Engineering Chemistry Textbook 2024 PDF

NEW EDITION A TEXT BOOK OF ENGINEERING CHEMISTRY For I Semester B.E. / B.Tech. Students (As per the SSN College of Engineering (Autonomous) Syllabus) COMMON TO ALL BRANCHES Dr. Arun Luiz T Dr. Davis Presley S I Dr. Murugesan A Dr. Tamilmani V Department of Chemistry SSN COLLEGE OF ENGINEERING Kalavakkam, Chennai DHANAM PUBLICATIONS Chennai - 600 042 Phone: 044 - 4303 6502 Mobile: 99406 41496 / 98401 25695 Email [email protected] © Reserved First Edition – August 2024 Copyright Warning: No part of this book may be reproduced or transmitted in any form of by any means electronic or mechanical including photocopying or recording or by any information storage and retrieval system without permission in writing from the author and publisher. Price: Rs. /- ISBN: 978 - 81 - 89843 - 76 - 2 For copies contact DHANAMPUBLICATIONS No.18, 3rd Main Road Dhandeeswaram Nagar, Velachery, Chennai - 600 042. Laser Typeset at: JV Computers, No.62, Pillaiyar Koil Street, Kanagam, Chennai 600 113. Printed at: Sekar Offset, Triplicane, Chennai 600 005. ii PREFACE This book has been written to cater the needs of first semester B.E./ B.Tech. students according to the new syllabus (Regulation 2024) prescribed followed by SSN College of Engineering (Autonomous affiliated to Anna University). This book comprises of five units which covers the entire syllabus. Topics like water technology, electrochemistry, corrosion, polymers and analytical techniques have been written in a simple and lucid manner. Most of these topics are traditionally taught in engineering chemistry courses in several universities and institutes. Figures and tables are incorporated wherever necessary to make the concept clearer. Authors hope that this book will be useful for both students and faculty alike. Despite all precautions and care, some error or misprints might have been left in advertently. The author welcomes comments, suggestions and criticisms for the improvement of the book. - Authors iii Course Code Course Title L T P C ENGINEERING CHEMISTRY XXXX 3 0 2 4 (THEORY CUM PRACTICE) OBJECTIVES:  To impart knowledge on various aspects of chemistry  To improve the ability of students to think logically and solve the problems in industries and day-to-day life efficiently Unit I WATER AND ITS TREATMENT 9+6 Theory: Hardness of water – types – expression of hardness – units – estimation of hardness of water by EDTA – numerical, Water quality parameters - WHO guidelines and BIS guidelines. Industry water treatment - External treatment – Ion exchange process, zeolite process. Domestic water treatment - Reverse Osmosis-Working of RO system – Advantages/limitations. Practical: Exp.1: Estimation of hardness of water Exp.2: Hardness of different sources of water Laboratory Demonstration:  Ion Exchange process Unit II ELECTROCHEMISTRY 9+6 Theory: Conductance- Conductometric titration and its applications - estimation of strong acid, estimation of mixture of strong and weak acids (numerical based on conductance). Electrochemistry-redox reaction - types of Electrode-Ion selective electrodes – Glass electrodes-measurement of pH-potentiometry Energy storage systems for electric vehicles – Principle & Electrochemistry of a H2–O2 fuel cell, Li-ion battery, Na-ion battery. Practical: Exp.3: Conductometric titrations- strength of mixture of acids Exp.4: Estimation of strong acid-pH Exp.5: Estimation of ferrous ion by potentiometry iv Unit III CORROSION 9+6 Theory: Corrosion - Types of corrosion - wet corrosion – mechanism –galvanic corrosion -differential aeration- Rate of corrosion, Corrosion control –Cathodic protection-electroless plating (Printed Circuit Board), Corrosion in different industries –concrete (reinforcing steel in concrete), boilers, electronic components Practical: Exp.6: Rate of corrosion Industrial visit:  Plating industry- Electroless plating Unit IV POLYMERS IN EVERY DAY LIFE 9+6 Theory: Polymers and Polymerization: types of polymerization: addition and condensation – Properties: Crystallinity, Glass Transition temperature (Tg), Average Molecular weight-viscosity method &PDI, tacticity, polymer recycling- biodegradable polymers. Practical: Exp.7 : Molecular weight of water soluble polymer by viscosity method Exp.8: Finding the Tg point of different polymer. Laboratory Demonstration: Chemical recycling of post consumed polymer Unit V ANALYTICAL TECHNIQUES 9+6 Theory: Spectroscopy: Beer-Lambert’s law. Colorimetric estimation of Fe3+, Principle, working and applications of IR, UV-Visible spectroscopy and Chromatography (TLC and column) Practical: Experiment 9: Colorimetric estimation of Fe3+ ions Laboratory Demonstration: Thin Layer Chromatography Theory Periods 45 Practical Periods 30 Total Periods 75 v Course Outcomes: Upon successful completion of the course, students will be able to CO1: Analyze the water samples, categorize based on the nature of impurities and suggest suitable method of treatment for domestic and industrial usage (BL:L3). CO2: Understand the principles of electrochemistry and apply the principles for zero-emission vehicles (BL: L3). CO3: Identify the type of corrosion and analyze different preventive methods of corrosion in various industries (BL: L3). CO4: Acquire sound knowledge on polymeric material and understand the need for sustainable polymeric materials (BL: L3) CO5: Apply analytical skills of techniques such as chromatography, spectroscopy to characterize materials to solve real life problems (BL: L3)  Text Books: 1. Jain P.C. and Monika Jain, ‘Engineering Chemistry’ 17th edition, Dhanpat Rai Publishing Company (P) Ltd, New Delhi, 2023. 2. S.S.Dara, ‘The Text Book of Engineering Chemistry, S.Chand & Co.Ltd, New Delhi, 2011 References: 1. N.F. Gray, ‘Water Technology-An Introduction for Environmental Scientists and Engineers’ Third Edition, Taylor & Francis, USA, 2010 2. S. Glasstone, ‘An Introduction to Electrochemistry’ East-West Press Pvt. Ltd., New Delhi, 2007 3. Bengt Sundén, ‘Hydrogen, Batteries and Fuel Cells’ Academic Press Inc, USA, 2019 4. R. Gowariker, N. V. Viswanathan, Jayadev Sreedhar, ‘Polymer Science’ New Age International (P) Ltd, New Delhi, reprint, 2005 4. R. Gopalan, K. Rangarajan, P.S. Subramanian, “Elements of Analytical Chemistry” Sultan Chand & Sons, 2003 5. B.Viswanath, B, P.S. Raghavan, ‘Practical Physical Chemistry’, ViVa Books PVT. Ltd., New Delhi, 2012 vi CO to PO Mapping COs POs PSOs 1 2 3 4 5 6 7 8 9 10 11 12 1 3 1 1 1 1 2 3 1 1 1 3 3 1 1 1 1 1 4 3 1 1 1 1 1 5 3 1 1 PO1: Engineering knowledge: Apply the knowledge of mathematics, science, engineering fundamentals, and an engineering specialization for the solution of complex engineering problems. PO4: Conduct investigations of complex problems: Use research-based knowledge and research methods including design of experiments, analysis and interpretation of data, and synthesis of the information to provide valid conclusions. PO6: The engineer and society: Apply reasoning informed by the contextual knowledge to assess societal, health, safety, legal, and cultural issues and the consequent responsibilities relevant to the professional engineering practice. PO7: Environment and sustainability: Understand the impact of the professional engineering solutions in societal and environmental contexts, and demonstrate the knowledge of, and the need for sustainable development. PO9: Individual and team work: Function effectively as an individual, and as a member or leader in diverse teams, and in multidisciplinary settings. PO12: Life-long learning: Recognize the need for, and have the preparation and ability to engage in independent and life-long learning in the broadest context of technological change. vii CONTENTS Unit – 1 Page Nos. Water and its Treatment ………………………………………. 