Chapter 33



Chapter 22, 23, 24 & 25

Electroanalytical Chemistry

[pic]

Why electroanalytical chemistry?

Electroanalytical methods have certain advantages over other analytical methods. Electrochemical analysis allows for the determination of different oxidation states of an element in a solution, not just the total concentration of the element. Electroanalytical techniques are capable of producing exceptionally low detection limits and an abundance of characterization information including chemical kinetics information. The other important advantage of this method is its low cost.

History:

Polarography was first discovered by a czechoslovavian chemist by the name of Heyrovsky in 1920. He won the Nobel prize for it in 1959. He proposed that the current recording generated by a oxidation or reduction in a cell as the A.P. is continuously increased:

Ox + ne ( Red

Oxygen Probe:

A.P. – 650 mV Ag|AgCl

Reactions:

O2 + 2H2O + 4e- ( 4OH-

4Ag + 4Cl- ( 4AgCl + 4e-

Hydrogen Probe:

A.P. +650 mV

Reactions:

H2O2 ( O2 +2H+ +2e-

2Ag+ + 2e- ( 2Ag

The following comprehensive chart below shows the applications of electrochemistry and its interactions with other branches of science and technology.

Electrochemical Cells

Extremely important classes of oxidation and reduction reactions are used to provide useful electrical energy in batteries. A simple electrochemical cell can be made from copper and zinc metals with solutions of their sulfates. In the process of the reaction, electrons can be transferred from the zinc to the copper through an electrically conducting path as a useful electric current. An electrochemical cell can be created by placing metallic electrodes into an electrolyte where a chemical reaction either uses or generates an electric current. Electrochemical cells that generate an electric current are called voltaic cells or galvanic cells, and common batteries consist of one or more such cells. In other electrochemical cells an externally supplied electric current is used to drive a chemical reaction, which would not occur spontaneously. Such cells are called electrolytic cells.

• Voltaic Cells: A voltaic cell is an electrochemical cell that external electrical current flow can be created using any two different metals since metals differ in their tendency to lose electrons. Zinc more readily electrons than copper, so placing zinc and copper metal in solutions of their salts can cause electrons to flow through an external wire which leads from the zinc to the copper. The following is a diagram of a voltaic cell.

[pic]

As a zinc atom provides the electrons, it becomes a positive ion and goes into aqueous solution, decreasing the mass of the zinc electrode. On the copper side, the two electrons received allow it to convert a copper ion from solution into an uncharged copper atom which deposits on the copper electrode, increasing its mass. In order for the voltaic cell to continue to produce an external electric current, there must be a movement of the sulfate ions in solution from the right to the left to balance the electron flow in the external circuit. The metal ions themselves must be prevented from moving between the electrodes, so some kind of porous membrane or other mechanism must provide for the selective movement of the negative ions in the electrolyte from the right to the left.

• Galvanic Cells: A galvanic cell consists of at least two half cells, a reduction cell and an oxidation cell. Chemical reactions in the two half cells provide the energy for the galvanic cell operations. The reactions always run spontaneously in the direction that produced a positive cell potential. The following shows a picture of an old fashioned galvanic cell:

|[pic] |

-the net reaction in a cell is the sum of the two half reactions

-the potential is the measure of the tendency of the cell to

move towards equilibrium

-Galvanic cells react in a way that produces electrical energy

-Electrolytic cells consume energy

-Chemically reversible cell exists when reversing the

direction of current reverses the reaction at the electrodes

[pic]

• The Daniell Cell

This cell is based on the overall reaction

[Cu(OH2)6]2+(aq) + Zn --> Cu + [Zn(OH2)6]2+(aq)

and functions by dissolution of Zn from the anode and deposition of Cu at the cathode. It is therefore very simply represented as

Zn | [Zn(OH2)6]2+(aq) || [Cu(OH2)6]2+(aq) | Cu

or just as

Zn | Zn(II)(aq) || Cu(II)(aq) | Cu

22A-1. Conduction in a Cell:

-charge is conducted by three distinct processes:

1) in the electrodes and the external conductor, electrons

serve as carriers

2)within the solution the flow of electricity involves the

migration of both the cations and the anions

3)oxidation and reduction occurs at the 2 electrode surfaces

22A-2 Solution Structure- The Double Layer

It is very important to realize that electrochemical measurements involve heterogeneous systems because an electrode can only donate or accept electrons from a species that is present in a layer of solution that is immediately adjacent to the electrode.

