Electrochemical Cells Introduction

From MyMCAT

1 Introduction
2 Redox Reactions
3 Half Reactions
4 Reduction Potentials
5 The Electrochemical Cell
6 Signs of the Cathode and Anode
7 See Also

[edit] Introduction

If anyone has ever seen a lemon powered clock, then they have seen an example of an electrochemical cell. In general, electrochemical cells involve reactions that transfer electrons between compounds, this transfer can often be harnessed to create current flowing electrons which can power electronic devices.

[edit] Redox Reactions

All electrochemical reactions can be classified as redox, or reduction-oxidation, reactions. This type of reaction involves a transfer of electrons between species. A reaction species that gains electrons is said to under go reduction, while a species that loses them undergoes oxidation. Assigning oxidation numbers allows one a simple process in determining when a species is oxidized or reduced when it is not obviously clear (see Oxidation Numbers for a review of this process).

[edit] Half Reactions

In a full electrochemical cell or reaction, some ions, atoms or molecules lose electrons (oxidation) while the other ions, atoms or molecules gain electrons (reduction) from their electrode. Together, they make up a complete reaction, however they can also be viewed independently as half-reactions or half-cells (when used in the electrochemical cell sense).

For instance, the most common electrochemical reaction, the Daniell Cell, consists of a reaction involving Zinc and Copper exchanging electrons (a redox reaction),

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

But we could also view it as two separate half reactions, one reduction, and one oxidation,

$\begin{align} \mbox{Zn} &\rightarrow \mbox{Zn}^{2+} + 2\mbox{e}^{-} \\ \mbox{Cu}^{2+} + 2\mbox{e}^{-} &\rightarrow \mbox{Zn} \end{align}$

Notice that in this reaction, it is clear that two electrons are transfered at a time.

Zn²⁺

Cu²⁺

→

In the above reaction, Zn is turning into Zn²⁺, while Cu²⁺ is turning into Cu. A gain in electrons is reduction, thus we need to identify the species which gains electrons, Cu²⁺ goes to Cu by neutralizing the positive charge with two electrons, thus Cu²⁺ is the correct answer.

Zn²⁺

Cu²⁺

→

Firstly, species with positive charge will tend not to want to lose electrons since they are already deficient in them, thus Zn²⁺ and Cu²⁺ would not lose electrons. Secondly, the reaction favours the forward direction because Zn overpowers the Cu when it comes to releasing electrons, thus Zn is the answer. This question was really just asking which species is oxidized (loss of electrons) in the forward direction (the other oxidation in the reverse direction can not be as strong since the reaction doesn't go that way!)

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[edit] Reduction Potentials

As a result of each species' ability to transfer electrons, half-cells have a characteristic voltage. Different choices of substances for each half-cell give different potential differences. Each reaction is undergoing an equilibrium reaction between different oxidation states of the ions—when equilibrium is reached the cell cannot provide further voltage.

Typically, the half-cell reactions written in the direction of reduction (the reverse being oxidation) and are assigned a reduction potential, which allows one to assess how the half-cell will behave. Large, positive numbers imply the half-cell favours reduction and wants to proceed in that direction, whereas large, negative numbers imply the half-cell dislikes the forward, reduction reaction, and would rather undergo oxidation.

A typical Standard Reduction Potential Table can be seen below, notice how all reactions are written in the direction of reduction.

Reaction	Reduction Potential(v)
Li⁺ + e^- --> Li	-3.04
Na⁺ + e^- --> Na	-2.71
Mg²⁺ + 2e^- --> Mg	-2.38
Zn²⁺ + 2e^- --> Zn	-0.76
Ni²⁺ + 2e^- --> Ni	-0.23
Cu²⁺ + 2e^- --> Cu	+0.34
Ag⁺ + e^- -->Ag	+0.80

Notice how some are positive while others are negative. A negative value implies that this species does not easily reduce and in fact more favorably undergoes the reverse reaction of oxidation.

[edit] The Electrochemical Cell

Taking the concept of a half-cell and its reduction potential, it should be evident that there is a possibility to combine half-cells into different redox reactions. Each combination, with its unique pair reduction potentials, could lead to reactions of varying strengths. A reaction in which one half-cell prefers reduction (a positive reduction potential) with another half-cell which prefers oxidation (a negative reduction potential) would rapidly react exchanging electrons as quickly as possible, however if we chose a different combination it may be that no reaction occurs at all because they both resist each other's reductive/oxidative needs.

Taking this one step further, this transfer of potential energy, could be harnessed to perform useful work if we controlled our reactions correctly. If one half-cell reactions gives up electrons easily and another accepts electrons, then it should be possible to generate a current between the two cells and perform useful work due to the flowing electrons.

Below is a design for a basic electrochemical cell,

[edit] Cathode and Anode

Electrochemical cells have two conductive electrodes (the anode and the cathode). The cathode is always the site of reduction. The cathode supplies electrons to the positively charged cations (positive ions) which flow to it from the electrolyte (solution). Because ions in solution are getting reduced and becoming solid, the electrode slowly gains mass with every molecule which attaches. In electroplating, as in gold plating and other commercial applications, items to be plated with pure metal are attached to and become part of the cathode in the electrolytic solution.