1.1 – 1.24 Unit – 2 Electrochemistry ……………………………………………… 2.1 – 2.42 Unit – 3 Corrosion ……………………………………………………... 3.1 – 3.27 Unit – 4 Polymers in Every Day Life ………………………………….. 4.1 – 4.32 Unit – 5 Analytical Techniques ………………………………………… 5.1 – 5.45 viii Water and its treatment 1.1 UNIT - 1 WATER AND ITS TREATMENT Hardness of water – types – expression of hardness – units – estimation of hardness of water by EDTA – numerical, Water quality parameters - WHO guidelines and BIS guidelines. Industry water treatment - External treatment – Ion exchange process, zeolite process. Domestic water treatment - Reverse Osmosis- Working of RO system – Advantages/limitations. 1.1. INTRODUCTION Earth is the only planet containing water. Hence it is known as “blue planet”. About 71% of the earth’s surface is covered by water, out of which only 0.014% can be used for domestic, agriculture and industrial purposes and rest is locked in oceans, polar ice caps, giant glaciers and rock crevices. Nearly 70% of human body is made up of water. Water is one of the most widely used and abundant chemical in nature. Water is so essential part of plant and animal life that without which life cannot survive. Hence, water is also known as “elixir of life”. Impurities in Water Water dissolves many salts and gases. So it is known as an “universal solvent”. So many undesirable impurities are present in water. These impurities can be classified as 1. Dissolved impurities: The following impurities may be present in the water in the dissolved form. a) Inorganic salts: Cations like Ca2+, Mg2+, Na+, K+, Fe2+, Al3+ and sometimes traces of Cu2+ and Zn2+ salts. Anions like Cl-, SO42-, HCO3-, NO3- and sometimes F- and NO2-. b) Organic compounds: Hydrocarbon, phenols etc. 1.2 Engineering Chemistry c) Dissolved gases: Like O2, CO2, N2, oxides of nitrogen and sulphur, and sometimes NH3 and H2S. 2. Colloidal impurities: Like clay, silica, Al(OH)3, Fe(OH)3, organic waste products, humic acids, amino acids, proteins etc. 3. Suspended impurities: Inorganic matter like clay and sand; organic matter like oil globules. 4. Biological impurities: Bacteria and other microorganism like virus, algae, fungi, phytoplanktons etc. These impurities may give water a bad taste, odour, turbidity, etc. and cause hardness, frothing, corrosiveness, staining etc. 1.2. HARDNESS Water which does not produce lather with soap is known as hard water. It is due to the presence of bicarbonates, chlorides and sulphates of calcium and magnesium. Salts of Fe, Mn and other heavy metals also can cause hardness. Soap is sodium or potassium salts of higher fatty acids like oleic, palmetic stearic acid etc. Sodium salt is soluble in water while corresponding calcium and magnesium salts are insoluble in water. When soap is added to soft water, it dissolves and forms lather easily. When a hard water sample is treated with soap, it does not form lather but it gets precipitated in the form of insoluble salts (scum/precipitate) of calcium and magnesium. CaCl2 + 2C17 H35 COONa → (C17 H35 COO)2Ca + 2NaCl CaSO4 + 2C17 H35 COONa → (C17 H35 COO)2Ca + 2Na2SO4 Mg(HCO3)2 + 2C17 H35 COONa → (C17 H35 COO)2Mg + 2NaHCO3 (From hard water) (Soap) (Insoluble ppt) In case of hard water, initially soap is used to precipitate out the hardness causing ions. No lather is formed until all the hardness causing ions are removed. So, large amounts of soap is consumed unnecessarily. Anions