22A-3 Faradaic and Nonfaradiac Currents

Two types of processes can conduct currents across and electrode solution interface. One kind involves a direct transfer of electrons via and oxidation reaction at one electrode and reduction reaction at the other. Processes of this type are called faradaic processes because they are govern by Faraday’s Law, which states that the amount of chemical reaction at an electrode is proportional to the current; the resulting currents are called faradaic currents.

22A-4 Mass Transfer in Cells with the Passage of Current

- three mechanisms bring about these mass transfer: convection, migration, and diffusion. Convection involves mechanical motion of the solution as a result of stirring or the flow of the solution past the surface of the electrode. Migration is the movement of ions through the solution brought about by electrostatic attraction between the ions and the charged electrode. Diffusion is the motion of species brought about by a concentration gradient.

22A-6 Anodes and Cathodes:

[pic]

-Cathode- the electrode where reduction occurs in an

electrochemical cell

-Anode- the electrode where oxidation occurs in an

electrochemical cell

-reactions at cathodes

-electrons supplied by external circuit via an inert

electrode (platinum or gold)

-some examples are:

Cu2+ + 2e- Cu(s)

Fe3+ + e- Fe2+

2H+ + 2e- H2(g)

AgCl(s) + e- Ag(s) + Cl-

-reactions at anodes

-some examples are:

Cu(s) Cu2+ + 2e-

Fe2+ Fe3+ + e-

H2(g) 2H+ + 2e-

Ag(s) + Cl- AgCl(s) + e-

22A-7 Cells without Liquid Junctions:

-liquid junction - the interface between 2 different

electrolytic solutions

-cell can contain more than one

-a small Junction Potential arises at these interfaces

-sometimes it is possible to prepare cells that share a

common electrolyte to avoid this problem

22A-8 Schematic Representation of Cells:

-anode and information on the solution it is contacting on left

-single vertical line indicates phase boundary where

potential might develop

-two vertical lines indicates liquid junction

-concentration or activity put in parentheses

-example:

Zn/ZnSO4(azn2+ = 1.00)//CuSO4(acu2+ + 1.00)/Cu

Principles of Electrolysis:

When an electric current is forced to pass through an electrolyte or electrolyte solution, chemical reactions take place both at the anode and at the cathode. The stoichiometry of these reactions obeys Faraday's laws of electrolysis. However, when several different reactions are possible at an electrode of an electrolytic cell then the process which actually does take place will be determined by the potential of the electrode. The reactions which take place in electrolytic cells are those reactions which require the least potential difference between the two electrodes.

The reactions which occur at the electrodes during electrolysis will be the oxidation and the reduction of the solvent unless some solute is more easily oxidized or reduced than is the solvent. Electrolysis of an aqueous solution of Na2SO4 will produce hydrogen at the cathode and oxygen at the anode because reduction of H+(aq) is easier than is reduction of Na+(aq) and oxidation of water is easier than is oxidation of aqueous sulfate ion. The actual potentials will be close to the standard potentials, and the standard potential of H+(aq)/H2(g) is less negative than is the standard potential of Na+(aq)/Na, while the standard potential of O2(g)/H2O is less positive than the standard potential of sulfate oxidation. The standard potential of sulfate oxidation is so positive that it is meaningless in aqueous solutions, in which no higher oxidation state than six has been found for sulfur.

Electrolysis of an aqueous solution of copper sulfate will produce copper at the cathode, because both the Cu2+/Cu+ and Cu+/Cu couples have standard potentials which are considerably less negative than the standard potential of hydrogen. The electrolysis will again produce oxygen at the anode, because Cu2+(aq) cannot be further oxidized in aqueous solution.

Electrolysis of an aqueous solution of NaBr will produce hydrogen at the cathode, as the aqueous sodium sulfate solution did, because neither Na+(aq) nor Br-(aq) can be reduced at potentials which are less negative than standard hydrogen. The electrolysis of aqueous NaBr solutions will produce bromine at the anode because the standard potential of bromide oxidation is less positive than that of the oxidation of water.

Electrolysis of a solution of aqueous CuBr2 will produce copper at the cathode and bromine at the anode because these are the least negative possible reduction and the least positive possible oxidation reactions in that solution. The anode reaction in any electrolysis will always be the oxidation which occurs at the least positive potential and the cathode reaction will always be the reduction which occurs at the least negative potential.

