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Unit 1 Introduction to Electro- Analytical Methods

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Unit 1 Introduction to Electro- Analytical Methods Introduction to UNIT 1 INTRODUCTION TO ELECTRO- Electroanalytical ANALYTICAL METHODS Methods Structure 1.1 Introduction Objectives 1.2 Basic Concepts Electrical Units Basic Laws of Electrochemistry Electrode Potential Liquid-Junction Potentials Electrochemical Cells The Nernst Equation Cell Potential 1.3 Classification and an Overview of Electroanalytical Methods Potentiometry Voltammetry Polarography Amperometry Electrogravimetry and Coulometry Conductometry 1.4 Classification and Relationships of Electroanalytical Methods 1.5 Summary 1.6 Terminal Questions 1.7 Answers 1.1 INTRODUCTION This is the first unit of this course. This unit deals with the fundamentals of electrochemistry that are necessary for understanding the principles of electroanalytical methods discussed in this Unit 2 to 9. In this unit we have also classified of electroanalytical methods and briefly introduced of some important electroanalytical methods. More details of these elecroanalytical methods will be discussed in the consecutive units. Objectives After studying this unit, you will be able to: • name the different units of electrical quantities, • define the two basic laws of electrochemistry, • describe the single electrode potential and the potential of a galvanic cell, • derive the Nernst expression and give its applications, • calculate the electrode potentials and cell potentials using Nernst equation, • describe the basis for classification of the electroanalytical techniques, and • explain the basis principles and describe the essential conditions of the various electroanalytical techniques. 1.2 BASIC CONCEPTS Before going in detail of different electroanalytical techniques, let’s recapitulate some basic concepts which you have studied in your undergraduate classes. 7 Electroanalytical 1.2.1 Electrical Units Methods -I Ampere (A): Ampere is the unit of current. This is so called in honour of the French Physicist and Chemist A.M. Ampere. It is abbreviated as ‘A’. One ampere is equal to the unvarying direct current which when passed through a silver nitrate solution, under certain specified conditions, will deposit silver at the rate of 0.00111800 g s -1. Ohm (Ω): The unit of electrical resistance is called the Ohm in honour of the German physicist G.S. Ohm. The resistance offered by a uniform column of mercury 106.300 cm long and with a mass of 14.4521 g with a direct current at 0 o C is equal to one Ohm. Volt (V): The unit of electromotive forces (emf) and potential difference are called the volt in honour of Italian physicist C.A. Volta. The unit, volt, is derived from the units of current and resistance via Ohm’s law, thus, One volt is equal to the electromotive force which when applied to a conductor whose resistance is one Ohm will produce a current of one ampere. Coulomb (C): The coulomb is the usual unit to express the quantity of electricity. The name has been given in honour of French physicist C.A. Coulomb. One coulomb corresponds to a constant current of one ampere flowing for one second. Faraday (F): The quantity of electricity associated with one equivalent of chemical change in an electrochemical process is called the Faraday. One Faraday is equal to 96494 coulombs. It is named in honour of English Scientist M. Faraday. Siemens (S): This is the unit of electric conductance. S = A/V 1.2.2 Basic Laws of Electrochemistry Ohm’s Law The mathematical relationship among three fundamental electrical quantities, namely, (1) electromotive force, E (in volts), (2) current strength, I (in amperes), and (3) resistance, R (in Ohms), is expressed by Ohm’s law. The law states that the current flowing in a conductor is equal to the potential difference between any two points divided by the resistance between them. That is, I = E/R or E = IR … (1.1) Faraday’s Law Faraday’s law states that the quantity of current (Coulombs) associated with an electron transfer process is directly proportional to the number of equivalents of the substance involved in chemical change at the electrode. The number of equivalents are the number of moles divided by the number of electrons taking part in electrons transformation reaction. It is expressed as: Q ∝ number of equivalents or Q= F × number of equivalents … (1.2) or QF= × number of moles/n … (1.3) Note that F could be defined as the quantity of where F is the Faraday constant and is equal to 96494 Coulombs, n is the electricity associated with number of electrons taking part in the electrical transformation reaction, such an Avogadro number of electrons. as considering the general reaction for reduction with n electron transfer: O + ne R where O is the species being reduced and R is the reduction product. 