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Making an electric cell

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Making an electric cell The zinc turns out to be the negative electrode (the black lead) as it is the more reactive of the two metals. You can check the order of reactivity in the reactivity series. Remember that metals react by losing electrons and turning into positive ions. It follows that a more reactive metal will lose electrons more readily than a less reactive one, and consequently be the negative electrode of the pair. Background The differing reactivities of metals When metals react, they give away electrons and form positive ions. This particular topic sets about comparing the ease with which a metal does this to form hydrated ions in solution - for example, Mg2+(aq) or Cu2+(aq). We might want to compare the ease with which these two changes take place: Everybody who has done chemistry for more than a few months knows that magnesium is more reactive than copper. The first reaction happens much more readily than the second one. What this topic does is to try to express this with some numbers. Looking at this from an equilibrium point of view Suppose you have a piece of magnesium in a beaker of water. There will be some tendency for the magnesium atoms to shed electrons and go into solution as magnesium ions. The electrons will be left behind on the magnesium. In a very short time, there will be a build-up of electrons on the magnesium, and it will be surrounded in the solution by a layer of positive ions. These will tend to stay close because they are attracted to the negative charge on the piece of metal. ...read more.


This stops too much mixing of the contents of the salt bridge with the contents of the two beakers. The electrolyte in the salt bridge is chosen so that it doesn't react with the contents of either beaker. What happens? These two equilibria are set up on the two electrodes (the magnesium and the porous platinum): Magnesium has a much greater tendency to form its ions than hydrogen does. The position of the magnesium equilibrium will be well to the left of that of the hydrogen equilibrium. That means that there will be a much greater build-up of electrons on the piece of magnesium than on the platinum. Stripping all the rest of the diagram out, apart from the essential bits: There is a major difference between the charge on the two electrodes - a potential difference which can be measured with a voltmeter. The voltage measured would be 2.37 volts and the voltmeter would show the magnesium as the negative electrode and the hydrogen electrode as being positive. This sometimes confuses people! Obviously, the platinum in the hydrogen electrode isn't positive in real terms - there is a slight excess of electrons built up on it. But voltmeters don't deal in absolute terms - they simply measure a difference. The magnesium has the greater amount of negativeness - the voltmeter records that as negative. The platinum of the hydrogen electrode isn't as negative - it is relatively more positive. The voltmeter records it as positive. Throughout the whole of this redox potential work, you have to think in relative terms. For example, +0.4 is relatively more negative than +1.2. ...read more.


The magnesium is being oxidised. Taking another example . . . When the copper(II) ions gain electrons to form copper, they are being reduced. Reducing agents and oxidising agents A reducing agent reduces something else. That must mean that it gives electrons to it. Magnesium is good at giving away electrons to form its ions. Magnesium must be a good reducing agent. An oxidising agent oxidises something else. That must mean that it takes electrons from it. Copper doesn't form its ions very readily, and its ions easily pick up electrons from somewhere to revert to metallic copper. Copper(II) ions must be good oxidising agents. Summarizing this on the electrochemical series Metals at the top of the series are good at giving away electrons. They are good reducing agents. The reducing ability of the metal increases as you go up the series. Metal ions at the bottom of the series are good at picking up electrons. They are good oxidising agents. The oxidising ability of the metal ions increases as you go down the series. Judging the oxidising or reducing ability from E� values The more negative the E� value, the more the position of equilibrium lies to the left - the more readily the metal loses electrons. The more negative the value, the stronger reducing agent the metal is. The more positive the E� value, the more the position of equilibrium lies to the right - the less readily the metal loses electrons, and the more readily its ions pick them up again. The more positive the value, the stronger oxidising agent the metal ion is. ...read more.

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