Introduction - Oxidation and Reduction
Chemical reactions are very common in our daily life. Many industrial processes (production of various materials), generation of power, synthesis of newer exotic materials etc. all use characteristic chemical reactions called as oxidation and reduction reactions.
Oxidizing Agents and Reducing Agents
The substance, which can bring about oxidation of other substances, is called an oxidizing agent. Thus, an oxidizing agent is the substance that provides oxygen or removes hydrogen from another substance. Important oxidizing agents are oxygen, ozone, hydrogen peroxide, chlorine, bromine, nitric acid, concentrated sulphuric acid, potassium permanganate, potassium dichromate etc.
Electronic Concept of Oxidation and Reduction
From the above reactions, it can be seen that the oxidation-reduction reactions involve transfer of electrons from one species to another. Thus, the oxidation-reduction reactions may also be defined as follows.
Redox Reactions in Aqueous Solution
Direct Redox Reaction
When the reacting species, capable of losing electrons (reducing agent) and that capable of gaining electrons (oxidizing agent) are present in the same solution a direct redox reaction takes place. In direct redox reactions, the electrons move randomly in all the directions through very short distances. So, in direct redox reactions the net movement of electrons in any direction is zero. The change in chemical energy in such reactions appears as heat.
Electrochemical Reactions
Here, each of the reactant is taken in a separate container in contact with a rod/sheet of a metallic-conductor (electronic conductor) called an electrode. Electrical contact between the two reactants is established by placing a conducting salt bridge in-between.
The Concept of Half-cell
A galvanic cell consists of two electrodes dipping into the same or two different electrolytes. No reaction takes place until a conducting wire joins the two electrodes. Once, a chemical reaction takes place at each electrode, electron transfer takes place. The electron-transfer reactions, which take place at the surface of the electrode are called electrode reactions.
Electrode Potential
It is important to understand the development of charges at the electrodes. When a strip of a metal M is placed in a solution of its ions Mn+, a metal - metal ion electrode is obtained.
Single Electrode Potential
It is not possible to measure the absolute potential of a single electrode because neither the oxidation nor the reduction reaction can occur by itself. Also, we need two electrodes to measure the potential difference between two points. This difficulty can be solved selecting one of the electrodes as a reference electrode and arbitrarily fixing the potential of that electrode as zero. Thus, it is possible to obtain the potential of an electrode, if the given electrode is coupled with another electrode having a potential of zero volts. The electrode whose potential can be taken as zero volt is called primary reference electrode.
E.M.F. or Cell Potential of a Cell
The electrochemical cell consists of two half cells where one of the half-cell has a higher value of reduction potential as compared to the other. As a result of this potential difference, there is a flow of electrons from the electrode with a lower reduction potential (or higher oxidation potential) to the electrode with higher reduction potential (or lower oxidation potential). The difference between the electrode potentials of the two electrodes in the electrochemical cell is known as electromotive force or cell potential of a cell. The electromotive force is commonly abbreviated as EMF (emf) and is expressed in volts. The emf of a cell may be expressed in terms of the difference in the reduction electrode potential.
Measurement of Electrode Potentials
The absolute value of a single electrode cannot be measured experimentally because a half-cell reaction cannot take place independently. One can measure only the difference between the electrode potentials of any two half-cell reactions. If the cell potential and the electrode potential for one of the half-cell reactions are known, the electrode potential of the other electrode can be calculated. Therefore, a standard hydrogen electrode is chosen and all other cell reactions are compared with this standard and a set of Eo values are obtained.
Dependence of Electrode Potential on Concentration and Temperature (Nernst equation)
When one measures the standard electrode potentials, the temperature of the cell is 298 K and the concentration of the electrolyte solutions is fixed as 1 M. However, in actual practice, the electrochemical cells are not always at 298 K nor do always have this concentration of the electrolyte solutions. Hence, the electrode potentials depend on the concentration and temperature of the electrolyte solutions. Nernst gave a relationship between electrode potentials and the concentration and temperature of electrolyte solutions.
