Copper electrochemical cell

Electrons in this zinc-copper electrochemical cell flow from the zinc strip to the copper strip, causing an electric current that powers the light bulb. As the spontaneous reaction continues, much of the zinc strip is oxidized to zinc ions and the copper ions are reduced to copper metal, which is d osited on the copper strip. If an outside voltage source is applied to reverse the flow of electrons, the original conditions of the cell are restored. 

Figure 20.19 The zinc-copper electrochemical cell can be a voltaic cell or an electrolytic cell. 

When the reaction between zinc and copper(II) sulphate was carried out in the form of an electrochemical cell (p. 94), a potential difference between the copper and zinc electrodes was noted. This potential resulted from the differing tendencies of the two metals to form ions. An equilibrium is established when any metal is placed in a solution of its ions. 

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 ... 

If the copper electrode is the indicator electrode in a potentiometric electrochemical cell that also includes a saturated calomel reference electrode... 

Coupling reactions and related fluoroalkylations with polytTuoioalkyl halides are induced by vanous reagents, among them metals such as copper and zinc, or by an electrochemical cell. More recently, examples of carbon-carbon bond forma tion by coupling of unsaturated fluorides have been reported Both acyclic and cyclic fluoroolefins of the type (Rp)2C=CFRp undergo reducUve dimerization on treatment with phosphines [42] (equation 33) The reaction shown in equation 33 IS also accompbshed electrocheimcally but less cleanly [43]... 

Figure 5-7 shows a simple electrometer. It consists of two spheres of very light weight, each coated with a thin film of metal. The spheres are suspended near each other by fine metal threads in a closed box to exclude air draffs. Each suspending thread is connected to a brass terminal. Next to the box is a battery —a collection of electrochemical cells. There are two terminal posts on the batteiy. We shall call these posts Pi and Pi. If post Pi is connected by a copper wire... 

Let s begin our investigation of an electrochemical cell by assembling one. Fill a beaker with a dilute solution of silver nitrate (about 0.1 M will do) and another beaker with dilute copper sulfate. Put a silver rod in the AgN03 solution and a copper rod in the CuSO< solution. With a wire, connect the silver rod to one terminal of an... 

The moles of silver deposited per mole of copper dissolved are the same whether reaction (J) is carried out in an electrochemical cell or in a single beaker, as in Experiment 7. If, in the cell, electrons are transferred from copper metal (forming Cu+2) to silver ion (forming metallic silver), then electrons must have been transferred from copper metal to silver ion in Experiment 7. 

Now let s take a more detailed look into the electrochemical cell. Figure 12-5 shows a cross-section of a cell that uses the same chemical reaction as that depicted in Figure 12-1. The only difference is that the two solutions are connected differently. In Figure 12-1 a tube containing a solution of an electrolyte (such as KNOa) provides a conducting path. In Figure 12-5 the silver nitrate is placed in a porous porcelain cup. Since the silver nitrate and copper sulfate solutions can seep through the porous cup, they provide their own connection to each other. 

Suppose water is added to each of the beakers containing copper sulfate in the two electrochemical cells shown in Figure 12-4 (p. 204). What change will occur in the voltage in each cdl Explain. 

Although iodides are more reactive than bromides, 2-(trifluoro-methyl)pyridine was obtained in 95% yield from 2-bromopyridine and CF3Br using an undivided electrochemical cell, DMF, and a sacrificial copper anode. CF3Cu was the reactive intermediate (92CC53). Photochem-... 

An electrochemical cell in which electrolysis takes place is called an electrolytic cell. The arrangement of components in electrolytic cells is different from that in galvanic cells. Typically, the two electrodes share the same compartment, there is only one electrolyte, and concentrations and pressures are far front standard. As in all electrochemical cells, the current is carried through the electrolyte by the ions present. For example, when copper metal is refined electrolytically, the anode is impure copper, the cathode is pure copper, and the electrolyte is an aqueous solution of CuS04. As the Cu2f ions in solution are reduced and deposited as Cu atoms at the cathode, more Cu2+ ions migrate toward the cathode to take their place, and in turn their concentration is restored by Cu2+ produced by oxidation of copper metal at the anode. 

Diagram of a copper/zinc electrochemical cell operating under standard conditions. 

For forced-convection studies, the cathodic reaction of copper deposition has been largely supplanted by the cathodic reduction of ferricyanide at a nickel or platinum surface. An alkaline-supported equimolar mixture of ferri- and ferrocyanide is normally used. If the anolyte and the catholyte in the electrochemical cell are not separated by a diaphragm, oxidation of ferrocyanide at the anode compensates for cathodic depletion of ferricyanide.3... 

For the in situ studies, an electrochemical cell was designed to hold the nearly perfect copper crystal in contact with a thin layer (20 to 50 /Am) of electrolyte. Figures 34 and 35 show the cells employed in the ex situ and in situ experiments, respectively. In addition, Fig. 34 shows the voltammetric traces obtained for the deposition of T1 in the presence and absence of oxygen. In the... 

Figure 34. Voltammograms for T1 deposition onto a copper single crystal in the presence (a) and absence (b) of traces of oxygen. Inset electrochemical cell. (From Ref. 120, with permission.)...

The electrochemical cell used by Flcischmann and co-workers (1986) employing the Bragg configuration is shown in Figure 2.67(b). The source is a copper anode X-ray tube employing a Ni filter to select out the Cu Ka line the detector is a PS PD. 

An electrochemical cell was constructed by connecting the copper wire attached at the back of the Ti02 electrode to the platinum black cathode through a load. The two compartments were connected through an agar salt bridge that allows the exchange of ionic... 

