Copper sulfate solution redox reaction

Because electrons are neither products nor reactants in chemical reactions, the two processes are interdependent and neither can occur alone. The zinc metal dissolution must furnish electrons for the copper metal plating. The reaction of zinc and copper sulfate solution is a spontaneous reaction involving a transfer of electrons, i.e., is a spontaneous redox process. The spontaneity of the reaction is commonly explained by saying that zinc loses electrons more readily than copper or, alternatively, that Cu2+ ions gain electrons more readily than Zn2+ ions. 

Oxidation and reduction reactions, also known as redox reactions, involve the transfer of electrons from one substance that is being oxidized (losing electrons) to another that is being reduced (gaining electrons). A very common example of a redox reaction, one that you may very well have done in the lab, is the reaction that occurs when you place clean, iron nails into a copper sulfate solution. In this reaction, shown below, iron is oxidized and copper is reduced ... 

As a result of this redox reaction between iron and copper sulfate solution, solid copper metal is deposited on the iron nail. To balance the equation given in the text for this reaction you could use the method of halfreactions. 

School-made misconceptions play a more significant role by advanced students, they do not differentiate between atoms and ions, atoms and substances, or substances and particles they are simply used as synonyms. Several students do not seem to be aware that redox reactions between atoms and ions are necessary to explain the formation of elemental copper in their mind copper sulfate solution already contains the copper atoms copper from copper sulfate forms a deposit on the iron nail, it connects to the iron atoms . Students definitely lack an understanding of the concept of ions in salt solutions (see Chap. 5). 

Problem In order to show further phenomena on the structure of complexes and complex equilibria, it should be shown that the central ion is solidly bound to the ligands and is not solely present in the solution, in the stable tetra ammine copper complex. In order to do this, an iron nail is dipped into the complex solution, respectively, diluted sodium hydroxide solution is added and this is compared to regular copper sulfate solution the iron nail does not show the copper deposit as usual, no precipitation of the copper hydroxide is deposited. The copper sulfate solution should be interpreted in comparison to the complex solution as a solution with free Cu2 + (aq) ions or very instable aqueous copper complexes. With the explanation of the copper deposit on iron a cross-linkage to redox reactions (see Chap. 8) is possible. 

The main processes occurring in electrochemical cells are simultaneous oxidation and reduction reactions, or redox reactions. At one electrode, the anode, a reduced species is oxidized here meaning to release electrons, while at the other electrode, the cathode, an oxidized species absorbs electrons and is reduced. It is common to think of an electrochemical cell as consisting of two half-cells. (one containing the anode and the second containing the cathode) and to describe the processes in terms of halfcell reactions. For example, one common cell consists of a copper cathode in a copper sulfate solution, and a zinc anode in a zinc sulfate solution. The overall reaction is... 

With a copper sulfate solution, copper is dissolved from the anode and deposited on the cathode. For the other solution (also known as redox couples solution), only current transfer occurs at the electrodes. The respective reactions at the electrodes are ... 

The energy of oxidation is given by the redox potential of the reaction. If a piece of copper wire is placed in a solution of zinc sulfate, nothing happens, but if zinc metal is placed in a solution of copper sulfate, the zinc metal is corroded. Simultaneously, the blue color of the copper sulfate solution disappears, and metallic copper is deposited on the surface of the remaining zinc metal ... 

A voltaic cell produces electrical energy through spontaneous redox chemical reactions. When zinc metal is placed in a solution of copper sulfate, an electron transfer takes place between the zinc metal and copper ions. The driving force for the reaction is the greater attraction of the copper ions for electrons ... 

When a strip of zinc metal is added to a solution of copper(II) sulfate, the blue color slowly fades, and the zinc metal is replaced by copper metal (Figure 4-13). As copper ions in the solution are reduced to copper metal, zinc atoms are oxidized to Zn cations. This is an example of a metal displacement reaction, in which a metal ion in solution (Cu ) is displaced by another metal (Zn) by means of a redox reaction. Figure 4-13 also shows molecular views of this displacement reaction. 

The reactivities of potassium and silver with water represent extremes in the spontaneity of electron-transfer reactions. The redox reaction between two other metals illustrates less drastic differences in reactivity. Figure 19-5 shows the reaction that occurs between zinc metal and an aqueous solution of copper(II) sulfate zinc slowly dissolves, and copper metal precipitates. This spontaneous reaction has a negative standard free energy change, as does the reaction of potassium with water ... 

Zinc metal reacts spontaneously with an aqueous solution of copper sulfate when they re placed in direct contact. Zinc, being a more reactive metal than copper (it s higher on the activity series of metals presented in Chapter 8), displaces the copper ions in solution. The displaced copper deposits itself as pure copper metal on the surface of the dissolving zinc strip. At first, the reaction may appecir to be a simple single replacement reaction, but it s also a redox reaction. 

A voltaic cell (also known as a galvanic cell) is a device that allows for the transfer of electrons (in a redox reaction) to be completed in a separate pathway from the reaction mixtures. In a voltaic cell, the two half-reactions are physically separated from each other by placing them into two separate reaction vessels. The electrons are transferred from one vessel to the other by a connecting wire (see Figure 18.1). In voltaic cells, the reactions in each vessel must be spontaneous. In figure 18.1, in the reaction on the left, a zinc strip is placed in a zinc sulfate solution, where zinc from the strip replaces zinc in solution (Zn —> Zn2+ + 2 c ). In the reaction vessel on the left, the zinc strip will lose mass over time. Electrons create an electric potential difference across the wire, which is also known as a voltage. The voltage across the wire will allow electrons to be forced from the zinc strip, across the wire, to the copper strip. However, an electric current cannot be established until the circuit is completed. 