Potentials which are not reversible have values which depend on the current which flows through them. The difference between the reversible potential of an electrode and its actual potential when current does flow through it is called overpotential. The overpotential depends upon the current and other electrode conditions and its determination is beyond our scope. However, when currents through aqueous electrodes are small or moderate the overpotentials at those electrodes are generally not more than a few mV unless the electrode reaction consumes or generates a gas.

Electrode reactions involving gas evolution can have overpotentials as high as a full volt. Overpotential can lead to different electrode processes in the electrolysis of aqueous solutions.

Example. An aqueous solution is 1.0 molar in Pb2+ and 1.0 molar in Sn2+, as well as 1.0 molar in H+; the other ions present are electrochemically inert. The reactions which will occur on electrolysis and the reversible cell potential under standard conditions can be determined as follows.

The anode reaction will be Sn2+ --> Sn4+ + 2e- at +0.1539 V, since oxidation of water to produce oxygen could occur only at a much more positive potential. The cathode reaction at first would appear to be 2H+ + 2e- --> H2(g) at 0.000 V, and on platinum electrodes this might be observed, but on most metals hydrogen overpotential is at least 0.5 V. With any significant overpotential the cathode reaction will be Pb2+ + 2e- --> Pb at -0.1266 V, which will occur in preference to Sn2+ + 2e- --> Sn at -0.1410 V. The usual cell potential difference will then be DE0 = 0.2805 V, the lead electrode being negative.

Cell Potential Differences From Electrode Potentials

The standard potential difference across a cell, DE0, is the difference between the two standard electrode potentials, E0, of the two electrodes in the cell. The actual potential difference across a cell, DE, is the difference between the two actual electrode potentials, E, of the two electrodes in the cell, and this will be true whether or not the actual electrode potentials are standard or even reversible.

Potentials which are not reversible have values which depend on the current which flows through them. The difference between the reversible potential of an electrode and its actual potential when current does flow through it is called overpotential. The overpotential depends upon the current and other electrode conditions and it is generally beyond our scope. However, when currents through aqueous electrodes are small or moderate the overpotentials at those electrodes are generally not more than a few mV unless the electrode reaction consumes or generates a gas. Electrode reactions involving gases can have overpotentials as high as a full volt. In later sections, we shall see how overpotential can lead to different electrode processes in the electrolysis of aqueous solutions.

Concentration Cells

A concentration cell is an electrochemical cell in which the electrode couple at both electrodes is the same but the concentrations of substances at the two electrodes may differ. The potential difference across a concentration cell can be calculated using the Nernst equation. An example of a concentration cell is shown in the Figure below.

[pic]

Example. A concentration cell is based on the Fe3+/Fe2+ couple whose value of E0 is +0.769 V. The electrode on the right has a 0.1 molar concentration of Fe3+ and a 0.01 molar concentration of Fe2+. The electrode on the left has a 0.1 molar concentration of Fe2+ and a 0.01 molar concentration of Fe3+. We can determine which electrode is the more positive as follows.

The spontaneous reaction in the cell must be to make the concentrations of all species equal, so on the left the spontaneous reaction is Fe2+ --> Fe3+ + e- while on the right the spontaneous reaction is Fe3+ + e- --> Fe2+. In the external circuit electrons must then leave the electrode on the left and flow into the electrode on the right. Since electrons will flow from a location where they are in surplus, the left, to a place where they are in deficiency, the right, the electrode on the right must be the more positive.

This qualitative answer matches the quantitative answer obtained from the Nernst equation. For the electrode on the left,

E = +0.769 - 0.05915 log ([0.1]/[0.01]) = +0.769 - 0.05915 log 10 = +0.7098 V

For the electrode on the right,

E = +0.769 - 0.05915 log ([0.01]/[0.1]) = +0.769 + 0.05915 = +0.8282 V

The potential difference across the cell is 0.1186 V; the electrode on the right is the more positive.

