Simpo PDF Merge and Split Unregistered Version - http://www.com Simpo PDF Merge and Split Unregistered Version - http://www.com 15 Electroanalytical Chemistry Electrochemistry is the area of chemistry that studies the interconversion of chemical energy and electrical energy. Electroanalytical chemistry is the use of electrochemical techniques to characterize a sample. The original analytical applications of electrochem- istry, electrogravimetry and polarography, were for the quantitative determination of trace metals in aqueous solutions. The latter method was reliable and sensitive enough to detect concentrations as low as 1 ppm of many metals.
Since that time, many different types of electrochemical techniques have evolved, each useful for particular applications in organic, inorganic, and biochemical analyses. A species that undergoes reduction or oxidation is known as an electroactive species. Electroactive species in general may be solvated or complexed, ions or molecules, in aqueous or nonaqueous solvents. Electrochemical methods are now used not only for trace metal ion analyses, but also for the analysis of organic compounds, for continuous process analysis, and for studying the chemical reactions within a single living cell.
Appli- cations have been developed that are suited for quality control of product streams in indus- try, in vivo monitoring, materials characterization, and pharmaceutical and biochemical studies, to mention a few of the myriad applications. Under normal conditions, concen- trations as low as 1 ppm can be determined without much difficulty. By using electrode- position and then reversing the current, it is possible to extend the sensitivity limits for many electroactive species by three or four orders of magnitude, thus providing a means of analysis at the ppb level. In practice, electrochemistry not only provides a means of elemental and mol- ecular analysis, but also can be used to acquire information about equilibria, kinetics, and reaction mechanisms from research using polarography, amperometry, conducto- metric analysis, and potentiometry.
The analytical calculation is usually based on the determination of current or voltage or on the resistance developed in a cell under conditions such that these are dependent on the concentration of the species under study. Electrochemical measurements are easy to automate because they are electrical signals. The equipment is often far less expensive than spectroscopy instru- mentation. Electrochemical techniques are also commonly used as detectors for LC, as discussed in Chapter 13.
Gale, Department of Chemistry, Louisiana State University, Baton Rouge, LA 919 920 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www. FUNDAMENTALS OF ELECTROCHEMISTRY Electrochemistry is the study of reduction –oxidation reactions (called redox reactions) in which electrons are transferred from one reactant to another. A chemical species that loses electrons in a redox reaction is oxidized. A species that gains electrons is reduced.
A species that oxidizes is also called a reducing agent because it causes the other species to be reduced; likewise, an oxidizing agent is a species that is itself reduced in a reaction. An oxidation – reduction reaction requires that one reactant gain electrons (be reduced) from the reactant which is oxidized. We can write the reduction and the oxidation reactions separately, as half-reactions; the sum of the half-reactions equals the net oxidation –reduction reaction. Examples of oxidation half-reactions include: Fe2þ ! Fe3þ þ e Cu(s) ! Cu2þ þ 2e AsH3 (g) ! As(s) þ 3Hþ þ 3e H2 C2 O4 ! 2CO2 (g) þ 2Hþ þ 2e Examples of reduction half-reactions include: Co3þ þ e ! Co2þ 1 (IO3 ) þ 6Hþ þ 5e ! I2 (s) þ 3H2 O 2 Cl2 (g) þ 2e ! 2Cl Agþ þ e ! Ag(s) If the direction of an oxidation reaction is reversed, it becomes a reduction reaction; that is, if Al3þ accepts 3 electrons, it is reduced to Al(s).
All of the reduction reactions are oxi- dation reactions if they are written in the opposite direction. Many of these reactions are reversible in practice, as we shall see. A net oxidation – reduction reaction is the sum of the appropriate reduction and oxi- dation half-reactions. If necessary, the half-reactions must be multiplied by a factor so that no electrons appear in the net reaction.
For example, the reaction between Cu(s), Cu2þ, Ag(s), and Agþ is: Cu(s) þ 2Agþ ! Cu2þ þ 2Ag(s) We shall see why the reaction proceeds in this direction shortly. The net reaction is obtained from the half-reactions as follows: Oxidation reaction: Cu(s) ! Cu2þ þ 2e Reduction reaction: Agþ þ e ! Ag(s) Each mole of copper gives up 2 moles of electrons, while each mole of silver ion accepts only 1 mole of electrons. Therefore the entire reduction reaction must be multiplied by 2, so that there are no electrons in the net reaction after summing the half-reactions: Oxidation reaction: Cu(s) ! Cu2þ þ 2e Reduction reaction: 2(Agþ þ e ! Ag(s)) Net reaction: Cu(s) þ 2Agþ ! Cu2þ 2Ag(s) The equal numbers of electrons on both sides of the arrow cancel out. Electroanalytical Chemistry 921 Simpo PDF Merge and Split Unregistered Version - http://www.com Electrochemical redox reactions can be carried out in an electrochemical cell as part of an electrical circuit so that we can measure the electrons transferred, the current, and the voltage.
Each of these parameters provides us with information about the redox reaction, so it is important to understand the relationship between charge, voltage, and current. The absolute value of the charge of one electron is 1.602 10219 coulombs (C); this is the fundamental unit of electric charge.602 10219 C is the charge of one electron, the charge of one mole of electrons is: (1:602 1019 C/e )(6:022 1023 e =mol) ¼ 96,485 C/mol (15:1) This value 96,485 C/mol is called the Faraday constant (F), and provides the relationship between the total charge, q, transferred in a redox reaction and the number of moles, n, involved in the reaction. q¼nF (15:2) In an electric circuit, the quantity of charge flowing per second is called the current, i. The unit of current is the ampere, A; 1 A equals 1 C/s.
