Cu2+ ions

Murakami et al. reported that a cyclophane 27 having two imidazole groups is activated by Cu2+ ions in the hydrolysis of p-nitrophenyl dodecanoate 25,26), although the activation seemed to be small. 

The Table 7 shows Cu2+ ion effects. In a CTAB micelle, the Cu2+ ion itself without ligand enhances the rate 1900 fold as compared to 14 fold of the Zn2+ ion (Table 4). 

Table 7. Pseudo-first-order rate constants, kc and K values in the Cu2 + ion catalyzed reaction of 1...

In an Evans diagram 89> the mixed potential can easily be found and also be verified by measuring the open circuit potential of a zinc-amalgam electrode in a Cu2+-ion solution. Even the complication by the simultaneous presence of another reducible species, e.g., Pbz+ can be graphically demonstrated for different limiting conditions... 

Fig. 3. Evans-diagram for the cementation of Cu2+ and Pb2 with zinc amalgam of different zinc content. If the zinc concentration in the mercury employed for this special extraction technique is low, the anodic zinc-dissolution current density may be diffusion controlled and below the limiting cathodic current density for the copper reduction. The resulting mixed potential will lie near the halfwave potential for the reaction Cu2+ + 2e j Cu°(Hg) and only Cu2 ions are cemented into the mercury.

From the values listed we see that the Cu2+ ion has a heat of formation greater than that of H+ by about 65 kJ/mol, and Cd2+ has a heat of formation about 76 kJ/mol less than that of H+. 

When ammonia is added to an aqueous solution of a copper(II) salt, a deep, almost opaque, blue color develops (Figure 15.1). This color is due to the formation of the Cu(NH3)42+ ion, in which four NH3 molecules are bonded to a central Cu2+ ion. The formation of this species can be represented by the equation... 

The nitrogen atom of each NH3 molecule contributes a pair of unshared electrons to form a covalent bond with the Cu2+ ion. This bond and others like it, where both electrons are contributed by the same atom, are referred to as coordinate covalent bonds. 

When a piece of zinc is added to a water solution containing Cu2+ ions, the following redox reaction takes place ... 

The Zn electrode must not come in contact with Cu2+ ions. Why ... 

In this reaction, copper metal plates out on the surface of the zinc. The blue color of the aqueous Cu2+ ion fades as it is replaced by the colorless aqueous Zn2+ ion (Figure 18.1). Clearly, this redox reaction is spontaneous it involves electron transfer from a Zn atom to a Cu2+ ion. 

To design a voltaic cell using the Zn-Cu2+ reaction as a source of electrical energy, the electron transfer must occur indirectly that is, the electrons given off by zinc atoms must be made to pass through an external electric circuit before they reduce Cu2+ ions to copper atoms. One way to do this is shown in Figure 18.2. The voltaic cell consists of two half-cells—... 

Electrons generated at the anode move through the external circuit (right to left in Figure 18.2) to the copper cathode, connected through the black wire to the voltmeter. At the cathode, the electrons are consumed, reducing Cu2+ ions present in the solution around the electrode ... 

As the above half-reactions proceed, a surplus of positive ions (Zn2+) tends to build up around the zinc electrode. The region around the copper electrode tends to become deficient in positive ions as Cu2+ ions are consumed. To maintain electrical neutrality, cations must move toward the copper cathode or, alternatively, anions must move toward the zinc anode. In practice, both migrations occur. 

A Zn-Cu + voltaic cell. In this voltaic cell, a voltmeter (left) is connected to a half-cell consisting of a Cu cathode in a solution of blue Cu2+ ions and a half-cell consisting of a Zn anode in a solution of colorless Zn2+ ions. The following spontaneous reaction takes place in this cell Zn(s) + Ctf+lag) — M+(atfl + Cu(s). 

The salt bridge allows ions to pass from one solution to the other to complete the circuit and prevents direct contact between Zn atoms and the Cu2+ ions. 

Oxidation of Ni by Cu2+. Nickel metal reacts spontaneously with Cu2+ ions, producing Cu metal and Ni2+ ions. Copper plates out on the surface of the nickel, and the blue color of Cu2+ is replaced by the green color of NP+. 

It should be noted that this method is only applicable to solutions containing up to 25 mg copper ions in 100 mL of water if the concentration of Cu2+ ions is too high, the intense blue colour of the copper(II) ammine complex masks the colour change at the end point. The indicator solution must be freshly prepared. 

Procedure (copper in crystallised copper sulphate). Weigh out accurately about 3.1 g of copper sulphate crystals, dissolve in water, and make up to 250 mL in a graduated flask. Shake well. Pipette 50 mL of this solution into a small beaker, add an equal volume of ca AM hydrochloric acid. Pass this solution through a silver reductor at the rate of 25 mL min i, and collect the filtrate in a 500 mL conical flask charged with 20 mL 0.5M iron(III) ammonium sulphate solution (prepared by dissolving the appropriate quantity of the analytical grade iron(III) salt in 0.5M sulphuric acid). Wash the reductor column with six 25 mL portions of 2M hydrochloric acid. Add 1 drop of ferroin indicator or 0.5 mL N-phenylanthranilic acid, and titrate with 0.1 M cerium(IV) sulphate solution. The end point is sharp, and the colour imparted by the Cu2+ ions does not interfere with the detection of the equivalence point. 

