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7.5: Oxidation–Reduction Reactions

  • Page ID
    52371
  • In contrast to acid–base reactions, oxidation–reduction (or redox) reactions obey a different pattern. In the simplest kinds of redox reactions, polar products are generated from non-polar reactants. You may have run into such reactions already (even if you did not know what they were called!) When iron is left in contact with oxygen (in air) and water, it rusts. The iron is transformed from a hard, non-polar metallic substance, Fe (solid), into a powdery substance, Fe2O3.nH2O(s). Rusting is mechanistically similar to the reactions that occur when copper turns green, when silver tarnishes and turns black, or (in perhaps the favorite reaction of chemists everywhere142) when sodium metal explodes in water.143

    All of these reactions start with a metal in its elemental form. Pure metals have no charge or permanent unequal distribution of charge (which makes them different from salts like NaCl). In fact we can use the synthesis of sodium chloride (NaCl) from its elements sodium (Na) and chlorine (Cl2) to analyze what happens during a redox reaction. The reaction can be written as:

    2Na(s) + Cl2(g) ⇄ 2NaCl(s)

    We have already looked at the structure of ionic compounds in Chapter 4 and know that the best way to think about them is to consider NaCl as a three-dimensional lattice of alternating positive (Na+) and negative (Cl–) ions. That is as the reaction proceeds the metal atoms becomes cations, and the chlorine molecules become anions. We could write this as two separate reactions: The Na loses an electron – a process that we define as oxidation.

    Na ⇄ Na+ +e(an oxidation reaction)

    The electrons must go somewhere (they cannot just disappear) and since chlorine is an electronegative element, it makes sense that the electrons should be attracted to the chlorine. We define the gain of electrons as a reduction.

    Cl + e ⇄ Cl (a reduction reaction).

    It turns out that all reactions in which elements react with each other to form compounds are redox reactions. For example, the reaction of molecular hydrogen and molecular oxygen is also a redox reaction:

    2H2(g) + O2(g) ⇄ 2H2O(l)

    The problem here is that there is no obvious transfer of electrons. Neither is there an obvious reason why these two elements should react in the first place, as neither of them has any charge polarity that might lead to an initial interaction. That being said, there is no doubt that H2 and O2 react. In fact, like sodium and water, they react explosively.144 When we look a little more closely at the reaction, we can see that there is a shift in electron density on individual atoms as they move from being reactants to being products. The reactants contain only pure covalent (H—H and O—O) bonds, but in the product (H2O) the bonds are polarized: Hδ+ and Oδ– (recall that oxygen is a highly electronegative atom because of its highly effective nuclear charge.) There is a shift in overall electron density towards the oxygen. This is a bit subtler than the NaCl case. The oxygen has gained some extra electron density, and so been reduced, but only partially – it does not gain the whole negative charge. The hydrogen has also been oxidized by losing some electron density. We are really talking about where the electron spends most of its time. In order to keep this straight, chemists have developed a system of oxidation numbers to keep track of the losses and gains in electron density.

    References

    142 This is based on the personal memories of one (and only one) of the authors.
    143 Visit http://www.youtube.com/watch?v=eCk0lYB_8c0 for an entertaining video of what happens when sodium and

    other alkali metals are added to water (yes, they probably faked the cesium).

    144 Hydrogen and oxygen can be used as rocket fuel, and the so-called “hydrogen economy” is based on the energy released when hydrogen reacts with the oxygen from the air.

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