Copper-hydrogen electrochemical cell

Despite its electrode potential (p. 98), very pure zinc has little or no reaction with dilute acids. If impurities are present, local electrochemical cells are set up (cf the rusting of iron. p. 398) and the zinc reacts readily evolving hydrogen. Amalgamation of zinc with mercury reduces the reactivity by giving uniformity to the surface. Very pure zinc reacts readily with dilute acids if previously coated with copper by adding copper(II) sulphate ... 

The reduction is usually made in a multi-compartment electrochemical cell, where the reference electrode is isolated from the reaction solution. The solvent can be water, alcohol or their mixture. As organic solvent A,A-dimethyl form amide or acetonitrile is used. Mercury is often used as a cathode, but graphite or low hydrogen overpotential electrically conducting catalysts (e.g. Raney nickel, platinum and palladium black on carbon rod, and Devarda copper) are also applicable. 

In bimetallic catalysts prepared by catalytic reduction of copper by hydrogen, copper is deposited as three-dimensional agglomerates which are located, at low copper loadings, on the edges, corners, and rims of the parent metallic particles. The mechanism of deposition can be transferred from that proposed in corrosion and involving a local electrochemical cell ... 

Figure 1.4 An electrochemical cell consisting of copper and hydrogen electrodes...

An electrochemical cell consists of a standard hydrogen electrode and a copper metal electrode. 

An electrochemical cell consists of a standard hydrogen electrode and a copper metal electrode. If the copper electrode Is placed in a solution of O.IOATNaOH that is saturated with Cu(OH)2, what is the cell potential at 25°C [For Cu(OH)2, = 1.6 X 10 .]... 

An example of an experiment is given by Barral et al. (1992). In this student experiment, a sandpapered zinc bar is put in a beaker with a diluted solution of sulphuric acid, and a clean copper bar is placed in another beaker with the same solution of sulphuric acid. Both experiments are well known to students. Subsequently, both bars are cleaned again, connected with a metal wire and placed in a third beaker of the solution. Within a minute, bubbles can be observed at each bar. The teacher should not explain that this is due to zinc losing electrons to the hydrogen ions via copper, nor point out that an electric current is flowing, that is, an electrochemical cell has been created. Instead, students are asked to write down their observations and explain the production of the bubbles by themselves. 

The reaction of a metal with an aqueous acid to yield hydrogen, a severe form of corrosion, involves oxidation of the metal and reduction of hydrogen ions in solution, H (aq), to H2 gas, and so can be thought of in terms of an electrochemical cell. The tendency for a reaction to occur follows the order of the electrochemical series (Table 9.1). Metals below H2 in the electrochemical series-those with a negative standard reduction potential-will react with aqueous acids to release hydrogen gas. Those above it will not react with acid. Thus, zinc will dissolve in acid to give hydrogen, whereas copper will not. 

During normal cell operation the anode contaminants that are more electrochemically active than copper (and less than hydrogen), mainly arsenic. 

When the compensation potential is determined using the Kenrick apparatus, the mercury streams down the center of the vertical tube and the HCl solution down its walls so that the Volta potential difference across the air gap is eliminated. In this system both the concentration of HCl and the pressure of hydrogen gas can be varied. It differs from cells (8.7.8) and (8.7.16) in that no attempt is made to balance the charge carriers on opposite sides of the cell. The Galvani potential difference measured between the two copper leads can be related to the electrochemical potentials of electrons in the mercury and platinum as follows ... 

---