Unit-1-ELECTROCHEMISTRY
Introduction – cells – types - representation of galvanic cell - electrode potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from single electrode potential - reference electrode: construction, working and applications of standard hydrogen electrode, standard calomel electrode - glass electrode – EMF series and its applications - potentiometric titrations (redox) - conductometric titrations - mixture of weak and strong acid vs strong base.
22CYT12-Unit_I_Electrochemistry - EMF Series & its Applications.pptKrishnaveniKrishnara1
Electrochemistry:Introduction – cells – types - representation of galvanic cell - electrode potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from single electrode potential - reference electrode: construction, working and applications (Determination of potential of the unknown electrode and pH of the unknown electrode) of standard hydrogen electrode, standard calomel electrode - glass electrode – EMF series and its applications - potentiometric titrations (redox) - conductometric titrations - mixture of weak and strong acid vs strong base.
1. The document provides information about a course on Engineering Chemistry taught by Dr. Suresh Siliveri. It includes the course outline, learning objectives, and topics that will be covered.
2. The course aims to teach students about electrochemistry, materials chemistry, energy sources, water technology, and engineering materials. After completing the course, students will be able to apply concepts in various applications and summarize manufacturing processes.
3. The topics covered include electrochemistry, batteries, corrosion, polymers, fuels, treated water, and advanced engineering materials. Learning objectives for electrochemistry include introduction, definitions, galvanic cells, the Nernst equation, and fuel cells.
This document provides an overview of electrochemistry. It defines electrochemistry as the branch of chemistry dealing with the transformation of electrical and chemical energy. It describes the key topics that will be covered, including conductors, electrochemical cells, electrode potentials, and how to predict spontaneity of reactions. It also summarizes the basic components and functions of electrolytic and galvanic (voltaic) cells, including how they convert between electrical and chemical energy.
The document discusses corrosion, including its types, why we prevent it, and various electrochemical principles related to corrosion. It defines corrosion as the destruction of metals by chemical and electrochemical attack from the environment. There are two main types of corrosion: chemical (dry) corrosion which occurs without moisture, and electrochemical (wet) corrosion, which is the most common and occurs in the presence of moisture via an electrochemical cell. The document outlines several electrochemical concepts like the Pilling Bedworth rule, electrochemical series, Nernst equation, standard electrodes, and Pourbaix and Ellingham diagrams which can be used to understand and predict corrosion reactions and products.
The document discusses corrosion, including its types, why we prevent it, and several related concepts from electrochemistry. It defines corrosion as the destruction of metals by chemical and electrochemical attack from the environment. There are two main types: chemical corrosion from oxygen and electrochemical corrosion, which occurs in the presence of moisture and produces an electrochemical cell. The document outlines several electrochemical concepts like the Pilling Bedworth rule, electrochemical series, Nernst equation, standard electrodes, and Pourbaix and Ellingham diagrams, which can be used to understand and predict corrosion reactions.
The document discusses corrosion, including its types, why we prevent it, and several related concepts from electrochemistry. It defines corrosion as the destruction of metals by chemical and electrochemical attack from the environment. There are two main types: chemical corrosion from oxygen and electrochemical corrosion, which occurs in the presence of moisture and produces an electrochemical cell. The document outlines several concepts to understand and control corrosion, including the Pilling-Bedworth rule, electrochemical series, Nernst equation, standard electrodes, and Pourbaix and Ellingham diagrams.
This document provides an overview of electrochemistry. It begins by defining electrochemistry as the study of chemical reactions at the interface of an electrode and electrolyte involving the interaction of electrical and chemical changes. The document then discusses the history and founders of electrochemistry, including Faraday's two laws of electrolysis. It explains key concepts such as oxidation-reduction reactions, balancing redox equations, and the Nernst equation. The document also covers applications including batteries, corrosion, electrolysis, and branches of electrochemistry like bioelectrochemistry and nanoelectrochemistry.
22CYT12-Unit_I_Electrochemistry - EMF Series & its Applications.pptKrishnaveniKrishnara1
Electrochemistry:Introduction – cells – types - representation of galvanic cell - electrode potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from single electrode potential - reference electrode: construction, working and applications (Determination of potential of the unknown electrode and pH of the unknown electrode) of standard hydrogen electrode, standard calomel electrode - glass electrode – EMF series and its applications - potentiometric titrations (redox) - conductometric titrations - mixture of weak and strong acid vs strong base.
1. The document provides information about a course on Engineering Chemistry taught by Dr. Suresh Siliveri. It includes the course outline, learning objectives, and topics that will be covered.
2. The course aims to teach students about electrochemistry, materials chemistry, energy sources, water technology, and engineering materials. After completing the course, students will be able to apply concepts in various applications and summarize manufacturing processes.
3. The topics covered include electrochemistry, batteries, corrosion, polymers, fuels, treated water, and advanced engineering materials. Learning objectives for electrochemistry include introduction, definitions, galvanic cells, the Nernst equation, and fuel cells.
This document provides an overview of electrochemistry. It defines electrochemistry as the branch of chemistry dealing with the transformation of electrical and chemical energy. It describes the key topics that will be covered, including conductors, electrochemical cells, electrode potentials, and how to predict spontaneity of reactions. It also summarizes the basic components and functions of electrolytic and galvanic (voltaic) cells, including how they convert between electrical and chemical energy.
The document discusses corrosion, including its types, why we prevent it, and various electrochemical principles related to corrosion. It defines corrosion as the destruction of metals by chemical and electrochemical attack from the environment. There are two main types of corrosion: chemical (dry) corrosion which occurs without moisture, and electrochemical (wet) corrosion, which is the most common and occurs in the presence of moisture via an electrochemical cell. The document outlines several electrochemical concepts like the Pilling Bedworth rule, electrochemical series, Nernst equation, standard electrodes, and Pourbaix and Ellingham diagrams which can be used to understand and predict corrosion reactions and products.
The document discusses corrosion, including its types, why we prevent it, and several related concepts from electrochemistry. It defines corrosion as the destruction of metals by chemical and electrochemical attack from the environment. There are two main types: chemical corrosion from oxygen and electrochemical corrosion, which occurs in the presence of moisture and produces an electrochemical cell. The document outlines several electrochemical concepts like the Pilling Bedworth rule, electrochemical series, Nernst equation, standard electrodes, and Pourbaix and Ellingham diagrams, which can be used to understand and predict corrosion reactions.
