Reduction potentials table

They are the basis of many products and processes, from batteries to photosynthesis and respiration. You know redox reactions involve an oxidation half-reaction in which electrons are lost and a reduction half-reaction in which electrons are gained. In order to use the chemistry of redox reactions, we need to know about the tendency of the ions involved in the half-reactions to gain electrons. This tendency is called the reduction potential. Tables of standard reduction potentials exist that provide quantitative information on electron movement in redox half-reactions. In this lab, you will use reduction potentials combined with gravimetric analysis to determine oxidation numbers of the involved substances. 

FIGURE 14.3 Mnemonic devices, the arrows on the left, for predicting which chemicals will participate in a redox reaction and which will not. A segment of the table of standard reduction potentials (Table 14.1) is presented on the right as a help to understand the use of the arrows. See text for an example. 

A voltaic cell converts chemical energy into electrical energy. It consists of two parts called half-cells. When two different metals, one in each half-cell, are used in the voltaic cell, a potential difference is produced. In this experiment, you will measure the potential difference of various combinations of metals used in voltaic cells and compare these values to the values found in the standard reduction potentials table. 

Applying Concepts Write the half-reactions for the anode and cathode in each of the voltaic cells in the data table. Look up the half-reaction potentials from the standard reduction potentials table (Table 21-1) and record these in the data table. 

Why is lithium metal becoming a popular electrode in modern batteries Use the standard reduction potentials table to help you answer this question. 

A The HiO MT) couple has a high positive reduction potential ( Table 6.I4) and because the O ICO potential is considerably smaller, hydrogen peroxide is unstable with respect to disproportionation ... 

Here n represents the number of electrons transferred in the reaction. With this equation we can calculate the free-energy change for any oxidation-reduction reaction from the values of E" in a table of reduction potentials (Table 13-7) and the concentrations of the species participating in the reaction. 

The following problem can be solved using standard reduction potentials (Table 6-8). Use E0 ... 

Table 20.2 lists standard potentials E° for oxidation of first-series transition metals. Note that these potentials are the negative of the corresponding standard reduction potentials (Table 18.1, page 775). Except for copper, all the E° values are positive, which means that the solid metal is oxidized to its aqueous cation more readily than H2 gas is oxidized to H+(aq). 

Once an ore has been concentrated, it is reduced to the free metal, either by chemical reduction or by electrolysis. The method used (Table 21.2) depends on the activity of the metal as measured by its standard reduction potential (Table 18.1). The most active metals have the most negative standard reduction potentials and are the most difficult to reduce the least active metals have the most positive standard reduction potentials and are the easiest to reduce. 

In crystal field theory calculations the direction of the axial distortion is along the z-axis. Therefore, the dz2 orbitals in iron atoms in Fig. 15 are along the line adjoining the two iron atoms. Remembering that the dz2 orbital lies lowest in this symmetry, the effect of reducing the complex is to add electron density to the dz% orbitals of the iron atoms. Since the dz2 iron-orbitals in Fig. 15 overlap, this structure results in an electron repulsion term between the iron atoms which increases as the iron atoms in these proteins are reduced. Thus, the negative reduction potentials (Table 1) of the plant-type ferredoxins can be accounted for by this model. 

Purines. In general, -OH readily adds to double bonds but undergoes ET reactions only very reluctantly (Chap. 3.2). This also applies to purines despite their relatively low reduction potentials (Table 10.2). Thus, G- which is formed in the reaction of OH with dGuo has a short-lived -OH-adduct rather than G-+ as precursor (Candeias and Steenken 2000), and the H-abstraction that could also lead to G (for theoretical calculations see Mundy et al. 2002) does not occur to any significant extent. 

At the present time, tables of E° values use reduction half reactions called standard reduction potentials. Because the value of E° is affected by the concentration of the electrolyte solution, these values are given for 1 molar solutions. On a table of standard reduction potentials, we would find the E° value for zinc to be -0.76 V, the negative value resulting because the zinc oxidation half reaction must be reversed for a reduction potential table When the half reaction is reversed, the sign must also be reversed. Some standard reduction potential half reactions include the following ... 

The M=0 bond energies (—AGBF) for these four oxene adducts have been estimated on the basis of their reduction potentials (Table 13.1)17,18 relative to... 

Because oxidation is simply the opposite of reduction, it is only necessary to create a table of one of the values. By convention, the reduction potential is used in tables, and the values are typically given for the standard reduction potential, E°, also written E°red, in units of volts, V. Because oxidation takes place at the anode, this is the value that will need to be reversed (since oxidation potential = -reduction potential). Therefore, we can rearrange Equation 18.1 so we can use values from the reduction potential tables in Equation 18.2 ... 

The diversity in oxidation-reduction potentials (Table 9) can not be accounted for as iron environmental differences among these non-heme iron proteins, since the fundamental structures of iron coordination must be identical (Section III-A and B). Therefore, it is of interest that the secondary environmental effect causes such a great variety in oxidation-reduction potentials. 

For the phenylchlorosilanes, the trend of lower reduction potentials (Table 2) is veryfied in the galvanostatic electrolysis, though additional overpotentials emerge. For Ph2SiCl2 we measured a cathodic potential of -28 V (vs SCE) at a current density of 0.3 mA/cm [5]. This effect above all is responsible for the low current yield in comparison with the methylchlorosilanes. 

The table below lists the cell potentials for the 10 possible galvanic cells assembled from the metals A, B, C, D, and E, and their respective 1.00 M 2+ ions in solution. Using the data in the table, establish a standard reduction potential table similar to Table 11.1 in the text. Assign a reduction potential of 0.00 V to the half-reaction that falls in the middle of the series. You should get two different tables. Explain why, and discuss what you could do to determine which table is correct. 

What we have informally called electron pressure is the tendency to undergo oxidation. This is formally expressed as an oxidation potential however, we usually tabulate reduction potentials (Table 21-2). 

Reactivity patterns in widely varying oxidants are seldom considered, the reduc-tant patterns being more often compared. Such studies can be approached in the same way as that of reductants, but because the [CoCNHjjjX]"" oxidants are so common, there may be more difficulties in determining both the self-exchange rate and reduction potential. Table 1 lists values of and for several oxidants, as well as the calculated oxidation factors (OF) [using (f) in 12.2.5.1.1]. These OF values can be corrected to give effective oxidation (actors, but because fewer reversals of trends appear, the effective oxidation factors are not included here. The OF values suggest a reactivity pattern with a reductant of = 0.3 of [Ru(bipy)3] > [Fe(l,10-phen)3 + > [IrBr ] ",... 

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