associated with these metal ions Water and its treatment 1.3 include chloride, sulphate, bicarbonate but these do not contribute to the hardness of water. TYPES OF HARDNESS: Temporary Hardness (Carbonate hardness or alkaline hardness): It is due to the presence of bicarbonates of calcium and magnesium. It can be easily removed by boiling. On boiling, bicarbonates are decomposed to insoluble carbonates and hydroxides. Ca(HCO3)2 → CaCO3  + H2O + CO2  Mg(HCO3)2 → Mg(OH)2  + 2 CO2  Permanent Hardness (Non-carbonate hardness or non-alkaline hardness): This is due to the presence of chlorides and sulphates of calcium and magnesium. This type of hardness cannot be removed by boiling. The hardness can be removed by lime soda process, zeolite process and ion exchange process. 1.3. CALCIUM CARBONATE EQUIVALENT (Degree of hardness) The concentration of hardness producing salts or ions such as bicarbonates, chlorides and sulphates of Ca, Mg, etc. is measured in terms of equivalent amounts of calcium carbonate. The weights of different salts causing hardness are converted to weights equivalents to that of CaCO3 CaCO3 eq. of hardness = Mass of hardness producing substance (in mg) x Mol. mass of CaCO3 Molecular mass of hardness producing saltCaCO3 is particularly selected to express the hardness of water because: 1. CaCO3 is the most insoluble salt which can be precipitated during the water treatment. 2. Since the molecular mass of CaCO3 is 100 (eq. weight = 50), the calculation becomes easy. 1.4 Engineering Chemistry Hardness causing Salt Molecular Weight Ca(HCO3)2 162 Mg(HCO3)2 146 CaSO4 136 CaCl2 111 MgSO4 120 MgCl2 95 Ca(NO3)2 164 Mg(NO3)2 148 Units of hardness 1. Parts per million (ppm) It is the parts of calcium carbonate equivalent hardness per 106 parts of water. i.e. 1ppm = 1 part of CaCO3 equivalent hardness in 106 parts of water. 2. Milligrams per litre (mg/L) It is the number of milligrams of calcium carbonate equivalent hardness present per litre of water. Thus, 1 mg/L = 1 mg of CaCO3 equivalent hardness per1 L of water But 1 L water = 1 kg = 1000g = 106 mg 1 mg/L = 1 mg of CaCO3 equivalent per 106 mg of water = 1 part of CaCO3 equivalent per 106 parts of water = 1 ppm 3. Degree Clark (0Cl) It is the number of grains of CaCO3 equivalent hardness per gallon of water (or) it is the parts CaCO3 equivalent hardness per 70,000 parts of water. Water and its treatment 1.5 1°Cl = 1 grain of CaCO3 equivalent hardness per gallon of water 1°Cl = 1 part of CaCO3 equivalent hardness per 70,000 parts of water 4. Degree French (0Fr) It is the parts of calcium carbonate equivalent hardness per 105 parts of water. i.e. 10Fr = 1 part of CaCO3 equivalent hardness in 105 parts of water. Relationship between various units of hardness 1ppm = 1 mg/L = 0.1 0Fr = 0.07 0Cl 1.4. ESTIMATION OF HARDNESS OF WATER Hardness in water is due to the presence of dissolved salts of calcium and magnesium. It is unfit for drinking, bathing, washing and it also forms scales in boilers. Hence it is necessary to estimate the amount of hardness producing substances present in the water sample. The estimation of hardness is based on complexometric titration. Aim: To estimate the amount of total hardness present in the given sample of water by EDTA titration method. Apparatus required: 50 ml Burette, 20 ml Pipette, 250 ml Conical flask, 100 ml Beaker, 250 ml beaker, Glass funnel. Reagents: EDTA solution, Standard CaCO3 solution, Eriochrome Black–T indicator, Buffer solution. Principle : EDTA ( Ethylene diamine tetra acetic acid) forms colourless stable complexes with Ca2+ and Mg2+ ions present in water at pH = 9-10. To maintain the pH of the solution at 9-10, buffer solution (NH4Cl + NH4OH) is used. Eriochrome