The document discusses corrosion, including its types, why we prevent it, and several related concepts from electrochemistry. It defines corrosion as the destruction of metals by chemical and electrochemical attack from the environment. There are two main types: chemical corrosion from oxygen and electrochemical corrosion, which occurs in the presence of moisture and produces an electrochemical cell. The document outlines several concepts to understand and control corrosion, including the Pilling-Bedworth rule, electrochemical series, Nernst equation, standard electrodes, and Pourbaix and Ellingham diagrams.
This document provides an overview of electrochemistry. It begins by defining electrochemistry as the study of chemical reactions at the interface of an electrode and electrolyte involving the interaction of electrical and chemical changes. The document then discusses the history and founders of electrochemistry, including Faraday's two laws of electrolysis. It explains key concepts such as oxidation-reduction reactions, balancing redox equations, and the Nernst equation. The document also covers applications including batteries, corrosion, electrolysis, and branches of electrochemistry like bioelectrochemistry and nanoelectrochemistry.
Electroanalytical methods provide several advantages for quantitative analytical chemistry. They involve measuring the electrical properties of analyte solutions in electrochemical cells. Some key points:
- Electroanalytical methods allow easy automation through electrical signal measurements. They can also determine low analyte concentrations without difficulty.
- Electrochemical processes involve the transfer of electrons between substances during redox reactions. This occurs at the interface between electrodes and solutions in electrochemical cells.
- Advantages include low cost compared to spectroscopy and the ability to easily automate measurements and detect low analyte concentrations through electrical signals.
CONTENTS
Electrochemistry: definition & importance
Conductors: metallic & electrolytic conduction,
Electrolytes, Electrochemical cell & electrolytic cell
A simple electrochemical cell: Galvanic cell or (Daniell Cell)
Cell reaction, cell representation, Salt bridge & its use,
Electrode potential, standard electrode potential, SHE,
Standard cell potential or standard electromotive force of a cell
Electrochemical series (Standard reduction potential values)
Nernst Equation, Relationship with Standard cell potential with Gibbs energy & also equilibrium constant
Resistance (R) & conductance (G) of a solution of an electrolyte
Conductivity (k) of solution, Cell constant (G*) & their units,
Molar conductivity (Λm) & its variation with concentration & temperature,
Debye Huckel Onsager equation & Limiting molar conductivity,
Kohlrausch’s law & its application & numerical problems.
Electrolytic cells & electrolysis.
Some examples of electrolysis of electrolytes in molten / aq. state.
Faraday’s laws of electrolysis: First & second law- numerical problems. Corrosion, Electrochemical theory of rusting.
Prevention of rusting.
Class XII Electrochemistry - Nernst equation.Arunesh Gupta
This document provides an overview of electrochemistry and some key concepts. It begins by defining electrochemistry as the study of how spontaneous chemical reactions can produce electricity and how electrical energy can drive non-spontaneous reactions. It then discusses several applications of electrochemistry including metal production, electroplating, and batteries. The document goes on to define conductors and the differences between metallic and electrolytic conduction. It also introduces concepts like galvanic cells, salt bridges, standard electrode potentials, and the electrochemical series. In summary, the document provides a broad introduction to fundamental electrochemistry topics and concepts.
1. The document discusses electrode potential and how it is measured. Electrode potential is the tendency of an electrode to gain or lose electrons when in contact with its own ions in solution.
2. Oxidation occurs at the anode where electrons are lost, and reduction occurs at the cathode where electrons are gained. The standard hydrogen electrode is used as a reference to measure other electrode potentials.
3. Electrode potentials can be oxidation potentials if the electrode loses electrons or reduction potentials if it gains electrons. Nernst theory explains how solution pressure and osmotic pressure determine electrode behavior.
Electrochemistry involves the study of electricity produced from spontaneous chemical reactions in galvanic cells and the use of electricity to drive non-spontaneous reactions in electrolytic cells. Galvanic cells produce electricity through spontaneous redox reactions, with oxidation occurring at the anode and reduction at the cathode. Electrolytic cells use electricity to carry out non-spontaneous reactions. The potential difference between electrodes in a galvanic cell is called the cell potential, which can be calculated using standard electrode potentials and concentrations based on the Nernst equation.
(1) Electrochemistry involves the transfer of electrons during chemical reactions and electrical changes brought about by chemical changes.
(2) Cell potential, measured in Volts, is the tendency of a species to lose or gain electrons compared to the Standard Hydrogen Electrode potential of 0.00V.
(3) The Standard Hydrogen Electrode consists of hydrogen gas bubbling over a platinum electrode in a solution of 1M hydrogen ions, and its reduction potential is defined as 0.00V.
1. The document discusses different types of electrochemical cells including galvanic/voltaic cells and electrolytic cells.
2. Galvanic cells are further classified as primary cells, which cannot be recharged, and secondary cells, which are rechargeable batteries.
3. The Nernst equation is derived, which relates the electrode potential to the standard electrode potential and the concentrations of the metal ions involved in the electrochemical cell reaction.
This document provides an overview of electrochemistry concepts including:
- Types of cells like reversible cells where equilibrium exists and irreversible cells where it does not.
- Different types of electrodes such as metal-metal ion electrodes containing a metal and its ions, gas electrodes containing a gas, and oxidation-reduction electrodes containing ions in different oxidation states.
- Key electrode reactions involving oxidation where electrons are removed from metals and reduction where electrons are gained by metal ions.
- Terminology used in electrochemistry like electrode, anode, cathode, and cell reactions.
This document discusses electrochemistry and electrochemical cells. It defines electrochemistry as the study of chemical reactions that produce electricity or use electricity to cause reactions. There are two types of electrochemical cells: galvanic cells that convert chemical energy to electrical energy, and electrolytic cells that use electrical energy to drive non-spontaneous reactions. Examples of galvanic cells include Daniell cells and concentration cells. The document explains concepts like standard electrode potentials, the electrochemical series, and how to represent cell diagrams according to IUPAC recommendations. It also discusses the functions of salt bridges and how junction potentials can affect cell potentials.
This document provides an overview of redox (oxidation-reduction) chemistry. It defines oxidation as the loss of electrons and reduction as the gain of electrons. Oxidation and reduction always occur simultaneously in redox reactions. The document discusses identifying oxidizing and reducing agents, balancing redox reactions using the half-reaction method, standard reduction potentials, galvanic (voltaic) cells that produce electricity from spontaneous redox reactions, and electrolytic cells that use electricity to drive nonspontaneous reactions.
