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REDOX REACTIONS 239 Fig. 7.2 Redox reaction between copper and aqueous solution of silver nitrate occurring in a beaker. At equilibrium, chemical tests reveal that both However, as we shall see later, the charge Ni2+(aq) and Co (aq) are present at moderate transfer is only partial and is perhaps better concentrations. In this case, neither the described as an electron shift rather than a reactants [Co(s) and Ni2+(aq)] nor the products complete loss of electron by H and gain by [Co 2(aq) and Ni (s)] are greatly favoured. O. What has been said here with respect This competition for release of electrons to equation (7.18) may be true for a good incidently reminds us of the competition for number of other reactions involving covalent compounds. Two such examples of this class release of protons among acids. The similarity suggests that we might develop a table in of the reactions are: which metals and their ions are listed on the H (s) + Cl (g) → 2HCl(g) (7.19) 2 2 basis of their tendency to release electrons and, just as we do in the case of acids to indicate the strength of the acids. As a matter of fact CH 4(g) + 4Cl 2g) → CCl (4) + 4HCl(g) (7.20) In order to keep track of electron shifts we have already made certain comparisons. By comparison we have come to know that in chemical reactions involving formation zinc releases electrons to copper and copper of covalent compounds, a more practical method of using oxidation number has releases electrons to silver and, therefore, the electron releasing tendency of the metals been developed. In this method, it is always is in the order: Zn>Cu>Ag. We would love to assumed that there is a complete transfer of electron from a less electronegative atom make our list more vast and design a metal activity series or electrochemical series. to a more electonegative atom. For example, The competition for electrons between various we rewrite equations (7.18 to 7.20) to show charge on each of the atoms forming part of metals helps us to design a class of cells, named as Galvanic cells in which the chemical the reaction : 0 0 +1 –2 reactions become the source of electrical energy. We would study more about these 2H 2g) + O2(g) → 2H O 2l) (7.21) cells in Class XII. 0 0 +1 –1 H2(s) + Cl2(g) → 2HCl(g) (7.22) 7.3 OXIDATION NUMBER – +1 0 + –1 +1 –1 A less obvious example of electron transfer is realised when hydrogen combines with oxygen CH (g) + 4Cl (g) → CCl (l) +4HCl(g) (7.23) 4 2 4 to form water by the reaction: It may be emphasised that the assumption 2H (g) + O (g) → 2H O (l) (7.18) of electron transfer is made for book-keeping 2 2 2 Though not simple in its approach, yet purpose only and it will become obvious at we can visualise the H atom as going from a a later stage in this unit that it leads to the simple description of redox reactions. neutral (zero) state in H2to a positive state in H 2, the O atom goes from a zero state in O 2 Oxidation number denotes the oxidation to a dinegative state in H O. It is assumed that state of an element in a compound 2 there is an electron transfer from H to O and ascertained according to a set of rules consequently H i2 oxidised and O is2reduced. formulated on the basis that electron pair Reprint 2026-27