السلام عليكم ورحمة الله وبركاته....
انا بحاجة الى معلومات عنelectrode potential وgalvanic cell و الفرق بينها وبينcell membrane وcopper electrode... رجاء من يعرف يرسلها مع مصادرها,ولكن بالانجليزية.... تحياتي
عرض للطباعة
السلام عليكم ورحمة الله وبركاته....
انا بحاجة الى معلومات عنelectrode potential وgalvanic cell و الفرق بينها وبينcell membrane وcopper electrode... رجاء من يعرف يرسلها مع مصادرها,ولكن بالانجليزية.... تحياتي
Galvanic CellGalvanic cell (also called voltaic cell) uses chemical reaction to produce electrical energy (flow of electrons).
http://web.fccj.org/~ethall/2046/ch18/zncu2.gifWhen zinc ****l placed in CuSO4 solution, following reaction take place:
Zn(s) + CuSO4(aq)http://web.fccj.org/~ethall/2046/ch18/arrow.gif ZnSO4(aq) + Cu(s)
Oxidation: Zn(s)http://web.fccj.org/~ethall/2046/ch18/arrow.gif Zn+2 + 2e-1
Reduction: Cu+2 + 2e-1http://web.fccj.org/~ethall/2046/ch18/arrow.gif Cu
Overall: Zn(s) + Cu+2http://web.fccj.org/~ethall/2046/ch18/arrow.gif Zn+2 + Cu(s)
Electrons will not flow in the following apparatus:
http://web.fccj.org/~ethall/2046/ch18/zncu3.gif
But if reaction carried out in the apparatus shown in Figure 21.2 (Cu/Ag system), electrons are transferred from Zn° to Cu+2 through a wire producing electrical energy.
http://web.fccj.org/~ethall/2046/ch18/zncu.gif
The salt bridge is necessary to complete the circuit and maintain charge neutrality.
Zn+2 http://web.fccj.org/~ethall/2046/ch18/arrow.gif cathode
SO4-2 http://web.fccj.org/~ethall/2046/ch18/arrow.gif anode
AnodeCathodeOxidation occursReduction occursEectrons producedElectrons are consumedAnions migrate towardCations migrate towardHas negative signHas positive sign
Shorthand Notation for:
Zn° + Cu+2http://web.fccj.org/~ethall/2046/ch18/arrow.gifZn+2 + Cu°
http://web.fccj.org/~ethall/2046/ch18/short.gif
Example
Write shorthand notation for: Fe(s) + 2Fe+3(aq)http://web.fccj.org/~ethall/2046/ch18/arrow.gif 3Fe+2(aq)
Fe°http://web.fccj.org/~ethall/2046/ch18/absolute.gifFe+2http://web.fccj.org/~ethall/2046/ch18/absolute.gifhttp://web.fccj.org/~ethall/2046/ch18/absolute.gifFe+3http://web.fccj.org/~ethall/2046/ch18/absolute.gifFe+2 Write shorthand notation for: 2Ag+1(aq) + Ni(s)http://web.fccj.org/~ethall/2046/ch18/arrow.gif 2Ag(s) + Ni+2(aq)
Ni°http://web.fccj.org/~ethall/2046/ch18/absolute.gifNi+2http://web.fccj.org/~ethall/2046/ch18/absolute.gifhttp://web.fccj.org/~ethall/2046/ch18/absolute.gifAg+1http://web.fccj.org/~ethall/2046/ch18/absolute.gifAg°
هذا ما توصت اليه عن galvanic cell......
Electrode Potential and Galvanic Corrosion
The potential difference between an anode and a cathode can be measured by a voltage measuring device. The absolute potential of the anode and cathode cannot be measured directly. Defining a standard electrode, all other potential measurements can be made against this standard electrode. If the standard electrode potential is set to zero, the potential difference measured can be considered as the absolute potential.
Standard Hydrogen Electrode
The half-cell in which the hydrogen reaction takes place is called the Standard Hydrogen Electrode - SHE
Standard Electrode Potential
The potential difference measured between ****l M, and the Standard Hydrogen Electrode - SHE
The electrochemical series consists of a list of ****ls which have been arranged in order of their standard electrode potentials.