Electrolysis is the process of using a direct electric current to drive nonspontaneous chemical reactions. It involves the decomposition of an electrolyte into its constituent ions by the removal or addition of electrons to the ions. During electrolysis, ions migrate to the electrodes where they undergo oxidation or reduction reactions. In the electrolysis of molten lead bromide, lead ions are reduced to metallic lead at the cathode, while bromide ions are oxidized to bromine gas at the anode. When an aqueous solution of copper sulfate is electrolyzed using copper electrodes, copper ions are reduced at the cathode to form metallic copper while oxygen gas forms at the anode. Electrolysis requires an electrolyte, electrodes, and a direct current power
This document provides a summary of key concepts in oxidation-reduction (redox) reactions:
1) Redox reactions involve the transfer of electrons between chemical species, either through the complete transfer of electrons to form ionic bonds, or partial transfer to form covalent bonds. Oxidation is the loss of electrons and reduction is the gain of electrons.
2) Redox pairs are couples of oxidized and reduced forms of elements that differ in their oxidation state. Common redox pairs include Fe3+/Fe2+, O2/H2O, and MnO4-/MnO2.
3) The standard electrode potential (E°) indicates the tendency of half-reactions to occur.
1. Redox reactions involve both oxidation and reduction processes occurring simultaneously, such as a displacement reaction where one element replaces another in a compound.
2. Metals can be ranked based on their ability to displace hydrogen from sources like water or acids in an activity series. More reactive metals displace hydrogen from less extreme sources.
3. Electrochemical cells consist of an anode where oxidation occurs, a cathode where reduction occurs, and an electrolyte for ion transport between electrodes. Different types of cells generate voltage based on factors like temperature, concentration, oxygen levels, or pH.
Electrochemistry is the study of the interchange between chemical change and electrical work. Electrochemical cells utilize redox reactions to produce or use electrical energy. Redox reactions involve the transfer of electrons between oxidizing and reducing agents. Voltaic or galvanic cells generate electrical energy from a spontaneous redox reaction, while electrolytic cells use an applied electrical current to drive a nonspontaneous reaction. Common components of cells include electrodes, electrolytes, and salt bridges. Oxidation occurs at the anode and reduction at the cathode. Batteries contain multiple connected galvanic cells and electrolysis uses a current to force a nonspontaneous redox reaction.
This document provides an overview of the key concepts in electrochemistry including oxidation-reduction reactions, galvanic cells, standard reduction potentials, the Nernst equation, electrolysis, batteries, corrosion, and commercial electrolytic processes. It defines important terms, describes experimental set ups and calculations for electrochemical cells, and summarizes fundamental electrochemical principles and laws such as Faraday's laws of electrolysis.
The document discusses conductivity (or specific conductance) of metal ions in solution, which is a measure of its ability to conduct electricity. It explains that conductivity is higher for strong electrolytes that nearly completely dissociate into ions in solution, while weak electrolytes only partially dissociate. Several factors influence conductivity, including the nature of the solute and solvent, concentration, and temperature. Conductivity measurements are used in various industrial applications like water treatment and leak detection.
This document discusses electrochemistry and provides details about electrochemical cells. It contains the following key points:
1. Electrochemistry is the study of production of electricity from chemical reactions and use of electrical energy to drive non-spontaneous reactions.
2. An electrochemical cell converts chemical energy to electrical energy (galvanic/voltaic cell) or electrical energy to chemical energy (electrolytic cell).
3. A Daniell cell is a voltaic cell that generates a voltage of 1.1V from the redox reaction of zinc and copper. Measurement of electrode potentials and the Nernst equation are also discussed.
There are several types of electrodes classified by their composition and function. Reference electrodes like the standard hydrogen electrode (SHE) and saturated calomel electrode (SCE) maintain a known and constant potential used for comparison. The SHE represents the standard reduction potential but is difficult to maintain at standard conditions. The SCE uses a mercury/mercury chloride mixture and is easier to construct and maintain compared to the SHE. Indicator electrodes like the glass electrode are used in titration analysis, with the glass electrode potential indicating pH. Electrodes can also be classified as anodes, which experience oxidation, or cathodes, which undergo reduction.
This document provides an overview of electrochemistry. It discusses key topics like what electrochemistry is, the history and founders of electrochemistry, oxidation-reduction reactions, balancing redox equations, standard electrode potential, the Nernst equation, batteries, corrosion, electrolysis, Faraday's laws of electrolysis, and more. The document serves as a high-level introduction to many fundamental concepts in electrochemistry.
Introduction- e - waste – definition - sources of e-waste– hazardous substances in e-waste - effects of e-waste on environment and human health- need for e-waste management– e-waste handling rules - waste minimization techniques for managing e-waste – recycling of e-waste - disposal treatment methods of e- waste – mechanism of extraction of precious metal from leaching solution-global Scenario of E-waste – E-waste in India- case studies.
Batteries -Introduction – Types of Batteries – discharging and charging of battery - characteristics of battery –battery rating- various tests on battery- – Primary battery: silver button cell- Secondary battery :Ni-Cd battery-modern battery: lithium ion battery-maintenance of batteries-choices of batteries for electric vehicle applications.
Fuel Cells: Introduction- importance and classification of fuel cells - description, principle, components, applications of fuel cells: H2-O2 fuel cell, alkaline fuel cell, molten carbonate fuel cell and direct methanol fuel cells.
More Related Content
Similar to 22CYT12 & Chemistry for Computer Systems-Unit_I_Electrochemistry.ppt
Electroanalytical methods provide several advantages for quantitative analytical chemistry. They involve measuring the electrical properties of analyte solutions in electrochemical cells. Some key points:
- Electroanalytical methods allow easy automation through electrical signal measurements. They can also determine low analyte concentrations without difficulty.
- Electrochemical processes involve the transfer of electrons between substances during redox reactions. This occurs at the interface between electrodes and solutions in electrochemical cells.
- Advantages include low cost compared to spectroscopy and the ability to easily automate measurements and detect low analyte concentrations through electrical signals.