Equilibrium Constants From Electrode Potentials:

One of the procedures used in quantitative analysis is the reduction of Fe(III) to Fe(II) by Sn(II); the reaction Sn2+ + 2Fe3+ --> 2Fe2+ + Sn4+ proceeds essentially quantitatively to the right. This is the direction we would expect it to go in an electrochemical cell made up of a standard tin and a standard iron half-cell. The tin is being oxidized at the anode and the iron is being reduced at the cathode. The chemical energy of the cell would be used to do electrical work in such a system, and the concentrations of Fe(III) and Sn(II) would decrease as this happened.

Now suppose that the two solutions in the separate half-cells were taken and mixed together, then the one solution divided into two halves and one half put in each half-cell. The potential difference between the two electrodes must now be zero since the two half-cells are identical. The energy which could have gone into electrical work has gone into heat instead. There is now equilibrium between the iron couple which would have been in one half-cell and the tin couple which would have been in the other half-cell. The concentrations of the forms of tin and iron present are the equilibrium concentrations and the potential of each electrode must be the equilibrium potential. In either half-cell, we could write the half-cell Nernst equation using either couple. So:

E(Sn) = +0.1539 - 0.05915/2 log ([Sn2+]/[Sn4+])

E(Fe) = +0.769 - 0.05915 log ([Fe2+]/[Fe3+])

We want to obtain the value of the equlibrium constant K for the reaction Sn2+ + 2Fe3+ --> Sn4+ + 2Fe2+, which by definition is

K = a(products)/a(reactants) = [Sn4+][Fe2+]2/[Sn2+][Fe3+]2.

Since at equilibrium there is only one E, that is, E = E(Sn) = E(Fe), we can then write:

0.1539 - 0.05915/2 log ([Sn2+]/[Sn4+]) = 0.769 - 0.05915 log ([Fe2+]/[Fe3+])

It is necessary to have [Fe2+] squared in the resulting equilibrium constant K, so the above equation is rewritten and then rearranged:

0.1539 - 0.05915/2 log ([Sn2+]/[Sn4+]) = 0.769 - 0.05915/2 log ([Fe2+]2/[Fe3+]2)

0.1539 - 0.769 = 0.05915/2 (log [Sn2+]/[Sn4+] - log [Fe2+]2/[Fe3+]2)

-0.6151 = 0.05915/2 log ([Sn2+][Fe3+]2/[Sn4+][Fe2+])2

0.6151 = 0.05915/2 log ([Sn4+][Fe2+]2/[Sn2+][Fe3+]2)

Then 0.6151 = 0.05915/2 log K, 1.2302 = 0.05915 log K, 20.80 = log K, and K = 6.28 x 10+20. This value is in fact very large as we know it would have to be in order for the reaction to proceed quantitatively as written.

Potentials in an Electroanalytical Cell:

Electroanalytical chemistry is a group of methods for qualitative and quantitative analysis based on the behavior of a solution of sample when it is made part of an electrochemical cell.

In an electrochemical cell, an electric potential is created between two dissimilar metals. This potential is a measure of the energy per unit charge that is available from the oxidation/reduction reactions to drive the reaction. It is customary to visualize the cell reaction in terms of two half-reactions, an oxidation half-reaction and a reduction half-reaction. The cell potential (often called the electromotive force or emf) has a contribution from the anode which is a measure of its ability to lose electrons - it will be called its "oxidation potential". The cathode has a contribution based on its ability to gain electeons, its "reduction potential". The cell potential can then be written:

Ecell = oxidation potential + reduction potential

It is important to understand that the potential of an electrochemical cell is related to the activities of the reactants and products of the cell reaction and indirectly to their molar concentrations.

If we could tabulate the oxidation and reduction potentials of all available electrodes, then we could predict the cell potentials of voltaic cells created from any pair of electrodes. Actually, tabulating one or the other is sufficient, since the oxidation potential of a half-reaction is the negative of the reduction potential for the reverse of that reaction. Two main hurdles must be overcome to establish such a tabulation

1. The electrode potential cannot be determined in isolation, but in a reaction with some other electrode.

2. The electrode potential depends upon the concentrations of the substances, the temperature, and the pressure in the case of a gas electrode.

