Standard electrode potentials chart

When water is electrolyzed with copper electrodes or using other common metals, the amount of 02(g) is less than when Pt electrodes are used, but the amount of H2(g) produced is independent of electrode material. Why does this happen In electrolysis, the most easily oxidized species is oxidized and the most easily reduced species is reduced. If we compare Cu and H20 by looking on the standard reduction potentials chart (data given below), we see that Cu is a stronger reducing agent than H20, because 0.337 V is less than 0.828 V. This means that Cu is more easily oxidized than water. 

The complete zinc-copper cell has a total potential of 1.10 volts (the sum of 0.76v and 0.34v). Notice that the sign of the potential of the zinc anode is the reverse of the sign given in the chart of standard electrode potentials (see Table 12-4) because the reaction at the anode is oxidation. 

In the chart of standard electrode potentials (Table 12-1), reactions are arranged in order of their tendency to occur. Reactions with a positive EMF occur more readily than those with a negative EMF. The zinc-copper cell has an overall EMF of + 1.10 volts, so the solution of zinc and deposition of copper can proceed. 

If you select any two half-reactions from the chart of standard electrode potentials, the half-reaction higher on the list will proceed as a reduction, and the one lower on the list will proceed in the reverse direction, as an oxidation. Beware Some references give standard electrode potentials for oxidation half-reactions, so you have to switch higher and lower in the rule stated in the preceding sentence, though this is not common. 

Problem 38 Considering only the elements in the chart of standard electrode potentials (Table 12-2), which pair can make a battery with the greatest voltage What would the voltage be ... 

This book contains extensive tables of standard electrode potentials covering the periodic chart. For each electrode reaction Is given the standard potential, the temperature and the pressure, the solvent, and a literature reference. Occasionally the temperature coefficient of the electrode potential Is given together with an estimate of uncertainty. Much of the tabulated data Is taken from secondary sources (such as Item [149]). 

Appendixes and Endpapers. Included in the appendixes are an updated guide to the literature of analytical chemistry, tables of chemical constants, electrode potentials, and recommended compounds for the preparation of standard materials sections on the use of logarithms and exponential notation, and on normality and equivalents (terms that are not used in the text itself) and a derivation of the propagation of error equations. The inside front and back covers of this book provide a full-color chart of chemical indicators, a table of molar masses of com-... 

In 1889 the German chemist Nernst gave the theoretical foundation for the use of electrode potential to measure the concentration of an ion in solution. This led S0rensen in 1900 to develop the pH scale and an early pH meter based on a hydrogen electrode in combination with a calomel reference electrode. At that time pH was measured using known concentrations of coloured indicator dyes which lost or changed colour by drop-wise addition of solutions. This allowed standardized measurements be made. Litmus paper and other indicator strips were also used, with colour comparisons against colour charts to determine pH. 

The measurement of pH. (a) A strip of test paper impregnated with indicator (a materiai that changes coior as the acidity of the surroundings changes) is put in contact with the soiution of interest. The resuiting coior is matched with a standard coior chart (coiors shown as a function of pH) to obtain the approximate pH. (b) A pH meter uses a sensor (a pH electrode) that develops an electrical potential that is proportional to the pH of the solution. 

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