Chemical reaction half-cell

The explicit aims of boiler and feed-water treatment are to minimise corrosion, deposit formation, and carryover of boiler water solutes in steam. Corrosion control is sought primarily by adjustment of the pH and dissolved oxygen concentrations. Thus, the cathodic half-cell reactions of the two common corrosion processes are hindered. The pH is brought to a compromise value, usually just above 9 (at 25°C), so that the tendency for metal dissolution is at a practical minimum for both steel and copper alloys. Similarly, by the removal of dissolved oxygen, by a combination of mechanical and chemical means, the scope for the reduction of oxygen to hydroxyl is severely constrained. 

Although it is attractive to directly convert chemical energy to electricity, PEM fuel cells face significant practical obstacles. Expensive heavy metals like platinum are typically used as catalysts to reduce energy barriers associated with the half-cell reactions. PEM fuel cells also cannot use practical hydrocarbon fuels like diesel without complicated preprocessing steps. Those significantly increase the complexity of the overall system. At this time, it appears likely that PEM fuel cells will be confined to niche applications where high cost and special fuel requirements are tolerable. 

We see that the overall chemical reaction that occurs in an electrochemical cell is conveniently described in terms of two types of half-reactions. In one, electrons are lost in the other, they are gained. To distinguish these half-reactions we need two identifying names. 

The value of this list is obvious. Any half-reaction can be combined with the reverse of another half-reaction (in the proportion for which electrons gained is equal to electrons lost) to give a possible chemical reaction. Our list permits us to predict whether equilibrium favors reactants or products. We would like to expand our list and to make it more quantitative. Electrochemical cells help us do this. 

Figure 9.3 The lead storage battery. The key to obtaining electrical energy from a redox chemical reaction is to physically separate the two half-cell reactions so that electrons are transferred from the anode through an external circuit to the cathode. In the process, electrical work is accomplished.

Haber, F., 357, 385 Hahn, O., 717 half-cell, 492 half-life (chemical), 543 half-life (radioactive), 712 half-reaction, 484 halide, nomenclature, 763 Hall, C., 598 Hall process, 598 haloakane, F36, 739, 756 nomenclature, 763 halogen, F20, 639... 

This is a quantitative problem, so we follow the standard strategy. The problem asks about an actual potential under nonstandard conditions. Before we determine the potential, we must visualize the electrochemical cell and determine the balanced chemical reaction. The half-reactions are given in the problem. To obtain the balanced equation, reverse the direction of the reduction half-reaction with the... 

The electric current connected with chemical conversion is termed the faradaic current in contrast to the non-faradaic current required to charge the electrical double layer. The equation for the electrode reaction is formulated similarly to that for the half-cell reaction (Section 3.1.4) for the cathodic reaction, the electrons are placed on the left-hand side of the equation and, for the anodic reaction, on the right-hand side. 

We can combine the half-cell potentials for any two half-reactions in the table to get a complete cell potential. The chemical reaction may proceed spontaneously if the complete cell potential is positive. Otherwise, the opposite reaction may proceed spontaneously. We combine half-cells by adding the chemical reactions and by adding the corresponding half-cell potentials. We must first get the correct chemical reactions and corresponding half-cell potentials for the half-reactions, as follows ... 

If you reverse the direction of the chemical reaction, change the sign of the potential. You can ensure getting a positive cell potential by reversing the lower half-reaction from Table 14.2. 

What difference does it make to the conclusions about the chemical reaction that may occur in a cell if you reverse the wrong equation for a half-reaction ... 

But the chemical reaction forming this coloured layer of oxide represents only one part of the cell. A cell contains a minimum of two electrodes, so a cell comprises two reactions - we call them half-reactions one describes the chemical changes at the positive electrode (the anode) and the other describes the changes that occur at the negative cathode. 

A concentration cell contains the same electroactive material in both half-cells, but in different concentration (strictly, with different activities). The emf forms in response to differences in chemical potential /r between the two half-cells. Note that such a concentration cell does not usually involve different electrode reactions (other than, of course, that shorting causes one half-cell to undergo reduction while the other undergoes oxidation). 