In practice, the first of these hurdles is overcome by measuring the potentials with respect to a standard hydrogen electrode. It is the nature of electric potential that the zero of potential is arbitrary; it is the difference in potential which has practical consequence. Tabulating all electrode potentials with respect to the same standard electrode provides a practical working framework for a wide range of calculations and predictions. The standard hydrogen electrode is assigned a potential of zero volts. The second hurdle is overcome by choosing standard thermodynamic conditions for the measurement of the potentials. The standard electrode potentials are customarily determined at solute concentrations of 1 Molar, gas pressures of 1 atmosphere, and a standard temperature which is usually 25°C. The standard cell potential is denoted by a degree sign as a superscript.

|E°Cell |Measured against standard hydroden electrode. |

| |Concentration 1 Molar |

| |Pressure 1 atmosphere |

| |Temperature 25°C |

The example below shows some of the extreme values for standard cell potentials.

|Cathode (Reduction) |Standard Potential |

|Half-Reaction |E° (volts) |

|Li+(aq) + e- -> Li(s) |-3.04 |

|K+(aq) + e- -> K(s) |-2.92 |

|Ca2+(aq) + 2e- -> Ca(s) |-2.76 |

|Na+(aq) + e- -> Na(s) |-2.71 |

|Zn2+(aq) + 2e- -> Zn(s) |-0.76 |

|Cu2+(aq) + 2e- -> Cu(s) |0.34 |

|O3(g) + 2H+(aq) + 2e- -> O2(g) + H2O(l) |2.07 |

|F2(g) + 2e- -> 2F-(aq) |2.87 |

The values for the table entries are reduction potentials, so lithium at the top of the list has the most negative number, indicating that it is the strongest reducing agent. The strongest oxidizing agent is fluorine with the largest positive number for standard electrode potential.

-we will deal mainly with activities rather than concentration

therefore:

ax = fx [X]

where fx= activity coefficient of solute X and [X] is molar

concentration

-equilibrium constant for a reaction (a + b c + d + e)

K =ac x ad x ae/pb x aa

where a is the activity and p is the pressure in atmospheres

-a pure solid at unity gives this equation:

K = ac x ad/pb

-the second quantity Q:

Q= (ac)i(ad)i/(pb)i

-the change in free energy for the cell reaction:

(G= RT ln Q -RT ln K

-cell potential:

(G =-nFEcell

where n = # of moles and F = faraday(96,485C/mole of electrons)

Ecell:

Ecell = -RT/nF lnQ + RT/nF ln K

= E0cell - RT/nF ln (ac)i(ad)i/(pb)i

-standard electrode potential:

E0cell = RT/nF ln K

Electrode potentials:

Ecell = E(cathode) - E(anode)

-potentials are measured by difference, they are relative not

absolute

-SHE =Standard hydrogen electrode-1 atm 0volts at all temps.

-NHE = normal hydrogen electrode

-IUPAC says electrode potential refers to reduction 1/2 reactions

-in dilute solutions molar concentration can be used for

computations, rather than activities

-shifts in equilibria cause a shift in cell potential

-to compensate for activity effects and side reactions we can

use (Ef) formal potential instead of standard electrode potential

Calculation of Cell Potential from Electrode Potentials:

Nernst Equation: The cell potential for a voltaic cell under standard conditions can be calculated from the standard electrode potentials. But real voltaic cells will typically differ from the standard conditions. The Nernst equation relates the cell potential to its standard cell potential.

|[pic] |R = gas constant |

| |T = temperature in Kelvins |

| |Q = thermodynamic reaction quotient |

| |F = Faraday's constant |

| |n = number of electrons transferred |

The quantity Q, the thermodynamic reaction constant, is like a dynamic version of the equilibrium constant in which the concentrations and gas pressures are the instantaneous values in the reaction mixture. For a reaction

[pic]

the reaction quotient has the form

[pic]

where [C] is understood to be the molar concentration of product C, or the partial pressure in atmospheres if it is a gas.

Applied to the Daniell cell where zinc and copper form the electrodes, the reaction is

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

the form of Q is

[pic]

since the concentrations of the pure metal solids are assigned the value 1. This implies that the departure of the cell potential from its standard value of 1.10 volts will be influenced by the temperature and the ion concentrations.