هذا ا توصلت اليه عن electrode potential......
The Copper-copper(II) sulfate electrode is a type of reference electrode, based on the redox reaction with participation of the ****l (copper) and its salt - copper(II) sulfate. It is used for measuring electrochemical potential and is the most commonly used reference electrode for testing cathodic protectioncorrosion control systems.Electrochemistry deals with the links between chemical reactions and electricity. This includes the study of chemical changes caused by the passage of an electric current across a medium, as well as the production of electric energy by chemical reactions. Electrochemistry also embraces the study of electrolyte solutions and the chemical equilibria that occur in them.
The corresponding equation can be presented as follow:
Cu2+ + 2e- → Cu0(****l)
This reaction characterized by fast electrode kinetics, meaning that a sufficiently high current can be passed through the electrode with the 100% efficiency of the redox reaction (dissolution of the ****l or cathodic deposition of the copper-ions).
The Nernst equation below shows the dependence of the potential of the copper-copper(II) sulfate electrode on the activity or concentration copper-ions:
Commercial reference electrodes consist of a plastic tube holding the copper rod and saturated solution of copper sulfate. A porous plug on one end allows contact with the copper sulfate electrolyte. The copper rod protrudes out of the tube. A voltmeter negative lead is connected to the copper rod.
The potential of a copper copper sulfate electrode is +0.314 volt with respect to the standard hydrogen electrode
Many chemical reactions require the input of energy. Such reactions can be carried out at the surfaces of electrodes in cells connected to external power supplies. These reactions provide information about the nature and properties of the chemical species contained in the cells, and can also be used to synthesize new chemicals. The production of chlorine and aluminum and the electroplating and electrowinning of ****ls are examples of industrial electrochemical processes. Electrochemical cells that produce electric energy from chemical energy are the basis of primary and secondary (storage) batteries and fuel cells. Other electrical phenomena of interest in chemical systems include the behavior of ionic solutions and the conduction of current through these solutions, the separation of ions by an electric field ( electrophoresis ), the corrosion and passivation of ****ls, electrical effects in biological systems (bioelectrochemistry), and the effect of light on electrochemical cells (photoelectrochemistry).
Electrochemical Cells
An electrochemical cell generally consists of two half-cells, each containing an electrode in contact with an electrolyte. The electrode is an electronic conductor (such as a ****l or carbon) or a semiconductor. Current flows through the electrodes via the movement of electrons. An electrolyte is a phase in which charge is carried by ions. For example, a solution of table salt (sodium chloride, NaCl) in water is an electrolyte containing sodium cations (Na + ) and chloride anions (Cl − ). When an electric field is applied across this solution, the ions move: Na + toward the negative side of the field and Cl − toward the positive side.
The half-cells are connected by a cell separator that allows ions to move between the half-cells but prevents mixing of the electrolytes. The separator can consist of a salt bridge, or tube of aqueous solution plugged at both ends with glass wool, or it can be an ion exchange membrane or a sintered-glass disk. In some cases both half-cells use the same electrolyte, so that the electrochemical cell consists of two electrodes in contact with a single electrolyte. Electrochemical cells are usually classified as either galvanic or electrolytic. In galvanic cells, reactions occur spontaneously at the electrode–electrolyte interfaces when the two electrodes are connected by a conductor such as a ****l wire. Galvanic cells convert chemical energy to electric energy and are the components of batteries, which usually contain several cells connected in series. In electrolytic cells, reactions are forced to occur at the electrode–electrolyte interfaces by way of an external source of power connected to both electrodes. Electric energy from the external source is converted to chemical energy in the form of the products of the electrode reactions.
The galvanic cell shown in Figure 1 is known as the Daniell cell and was used as an early source of energy. It consists of a zinc (Zn) electrode in contact with an aqueous zinc sulfate solution and a copper (Cu) electrode in contact with an aqueous copper sulfate solution. When the external switch is closed, an atom of zinc on the zinc electrode is oxidized to zinc ion, liberating two electrons.