CONTENTS
Electrochemistry: definition & importance
Conductors: metallic & electrolytic conduction,
Electrolytes, Electrochemical cell & electrolytic cell
A simple electrochemical cell: Galvanic cell or (Daniell Cell)
Cell reaction, cell representation, Salt bridge & its use,
Electrode potential, standard electrode potential, SHE,
Standard cell potential or standard electromotive force of a cell
Electrochemical series (Standard reduction potential values)
Nernst Equation, Relationship with Standard cell potential with Gibbs energy & also equilibrium constant
Resistance (R) & conductance (G) of a solution of an electrolyte
Conductivity (k) of solution, Cell constant (G*) & their units,
Molar conductivity (Λm) & its variation with concentration & temperature,
Debye Huckel Onsager equation & Limiting molar conductivity,
Kohlrausch’s law & its application & numerical problems.
Electrolytic cells & electrolysis.
Some examples of electrolysis of electrolytes in molten / aq. state.
Faraday’s laws of electrolysis: First & second law- numerical problems. Corrosion, Electrochemical theory of rusting.
Prevention of rusting.
Class XII Electrochemistry - Nernst equation.Arunesh Gupta
This document provides an overview of electrochemistry and some key concepts. It begins by defining electrochemistry as the study of how spontaneous chemical reactions can produce electricity and how electrical energy can drive non-spontaneous reactions. It then discusses several applications of electrochemistry including metal production, electroplating, and batteries. The document goes on to define conductors and the differences between metallic and electrolytic conduction. It also introduces concepts like galvanic cells, salt bridges, standard electrode potentials, and the electrochemical series. In summary, the document provides a broad introduction to fundamental electrochemistry topics and concepts.
1. The document discusses electrode potential and how it is measured. Electrode potential is the tendency of an electrode to gain or lose electrons when in contact with its own ions in solution.
2. Oxidation occurs at the anode where electrons are lost, and reduction occurs at the cathode where electrons are gained. The standard hydrogen electrode is used as a reference to measure other electrode potentials.
3. Electrode potentials can be oxidation potentials if the electrode loses electrons or reduction potentials if it gains electrons. Nernst theory explains how solution pressure and osmotic pressure determine electrode behavior.
Electrochemistry involves the study of electricity produced from spontaneous chemical reactions in galvanic cells and the use of electricity to drive non-spontaneous reactions in electrolytic cells. Galvanic cells produce electricity through spontaneous redox reactions, with oxidation occurring at the anode and reduction at the cathode. Electrolytic cells use electricity to carry out non-spontaneous reactions. The potential difference between electrodes in a galvanic cell is called the cell potential, which can be calculated using standard electrode potentials and concentrations based on the Nernst equation.
(1) Electrochemistry involves the transfer of electrons during chemical reactions and electrical changes brought about by chemical changes.
(2) Cell potential, measured in Volts, is the tendency of a species to lose or gain electrons compared to the Standard Hydrogen Electrode potential of 0.00V.
(3) The Standard Hydrogen Electrode consists of hydrogen gas bubbling over a platinum electrode in a solution of 1M hydrogen ions, and its reduction potential is defined as 0.00V.
1. The document discusses different types of electrochemical cells including galvanic/voltaic cells and electrolytic cells.
2. Galvanic cells are further classified as primary cells, which cannot be recharged, and secondary cells, which are rechargeable batteries.
3. The Nernst equation is derived, which relates the electrode potential to the standard electrode potential and the concentrations of the metal ions involved in the electrochemical cell reaction.
This document provides an overview of electrochemistry concepts including:
- Types of cells like reversible cells where equilibrium exists and irreversible cells where it does not.
- Different types of electrodes such as metal-metal ion electrodes containing a metal and its ions, gas electrodes containing a gas, and oxidation-reduction electrodes containing ions in different oxidation states.
- Key electrode reactions involving oxidation where electrons are removed from metals and reduction where electrons are gained by metal ions.
- Terminology used in electrochemistry like electrode, anode, cathode, and cell reactions.
This document discusses electrochemistry and electrochemical cells. It defines electrochemistry as the study of chemical reactions that produce electricity or use electricity to cause reactions. There are two types of electrochemical cells: galvanic cells that convert chemical energy to electrical energy, and electrolytic cells that use electrical energy to drive non-spontaneous reactions. Examples of galvanic cells include Daniell cells and concentration cells. The document explains concepts like standard electrode potentials, the electrochemical series, and how to represent cell diagrams according to IUPAC recommendations. It also discusses the functions of salt bridges and how junction potentials can affect cell potentials.
This document provides an overview of redox (oxidation-reduction) chemistry. It defines oxidation as the loss of electrons and reduction as the gain of electrons. Oxidation and reduction always occur simultaneously in redox reactions. The document discusses identifying oxidizing and reducing agents, balancing redox reactions using the half-reaction method, standard reduction potentials, galvanic (voltaic) cells that produce electricity from spontaneous redox reactions, and electrolytic cells that use electricity to drive nonspontaneous reactions.
Electrolysis is the process of using a direct electric current to drive nonspontaneous chemical reactions. It involves the decomposition of an electrolyte into its constituent ions by the removal or addition of electrons to the ions. During electrolysis, ions migrate to the electrodes where they undergo oxidation or reduction reactions. In the electrolysis of molten lead bromide, lead ions are reduced to metallic lead at the cathode, while bromide ions are oxidized to bromine gas at the anode. When an aqueous solution of copper sulfate is electrolyzed using copper electrodes, copper ions are reduced at the cathode to form metallic copper while oxygen gas forms at the anode. Electrolysis requires an electrolyte, electrodes, and a direct current power
This document provides a summary of key concepts in oxidation-reduction (redox) reactions:
1) Redox reactions involve the transfer of electrons between chemical species, either through the complete transfer of electrons to form ionic bonds, or partial transfer to form covalent bonds. Oxidation is the loss of electrons and reduction is the gain of electrons.
2) Redox pairs are couples of oxidized and reduced forms of elements that differ in their oxidation state. Common redox pairs include Fe3+/Fe2+, O2/H2O, and MnO4-/MnO2.
3) The standard electrode potential (E°) indicates the tendency of half-reactions to occur.
1. Redox reactions involve both oxidation and reduction processes occurring simultaneously, such as a displacement reaction where one element replaces another in a compound.
2. Metals can be ranked based on their ability to displace hydrogen from sources like water or acids in an activity series. More reactive metals displace hydrogen from less extreme sources.
3. Electrochemical cells consist of an anode where oxidation occurs, a cathode where reduction occurs, and an electrolyte for ion transport between electrodes. Different types of cells generate voltage based on factors like temperature, concentration, oxygen levels, or pH.