|One implication is that the cell potential will be reduced from the standard value if |[pic] |

|the concentration of Zn2+(aq) is greater than that of Cu2+(aq) at the standard | |

|temperature. An excess concentration of Cu2+(aq) will give a higher voltage. The graph | |

|at right shows the increase in cell voltage with increasing concentration of the cation.| |

|Note that the horizontal axis is logarithmic, and that the straight line variation of | |

|the voltage represents an logarithmic variation with Q. Note that the cell potential is | |

|equal to the standard value if the concentrations are equal even if they are not equal | |

|to the standard value of 1M, since the logarithm gives the value zero. | |

|Consider a concentration of 10-5Molar for Zn2+(aq) and 0.1 Molar for Cu2+(aq) as a test |[pic] |

|case for temperature dependance. We can see that the cell potential tends to increase | |

|with temperature, or that a colder cell prodices less voltage - a commonly observed | |

|phenomenon with dry cell batteries. The variation with temperature is linear with | |

|temperature, but quite small for this cell. The large variations of practical output | |

|voltage with temperature for dry cells does not arise from the Nernst equation alone. | |

-calculated potentials are sometimes called thermodynamic

potentials

Ecell = E(cathode) - E(anode)

-negative Ecell indicates non-spontaneity of reaction

-Liquid Junction Potentials can be calculated from the knowledge

of the mobility of the two ions involved, but it is rare that the

system in question is simple enough for this computation

Currents in Electrochemical Cells:

-Ohms law is usually obeyed:

E=IR

where E is the potential difference in volts responsible for the

movement of the ions, I is the current in amps, and R is the

resistance in ohms of the electrolyte to the current

Electrochemical Power Sources:

There are three significant types of power sources which produce electricity by reaction within electrochemical cells. The two types which use reactants stored within them are called primary cells and secondary cells. Groups of primary or secondary cells are called batteries, although the term battery has been extended to include also a single cell used as a power source. Secondary cells, unlike primary cells, can be driven in reverse or charged by external electrical power. The third type, fuel cells, employ reactants which are continuously supplied to the cell; products are also continuously removed. In primary and secondary cells, the reactants and products are contained within the cell.

Primary Cells

Primary cells are electrochemical cells in which a spontaneous electrochemical reaction serves as a source of electrical power. A primary cell is not designed to be recharged after use and must be discarded when its internal supply of reactants is exhausted. Attempts to recharge sealed primary cells can produce explosions because the cells are not designed to withstand the pressure of gases generated within the cell by electrolysis. Primary cells are most useful in providing low-current DC power, particularly when the demand is intermittent. The oldest practical primary cell is the now-obsolete Daniell cell or wet cell which was used to power electrical telegraphs at the turn of the century. These cells consisted of a large glass jar (now called a battery jar) filled with two solutions of aqueous sulfuric acid. The lower half of the jar was filled with a denser solution of CuSO4 in the acid, while the upper half was filled with the acid alone. A zinc electrode was placed in the upper half of the cell and a copper electrode was placed in the lower half. This cell provided a voltage of about one volt. If the cell was connected constantly and not agitated, the difference in density between the solutions kept them from mixing for months. The overall cell reaction is Zn(s) + Cu2+(aq) --> Zn2+(aq) + Cu(s) in the cell:

(-) Zn(s)/Zn2+(aq),H2SO4(aq)//Cu2+(aq),H2SO4(aq)/Cu(s) (+)

The common flashlight cell or dry cell is the acidic Zn/MnO2 cell designed by Leclanche' in 1887. The structure of this cell is shown in the Figure below.

The zinc casing serves as the anode and is consumed in the anodic electrode reaction Zn(s) --> Zn2+ + 2e-; the zinc ion dissolves in the moist ZnCl2-NH4Cl electrolyte. A carbon rod serves as the cathode, but it is chemically inert. The cathodic electrode reaction which consumes MnO2 is best written as:

[Mn4+ + 2O2-] + H2O + e- --> [Mn3+ + O2- + OH-] + OH-

where the square brackets indicate the species present in the solid phase at the cathode. The cathode reaction actually occurs within the solid structure; the carbon rod serves only to transfer electrons from the external circuit.

The dry cell has a potential difference of about 1.25 V; the zinc electrode is negative. It is a good source of electrical power and the materials of construction are relatively cheap. The cell voltage during discharge falls off rather badly and the dry cell is not a good source of power when a constant voltage is needed.