Zn → Zn 2+ + 2 e − (1)
The electrons pass through the external wire and reduce a copper ion to an atom of copper ****l on the surface of the copper electrode.
Figure 1. Daniell Cell.
Cu 2+ + 2 e − → Cu (2)
The electron flow in the external circuit represents an electric current produced by the cell. Ions flow within the electrolytes and across the salt bridge, as shown in the figure, to prevent an imbalance of ionic charge in the solutions that could result from the occurrence of these two electrode reactions. The overall cell reaction is the reduction of copper ion by zinc.
Cu 2+ + Zn → Cu + Zn 2+ (3)
The electrolytic cell shown in Figure 2 is the industrial chloralkali cell in which brine (an aqueous sodium chloride solution) is electrolytically converted to chlorine and caustic soda (sodium hydroxide, NaOH). The external power source supplies electric energy to drive the overall reaction.
2Cl − + 2H 2 O → Cl 2 + H 2 + 2OH − (4)
Chloride ion is oxidized to chlorine gas at the carbon electrode, and water is reduced to hydrogen gas (H 2 ) and hydroxide ion (OH − ) at the iron electrode. The electrolytes are maintained as electrically neutral by a flow of sodium ions through the separator (such as an ion exchange membrane).
The electrode where oxidation occurs, the zinc electrode in Figure 1 and the carbon electrode in Figure 2, is called the anode, while the electrode where reduction occurs is called the cathode. Reactions (1) and (2) are known as half-reactions, whereas reactions (3) and (4) are called oxidation-reduction (redox) reactions.
Electrode Potentials
Current and potential (or voltage ) are the two electrical variables of greatest interest in electrochemical cells. Current is related to the rate of the electrode
reactions, and the potential, to the cell energetics. Current is measured in amperes (A), or the amount of electricity in coulombs (C) that passes across a medium per second(s). Potential between the two electrodes is measured in volts (V) with a voltmeter. Potential (V) has units of energy or work (joules, J) per amount of electric charge (C). That is, 1 V = 1 J/C, so that the cell potential is a measure of the energy of the cell reaction. The cell is said to be at open circuit when no current flows; that is, when there are no external connections to the electrodes. Under these conditions, no electrode reactions occur.
Figure 2. Electrolytic Cell.
Measurements of the potentials of galvanic cells at open circuit give information about the thermodynamics of cells and cell reactions. For example, the potential of the cell in Figure 1, when the solution concentrations are 1 molar (1 M) at 25°C, is 1.10 V. This is called the standard potential of the cell and is represented by E°. The available energy (the Gibb’s free energy Δ G °) of the cell reaction given in equation (3) is related to E° by
Δ G ° = − nF E° (5)
where n is the number of electrons transferred in the reaction (in this case two) and F is a proportionality constant, called the Faraday (96,485 coulombs/equivalent). The cell potential is the difference in potential of the two half-cells. Tables of standard electrode potentials of half-reactions have been compiled; representative values are given in Table 1. These are frequently tabulated with respect to the standard or normal hydrogen electrode (SHE or NHE), which is arbitrarily assigned a half-cell potential of zero. Thus the value, +0.34 V, is assigned to the half-reaction Cu 2+ + 2 e − → Cu and
E° = +0.34 V vs NHE (6)
Similarly, the standard potential for the Zn/Zn 2+ cell yields Zn 2+ + 2 e − → Zn and
Table 1.
REPRESENTATIVE STANDARD POTENTIALS
Half Reaction
E o vs NHE
SOURCE: Bard, Allen J. et al., eds. (1985). Standard Potentials in Aqueous Solutions. New York: Marcel Dekker.
Li + + e − → Li
−3.045
Mg 2+ + 2e − → Mg
−2.356
Al 3+ + 3e − → Al
−1.67
Zn 2+ + 2e − → Zn
−0.7626
Cr 3+ + e − → Cr 2+
−0.424
2H + + 2e − → H 2
0.000
Cu 2+ + 2e − → Cu
0.340
O 2 + 4H + + 4e − → 2H 2 O
1.229
F 2 + 2e − → 2F −
2.87
E° = −0.76 V vs NHE (7)
The difference between these two half-cell potentials yields the standard potential of the Zn-Cu cell.