The potential difference, E, between two points in the cell is the amount of energy required to move the charged electrons between the two points. If the electrons are attracted from the first point to the second point, the electrons can do work. If the second point repels the electrons, work must be done to force them to move. Work is expressed in joules, J, and the potential difference, E, is measured in volts.
The relationship between work and potential difference is: w (in joules) ¼ E (in volts) q (in coulombs) (15:3) Since the unit of charge is the coulomb, 1 V equals 1 J/C. The relationship between current and potential difference in a circuit is expressed by Ohm’s Law: E i¼ (15:4) R where i is the current; E, the potential difference, and R, the resistance in the circuit. The units of resistance are V/A or ohms, V. ELECTROCHEMICAL CELLS At the heart of electrochemistry is the electrochemical cell.
We will consider the creation of an electrochemical cell from the joining of two half-cells. When an electrical conductor such as a metal strip is immersed in a suitable ionic solution, such as a solution of its own ions, a potential difference (voltage) is created between the conductor and the solution. This system constitutes a half-cell or electrode (Fig. The metal strip in the solution is called an electrode and the ionic solution is called an electrolyte.
We use the term elec- trode to mean both the solid electrical conductor in a half-cell (e., the metal strip) and the complete half-cell in many cases, for example, the standard hydrogen electrode, the calomel electrode. Each half-cell has its own characteristic potential difference or elec- trode potential. The electrode potential measures the ability of the half-cell to do work, or the driving force for the half-cell reaction. The reaction between the metal strip and the ionic solution can be represented as M0 ! Mnþ þ ne (15:5) 922 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.1 A half-cell composed of a metal electrode M0 in contact with its ions, Mþ, in solution.
The salt bridge or porous membrane is shown on the lower right side. where M0 is an uncharged metal atom, Mnþ is a positive ion, and e2 is an electron. The number of electrons lost by each metal atom is equal to n, where n is a whole number. This is an oxidation reaction, because the metal has lost electrons.
It has been oxidized from an uncharged atom to a positively charged ion. In the reaction, the metal ions enter the sol- ution (dissolve). By definition, the electrode at which oxidation occurs is called the anode. We say that at the anode, oxidation of the metal occurs according to the reaction shown in Eq.
Some examples of this type of half-cell are: Cd(s) ! Cd2þ þ 2e Ag(s) ! Agþ þ e Cr(s) ! Cr3þ þ 3e Note that in normal usage, the zero oxidation state of the solid metal is understood, not shown with a zero superscript. It has been found that with some metals the spontaneous reaction is in the opposite direction and the metal ions tend to become metal atoms, taking up electrons in the process. This reaction can be represented as Mnþ þ ne ! M0 (15:6) This is a reduction reaction because the positively charged metal ions have gained elec- trons, lost their charge, and become neutral atoms. The neutral atoms deposit on the electrode, a process called electrodeposition.
This electrode is termed a cathode. At the cathode, reduction of an electroactive species takes place. An electroactive species is one that is oxidized or reduced during reaction. Electrochemical cells also contain nonelectroactive (or inert) species such as counterions to balance the charge, or electri- cally conductive electrodes that do not take part in the reaction.
Often these inert electrodes are made of Pt or graphite, and serve only to conduct electrons into or out of the half-cell. It is not possible to measure directly the potential difference of a single half-cell. However, we can join two half-cells to form a complete cell as shown in Fig. In this example, one half-cell consists of a solid copper electrode immersed in an aqueous solution of CuSO4 ; the other has a solid zinc electrode immersed in an aqueous solution Electroanalytical Chemistry 923 Simpo PDF Merge and Split Unregistered Version - http://www.2 A complete Zn/Cu galvanic cell with a salt bridge separating the half-cells.
The two half-cell reactions and the net spontaneous reaction are shown: Anode (oxidation) reaction: Zn(s) ! Zn2þ þ 2e Cathode (reduction) reaction: Cu2þ þ 2e ! Cu(s) Net reaction: Zn(s) þ Cu2þ ! Zn2þ þ Cu(s) No reaction will take place, and no current will flow, unless the electrical circuit is com- plete. As shown in Fig.2, a conductive wire connects the electrodes externally through a voltmeter (potentiometer). A salt bridge, a glass tube filled with saturated KCl in agar gel, physically separates the two electrolyte solutions. The salt bridge permits ionic motion to complete the circuit while not permitting the electrolytes to mix.
The reason we need to prevent the mixing of the electrolytes is that we want to obtain information about the electrochemical system by measuring the current flow through the external wire. If we had both electrodes and both ionic solutions in the same beaker, the copper ions would react directly at the Zn electrode, giving the same net reaction but no current flow in the external circuit. In the electrolyte solution and the salt bridge the current flow is ionic (ion motion), and in the external circuit the current flow is electronic (electron motion). This cell and the one in Fig.3 show the components needed for an electrochemical cell: two electrical conductors (electrodes), suitable electrolyte solutions and a means of allowing the movement of ions between the solutions (salt bridge in Fig.2, a semipermeable glass frit or membrane in Fig.