Electrochemistry is the study of the interchange between chemical change and electrical work. Electrochemical cells utilize redox reactions to produce or use electrical energy. Redox reactions involve the transfer of electrons between oxidizing and reducing agents. Voltaic or galvanic cells generate electrical energy from a spontaneous redox reaction, while electrolytic cells use an applied electrical current to drive a nonspontaneous reaction. Common components of cells include electrodes, electrolytes, and salt bridges. Oxidation occurs at the anode and reduction at the cathode. Batteries contain multiple connected galvanic cells and electrolysis uses a current to force a nonspontaneous redox reaction.
This document provides an overview of the key concepts in electrochemistry including oxidation-reduction reactions, galvanic cells, standard reduction potentials, the Nernst equation, electrolysis, batteries, corrosion, and commercial electrolytic processes. It defines important terms, describes experimental set ups and calculations for electrochemical cells, and summarizes fundamental electrochemical principles and laws such as Faraday's laws of electrolysis.
The document discusses conductivity (or specific conductance) of metal ions in solution, which is a measure of its ability to conduct electricity. It explains that conductivity is higher for strong electrolytes that nearly completely dissociate into ions in solution, while weak electrolytes only partially dissociate. Several factors influence conductivity, including the nature of the solute and solvent, concentration, and temperature. Conductivity measurements are used in various industrial applications like water treatment and leak detection.
This document discusses electrochemistry and provides details about electrochemical cells. It contains the following key points:
1. Electrochemistry is the study of production of electricity from chemical reactions and use of electrical energy to drive non-spontaneous reactions.
2. An electrochemical cell converts chemical energy to electrical energy (galvanic/voltaic cell) or electrical energy to chemical energy (electrolytic cell).
3. A Daniell cell is a voltaic cell that generates a voltage of 1.1V from the redox reaction of zinc and copper. Measurement of electrode potentials and the Nernst equation are also discussed.
There are several types of electrodes classified by their composition and function. Reference electrodes like the standard hydrogen electrode (SHE) and saturated calomel electrode (SCE) maintain a known and constant potential used for comparison. The SHE represents the standard reduction potential but is difficult to maintain at standard conditions. The SCE uses a mercury/mercury chloride mixture and is easier to construct and maintain compared to the SHE. Indicator electrodes like the glass electrode are used in titration analysis, with the glass electrode potential indicating pH. Electrodes can also be classified as anodes, which experience oxidation, or cathodes, which undergo reduction.
This document provides an overview of electrochemistry. It discusses key topics like what electrochemistry is, the history and founders of electrochemistry, oxidation-reduction reactions, balancing redox equations, standard electrode potential, the Nernst equation, batteries, corrosion, electrolysis, Faraday's laws of electrolysis, and more. The document serves as a high-level introduction to many fundamental concepts in electrochemistry.
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Introduction – cells – types - representation of galvanic cell - electrode potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from single electrode potential - reference electrode: construction, working and applications of standard hydrogen electrode, standard calomel electrode - glass electrode – EMF series and its applications - potentiometric titrations (redox) - conductometric titrations - mixture of weak and strong acid vs strong base.
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22CYT12 & Chemistry for Computer Systems-Unit_I_Electrochemistry.ppt
1. DEPARTMENT OF CHEMISTRY
WELCOMES YOUALL
22CYT12 & Chemistry for Computer Systems
2022R
Unit-I-Electrochemistry
Prepared By
Krishnaveni K
Assistant Professor
Department of Chemistry
Kongu Engineering
College, Perundurai,
Erode
Course Outcome: Apply the principle of
electrochemistry for various applications
2. APPLIED CHEMISTRY
• The development of science and technology has
been giving us a lot of benefits. The advanced
technology has often required the basic research.
• Applied Chemistry is the scientific field for
understanding basic chemical properties of
materials and for producing new materials with
well-controlled functions.
• Applied chemistry is increasingly important in
solving environmental problems and contributing
to the development of new materials, both of which
are key issues in the 21st century.
5. Introduction – cells – types - representation of galvanic cell - electrode
potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from
single electrode potential - reference electrode: construction, working and
applications of standard hydrogen electrode, standard calomel electrode - glass
electrode – EMF series and its applications - potentiometric titrations (redox) -
conductometric titrations - mixture of weak and strong acid vs strong base.
UNIT-II
ELECTROCHEMISTRY
6. History of Electrochemistry
16 th Century - William Gilbert –Father of Magnetism
18 th Century – William Nicholson & Wilhelm Ritter – Decomposition of water – Electrolysis
Svante Arrhenius - Dissociation of electrolytes
Walther Hermann Nernst – Theory of Electromotive Force
Conductance? Ability to conduct current , mho
7. ELECTROCHEMISTRY
INTRODUCTION
It is a branch of chemistry
The study of process involved in the interconversion of
chemical and electrical energy.
KEY TERMS IN ELECTROCHEMISTRY
Conductor: Material which conduct electric current
Non conductor: Material which do not conduct electric current
Current: The flow of electrons through a wire or any conductor
Oxidation: Loss of electrons
Reduction: Gain of electrons
Redox reaction: oxidation and reduction reactions occur simultaneously
Reducing agent: A reactant in which donates an electron to the reduced species. (The reducing agent
is oxidized)
22-Feb-24
8. Oxidizing agent: A reactant in which accepts an electron from the oxidized species. (The oxidizing agent
is reduced)
Anode: The electrode at which oxidation occurs
Cathode: The electrode at which reduction occurs
Electrolyte: A water soluble substance and conduct an electric current
Half cell: A single electrode immersed in an electrolytic solution and developing a definite potential
difference.
Cell: Two half cells are connected through one wire
Oxidation Potential : It is the tendency of an electrode to loss electrons
Reduction potential: It is the tendency of an electrode to gain electrons
Electrode Potential: It is the tendency of an electrode to loss or gain electrons
Single Electrode Potential: It is the tendency of an electrode to loss or gain electrons when it is dipped in
its own salt solution. (Standard- 1M concentration at 250C).
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12. ELECTROCHEMICAL CELL
Introduction
An electrochemical cell is a device in
which a redox reaction is utilized to get
electrical energy.
An electrochemical cell is also commonly
referred to as voltaic or galvanic cell.
The electrode where reduction occurs is
called cathode.
The electrode where oxidation occurs is
called anode.
