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Galvanic cell

 
Sci-Tech Dictionary: galvanic cell
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(gal′van·ik ′sel)

(electricity) An electrolytic cell that is capable of producing electric energy by electrochemical action.

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The Galvanic cell, named after Luigi Galvani, is a part of a battery consisting of an electrochemical cell with two different metals connected by a salt bridge or a porous disk between the individual half-cells. It is sometimes also called a Voltaic cell.

Common usage of the word battery has evolved to include a single Galvanic cell, but the first batteries had many Galvanic cells.[1][2] A Galvanic cell is a source of electrical power, obtained from chemical reactions, whereas an electrolytic cell, uses electrical power to perform chemical reactions.

Contents

History

In 1780, Luigi Galvani discovered that when two different metals (copper and zinc for example) were connected together and then both touched to different parts of a nerve of a frog leg at the same time, they made the leg contract.[3] He called this "animal electricity". The Voltaic pile invented by Alessandro Volta in the 1800s is similar to the galvanic cell. These discoveries paved the way for electrical batteries.

Description

Schematic of Zn-Cu galvanic cell

A Galvanic cell consists of two half-cells. In its simplest form each half-cell consists of a metal and a solution of a salt of the metal. The salt solution contains a cation of the metal and an anion to balance the charge on the cation. In essence the half-cell contains the metal in two oxidation states and the chemical reaction in the half-cell is an oxidation-reduction (redox) reaction, written symbolically as

M (reduced species) \leftrightarrow Mn+ (oxidised species) +n e-

In a galvanic cell one metal is able to reduce the cation of the other and, conversely, the other cation can oxidise the first metal. The two half-cells must be physically separated so that the solutions do not mix together. A salt bridge or porous plate is used to separate the two solutions.

The number of electron transferred in both directions must be the same, so the two half-cells are combined to give the whole-cell electrochemical reaction. For two metals A and B

A \leftrightarrow A^{n+} + n e^-
B \leftrightarrow B^{m+} + m e^-
mA + nB^{m+} \leftrightarrow nB + mA^{n+}

This is not the whole story as anions must also be transferred from one half-cell to the other. When a metal in one half-cell is oxidised anions must be transferred into that half-cell to balance the electrical charge of the cation produced. The anions are released from the other half-cell where a cation is reduced to the metallic state. Thus, the salt bridge or porous membrane serves both to keep the solutions apart and to allow the flow of anions in the direction opposite opposite to the flow of electrons in the wire connecting the electrodes.

The voltage of the Galvanic cell is the sum of the voltages of the two half-cells. It is measured by connecting a voltmeter to the two electrodes. The voltmeter has very high resistance, so the current flow is effectively negligible. When a device such as an electric motor is attached to the electrodes a current flows and redox reactions occur in both half-cells. This will continue until the concentration of the cations that are being reduced goes to zero.

For the Daniell cell, depicted in the figure, the two metals are zinc and copper and the two salts are sulfates of the respective metal. Zinc is the more reducing metal so when a device is connected to the electrodes, the electrochemical reaction is

Zn + Cu2+ → Zn2+ + Cu

The zinc electrode is dissolved and copper is deposited the copper electrode. By definition, the cathode is the electrode where reduction (gain of electrons) takes place, so the copper electrode is the cathode. The cathode attracts cations, so has a negative charge. In this case copper is the cathode and zinc the anode.

Galvanic cells are typically used as a source of electrical power. By their nature they produce direct current. For example, a lead-acid battery contains a number of galvanic cells. The two electrodes are effectively lead and lead oxide.

The Weston cell was adopted as an International Standard for voltage in 1911. The anode is a cadmium mercury amalgam, the cathode is made of pure mercury, the electrolyte is a (saturated) solution of cadmium sulfate and the depolarizer is a paste of mercurous sulfate. When the electrolyte solution is saturated the voltage of the cell is very reproducible, hence its use as a standard.

Cell voltage

The standard electrical potential of a cell can be determined by use of a standard potential table for the two half cells involved. The first step is to identify the two metals reacting in the cell. Then one looks up the standard electrode potential, EStrikeO.png, in volts, for each of the two half reactions. The standard potential for the cell is equal to the more positive EStrikeO.png value minus the more negative EStrikeO.png value.