The ellipsometer used in this study is described elsewhere(3). It consists of a Xenon light source, a monochromator, a polarizer, a sample holder, a rotating analyzer and a photomultiplier detector (Figure 1). An electrochemical cell with two windows is mounted at the center. The windows, being 120° apart, provide a 60° angle of incidence for the ellipsometer. A copper substrate and a platinum electrode function as anode and cathode respectively. Both are connected to a DC power supply. The system is automated with a personal computer to collect all experimental data during the deposition. Data analysis is carried out by a Fortran program run on a personal computer. 

Figure 1 A schematic diagram of the experimental set-up consisting of a rotating analyzer ellipsometer, an electrochemical cell with a copper substrate and a platinum electrode connected to a DC power supply. 

Hybinette A process for extracting nickel from sulfide ores. The nickel ore that occurs in Canada is a mixture of the sulfides of nickel, copper, and iron. Several methods have been used to separate these metals. In the Hybinette process, the ore is first smelted in a blast furnace, yielding a nickel-copper matte (i.e., a mixture of their lower sulfides). This is roasted to remove sulfur and leached with dilute sulfuric acid to remove copper. The resulting crude nickel oxide is used as the anode of an electrochemical cell. The nickel deposits on the cathode, which is contained in a cloth bag. Precious metals collect in the anode slime. The process was invented by N. V Hybinette in 1904 and operated at the Kristiansand refinery, Norway, from 1910. 

A basic electrochemical cell is depicted in Figure 9.3 and is made of a copper wire in one container with a solution of copper sulfate and a zinc rod in a different container with a zinc sulfate solution. There is a salt bridge containing a stationary saturated KC1 solution between the two containers. Electrons flow freely in the salt bridge in order to maintain electrical neutrality. A wire is connected to each rod and then to a measuring device such as a voltmeter to complete the cell. 

In these redox reactions, there is a simultaneous loss and gain of electrons. In the oxidation reaction part of the reaction (oxidation half-reaction), electrons are being lost, but in the reduction half-reaction, those very same electrons are being gained. Therefore, in redox reactions there is an exchange of electrons, as reactants become products. This electron exchange may be direct, as when copper metal plates out on a piece of zinc or it may be indirect, as in an electrochemical cell (battery). 

The rules of stoichiometry also apply in this case. In electrochemical cells, we must consider not only the stoichiometry related to chemical formulas, but also the stoichiometry related to electric currents. The half-reaction under consideration not only involves 1 mol of each of the copper species, but also 2 mol of electrons. We can construct a mole ratio that includes moles of electrons or we could construct a mole ratio using faradays. A faradav (F) is a mole of electrons. Thus, we could use either of the following ratios for the copper half-reaction ... 

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. 

For example, in Chapter 12, Section 4, we have examined the electrochemical response of azurin (from Pseudomonas aeruginosa), the only cupredoxin in which the copper(II) ion is pentacoordinate. Its reversible Cu(II)/Cu(I) reduction occurs at Eol= +0.31 V, vs. NHE, at 25° C. Measurements of the variation of the formal electrode potential with temperature in a non-iso thermic electrochemical cell gives the two diagrams illustrated in Figure ll.20... 

In many STM studies little effort has been made to control the atmosphere within the electrochemical cell. Yet oxygen is known to exert a major role in the chemistry and corrosion of many transition metals. For example, several STM studies have used the copper/copper ion reference electrode, yet the electrode is known to be polarized from its reversible condition by oxygen, leading to significant dissolution [154]. These effects become particularly significant in the smdy of metal deposition and dissolu-... 

Let us continue with the example of copper ions in contact with copper metal and zinc ions in contact with zinc metal. This combination is usually referred to as the Darnell cell or zinc/copper couple(Fig. 6.5a). For this electrochemical cell the reduction and oxidation processes responsible for the overall reaction are separated in space one half reaction taking place in one electrode compartment and the other takes place in the other compartment. 

A simple electrochemical cell can be made from two test tubes connected with a third tube (the crossbar of the H ), as shown in Figure 12-1. The hollow apparatus is filled by simultaneously pouring different solutions into the two test tubes, an aqueous solution (aq) of zinc sulfate into the left tube and a copper sulfate solution into the one on the right. Then a strip of zinc metal is dipped into the ZnS04 solution, a piece of copper is inserted into the CUSO4 solution, and the two ends of the metal strips are connected by wires to an voltmeter. The lateral connecting tube allows ionic migration necessary for a closed electrical circuit. The voltmeter will show the electrical potential of 1.10 volts, which leads to the movement of electrons in the wire from the zinc electrode toward the copper electrode. 

The electrochemical cell with zinc and copper electrodes had an overall potential difference that was positive (+1.10 volts), so the spontaneous chemical reactions produced an electric current. Such a cell is called a voltaic cell. In contrast, electrolytic cells use an externally generated electrical current to produce a chemical reaction that would not otherwise take place. 

The left half of this electrochemical cell, containing a zinc electrode and zinc sulfate solution, engages in an oxidation reaction. This reaction liberates electrons and turns zinc atoms into zinc ions (Zn, which is a zinc atom that has lost two electrons and therefore has a net positive charge of 2). The zinc atoms come from the electrode, which is gradually depleted the zinc ions that are produced in the reaction enter the solution. In the right half of the cell, a reduction reaction occurs electrons combine with copper ions to produce neutral atoms of copper. Copper ions leave the solution in the process and collect at the copper electrode. Over time, the zinc electrode and copper solution will run out of material, causing the reaction to cease unless the material is replenished. 

Figure 14.5 shows the basic arrangement of a electrochemical cell called the Daniell cell. This cell is named for John Frederick Daniell (1790-1845) who constructed this type of cell in 1836. The Daniell cell components include zinc and copper solutions in separate containers. Between the solutions is a salt bridge... 

---