Black-T (E.B.T) is used as an indicator. The sample of hard water is first treated with buffer solution and EBT indicator which forms unstable, wine-red coloured complexes with Ca2+ and Mg2+ present in water. Now this wine red complex is 1.6 Engineering Chemistry titrated against EDTA solution (of known strength) the colour of the sample changes from wine red to steel blue showing the endpoint. Preparation of Solutions Following solutions are prepared for estimation of various types of hardness 1. Standard Hard Water (SHW): Dissolve 1 gram of pure CaCO3 in minimum quantity of dil. HCl and dilute to 1L with distilled water. 1mL of SHW  1mg of CaCO3 equivalent hardness 2. EDTA solution: 4 gram of pure EDTA crystal along with 0.1 gram of MgCl2 in 1L of distilled water 3. Indicator: 0.5 gram of EBT in 100 mL of alcohol. 4. Buffer solution: Add 67.5 gram of NH4Cl in 570 mL of concentrated ammonia solution and further dilute to 1L with distilled water Water and its treatment 1.7 PROCUDERE Titration I: Standardization of EDTA 20 mL of standard hard water (SHW) is pipetted out into a clean conical flask. 5 mL of ammonia buffer and few drops of EBT indicator are added. The solution turned wine red in colour and is titrated against EDTA taken in the burette. The end point is the change in colour from wine red to steel blue. The titration is repeated for concordant value and readings are tabulated. Let the volume of EDTA consumed by SHW be V1 mL. From the titre value, EDTA equivalence is calculated. Titration II: Estimation of Total Hardness 20 mL of given water sample is pipette out into a conical flask and 5 mL of ammonia buffer and few drops of EBT indicator are added and the titration is carried out as before. Let the volume of EDTA consumed by sample water be V2 mL. From the titre value the total hardness of the given sample(s) can be calculated. Titration III: Estimation of Permanent Hardness 20 mL of boiled water sample is pipette out into a conical flask and 5 mL of ammonia buffer and few drops of EBT indicator are added and the titration is carried out as before. Let the volume of EDTA consumed by boiled sample water be V3 mL. From the titre value the permanent hardness of the given sample can be calculated. The difference between the Total and Permanent hardness gives the Temporary Hardness of the water sample. Calculation for Titration I 1mL of SHW  1mg of CaCO3 equivalent hardness 20 mL of SHW  20 mg of CaCO3 equivalent hardness 20 mL of SHW = V1 mL of EDTA  20 mg of CaCO3 = V1 mL of EDTA  1 mL of EDTA = 20/ V1 mg of EDTA (This is known as EDTA equivalence) 1.8 Engineering Chemistry Calculation for Titration II and III 20 mL of given water sample require V2 mL of EDTA V2  20 Total Hardness in 20 mL of water sample = mg of CaCO3 V1 V2  20 1000 1L (1000 mL) of given water contains = mg of CaCO3 V1  20 V2 ×1000 Total hardness = ppm of CaCO3 equivalents V1 Similarly, 20 mL of boiled water sample require V3 mL of EDTA V  20 Permanent Hardness in 20 mL of water sample = 3 mg of CaCO3 V1 V3  20 1000 1L (1000mL) of given water contains = mg of CaCO3 V1  20 V3 ×1000 Permanent hardness = ppm of CaCO3 equivalents V1 Temporary Hardness = Total Hardness – Permanent Hardness 1000(V2  V3 ) = ppm of CaCO3 eq. V1 Water and its treatment 1.9 PROBLEMS 1. Calculate the total hardness of water containing 16.8mg/L of Mg(HCO3)2,19.0mg/L of MgCl2,24.0mg/L of MgSO4 and 22.2 mg/L of CaCl2. Multiplication Constituent Amount CaCO3 Equivalent Factor Mg(HCO3)2 16.8mg/L 100/146 16.8x100/145=11.5mg/L MgCl2 19.0mg/L 100/95 19.0X100/95=10.0mg/L MgSO4 24.0mg/L 100/120 24.0X100/120=20.0mg/L CaCl2 22.2mg/L 100/111 22.2X100/111= 20.0mg/L Total Hardness = (11.5+10.0+20.0+20.0)mg/L = 61.5mg/L 2. A sample of water on