The standard potential applies to a half-cell when all the reactants are present at unit activity; that is, when the solution species are near a concentration of 1 molar. The actual half-cell potential E is a function of the solution concentrations and is related to these and to the standard potential E° by the Nernst equation. The Nernst equation for the generalized half-reaction
Oxidized Species + ne − → Reduced Species (8)
is
(9)
where R is the universal gas constant, T is the temperature on the Kelvin (absolute) scale, and the terms [Oxidized Species] and [Reduced Species] denote the activities of the species in the half-cells. The activities of pure solids and liquids are taken as unity. At 25°C, ( RT/F ) = 0.025 V.
Common Batteries
In most flashlights, toys, and remote controllers for televisions, primary batteries are used. The cell reactions in primary batteries are irreversible. During use, reactants are converted to products, and when the reactants are used up, the battery is "dead." The inexpensive flashlight batteries sold in retail stores use a design called a Leclanche dry cell. The body of the battery is made of zinc, which acts as the anode. A carbon rod in the center of the cell serves as the cathode. It is surrounded by a moist paste of graphite powder (carbon), manganese dioxide (MnO 2 ), and ammonium chloride (NH 4 Cl). The anode reaction is the oxidation of the zinc cylinder to zinc ions. The cathode reaction involves the reduction of manganese dioxide. A simplified version of the overall reaction is
Zn + 2MnO 2 + H 2 O → Zn 2+ + Mn 2 O 3 + 2OH − (10)
Alkaline cells are similar, except that the zinc case is porous and the paste around the carbon cathode is moist manganese dioxide and potassium hydroxide. These are more expensive than ordinary zinc-carbon cells, but they maintain a high voltage longer.
The lead-acid storage battery used in automobiles is a secondary battery; it is rechargeable. That is, the automobile battery operates as a galvanic cell when used to start the engine (when discharging), and as an electrolytic cell when it is charged by the alternator or by an external battery charger. The anode consists of porous lead plates in contact with a sulfuric acid (H 2 SO 4 ) solution. The cathode consists of lead dioxide (PbO 2 ) plates, also in sulfuric acid. Electrons flow from the lead plates to the lead oxide plates. As lead (Pb) loses electrons, it forms lead ions (Pb 2+ ) that react with sulfate ions (SO 4 2− ) in solution to form insoluble lead sulfate (PbSO 4 ). When PbO 2 gains electrons, it too reacts with SO 4 2− ions in solution to form solid PbSO 4 . The cell reaction is
Pb + PbO 2 + 4H + + 2 SO 4 2− → 2PbSO 4 + 2H 2 O (11)
and proceeds from left to right when the battery is discharging and from right to left when charging.
The rechargeable nickel-cadmium (Ni-Cad) batteries are used in a variety of cordless appliances such as telephones, battery operated tools, and portable computers. During discharge, cadmium ****l (Cd) acts as the anode, and nickel dioxide (NiO 2 ) as the cathode. Both ****ls form insoluble hydroxides due to the presence of the potassium hydroxide electrolyte. The cell reaction during discharge is
NiO 2 + Cd + 2H 2 O → Ni(OH) 2 + Cd(OH) 2 (12)
The reaction is reversed during charging.
SEE ALSO A LUMINUM ; F ARADAY , M ICHAEL .
Cynthia G. Zoski
Bibliography
Bard, Allen J.; Parsons, Roger; and Jordan, Joseph, eds. (1985). Standard Potentials in Aqueous Solutions. New York: Marcel Dekker.
Bard, Allen J., and Faulkner, Larry R. (2000). Electrochemical Methods: Fundamentals and Applications. New York: Wiley.
Bockris, J. M., and Reddy, A. K. N. (1998). Modern Electrochemistry. New York: Plenum.
Oldham, Keith B., and Myland, Janice C. (1994). Fundamentals of Electrochemical Science. New York: Academic Press.
Rieger, Philip H. (1994). Electrochemistry, 2nd edition. New York: Chapman and Hall.