Chemistry_Grade_10-12 (1).pdf

1742 equilibrium reactions in half cells lets think

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17.4.2 Equilibrium reactions in half cells Let’s think back to the Zn-Cu electrochemical cell. This cell is made up of two half cells and the reactions that take place at each of the electrodes are as follows: Zn Zn 2+ + 2 e Cu 2+ + 2 e Cu At the zinc electrode, the zinc metal loses electrons and forms Zn 2+ ions. The electrons are concentrated on the zinc metal while the Zn 2+ ions are in solution. But some of the ions will be attracted back to the negatively charged metal, will gain their electrons again and will form zinc metal. A dynamic equilibrium is set up between the zinc metal and the Zn 2+ ions in solution when the rate at which ions are leaving the metal is equal to the rate at which they are joining it again. The situation looks something like the diagram in figure 17.1. - - - - - - - - - - - 2+ 2+ 2+ 2+ 2+ Zn 2+ ions in solution zinc metal concentration of electrons on metal surface Figure 17.1: Zinc loses electrons to form positive ions in solution. The electrons accumulate on the metal surface. The equilibrium reaction is represented like this: Zn 2+ ( aq ) + 2 e Zn ( s ) (NOTE: By convention, the ions are written on the left hand side of the equation) In the zinc half cell, the equilibrium lies far to the left because the zinc loses electrons easily to form Zn 2+ ions. We can also say that the zinc is oxidised and that it is a strong reducing agent . At the copper electrode, a similar process takes place. The difference though is that copper is not as reactive as zinc and so it does not form ions as easily. This means that the build up of electrons on the copper electrode is less (figure 17.2). The equilibrium reaction is shown like this: 329
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17.4 CHAPTER 17. ELECTROCHEMICAL REACTIONS - GRADE 12 - - - - - - 2+ 2+ Cu 2+ ions in solution copper metal concentration of electrons on metal surface Figure 17.2: Zinc loses electrons to form positive ions in solution. The electrons accumulate on the metal surface. Cu 2+ ( aq ) + 2 e Cu ( s ) The equation lies far to the right because most of the copper is present as copper metal rather than as Cu 2+ ions. In this half reaction, the Cu 2+ ions are reduced . 17.4.3 Measuring electrode potential If we put the two half cells together, a potential difference is set up in two places in the Zn-Cu cell: 1. There is a potential difference between the metal and the solution surrounding it because one is more negative than the other. 2. There is a potential difference between the Zn and Cu electrodes because one is more negative than the other. It is the potential difference (recorded as a voltage) between the two electrodes that causes electrons, and therefore current, to flow from the more negative electrode to the less negative electrode. The problem though is that we cannot measure the potential difference (voltage) between a metal and its surrounding solution in the cell. To do this, we would need to connect a voltmeter to both the metal and the solution, which is not possible. This means we cannot measure the exact electrode potential (E o V) of a particular metal. The electrode potential describes the ability of a metal to give up electrons.
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