analysis has been found to contain following in ppm. Ca(HCO3)2= 10.5, Mg(HCO3)2= 12.5,CaSO4=7.5 CaCl2= 8.5,MgSO4=2.6 Calculate the temporary and permanent hardness. Multiplication Constituent Amount CaCO3 Equivalent Factor Ca(HCO3)2 10.5 ppm 100/162 10.5x100/162=6.48 ppm Mg(HCO3)2 12.5 ppm 100/146 12.5X100/146=8.56 ppm CaSO4 7.5 ppm 100/136 7.5X100/136=5.15ppm CaCl2 8.2ppm 100/111 8.2.2X100/111= 7.38ppm MgSO4 2.6 ppm 100/120 2.6X100/120= 2.16 ppm Temporary hardness = 6.48+8.56= 15.04ppm Permanent hardness = 5.15+7.38+2.16= 14.69ppm 1.10 Engineering Chemistry 3. A sample of water on analysis was found to contain the following impurities expressed in mg/litre Ca(HCO3)2 =10.0mg/L , Mg(HCO3)2 = 8.5mg/L, CaSO4 = 12.0 mg/L, MgSO4 = 14.0 mg/L. Calculate the temporary and permanent hardness and express the total hardness in degree French. Multiplication CaCO3 Equivalent Constituent Amount Factor Ca(HCO3)2 10.0 mg/L 100/162 10.0x100/162=6.17mg/L Mg(HCO3)2 8.5 mg/L 100/146 8.5X100/146=4.45mg/L CaSO4 12.0mg/L 100/136 12.0X100/136=8.82mg/L MgSO4 14.0mg/L 100/120 14.0X100/120= 11.67mg/L Temporary hardness = 6.17 +4.45= 10.62mg/L Permanent hardness = 8.82+11.67 = 20.49 mg/L Total hardness = 10.62 + 20.49 = 31.11 mg/LX 0.1OFr = 3.111 OFr 1.5. WATER QUALITY PARAMETERS The water we consume must adhere to specific quality standards. In industrial settings, water often requires treatment to achieve a quality suitable for various essential processes. To assess water quality, physical, chemical, and biological parameters are checked. Physical parameters include colour, taste, odour, and turbidity. Chemical parameters include pH, acidity, alkalinity, chloride, and hardness. Biological parameters involve the presence of bacteria, algae, and viruses. Colour Pure water contains no colour because it is essentially colourless. Colour of the water is affected by materials that decay from organic matter and inorganic matter such as rock, soil and stones. Common sources of contamination comes from organic Water and its treatment 1.11 dyes and inorganic compounds like Fe, Mn, and Cr etc. The colour of the water is measured in terms of Hazen units. Yellowish tinge in water indicates the presence of chromium whereas yellowish red indicates the presence of iron. Colour can be removed by coagulation, settling and filtration. Taste and Odour The taste of water can change and odours can develop when foreign substances are introduced. These substances may include organic materials, dissolved gases, and inorganic compounds, often originating from agricultural, natural, and domestic sources. Water can have a salty taste due to high levels of sodium, potassium and magnesium. Bacteria growing in your sink drain or hot water heater may cause odour. Naturally occurring hydrogen sulfide in your water supply can also cause rotten egg odour.Taste and odour can be improved by aeration or by activated charcoal treatment. Turbidity Turbidity refers to how cloudy a water is. Turbidity occurs due to higher concentration of silt, clay, and organic materials. Additional problems that are caused due to turbidity include high water treatment costs, affects the aquatic life, increases the water temperature because suspended particles absorb more heat. Turbidity of water can be removed by coagulation, settling and filtering. Total Dissolved Solids (TDS) TDS represents the total concentration of dissolved substances in water. TDS is made up of inorganic salts as well as small amounts of organic matter. Common inorganic salts that can be found in water include calcium, magnesium, potassium and sodium, which are all cations, and carbonates, nitrates, bicarbonates, chlorides and sulfates, which are all anions. Cations are positively charged