8 1.2.3 Electrode Potential Introduction to Electroanalytical The understanding of electrode potentials is essential in electroanalytical chemistry; Methods therefore, this will be discussed first. Whenever two dissimilar conducting phases are brought into contact an electric potential is developed across the interface. In order to understand this effect let us consider first a metal solution interface, which gives the origin of electrode potential. Development of Electrode Potential Consider a metal M that is placed in a solution containing its ions M n+. The metal may be looked upon as being composed of metal ions and electrons. Both the phases, the metal and the solution contain metal ions M n+ but the activity of M n+ in the metal will be different from that in the solution and distribution of metal ions takes place in the two phases in order to get the position of equilibrium. We will consider, for two types of metals, one less active, for example copper, and the other more active say zinc by placing them in contact with their salt solutions. i) In case one, when a less active metal, say a piece of copper is placed in a solution of copper sulphate. Some of the copper ions may deposit on the copper metal, accepting electrons from the metal conduction band and leaving the metal with a small positive charge and the solution with a small negative charge. ii) In the second situation with a more active metal it will be the other way around, a few metal ions from the metal surface may pass into the solution phase, giving the metal a small negative charge and the solution a small positive charge. In both the above two cases, the positive and negative charges will be located at the surface of the metal solution phases, called the interface (see Fig. 1.1). Fig. 1.1: Metal – Metal ion Interface As a result an electrical double layer is established with a corresponding potential difference between metal and solution, and is called the electrode potential . Measurement of Electrode Potentials Unfortunately, there is no way of measuring directly the potential difference between an electrode and a solution. However, it can be measured, with respect to an arbitrarily defined reference electrode. Such a reference electrode was first proposed by Nernst, known as standard hydrogen electrode (SHE) and was given arbitrarily zero potential (at all temperatures). By universal agreement among chemists the standard hydrogen electrode was chosen as the reference electrode. Thus, the electrode, whose potential is to be measured, is coupled with a standard hydrogen electrode and the electromotive force (emf) of the resulting cell is the 9 Electroanalytical electrode potential of the electrode being studied, the experimental conditions being Methods -I such that the liquid-junction potential is negligible. 1.2.4 Liquid-Junction Potentials When two different electrolyte solutions are brought into contact an electrical potential difference arises at the zone of contact. This potential difference is termed liquid- junction potential ( Ej) or diffusion potential. It is caused due to the diffusion of ions from regions of higher to lower concentrations depending on the concentration gradient and the individual mobility of each ion. Various types of liquid-junction potentials are possible, one simple type of Ej is illustrated below. Suppose we could prepare a quiet interface between two solutions containing the same electrolyte but at different concentrations, such as 0.1 M HCl and 0.01 M HCl. On making a contact of these two solutions, immediately, both H+ and Cl− ions diffuse from left to right (see Fig. 1.2) due to concentration gradient. However, hydrogen ions move much more rapidly than do chloride ions and is indicated by the longer arrow for H+ in Fig. 1.2. Thus, H+ outruns Cl − and there is a slight tendency for a charge separation with the right side of the junction acquiring a positive charge and the left side a negative charge. Fig. 1.2: Liquid-Junction Potential The liquid-junction potentials may vary over a considerable range depending on the conditions of the cell. These potentials can be minimized by using a salt bridge containing a concentrated electrolyte solution of cation and anion having comparable mobilities. For example, potassium and chloride ions have comparable mobilities, and salt bridges of saturated aqueous potassium chloride, with agar gel, are widely used to minimize liquid junction potentials. 1.2.5 Electrochemical Cells An electrochemical cell consists of two metallic electrodes immersed in either the same electrolyte solution or in two different solutions that are in electrolytic contact. An electrochemical cell can operate to convert chemical energy into electrical energy or vice-versa depending on whether the cell reaction is spontaneous or force to occur in the non-spontaneous direction. The cell in which the electrode reaction occurs spontaneously when the electrodes are externally connected by a conductor and it serves as a means of converting chemical energy into electrical energy is called the Galvanic cell or voltaic cell . Alternatively, the cell in which the cell reaction is force to occur in the non-spontaneous direction by passing the current through the cell from an external source to affect a chemical transformation is called the e lectrolytic cell .
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