the faradic yield of each half-reaction drops lower than 100%. 

As in electrolysis mode, the main half-reaction usually offers the highest faradic yield. The current-potential curves define a potential range for each electrode such that the faradic yield remains close to 100%, as illustrated in figure 2.38. 

The chemical changes that occur at each electrode are written as chemical equations including electrons the corresponding chemical reaction is termed a half reaction and the equation a half equation. Summing up the two chemical half equations and balancing so that the electrons cancel out, one gets a total equation, which represents the overall reaction that has occurred in the electrolysis. With the electrolysis of molten NaCl as an example, we have ... 

In his creation of the first electric battery in 1799, Alessandro Volta laid the foundation of electrochemistry. His battery consisted of a zinc and a copper electrode dipped into an aqueous solution of sulfuric acid and connected externally by copper wires. The procedure for setting up and running this experiment is similar to the earlier practical of the electrolysis of aqueous copper sulfate solution (page 265), but without the external power supply. If a voltmeter is connected in parallel to the external circuit, it can be used to measure the potential difference. The half reactions and the overall reaction are as follows ... 

In a galvanic cell, there is a reversal of the anode and the cathode in relation to electrolysis. In this case, the anode is the negative electrode, while the cathode is the positive electrode. The reason is that here electrons are freed on the anode, from where they move up in the external metallic circuit and then move down onto the cathode where they are taken up by chemical species in the electrolyte. In common with electrolysis, an oxidation half reaction occurs at the anode and a reduction half reaction occurs at the cathode. Figure 8.9 shows an electrolytic cell which draws electric current from a galvanic cell. Electron flow, electrode polarity, anodes and cathodes, and the half reactions (oxidations and reactions) that take place are shown. 

Electrolysis - external current used to drive a non-spontaneous chemical reaction in the cell Positive - through the external metallic circuit, it drives away current (electrons) to the positive terminal of an external power supply. These electrons are freed on the electrode and come from an oxidation half reaction that occurs at the electrode-electrolyte interface. Negative - through the external metallic circuit, it attracts current (electrons) from the negative terminal of an external power supply. These electrons cause a reduction half reaction to occur at the electrode-electrolyte interface. 

Electrolysis of AgF(aq) in an acidic solution leads to the formation of silver metal and oxygen gas. (a) Write the half-reaction that occurs at each electrode, (b) Calculate the minimum external emf required for this process under standard conditions. 

In the electrolysis of aqueous sodium chloride, NaCl, shown below, what are the half-reactions that occm at each electrode ... 

Bioelectrochemical Hydrogen Production, Fig.1 Schematic layout of a microbial electrolysis cell (MEC) showing the anode and cathode chamber, the electrodes with attached biocatalysts, the membrane separator and the power supply, as well as the anodic and cathodic half reactions. AEM anion exchange membrane, CEM cation exchange membrane... 

Calculate the number of grams of magnesium that can be produced by supplying 1.00 F to the electrode. Consider the electrolysis of molten barium chloride, BaCl2. (a) Write the half-reactions, (b) How many grams of barium metal can be produced by supplying 0.50 A for 30 min ... 

In the electrolysis of a molten salt, the possible half-reactions are usually limited to those involving ions from the salt. When you electrolyze an aqueous solution of an ionic compound, however, you must consider the possibility that water is involved at one or both electrodes. Let us look at the possible half-reactions involving water. 

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