ions and anions are negatively charged ions. These minerals can originate from a number of sources, both 1.12 Engineering Chemistry natural and as a result of human activities. TDS can be removed by Reverse Osmosis. Permissible limits of TDS is given below. Level of TDS( mg/L) Rating Less than 300 Excellent 300-600 Good 600-900 Fair 900-1200 Poor Above 1200 Unacceptable Fluoride (F-) Fluoride occur normally in water, its source is only natural, soluble fluoride minerals from the soil get dissolved in water. F- prevents tooth decay at levels around 1mg/L in drinking water. However when the fluoride content of drinking water is high greater than 10mg/L it causes fluorosis, which damages teeth and bones, the teeth become stained with colours. Fluoride contents in water can be removed by a process called Defluoridation. This process involves treating water with Ca2+ (aq) which combines with F- to form insoluble CaF2. pH pH value of a water is a measure of its acidity and alkalinity. The pH level is a measurement of the activity of the hydrogen atom, because the hydrogen activity is a good representation of the acidity or alkalinity of the water. The hydrogen ion concentration is represented by pH = - log10[H+] The pH scale, as shown below, ranges from 0 to 14, with 7.0 being neutral. Water with a low pH is said to be acidic, and water with a high pH is basic, or alkaline Water and its treatment 1.13 Fig. 1.1. The pH scale The pH scale is logarithmic, which means that each step on the pH scale represents a ten-fold change in acidity. For example, a water body with a pH of 5.0 is ten times more acidic than water with a pH of 6.0. And water with a pH of 4.0 is 100 times more acidic than water with a pH of 6.0. Alkalinity Alkalinity is a chemical measurement of a water’s ability to neutralize acids. It can measure the bicarbonate, hydroxide ions, and carbonate naturally present in the water. The majority of alkalinity in the surface water is from calcium carbonate, leached from soil and rocks. High alkalinity values are harmful to aquatic organisms where as low alkalinity values can contribute to the corrosion of pipes. Chemical oxygen Demand (COD) Chemical Oxygen Demand (COD) measures the amount of oxygen, in milligrams, needed to oxidize all the organic pollutants in water into carbon dioxide and water. The test involves boiling the water sample with potassium dichromate in a 1.14 Engineering Chemistry strongly acidic medium under reflux, which oxidizes the organic compounds present. The COD test is commonly used to assess the efficiency of water treatment plants and to quantify pollutants in water, wastewater, and aqueous hazardous wastes Biological Oxygen Demand (BOD) Biochemical Oxygen Demand (BOD) is the amount of dissolved oxygen needed by bacteria to biologically oxidize organic matter under aerobic conditions at 20°C over a period of 5 days. BOD is calculated by measuring the oxygen consumed by bacterial and chemical action in a closed sample of water maintained at 20°C over a period of 5 days. A lower BOD value indicates better water quality. Water with a BOD of less than 3 ppm is considered pure, while a BOD greater than 4 ppm suggests the water is polluted. 