Electrochemical Series
Different electrodes have different standard electrode potentials. The standard electrode potentials (E°) for some electrodes are negative, while for some others E° values are positive. The E° values of many electrodes have been measured and their standard reduction potentials are arranged in a sequential order. This arrangement of elements in order of increasing reduction potential values is called electrochemical series or activity series.
Applications of Electrochemical Series
Oxidizing and Reducing Strengths
The electrochemical series helps to pick out substances that are good oxidizing agents and those which are good reducing agents. For example, a very high negative reduction potential of lithium electrode indicates that it is very difficult to reduce Li+ ions to Li atoms. Therefore, Li+ cannot accept electrons easily and so loses electrons to behave as a reducing agent. As the reduction potential increases (negative value decreases), the tendency of the electrode to behave as reducing agent decreases. Thus, all the substances appearing on the top of the series behave as good reducing agents. For example Li and K are good reducing agents while F- and Au are the poorest reducing agents.
Oxidation Number
In ionic substances, the redox reactions can be explained on the basis of electron transfer. However, the redox reactions of covalent compounds cannot be explained in terms of electron transfer. The term oxidation number explains the phenomenon of oxidation-reduction in covalent and ionic substances.
Redox Reactions in Terms of Oxidation Number
Oxidation is a chemical change in which there is an increase in oxidation number while reduction is a chemical change in which there is a decrease in the oxidation number. For example, in the reaction between manganese dioxide and hydrochloric acid, the oxidation number of manganese decreases from +4 (MnO2) to +2 (in MnCl2) indicating that manganese dioxide undergoes reduction. On the other hand, the oxidation number of chlorine increases from -1 (in HCl) to 0 (in Cl2) indicating that hydrochloric acid undergoes oxidation.
Oxidation Number and Nomenclature
When an element forms two monoatomic cations (representing different oxidation states), the two ions are distinguished by using the ending 'ous' and 'ic' at the end. The suffix 'ous' is used for the cation with lower oxidation state and the suffix 'ic' is used for the cation with higher oxidation state.
Balancing Oxidation-Reduction Reactions
There are two very important methods for balancing oxidation-reduction reactions. These are:
* Oxidation number method
* Ion-electron method
Stoichiometry of Redox Reactions
The quantitative relationships between different species in the redox reactions gives its stoichiometry. These relationships are derived in the same manner as derived for general chemical equations after writing the balanced redox equation.
Applications of Redox Reactions
Metallurgical processes
Metal oxides are reduced to metals using suitable reducing agents. For example Fe2O3 is reduced to iron in a blast furnace using coke. Al2O3 is reduced to aluminium by cathodic reduction in an electrolytic cell.
Electrolysis (Additional)
Chemical energy can be converted into electrical energy with the help of a galvanic cell. Conversely, the phenomenon of chemical changes taking place by the passage of electricity through an electrolyte is called electrolysis. Humphrey Davy (1807) showed this in his experiments, when he isolated potassium by passing electricity through molten potassium hydroxide.
Faraday's Laws of Electrolysis
In 1833 M. Faraday studied the quantitative aspects of electrolysis, and postulated two laws named after him.
Equivalent Mass and Number of Electrons
Equivalent mass is equal to the molecular or atomic mass divided by the number of electrons involved in the reaction per molecule, atom or ion. For example in the reaction,
two electrons are needed to produce one molecule of hydrogen gas. So, 2 Faraday of electricity is needed to produce one mole of hydrogen gas.
Current Efficiency
In many industrial processes, the process of electrolysis is less efficient than expected from the Faraday's laws. The amount of material obtained during electrolysis is generally less than that expected due to:
* Loss of energy during its flow through the system.
* Some other side-reactions taking place during electrolysis.
Electrolysis and Electrode Processes
The chemical reactions, which take place at the surface of electrodes are called electrode reaction (or electrode processes). According to the theory of ionization, electrolytes are present as ions in solution. These ions are directed towards the respective electrodes by the electricity supplied. The electrolytes can be electrolyzed only in the dissolved or molten state. Various types of electrode reactions are described below.
Applications of Electrolysis
The principle of electrolysis can be applied to many important processes of industrial or commercial importance. Some of these are described below.