22-Feb-24
13. Construction
Electrochemical Cells are made up of two half-cells, each consisting of an electrode
which is dipped in an electrolyte. The same electrolyte can be used for both half cells.
These half cells are connected by a salt bridge which provides the platform for ionic
contact between them. A salt bridge minimizes or eliminates the liquid junction
potential.
The practical application of an electrochemical or galvanic cell is the Daniel cell.
It consists of a Zn electrode dipping in ZnSO4 solution and a Cu electrode dipping in
CuSO4 solution.
EMF= Eoxi + E Red
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14. Cell reaction
Anode : Zn → Zn2+ + 2e- (Oxidation) {0.76V}
Cathode : Cu2+ + 2e- → Cu (reduction) {0.34V}
Overall : Zn + Cu2+ → Zn2+ + Cu (Redox)
Representation of Daniel cell : Zn / Zn2+ || Cu2+ / Cu
Zn / ZnSO4 (1M) // CuSO4 (1M) / Cu
Cell EMF : 1.1 V
EMF= Eoxi + E Red
= EZn + Ecu = 0.76+0.34
CuSO4 - Cu2+ + SO4
2-
15. Anode : Zn Cathode : Cu
Zn Zn2+ ZnSO4 CuSO4 Cu2+ Cu
, / ;
Zn / Zn2+ Cu2+ / Cu
Zn / ZnSO4 (1M) CuSO4 (1M) / Cu
Zn , Zn2+ Cu2+ , Cu
Zn / Zn2+ || or // Cu2+ / Cu
17. Electrochemical Series
The standard electrode potentials of a number of electrodes are arranged in the
increasing order of reduction potential at 25°C is referred to as emf or electrochemical
series.
Characteristics of electrochemical series:
Lithium is the first member of the series.
Highly reactive metal systems are at the top of the series.
In other words, good reducing agents are at the top of the series, having the negative sign and act as
anode.
All good oxidizing agents are at the bottom of the series , having the positive sign and act as cathode.
Hydrogen system is at the middle of the series. All the elements which displace hydrogen from dilute
acids are placed above it.
18.
19.
20. Applications of Electrochemical Series
To Find Reactivity of Metals
As we move down in the electrochemical series reactivity of metal
decreases
Alkali metals and alkaline earth metals at the top are highly reactive.
They can react with cold water and evolve hydrogen. They dissolve in
acids forming salts.
Metals like Fe, Pb, Sn, Ni and Co which lie a little down in the series,
do not react with cold water but react with steam and evolve hydrogen.
Metals like Cu, Ag and Au which lie below the hydrogen are less
reactive and do not evolve hydrogen from water.
21. Calculation of standard EMF of the cell
EMF= Eoxi + E Red
Zn & Cu Couple
EMF= Eoxi + E Red
= EZn + E Cu
= 0.76+ + 0.34
= 1.1V
Fe & H2
EMF= Eoxi + E Red
EMF= EFe + E H2
= 0.441+ 0
0.441V
22. Ni & Hg Couple
Ni – Anode
Hg - Cathode
EMF= Eoxi + E Red
= ENi + E Hg
= 0.236 + 0.61= 0.846V
EMF = Standard reduction potential of R.H.S electrode- Standard reduction potential of L.H.S
electrode
E0 = E0
RHS - E0
LHS
= E0
Hg- E0
Ni
= 0.61 – (-0.236)
= 0.61+0.236 = 0.846V
= ENi + E Hg
= 0.236+0.61
= 0.846V
23. Cr & Sn Couple
Cr – Anode
Sn - Cathode
EMF= Eoxi + E Red
EMF= ECr + E Sn
= -0.74+(-0.14)
= 0.60V
EMF = Standard reduction potential of R.H.S electrode-
Standard reduction potential of L.H.S electrode
E0 = E0
RHS - E0
LHS
= E0
Sn - E0
Cr
= – 0.14 -(-0.74)
= -0.14+0.74 = 0.60V
24. Zn + CuSO4 ZnSO4 + Cu
Cu + ZnSO4 No reaction
Zn + H2SO4 ZnSO4 + H2
Ag + H2SO4 no reaction
25. For Studying displacement reaction
Elements having higher reduction potential will gain electrons and that having lower
reduction potential will lose electrons. Hence element higher in electrochemical series
can displace an element placed lower in electrochemical series from its salt solution.
Example
Can zinc displaces copper from its salt solution?
Zn displaces Cu from CuSO4, because, zinc is placed higher in electrochemical series
while Cu is placed lower in electrochemical series. Hence zinc can easily displace
copper from CuSO4.
Zn+CuSO4 --------> ZnSO4 + Cu
26. For choosing elements as Oxidizing Agents
The elements which have more electron-accepting tendency are oxidizing agents. The
strength of an oxidizing agent increases as the value of reduction potential becomes more
and more positive. Elements at the bottom of the electrochemical series have higher (+ve)
reduction potential. So they are good oxidizing agents. Thus, oxidizing power increases
from top to bottom in the series.
Example- F2 is a stronger oxidant than Cl2, Br2 and I2.
Cl2 is a stronger oxidant than Br2 and I2.
27. For choosing elements as Reducing Agents
The elements which have more electron losing tendency are reducing agents. The
power of reducing agent increases as the value of reduction potential becomes more and
more negative. Elements at the top of the electrochemical series have higher (-ve)
reduction potential. So they are good reducing agents. Thus, reducing power decreases
from top to bottom in the series.
Example-
The element like Zn, K, Na, Fe, etc. are good reducing agent.
28. Displacement of hydrogen from dilute acids by metals
The metal which can provide electrons to H+ ions present in dilute acids for reduction evolve hydrogen
from dilute acids. The metal having negative values of reduction potential possesses the property of
losing an electron or electrons.
Thus, the metals occupying top positions in the electrochemical series readily liberate hydrogen from
dilute acids and on descending in the series, tendency to liberate hydrogen gas from dilute acids
decreases.
The metals which are below hydrogen in the electrochemical series like Cu, Hg, Au and Pt do not evolve
hydrogen from dilute acids.
Example
Zinc reacts with dil.H2SO4 to give H2 but Ag does not. Why?
Zn+H2SO4 --------> ZnSO4 + H2 ; E0
Zn = -0.76 volts
Ag+H2SO4 --------> No reaction; E0
Ag = +0.80 volts
The metal with a positive reduction potential will not displace hydrogen from an acid solution.