The alkaline manganese cell, or alkaline cell, is a variant of the dry cell in which the acidic NH4Cl(aq) electrolyte is replaced by a concentrated solution of NaOH. The carbon electrode is replaced by a steel rod or brass tube. In this battery the anode reaction is Zn + 2OH- --> ZnO(c) + H2O + 2e-, since ZnO is insoluble in the alkaline electrolyte solution. Alkaline cells cost two to three times as much as their counterpart dry cells because a more elaborate internal construction is required to prevent leakage of the caustic electrolyte. They do, however, have almost twice the charge capacity and a higher available current. The voltage of alkaline manganese cells is the same as the voltage of dry cells, but under heavy current drain their output voltage decreases less. These advantages have led to their increasing use in battery-powered electronic equipment.

The mercury cell, also known as the Ruben cell, was developed in 1947. Its reaction on discharge is Zn(s) + HgO(s) --> ZnO(s) + Hg(l), and it takes place in the cell

(-) Zn(s)/KOH,ZnO(s),H2O/HgO(s),Hg(l) (+)

Mercury cells are much more expensive to manufacture than are dry cells or alkaline cells. This is due to the high cost of the mercury and mercury (II) oxide they contain. However, the discharge curve of mercury cells is very flat, which is to say that the potentials of mercury cells do not drop off significantly as the cell discharges, and the temperature coefficient of the mercury cells is also less. At 250C the cell voltage is 1.344 V with the zinc electrode negative. The constancy of the mercury cell voltage makes it useful in sensitive instruments but the high cost prevents its use as a general power source.

Other types of primary cells, such as the lithium cell, have been developed for special uses where long life, small size, and very low power are required. Cameras, heart pacemakers, and electronic wrist watches are powered by these cells.

Secondary Cells

The familiar automobile battery is a group of secondary cells connected in series. A 12-volt automobile battery consists of six lead-acid cells in series. Each lead-acid cell is represented as follows:

(-) Pb(s)/PbSO4(s),H2SO4(aq)//H2SO4(aq),PbO2(s)/Pb(s) (+)

Modern automobile batteries are remarkably rugged in both the physical and electrical sense. A typical automobile battery rated at 54 ampere-hours can provide currents as high as 300 A for short periods and a continuous drain of 25 A for over two hours.

The electrode reaction at the negative electrode of a lead-acid cell is PbSO4(s) + 2e- --> Pb(s) + SO42-(aq), for which E0 = -0.356 V, as measured by Kolthoff in 1931. At the positive electrode the reaction is PbO2(s) + SO42- + 4H+ + 2e- --> PbSO4(s) + 2H2O, for which E0 = +1.685 V, as measured by Harned and Hamer, also in 1931. The standard cell potential difference is then 2.041 V, and the spontaneous direction is for the PbO2 side to be positive. There is therefore in the external circuit a spontaneous electron flow from the lead metal side to the lead dioxide side which means that the spontaneous reaction must be, at the less positive or more negative electrode, Pb --> PbSO4(s), giving electrons. This is an oxidation process and so this electrode must be the anode. At the more positive electrode, the spontaneous reaction must be PbO2(s) --> PbSO4(s), taking electrons; this is a reduction process and so this electrode must be the cathode.

As the battery is being charged, the reaction is forced to proceed in the nonspontaneous direction Pb(II) --> Pb(IV) + Pb(0) and therefore the amount of PbSO4(s) decreases, the amount of H2SO4(aq) increases, the amount of PbO2(s) increases, and the amount of lead metal increases.

Another rechargeable secondary cell, used for hand-held devices such as electric razors, is the nickel-cadmium cell (NiCd cell or nicad cell). This cell contains an alkaline electrolyte. The spontaneous reactions on cell discharge at the anode and cathode are to the right:

2NiOOH(s) + 2H2O 2Ni(OH)2(s) + Cd(OH)2(s)

These reactions take place in the cell

(-) steel/Cd(s),Cd(OH)2(s)/LiOH(aq)/NiOOH(s),Ni(OH)2(s)/steel (+)

The discharge reactions are reversed when the cell is recharged. A fully charged nicad cell has a potential of about 1.4 V, and the cadmium terminal is negative. Nicad cells are available in standard flashlight cell sizes and can replace them in many applications. Nicad cells do lose stored charge (self-discharge) over a period of a few months and so they are not suitable for very intermittent use.

Fuel Cells

A fuel cell is an electrochemical cell to which reactants are continually supplied and from which products are continually removed so that there is a continuous production of electrical power. The best known fuel cell, and the most highly developed, is the hydrogen/oxygen fuel cell known as the Bacon cell, which is used in the United States space program. An idealized drawing of this cell is shown in the Figure below.

