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In the redox reaction chapter, we learn that oxidation and reduction go hand in hand. The oxidising agent oxidises by accepting the electron released by the reducing agent and gets reduced in the process.

However, in the half cell electrode (say, $\ce{Zn}$/$\ce{Zn^{2+}}$) only oxidation is taking place and no simultaneous reduction is happening in the scene. Also, in a $\ce{Cu}$/$\ce{Cu^{2+}}$ half cell electrode, only reduction is taking place. Until the half cells are joined by a saltbridge, oxidation and reduction go on violating the basic principle of redox reactions.

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    $\begingroup$ Half reactions cannot happen independently and (as you have already stated in the first line) a reduction is simultaneous with an oxidation. Half cells are just representations. Nothing violated. Reaction occurs only when 2 half reactions are present, and the sum of their potentials is positive. $\endgroup$ – William R. Ebenezer Apr 21 at 10:56
  • $\begingroup$ Is that true? In liquid ammonia solutions you can get electrons pulled out at the anode but they just pour into the liquid through the cathode. $\endgroup$ – Oscar Lanzi Nov 19 at 20:28
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In a (standard) Zn/Zn++ half cell, no oxidation (and no reduction) is taking place. A piece of zinc is sitting in a 1M solution of Zn++ ions (with their accompanying anions). It is not imaginary, and on a micro scale, an equilibrium is occurring (could be occurring!) where Zn + Zn++ --> Zn++ + Zn. No overall change occurs; this is a reference point. If it were not stable, some process would be taking place that degrades the half cell and it would be reacting itself away. In reality, half cells are only useful to the extent that they don't react themselves away until they are connected to the other half of the overall cell.

Same for Cu/Cu++. However, when the half cells are connected (thru a salt bridge so ions can flow, and a wire so electrons can flow), two processes occur: electrons leave the zinc (the zinc oxidizes) and electrons flow to the copper (reducing it). No "process" or reaction occurs in the half cells until they are connected.

The reality of oxidation and reduction require both to occur simultaneously - WHEN THEY OCCUR - but we can speak of each process independently of the other. When electrons leave an atom (oxidation), they have to go somewhere (reduction); if electrons are not flowing (no actual reaction) there is no oxidation or reduction. The half cell potentials then define places from which electrons will flow, or locations to which they will flow.

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Redox reactions consist of oxidation and reduction processes occurring simultaneously. What does this actually mean?

Redox reaction involves the transfer of electrons between species (Easy way to remember: increase in oxidation number = oxidation, hence decrease in oxidation number = reduction). These processes occur together at the same time at different interfaces (boundaries).

Considering your cell (Galvanic Cell): The main reacting species (excluding the spectator ions) are Copper and Zinc. The overall reaction that is taking place is:

$$ \ce {Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu}$$

Applying the previously mentioned trick, you can see that zinc is being oxidised from 0 oxidation number to 2+ (increase in oxidation number = oxidised), whereas copper ion is reduced from 2+ to 0 oxidation number (decrease in oxidation number = reduced). You can think of a half-cell equation as small components that make the overall reaction in redox reaction of Galvanic cell. They provide vital information on what is happening to the species in terms of electrons involved and the standard potential can be used to identify if the species is capable of operating as reducing agent or oxidising agent.

The half-cell equations is usually tabulated against standard hydrogen electrode as being the reference.(Note: half-cell equations are usually written as reductions). In order to identify which species is oxidising agent, you can think of it this way: the more positive the potential $\Rightarrow$ stronger oxidising agent $\Rightarrow$ oxidising agents by definition are reduced (reduction occurs at the cathode).

The reactions are only taking place when the circuit is closed (cells are connected). This leads to the flow of the electrons around the circuit and redox reaction can take place. To qualitatively prove that your reaction occurs, you have to consider the cell potential:

$$E_{cell} = E_{cathode} -E_{anode} $$

applying the concept of electrical work = Gibbs energy, thus

$$G=-nFE_{cell}$$

(Important: provided the concentration 1 M, if they are not you have to apply Nernst Equation to identify new cell potential) .

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