The sodium potassium tartrate initially forms the soluble blue copper(II) tartrate complex witli the copper sulfate the color of this complex is a deeper blue than that of the hydrated ion. At the temperature of the experiment reduction to the beautiful golden orange CuzO occurs, and at the same time oxygen and (rather less) CO2 are formed in the redox proce.ss these are responsible lor the foaming of the solution. This reversible color reaction can be repeated several times when further hydrogen peroxide is added. The pH of the system increases from 5 to about 9. ... 

Another redox reaction that doesn t involve oxygen occurs when a strip of zinc metal is placed in a solution of copper(II) sulfate. The progress of this reaction can be followed easily because a readily observable change takes place. As shown in Figure 16.5, copper metal quickly begins to form on the zinc strip. 

Write the equation for the redox reaction that occurs when a piece of iron metal is dipped in a solution of copper(II) sulfate. 

The same overall redox reaction occurs if the magnesium metal is placed directly into a solution of copper sulfate, Figure 17.12. However, this is not a galvanic cell because the electrons do not flow through an external circuit. Instead, the electrons move directly from the magnesium metal to the copper ions, forming copper metal. This is a way to make copper metal from copper ions, but it is not a way to make electrical power. 

A voltaic cell, then, is composed of two separate half-cells. In the case of the zinc and copper(II) redox reaction, the first half-cell consists of a strip of zinc metal in a solution of zinc sulfate. The second half-cell consists of a strip of copper metal in a solution of copper(II) sulfate. In the Zn/Zn + half-cell, zinc atoms lose two electrons and are converted to zinc ions. The two electrons pass through the wire to the Cu/Cu half-cell, where Cu ions gain the two electrons to form uncharged copper atoms. The zinc metal and copper metal strips are called electrodes, the general name for electrical conductors placed in half-cells of voltaic cells. 

The occurrence of electron transfer is more apparent in some redox reactions than others. When metallic zinc is added to a solution containing copper(ll) sulfate (CUSO4), zinc reduces by donating two electrons to it ... 

I hroughout history, we have suffered from our ignorance of basic electrochemical principles. For example, during the Middle Ages, our chemistry ancestors (alchemists) placed an iron rod into a blue solution of copper sulfate. They noticed that bright shiny copper plated out onto an iron rod and they thought that they had changed a base metal, iron, into copper. What actually happened was the redox reaction shown in Equation 1. 

What do you think would happen if you separated the oxidation half-reaction from the reduction half-reaction Can a redox reaction occur Consider Figure 20.1a, in which a zinc strip is immersed in a solution of zinc sulfate and a copper strip is immersed in a solution of copper(II) sulfate. 

Conversely, some spontaneous redox reactions can be made to supply useful amounts of electrical energy. When a piece of zinc is put in a copper(II) sulfate solution, the zinc quickly becomes coated with metallic copper. We expect this coating to happen because zinc is above copper in the activity series copper(II) ions are therefore reduced by zinc atoms ... 

This reaction is clearly a spontaneous redox reaction, but simply dipping a zinc rod into a copper(II) sulfate solution will not produce useful electric current. However, when we carry out this reaction in the cell shown in Figure 17.5, an electric current is produced. The cell consists of a piece of zinc immersed in a zinc sulfate solution and connected by a wire through a voltmeter to a piece of copper immersed in cop-per(II) sulfate solution. The two solutions are connected by a salt bridge. Such a cell produces an electric current and a potential of about 1.1 volts when both solutions are 1.0 M in concentration. A cell that produces electric current from a spontaneous chemical reaction is called a voltaic cell. A voltaic cell is also known as a galvanic cell. 

In the Redox Chemistry of Iron and Copper movie (eChapter 20.5) an iron naU is placed in an aqueous solution of copper sulfate, (a) Write the half-reaction for the formation of the copper coating on the nail, (b) Is there another half-reaction occurring in the solution If so, what is it (c) What is the sign of AG for the overall process that occurs ... 

Redox reactions sometimes involve direct electron transfer, in which one substance immediately picks up the electrons another has lost. For instance, if you put a piece of zinc metal into a copper(ll) sulfate solution, zinc gives up two electrons (becomes oxidized) to the Cu ion that accepts the electrons (reducing it to copper metal). The copper metal begins spontaneously plating out on the surface of the zinc. The equation for the reaction is... 

A simple but important example is the reaction that occurs if you drop a piece of zinc into a solution of copper sulfate. You will find that the zinc becomes coated in a layer of copper and that some of the zinc dissolves (as ions). You can conclude that the zinc atoms, Zn, in the lump, having lost electrons, have been oxidized to Zn ions. Moreover, because the Cu ions in the solution have been converted to copper atoms, Cu, by gaining electrons, the ions have been reduced (Figure 5.3). This redox reaction might seem rather dull and a waste of zinc, but you will see in Reaction 7 that it is the foundation of the emergence of modem communication systems. 

You will learn more about the importance of half-reactions when you study electrochemistry in Chapter 21. For now, however, you can learn to use halfreactions to balance a redox equation. First, look at an unbalanced equation taken from Table 20-3 to see how to separate a redox equation into half-reactions. For example, the following unbalanced equation represents the reaction that occurs when you put an iron nail into a solution of copper(II) sulfate, as shown in Figure 20-8. Iron atoms are oxidized as they lose electrons to the copper(ll) ions. 

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