323063903 Electroanalytical Chemistry Nadya Ver 2 Not Mine PDF

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Chapter 22, 23, 24 & 25 Electroanalytical Chemistry

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-  4OH4Ag + 4Cl-  4AgCl + 4eHydrogen Probe: A.P. +650 mV

Reactions: H2O2  O2 +2H+ +2e2Ag+ + 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.

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:

-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



The Daniell Cell

This cell is based on the overall reaction [Cu(OH ) ] + + Zn --> Cu + [Zn(OH ) ] + 2 6

2

(aq)

2 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(OH ) ] + || [Cu(OH ) ] + | Cu or just as Zn | Zn(II) || Cu(II) | Cu 2 6

2

(aq)

(aq)

2 6

2

(aq)

(aq)

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:

-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+ + 2e2+ 3+ Fe Fe + eH2(g) 2H+ + 2eAg(s) + Cl- AgCl(s) + e22A-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.

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 halfcells 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 halfreaction. 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. 1. Measured against standard hydroden electrode.

E

2. Concentration 1 Molar

° Cell

3. Pressure 1 atmosphere 4. Temperature 25°C

The example below shows some of the extreme values for standard cell potentials. Cathode (Reduction) Half-Reaction

Standard Potential E° (volts)

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

-3.04

+

-...