In redox reactions, electrons are transferred from one species to another. If the reaction is spontaneous, energy is released, which can then be used to do useful work. To harness this energy, the reaction must be split into two separate half reactions: the oxidation and reduction reactions. The reactions are put into two different containers and a wire is used to drive the electrons from one side to the other. In doing so, a Voltaic/ Galvanic Cell is created.

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Introduction

When a redox reaction takes place, electrons are transferred from one species to the other. If the reaction is spontaneous, energy is released, which can be used to do work. Consider the reaction of a solid copper (Cu(s)) in a silver nitrate solution (AgNO3(s)).

<2Ag^+_(aq) + Cu_(s) leftrightharpoons Cu^2+_(aq) + 2Ag_(s)>

The (AgNO_3;(s)) dissociates in water to produce (Ag^+_(aq)) ions and (NO­­^-_3;(aq)) ions. The NO3-(aq) ions can be ignored since they are spectator ions and do not participate in the reaction. In this reaction, a copper electrode is placed into a solution containing silver ions. The Ag+(aq) will readily oxidize Cu(s) resulting in Cu2+(aq), while reducing itself to Ag(s).

This reaction releases energy. When the copper electrode solid is placed directly into a silver nitrate solution, however, the energy is lost as heat and cannot be used to do work. In order to harness this energy and use it do useful work, we must split the reaction into two separate half reactions; The oxidation and reduction reactions. A wire connects the two reactions and allows electrons to flow from one side to the other. In doing so, we have created a Voltaic/ Galvanic Cell.

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When an electrode is oxidized in a solution, it is called an anode and when an electrode is reduced in solution. it is called a cathode.

Anode: The anode is where the oxidation reaction takes place. In other words, this is where the metal loses electrons. In the reaction above, the anode is the Cu(s) since it increases in oxidation state from 0 to +2. Cathode: The cathode is where the reduction reaction takes place. This is where the metal electrode gains electrons. Referring back to the equation above, the cathode is the Ag(s) as it decreases in oxidation state from +1 to 0.

Remembering Oxidation and Reduction

When it comes to redox reactions, it is important to understand what it means for a metal to be “oxidized” or “reduced”. An easy way to do this is to remember the phrase “OIL RIG”.

OIL = Oxidization is Loss (of e-)

RIG = Reduction is Gain (of e-)

In the case of the example above (Ag^+_(aq)) gains an electron meaning it is reduced. (Cu_(s)) loses two electrons thus it is oxidized.



Flow of Electrons

Electrons always flow from the anode to the cathode or from the oxidation half cell to the reduction half cell. In terms of Eocell of the half reactions, the electrons will flow from the more negative half reaction to the more positive half reaction. A cell diagram is a representation of an electromslsec.comical cell. The figure below illustrates a cell diagram for the voltaic shown in Figure (PageIndex1) above.

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cell voltage or potential difference between it"s two two half-cells. Cell voltage is also known as cell potential or electromotive force (emf) and it is shown as the symbol (E_cell).

Standard Cell Potential:

The Eo values are tabulated with all solutes at 1 M and all gases at 1 atm. These values are called standard reduction potentials. Each half-reaction has a different reduction potential, the difference of two reduction potentials gives the voltage of the electromslsec.comical cell. If Eocell is positive the reaction is spontaneous and it is a voltaic cell. If the Eocell is negative, the reaction is non-spontaneous and it is referred to as an electrolytic cell.


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