The anode on the other hand, is always the site of oxidation. At the anode, anions (negative ions) are forced by the electrical potential to react chemically and give off electrons (oxidation) which then flow through the wire to the cathode. As the metal is oxidized it ionizes and is lost into solution, thus over time the anode loses mass.

Electrodes can be made from any sufficiently conductive materials, such as metals, semiconductors, graphite, and even conductive polymers. In between these electrodes is the electrolyte, which contains ions that can freely move.

[edit] Galvanic Cells and Electrolytic Cells

The shown above diagram is an example of a galvanic cell or voltaic cell. It is considered a galvanic cell because when the wire is connected to complete the circuit between the two electrodes, it will spontaneously produce a current through reduction and oxidation reactions in the solutions. Galvanic cells are simply just another name for a battery which uses two half-cells.

Electrolytic cells on the other hand, do not spontaneously react to produce current. In fact, they require current to be supplied to them (from another source) to even function. An electrolytic cell can be thought of as the opposite of a galvanic cell, its two half-cell reactions do not like to proceed and so power is supplied to force the electrochemical reaction to occur.

[edit] Electromotive Force

The electromotive force dictates how a pair of half-cells will behave when combined as it describes what the overall potential is for a reacting pair of half-cells (a complete reaction). If the electromotive force is positive, then it implies that both half-reactions will proceed to react sponatneously (and so we have a galvanic cell), if however the electromotive force is negative, then it implies that reaction cannot spontaneously occur and energy will have to be supplied if we wish it to proceed (and so it is an electrolytic cell).

There are two forms of the equation to determine electromotive force, but both imply the exact SAME thing.

$\begin{align} E^{\circ}_{\mbox{emf}} &= E^{\circ}_{\mbox{cathode}} - E^{\circ}_{\mbox{anode}} \\ \\ E^{\circ}_{\mbox{emf}} &= E^{\circ}_{\mbox{reduction}} + E^{\circ}_{\mbox{oxidation}} \end{align}$

The first says that the electromotive force of a electrochemical cell can be found by simply subtracting the cathode's reduction potential from the anode's reduction potential. Keeping in mind that the cathode is the electrode that undergoes reduction and the anode is the one that undergoes oxidation, then the second equation should become evident. If we know what the potential for reduction of one half reaction is and we know what the potential for oxidation of the other half reaction is (the reverse sign of the reduction potential), then the sum of them would tell us what the potential for the whole reaction is.

When using these formula's remember that normally a list of reduction potentials is given, thus the first formula can be used directly if we know how the electrochemical cells are set up (where the cathode and anode are).

Unlike the normal procedures for balancing reactions, if say for example one half-cell donates one e^- and the other accepts two (or vice versa), reduction potentials do not need to be adjusted by multiplication factors to balance the electrons. Since reduction potentials state how likely is the species willing to donate electrons, it does not make sense to think that because there is twice as much of that species it will donate twice as easily (yes, more will be donated to match the other half-cell's reaction needs, but the potential wont change!)

Li and Na

Li and Ag

Cu and Ag

Ni and Cu

→

We are given all reduction potentials and simply want to know which pair will produce the largest EMF. Thus to answer this question we need to use E_emf=E_cathode-E_anode for each pair, and also decide which should be the cathode and anode to maximize the value. A larger E_emf would mean a better potential, and thus a better battery. Li has a reduction potential of -3.04 and Ag has a reduction potential of +.80. If we set Ag to the cathode and Li to the anode we get (+0.80)-(-3.04)=+3.84. This is infact the largest number that can be made from any of the pairs of reduction potentials given, thus it is the answer. Note that flipping which was anode and cathode does change the potential!

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[edit] Signs of the Cathode and Anode

A significant difference in Galvanic and Electrolytic cells is the designation of signs to the cathode and anode. In a galvanic cell, the cathode and anode are labeled as (+) and (-) respectively, while in an electrolytic cell, the signs are reversed. Besides this nothing else changes! The cathode is always the site of reduction, cations always flow to it, and electrons always to it from the wire. The anode is always the site of oxidation, anions always flow to it, and electrons always move up the wire from it. This is true for both galvanic and electrolytic cells.

The sign difference comes purely from how they behave. In a galvanic cell, the anode is giving off electrons (oxidation), so it has a negative charge and as such, the cathode must have the positive charge as that is where electrons are flowing to. In an electrolytic cell on the other hand, external power is supplied to force the reaction to occur, and this is done by applying a stronger potential to the electrodes in a reverse orientation, thus the signs are designated to be in reverse.

[edit] See Also

A cheat sheet on electrochemical cells can be found here, electrochemical cells cheatsheet, which summaries most of what can be found on this page.

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Category: Chemical Reactions

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