As the cell is discharged, Zn2+ ions are produced at the anode while Cu2+ ions are used up at the cathode. To maintain electrical neutrality, SO4- ions must migrate through the porous membrane,dd which serves to keep the two solutions from mixing. The result of this migration is a potential difference across the membrane. This junction potential works in opposition to the cell voltage E and affects the value obtained. Junction potentials are usually small, and in some cases, corrections can be made to E if the transference numbers of the ions are known as a function of concentration.ee It is difficult to accurately make these corrections, and, if possible, cells with transference should be avoided when using cell measurements to obtain thermodynamic data. 

FIGURE K.5 When a strip of zinc is placed in a solution that contains Cu2 t ions, the blue solution slowly becomes colorless and copper metal is deposited on the zinc. The inset shows that, in this redox reaction, the zinc metal is reducing the Cu2+ ions to copper metal and the Cu2+ ions are oxidizing the zinc metal to Zn2 ions. 

Aqueous solutions of compounds containing the Cu2+ ion are blue as a result of the presence of the Cu(H,0)62+ complex. Does this complex absorb in the visible region Suggest an explanation. Refer to Major Technique 2, which follows these exercises. 

In the process of separating Pb2 ions from Cu2+ ions as sparingly soluble iodates, what is the Pb2+ concentration when Cu2+ just begins to precipitate as sodium iodate is added to a solution that is initially 0.0010 m Pb(N03)2(aq) and 0.0010 M Cu(N03)2(aq) ... 

If we were to place a piece of zinc metal into an aqueous copper(II) sulfate solution, we would see a layer of metallic copper begin to deposit on the surface of the zinc (see Fig. K.5). If we could watch the reaction at the atomic level, we would see that, as the reaction takes place, electrons are transferred from the Zn atoms to adjacent Cu2 r ions in the solution. These electrons reduce the Cu2+ ions to Cu atoms, which stick to the surface of the zinc or form a finely divided solid deposit in the beaker. The piece of zinc slowly disappears as its atoms give up electrons and form colorless Zn2+ ions that drift off into the solution. The Gibbs free energy of the system decreases as electrons are transferred and the reaction approaches equilibrium. However, although energy is released as heat, no electrical work is done. 

As Cu2+ ions are reduced, the solution at the cathode becomes negatively charged and the solution at the anode begins to develop a positive charge as the additional Zn2+ ions enter the solution. To prevent this charge buildup, which would quickly stop the flow of electrons, the two solutions are in contact through a porous wall ions provided by the electrolyte solutions move between the two compartments and complete the electrical circuit. 

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. 

FIGURE 12.14 A schematic representation showing the electrolytic process for refining copper. The anode is impure copper. The Cu2 ions produced by oxidation of the anode migrate to the cathode, where they are reduced to pure copper metal. A similar arrangement is used for electroplating objects. 

A solution is prepared by dissolving 1 mol each of Cu(N03)2, Ni(N03)2, and AgNO in 1.0 L of water. Using only data from Appendix 2B, identify the metals (if any) that when added to these solutions (a) will leave the Ni2 ions unaffected but will cause Cu and Ag to plate out of solution (b) will leave the Ni2+ and Cu2+ ions in solution but will cause Ag to plate out of solution (c) will leave all three metals ions in solution ... 

The copper product is known as blister copper because of the appearance of air bubbles in the solidified metal. In the hydrometallurgical process, soluble Cu2+ ions are formed by the action of sulfuric acid on the ores. Then the metal is obtained by reducing these ions in aqueous solution either electrolytically or chemically, by using an inexpensive reducing agent that has a more negative standard potential than that of copper, such as hydrogen or iron (see Section 14.3) ... 

Many of the d-block elements form characteristically colored solutions in water. For example, although solid copper(II) chloride is brown and copper(II) bromide is black, their aqueous solutions are both light blue. The blue color is due to the hydrated copper(II) ions, [Cu(H20)fJ2+, that form when the solids dissolve. As the formula suggests, these hydrated ions have a specific composition they also have definite shapes and properties. They can be regarded as the outcome of a reaction in which the water molecules act as Lewis bases (electron pair donors, Section 10.2) and the Cu2+ ion acts as a Lewis acid (an electron pair acceptor). This type of Lewis acid-base reaction is characteristic of many cations of d-block elements. 

The composition and properties of the ions contained in the solution are not the same as those of ions contained in the ionic crystal lattice. It is already known that anhydrous copper sulfate (CuS04) is colorless. This implies that Cu2+ and SCT ions that make up the crystal lattice of the sulfate are colorless. When the Cu2+ ions combine with water molecules during dissolution they turn blue (the color characteristic of copper salt). This color is therefore due to hydrated ions of copper, i.e., ions connected with the water molecules. 

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