29. Displacement of hydrogen from water
Iron and the metals above iron are capable of liberating hydrogen from water. The tendency
decreases from top to bottom in the electrochemical series.
Alkali metals and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe
liberate hydrogen from hot water or steam.
For Calculation of Standard emf of the cell
Standard reduction potential values are given in emf series. From the values E0
cell is calculated
using formula
E0
cell or standard emf of a cell = E0
oxi(cathode) - E0
red(anode)
30. For predicting spontaneity of the cell reaction
E0
cell > 0 cell reaction is spontaneous
E0
cell < 0 cell reaction is non-spontaneous
E0
cell = 0 cell reaction is in equilibrium
For determination of equilibrium constant for a reaction
We know that
-∆G0 = RTlnK
= 2.303RT logK
log K =
log K = (-∆G0 = nFE0)
Thus, from the value of E0 for a cell reaction, its equilibrium constant can be calculated.
32. Reference electrode
The electrode of standard potential with which we can compare the potentials of other
electrode is called a reference electrode.
It can acts both as anode or cathode depending upon the nature of other electrode.
Classification:
i) Primary reference electrodes Ex : Standard Hydrogen Electrode (SHE)
ii) Secondary reference electrodes Ex: Calomel, Ag/AgCl electrodes and
Quinehydrone electrodes
v
Reference
Electrode
Working or Indicator
Electrode
The part of
the cell that
is kept
constant
The part of
the cell that
contains the
unknown
solution
33. Construction and Working of Standard Calomel Electrode (SCE)
A common reference electrode.
It consists of a wide glass tube.
Mercury is placed at the bottom of the
glass tube.
A paste of mercury and mercurous
chloride(Calomel) is placed above the
mercury. The remaining portion above the
paste is filled with a KCl solution of
known concentration (0.1N, 1.0N and
saturated) .
A platinum wire is immersed into the
mercury to obtain electrical contact.
The side arm is provided for making
electrical contact through a salt bridge.
Pt wire
34. Electrode representation:
Hg, Hg2Cl2(s)// KCl(satd. solution)
Working of the electrode:
If it acts as Cathode :
Hg2Cl2 Hg2
2+ +2Cl-
Hg2
2+ +2e- 2Hg
Hg2Cl2+2e- 2Hg+2Cl-
if it acts as anode :
2Hg Hg2
2+ +2e-
Hg2
2+ +2Cl-
2Hg+2Cl- Hg2Cl2+2e-
Hg2Cl2
KCl E in Volts
saturated 0.2422V
1.0N 0.2800 V
0.1N 0.3338V
Electrode potential
36. Measurement of pH using Calomel electrode
Hydrogen electrode containing a solution of unknown pH combine with the
calomel electrode to set up a complete cell.
We use a saturated calomel electrode as the reference and the complete cell can be represented as :
Pt, H2 (1atm)/ H+ (C=?)/ /KCl(satd. solution)/Hg2Cl2(s), Hg
37.
38. Merits
It is easy to construct and easy to transport.
It provides almost a constant potential value with varying temperature and finds
application in laboratories for measuring potential of an electrode.
It is used in corrosion studies.
39. Ion selective electrode- Glass electrode
In Ion selective electrode, a membrane is in contact with a solution, with which it
can exchange ions. This ISE is responsive towards H+ and extensively used to
measure the pH of solution.
In 1906, cremer found that a thin bulb of glass conducted electricity when he put
two solutions of different acid strengths inside and outside the bulb.
The potential developed at the glass was in accordance with Nernst equation.
Ion selective electrode is one which selectively responds to a specific ion in a
mixture and the potential developed at the electrode is a function of the
concentration of that ion in solution
Example- Glass electrode
40. Ion selective electrode- Glass electrode
Construction:
It consists of a thick walled glass tube with a very thin glass bulb placed at the bottom.
The thickness of the bulb is 0.01-0.03mm. Glass have high electrical resistivity.
In glass electrode potential depends upon the pH of the medium
The glass electrode consists of a glass bulb made up of special type of glass called Corning-
015 contains Na2O(22%), CaO(6%) and SiO2(72%) with high electrical conductance and
high hygroscopic in nature..
The mixtures of the oxides is melted and cooled to form the glass. By altering the
composition of the glass, it is possible to make the electrode selective for different ions.
The glass bulb is filled with a solution of constant pH(0.1M HCl) and insert with a Ag-
AgCl electrode, which is the Internal reference electrode and also serves for the external
electrical contact.
Thin
walled
glass bulb
0.1M HCl
AgCl
coated Ag
wire
Glass electrode
41.
42. Electrode representation
Ag-AgCl /(0.1M) HCl/ Glass
Working:
The glass electrode works on the principle that when a thin glass membrane is placed between two
different concentration of a solution, a potential is developed at layers of the glass membrane. This
potential arises due to difference in the concentration of H+ ion inside and outside the membrane.
The potential developed is known as glass electrode potential EG and can be expressed as
H+ + e- 1/2H2
EG = E0
G --- 0.0591 Log[ H+]
--------
n
EG = E0
G - 0.0591 pH
potentiometer
43.
44.
45. The glass electrode is placed in the solution under test and coupled with a saturated calomel electrode.
Cell representation
Ag-AgCl /(0.1M) HCl/ Glass/ solution of unknown pH//saturated calomel electrode
The EMF of the cell is determined experimentally. From the emf, pH of the solution is calculated as
follows.
E cell = Ecalomel- Eglass
E cell = 0.2422 - (E0
G + 0.0591 pH)
E cell = 0.2422- E0
G -0.0591 pH
pH =
The value of E0
G can be determined by using a solution of known pH.
0.2422- Ecell -E0
G
0.0591
46. Advantages of Glass electrode
It is very easy to construct and simple to operate.
The potential developed remains constant for long time.
This electrode can be used with very small amount of the test solution.
This electrode can be used even in the presence of oxidized impurities, reducing impurities ,poison
molecules etc.,
It can be used in turbid coloured and colloidal solutions.
Limitations of Glass electrode
Since the glass membrane offers very high resistance, ordinary potentiometer cannot be used. It is
necessary to use electronic potentiometers.
This electrode cannot be used to determine the pH above 12.
48. It is similar to direct volumetric titration.
Instead of indicator, potential is measured across the analyte
Two electrodes are used – an indicator electrode and reference electrode
Since the potential of reference electrode is constant and with the potential of indicator electrode, the
concentration of ion in the analyte can be measured.