[pic]

The spontaneous reactions in this cell are oxidation of hydrogen at the anode and reduction of oxygen at the cathode:

Pt,H2(g) --> 2H+(aq) + e-; E0 = 0.0000 V

Pt,O2(g) + 4H+(aq) + 4e- --> 2H2O; E0 = +1.2288 V

The overall cell reaction, which yields 1.2288 V under standard conditions, is the formation of water, O2(g) + 2H2(g) --> 2H2O. Cost is a major problem with this cell; pure hydrogen, pure oxygen, and platinum are expensive. The platinum electrodes tend to pick up any trace of impurity in the oxygen or hydrogen and are poisoned by it, so that the cell efficiency is degraded severely.

Fuel cells have been suggested as future power sources, because they are not as restricted in efficiency as are power sources based on heat engines. The amount of heat obtainable and work obtainable from a fuel are roughly the same. However, the conversion of heat into work with 100% efficiency is not possible; the maximum efficiency possible depends on the temperature of the input and output heat. Efficiencies of 30 - 40% are commonly found in practice. The theoretical advantage in avoiding the heat step is about a factor of three, a considerable advantage if electrical power rather than heat is desired.

Example. Suppose we used methanol as a fuel in a fuel cell and in a heat engine. The net reaction in either case would be 2CH3OH(l) + 3O2(g) --> 2CO2(g) + 4H2O(g)

DG0 = 2(-394.359) + 4(-228.572) - 2(-166.36) - 3(0) = -1370.29 kJ

A total of -685.14 kJ/mole methanol of work are possible. The heat which could be obtained is -638.49 kJ/mole methanol:

DH0 = 2(-393.509) + 4(-241.818) - 2(-238.66) - 3(0) = -1276.97 kJ

Practical fuel cells using the oxidation of methane (natural gas) operate at 800oC to 1000oC with a molten carbonate salt serving as the electrolyte. Long-term operation of these cells is compromised by impurities in natural gas which react irreversibly with the molten carbonate electrolyte. For example, sulfur-containing impurities such as H2S are oxidized to sulfate ion which accumulates in the cell electrolyte.

Effect of current on cell potential:

-can result in

1) reduced potential of galvanic cell

2) increased potential needed to develop current in an

electrolytic cell:

-this is due to ohmic resistance and polarization effects such

as: charge transfer over-voltage, crystallization over-voltage

-to account for an ohmic potential, or IR drop:

Ecell = E(cathode) - E(anode) – IR

Polarization:

Polarization has occurred if the relationship Eapp = Ec – Ea – IR does not hold true. Cells that exhibit a nonlinear increase of current with potential are polarized. Complete polarization has occurred when I is independent of potential no matter how much more negative one makes the Eapp, I won’t change. Metallic surfaces can be polarized by the application of an external voltage or by the spontaneous production of a voltage away from equilibrium. This deviation from equilibrium potential is called polarization. The magnitude of polarization is usually described as an overvoltage (h) that is a measure of polarization with respect to the equilibrium potential (Eeq) of an electrode.

This polarization is said to be either anodic, when the anodic processes on the electrode are accelerated by changing the specimen potential in the positive (noble) direction or cathodic when the cathodic processes are accelerated by moving the potential in the negative (active) direction. There are three distinct types of polarization in any electrochemical cell, the total polarization across an electrochemical cell being the summation of the individual elements:

Eapp - Eeq = htotal = hact +hconc +iR

• Concentration Polarization- When the mass transfer limits

the rate of the reaction and therefor the current. Concentration polarization occurs when the concentration at the surface is not the same as that in the bulk, get less current than predicted. Stirring will decrease concentration polarization. As the voltage becomes much more negative all the reagent is reduced as soon as it reaches the surface, therefore the surface concentration of the reagent is zero.

• Reaction Polarization- When the rate of formation of the

intermediate limits the current

• When a physical process limits the reaction it is said to be

adsorption, desorption, crystallization etc. polarization

• Charge Transfer Polarization- When rate of electron transfer

from electrode to oxidized species or reduced species to

electrode limits the current. An overvoltage is required to observe the reaction.

• When the degree of polarization is measured by the overvoltage n

n = E - Eeq

where E is electrode potential and Eeq is thermodynamic or

equilibrium potential and E ................
................

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