On the Standard Reduction Potentials of Half-Cells table, why does the Oxidizing agent have to be higher than Reducing agent for a redox reaction to spontaneously occur?

  • $\begingroup$ I thought that the only requirement was: if $E^{\circ}$ is positive = the redox is spontaneous. I thought that it did not matter which one is more positive. $\endgroup$ – CoffeeIsLife Jan 15 '17 at 20:46
  • $\begingroup$ You're right. It has to be positive and I get that. However, that's the mathematical explanation. But how can you explain it scientifically? $\endgroup$ – A.AK Jan 16 '17 at 3:22
  • $\begingroup$ Need some clarification here based on your comment. Are you interested in why a positive cell potential means a negative deltaG or are you interested in why a negative deltaG means spontaneous reaction? $\endgroup$ – Burak Ulgut Jan 16 '17 at 5:48
  • $\begingroup$ The sign convention may be considered to be 'back to front' since $\Delta G=-nF\Delta E$ with $\Delta E$ as the difference in standard redox potentials for the reaction. $\Delta G$ is negative for a spontaneous reaction. Oxidising agents have more positive redox $E^0$ in tables of redox potentials than reducing agents do. It is the difference in redox that determines if a reaction is spontaneous. 'Scientifically' a reducing agent is an electron donor and so an oxidising agent is an electron acceptor. $\endgroup$ – porphyrin Jan 16 '17 at 10:24

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