For example, in the picture above the solutions are CuSO4 and ZnSO4. Each solution has a corresponding metal strip in it, and a salt bridge or porous disk connecting the two solutions and allowing SO42− ions to flow freely between the copper and zinc solutions. In order to calculate the standard potential one looks up copper and zinc's half reactions and finds:

Cu2+ + 2 e− is in equilibrium with Cu: EStrikeO.png = +0.34 V
Zn2+ + 2 e− is in equilibrium with Zn: EStrikeO.png = −0.76 V

Thus the overall reaction is:

Cu2+ + Zn is in equilibrium with Cu + Zn2+

The standard potential for the reaction is then +0.34 V - -0.76 V = 1.10 V. The polarity of the cell is determined as follows. Zinc metal is more strongly reducing than copper metal as shown by the fact that the standard (reduction) potential for zinc is more negative than that of copper. Therefore zinc metal will tend to loose electrons to copper ions and develop a positive electrical charge. The equilibrium constant, K, for the cell is given by

\ln K= \frac{nFE^\Theta}{RT}

where F is the Faraday, R is the gas constant and T is the temperature in Kelvin. For the Daniell cell K is approximately equal to 1.5×1037. Thus, at equilibrium, a few electrons are transferred, enough to cause the electrodes to be charged.[4]

Actual cell potentials must be calculated by using the Nernst equation as the solutes are unlikely to be in their standard states,

E_{\text{half-cell}} = E^\Theta - \frac{RT}{nF}\ln_e Q

where Q is the reaction quotient. This simplifies to

E_{\text{half-cell}} = E^\Theta - 2.303 \frac{RT}{nF} \log_{10} \{ M^{n+}\}

where {Mn+} is the activity of the metal ion in solution. The metal electrode is in its standard state so by definition has unit activity. In practice concentration is used in place of activity. The potential of the whole cell is obtained by combining the potentials for the two half-cells, so it depends on the concentrations of both dissolved metal ions. The value of 2.303R/F is 0.19845, so at 25°C (298.15 K) the cell potential will change by 59.18/n mV if the concentration of a metal ion is increased or decreased by a factor of 10.

These calculations are based on the assumption that all chemical reactions are in equilibrium. When a current flows in the circuit, equilibrium conditions do not obtain and the cell potential will usually be reduced by various mechanisms, such as the development of overpotentials.[5] Also, since chemical reactions occur when the cell is producing power, the electrolyte concentrations change and the cell voltage is reduced. The voltage produced by a galvanic cell is temperature dependent because standard potentials are temperature-dependent.

Notation

The galvanic cell, as the one shown in the figure, are conventionally described using the following notation:

Zn(s) | ZnSO4(aq) || CuSO4(aq) | Cu(s)
(anode)----------------------------------(cathode)

An alternate notation for this cell would be:
Zn(s) | Zn+2(aq) || Cu+2(aq) | Cu(s)

Where the following applies:

  • (s) denotes solid.
  • (aq) means aqueous solution.
  • The vertical bar, |, denotes a phase boundary.
  • The double vertical bar, ||, denotes a liquid junction for which the junction potential is near zero, such as a salt bridge.[6]

Galvanic corrosion

Main article: Galvanic corrosion

Galvanic corrosion is a process that degrades metals electrochemically. This corrosion occurs when two dissimilar metals are placed in contact with each other in the presence of an electrolyte, such as salt water, forming a galvanic cell. A cell can also be formed if the same metal is exposed to two different concentrations of electrolyte. The resulting electrochemical potential then develops an electric current that electrolytically dissolves the less noble material.

Cell types

See also

References

  1. ^ Merriam-Webster Online Dictionary: "battery"
  2. ^ "battery" (def. 4b), Merriam-Webster Online Dictionary (2008). Retrieved 6 August 2008.
  3. ^ Keithley, Joseph F. (1999). Daniell Cell. John Wiley and Sons. pp. 49–51. ISBN 0780311930. 
  4. ^ Atkins, P; de Paula (2006). Physical Chemistry. J. (8th. ed.). Oxford University Press. ISBN 9780198700722.  Chapter 7, sections on "Equilibrium electrochemistry"
  5. ^ Atkins, P; de Paula (2006). Physical Chemistry. J. (8th. ed.). Oxford University Press. ISBN 9780198700722.  Section 25.12 "Working Galvanic cells"
  6. ^ Atkins, P., "Physical Chemistry", 6th edition, W.H. Freeman and Company, New York, 1997

External links


v • d • e
Galvanic cells
Non-rechargeable:
primary cells		 Galvanic cell
Rechargeable:
secondary cells
Kinds of cells
Parts of cells

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