Ecell is recorded at intervals as the titrant is added.
A graph of potential against volume added can be drawn and the end point of the reaction is halfway
between the jump in voltage.
Ecell depends on the concentration of the interested ions with which the indicator electrode is in contact.
For example, the electrode reaction may be
Mn++ ne−-----> M
As the concentration of Mn+ changes, the Ecell changes correspondingly. Thus the potentiometric
titration involve measurement of Ecell with the addition of titrant.
49. Theory
Potential of an Electrode dipping in solution of eletrolyte depends upon the concentration of active ions.
E= E⁰ + (RT/nF) log C
Small Change in active ion concentration in the solution changes the electrode potential correspondingly
Concentration of Active ion decreases electrode potential of indicator electrode decreases
The potential of Indicator electrode is measured potentiometrically by connecting with a reference electrode
(Saturated Calomel Electrode)
22-Feb-24
50. Determination of End point
The emf of a cell changes by the addition of a small amount of titrant. So concentration
of reversible ion in contact with indicator electrode changes.
Record the change in emf with every small addition
The changes of potential will be slow at first, but at equivalence, the point change will
be sharp
The values are plotted against corresponding volume changes.
Change in emf with addition of titrant (⧍E/⧍V) is plotted against volume (V)
The maximum of the curve gives the end point.
22-Feb-24
51. Fig (a)– Volume of Titrant Vs Emf
Fig(b) Volume of titrant Vs (⧍E/⧍V)
22-Feb-24
53. When it is titrated against K2Cr2O7 the following redox reaction takes place
Fe2+ is converted to Fe3+ and its concentration increases.
The potential is determined by the ratio of [Fe2+] / [Fe3+]
Till the end point, there is variation in potential with respect to the ratio of [Fe2+] / [Fe3+]
and after end point there’s no much variation in potential.
(Oxidised)
(reduced)
54.
55. 22-Feb-24
Advantages of Potentiometric Titrations
Potentiometric titrations can be carried out in colored solutions, where indicators
cannot be used
There is no need of prior information about the relative strength of titrant before the
titration
56. CONDUCTOMETRIC TITRATION
Volumetric method based on the measurement of conductance of the solution during the titration
22-Feb-24
Conductance
Number and Charge on the free
ions
Mobility of the ions
57. Measurement of conductance using cells
A conductance of the solution is measured using conductance cell.
Which is made of a glass tube in which the platinized thin foils of
platinum electrodes are firmly fixed by sealing on a glass base.
Polarization is removed when the electrodes are coated with finely
divided platinum black( chloroplatinic acid + lead acetate )
The electrode are then washed repeatedly with distilled water and
finally with conductivity water.
After usage the electrodes should be kept in conductivity water.
The conductance of solution may be determined by measuring the
resistance of solutions into which a conductance cell is dipped.
Conductance measurements are used extensively in chemistry and
in chemical industries. The use of the method is based on the
information from the behavior of electrolytes.
58. Process
Taking a solution to be titrated in a beaker kept in a water
bath at a constant temperature.
Conductivity cell is dipped and connected to a conductivity
bridge.
The titrant is added from the burette(Fig)
Conductance is measured each addition of solution.
Recorded value is plotted the value of conductance against
the volume of the titrant.
Since the conductance of solution is proportional to the
concentration of ions present, the conductance first
decreases with increase in volume of titrant, it reach the
saturation point it increase with the addition of titrant.
From the graph end point is noted.
22-Feb-24
59. Procedure
Calibrate the instrument by releasing the calibration knob
Standard Sodium Hydroxide is taken in the burette
The given acids is made upto 100ml in the standard measuring flask (SMF)
20 ml of made up acids + 20 ml of conductivity water are added in 100 ml beaker
Conductance is noted for addition of every addition of 1ml of Standard Sodium
Hydroxide
Plot a graph between Volume of Standard Sodium Hydroxide Vs Conductance
End Points are noted from the graph
Equivalent Weight of Hydrochloric Acid = 36.5
Equivalent Weight of Acetic Acid = 60
22-Feb-24
60. Types of Conductometric Titrations
Acid –Base titration
Strong Acid Vs Strong Base
Weak Acid Vs Strong Base
Mixture of Weak and Strong Acid Vs Strong Base
Precipitation titration
Replacement titration
Redox titration
Complexometric titration
22-Feb-24
61. Strong Acid Vs Strong Base(HCl Vs NaOH)
Solution of electrolytes conducts electricity due to the presence of ions. The specific conductance of
solution is proportional to the concentration of ions in it. The reaction between HCl and NaOH may be
represented as
• H+ + OH------ H2O
When a solution of hydrochloric acid is titrated with NaOH, the fast moving hydrogen ions are
progressively replaced by slow moving sodium ions. As a result conductance of the solution decreases.
This decrease in conductance will take place until the end point is reached. Further addition of alkali
raises the conductance sharply as there is an excess of hydroxide ions.
A graph is drawn between volume of NaOH added and the conductance of solution. The exact end
point is the point of intersection of the two straight lines.
22-Feb-24
HCl + NaOH NaCl + H2O
63. Weak Acid Vs Strong Base
( CH3COOH Vs NaOH)
22-Feb-24
CH3COOH+NaOH 3COO‾ + Na+ +H2O
H+ +OH---> H20
CH3COO- + Na-- CH3COONa
64. Mixture of Weak and Strong Acid Vs Strong Base
(HCl and CH3COOH Vs NaOH)
22-Feb-24
HCl + NaOH NaCl +H2O
CH3COOH+NaOH CH3COO
-Na+ +H2O
65. Advantages
This method can be used with very diluted Solution
This method can be used with Coloured and Turbid solution in which the end point
cannot be seen clearly
This method can be used in which there is no suitable indicator
Used for acid-base, redox, precipitation titration etc.,
22-Feb-24
66. REFERENCES:
1.Palanisamy P.N., Manikandan P., Geetha A.& Manjula Rani K, “Applied
Chemistry”, 6th Edition, Tata McGraw Hill Education Private Limited, New
Delhi, 2019.
2 .Paya Payal B.Joshi, Shashank Deep., “Engineering Chemistry”, Oxford
University Press, New Delhi, 2019.
3.Palanna O., “Engineering Chemistry”, McGraw Hill Education, New
Delhi, 2017.
22-Feb-24