To understand potentiometric methods, those that measure electrical potentials and determine analyte concentrations from these potentials, it is necessary that numerical values for these tendencies be known under conventional standard modes and conditions. What are these modes and conditions First, all halfreactions must be written as either reductions or oxidations. Scientists have decided to write them as reductions. Second, the tendencies for half-reactions to proceed depend on the temperature, the concentrations of the chemical species involved, and, if gases are involved, the pressure in the half-cell. Scientists have defined standard conditions to be a temperature of 25°C, a concentration of exactly 1 M for all dissolved chemical species involved, and a pressure of exactly 1 atm. Third, because every cell consists of two half-cells, it is not possible to measure the value directly. However, if we were to assign the tendency of a certain half-reaction to be zero, then the tendencies of all other half-reactions can be determined relative to this reference half-reaction. 

The purpose of the silver-silver chloride combination is to prevent the potential that develops from changing due to possible changes in the interior of the electrode. The potential that develops is a membrane potential. Since the glass membrane at the tip is thin, a potential develops due to the fact that the chemical composition inside is different from the chemical composition outside. Specifically, it is the difference in the concentration of the hydrogen ions on opposite sides of the membrane that causes the potential—the membrane potential—to develop. There is no half-cell reaction involved. The Nernst equation is... 

A half-cell describes a single electrode and its associated chemical reaction which forms part of a voltaic cell. 

There can be no chemical reaction in such a system without a complementary electron donor or acceptor to complete the process. Each of these electrode systems is known as a half-cell and the potential developed by a halfcell cannot be measured in absolute terms but only compared with that of another half-cell. The chemical reaction occurring at each half-cell is known as a half-reaction. 

The reduction half-reaction does not include a solid conductor of electrons, so an inert platinum electrode is used in this half-cell. The platinum electrode is chemically unchanged, so it does not appear in the chemical equation or half-reactions. However, it is included in the shorthand representation of the cell. 

Represent one example of a galvanic cell, and one example of an electrolytic cell, using chemical equations, half-reactions, and diagrams. 

The reduction-oxidation potential (typically expressed in volts) of a compound or molecular entity measured with an inert metallic electrode under standard conditions against a standard reference half-cell. Any oxidation-reduction reaction, or redox reaction, can be divided into two half-reactions, one in which a chemical species undergoes oxidation and one in which another chemical species undergoes reduction. In biological systems the standard redox potential is defined at pH 7.0 versus the hydrogen electrode and partial pressure of dihydrogen of 1 bar. 

A fuel cell is a device that converts the free energy change of a chemical reaction directly into electrical energy. This conversion occurs by two electrochemical half cell reactions. 

The collision between reacting atoms or molecules is an essential prerequisite for a chemical reaction to occur. If the same reaction is carried out electrochemically, however, the molecules of the reactants never meet. In the electrochemical process, the reactants collide with the electronically conductive electrodes rather than directly with each other. The overall electrochemical Redox reaction is effectively split into two half-cell reactions, an oxidation (electron transfer out of the anode) and a reduction (electron transfer into the cathode). 

This section addresses the role of chemical surface bonding in the electrochemical oxidation of carbon monoxide, CO, formic acid, and methanol as examples of the electrocatalytic oxidation of small organics into C02 and water. The (electro)oxidation of these small Cl organic molecules, in particular CO, is one of the most thoroughly researched reactions to date. Especially formic acid and methanol [130,131] have attracted much interest due to their usefulness as fuels in Polymer Electrolyte Membrane direct liquid fuel cells [132] where liquid carbonaceous fuels are fed directly to the anode catalyst and are electrocatalytically oxidized in the anodic half-cell reaction to C02 and water according to... 

A chemical reaction that can occur within one half-cell win reach equilibrium and is assumed to remain at equilibrium. Such a reaction is not the net cell reaction. 

We allow half-cells to stand long enough to come to chemical equilibrium within each half-cell. For example, in the right-hand half-cell in Figure 14-9, the reaction... 

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