1.6. DRINKING WATER SPECIFICATIONS World Health organization (WHO) and Bureau of Indian standards (BIS) produces norms on water quality in form of guidelines that are used as basis for regulation. Assurance of drinking water safety is the foundation for the prevention and control of water borne diseases. Table 1 gives the guidelines for drinking water WHO and BIS guidelines for drinking water Standard limits as per Standard limit as per Parameter WHO in (mg/L) BIS in (mg/L) Colour in Hazen units No visible colour 5 Odour and taste Odourless and tasteless Unobjectionable and agreeable Turbidity ---- 5 pH Value ---- 6.5-8.5 Total hardness as CaCO3 ---- 300 Water and its treatment 1.15 Alkalinity ---- 200 Iron as Fe ---- 0.3 Chlorides as Cl 200-300 250 Total dissolved solids ---- 500 Nitrate 50 45 Fluoride 1.5 1.0 Lead 0.01 0.05 Arsenic 0.01 0.05 Cyanide 0.07 0.01 Mercury 0.006 0.001 1.7. INDUSTRY WATER TREATMENT Industrial water treatment addresses four main issues: scaling, corrosion, microbial activity, and the disposal of residual wastewater. Industrial waste water treatment Involves both Internal and external treatment. Internal treatment is done by adding chemicals to the water and the impurities are removed in the form of precipitates. External treatment of water External Treatment is the reduction or removal of impurities from feed water used in boilers. Common Industrial water treatment methods are 1. Filtration 2. Softening 3. Reverse osmosis 4. Ion exchange 5. Zeolite Process 1.16 Engineering Chemistry 1.8. DEMINERALIZATION OR ION EXCHANGE OR DEIONIZATION METHOD Demineralization is an external conditioning method. In this method cations are exchanged by H+ and anions are exchanged by OH-. Ion exchange resins are insoluble, cross linked, organic polymers with micro porous structure. Functional groups attached to the polymer chains are responsible for ion exchange properties. The ion exchange resins may be classified as cation exchange resins and anion exchange resins. Both have a network of divinyl benzene co-polymer.  Cation exchange resins (R-H): Cation exchange resins have –COOH or –SO3H group in their structure so that they can easily exchange H+ ions with cations of the water. Cation exchange resins can be represented as R-H Eg: Sulphonated or carboxylated styrene divinyl benzene copolymer. Commercially available cation exchangers are Amberlite IR-120 and Dowex-50.  Anion exchange resins (R-OH): Anion exchange resins have quarternary ammonium hydroxide [–NH3+OH-]. They are capable of exchanging OH- with anions present in hard water. Anion exchange resins can be represented as R-OH. Eg: Styrene divinyl benzene copolymer containing amino or quarternary ammo- nium salts. Commercially available anion exchangers are AG-1, AG-2. Fig. 1.2. Structures of cation and anion exchange resins Water and its treatment 1.17 Method: Cation and anion exchange resins are packed in separate cylindrical containers and hard water is passed through each of these consecutively. Hard water is first passed through cation exchange resin. The cations (Ca2+, Mg2+, Na+, K+ etc) present in the water get exchanged with H+ ions of the resin as shown by the reaction given below. Ca2+ + 2 R-H R2 Ca + 2 H+ Mg2+ + 2 R-H R2 Mg + 2 H+ (Hard water) (Cation exchanger) The water coming out of the first chamber contains free H+ ions and hence acidic in nature. Now it is passed through the anion exchange resin. Anion exchange resin exchanges anions (Cl-, SO42-, HCO3- etc) with OH- ions. Cl- + R-OH RCl + OH- HCO3- + R-OH R HCO3 + OH- (Anion exchanger) The OH- ions liberated from the chamber now react with free H+ ions (coming from the first chamber) to form water. H+ + OH- H2 O Thus the water coming out from the second chamber is free from all cations and anions that were initially present in hard water. Hence it is generally known as demineralized water or deionized water. Thus water becomes soft after this process. Regeneration of resin: As the continuous use of demineralization makes the resin exhausted, therefore for the future use, resins must be regenerated. Exhausted cation exchange resin is regenerated by passing a moderately concentrated (10-20%) hydrochloric acid or sulphuric acid solution through the chamber. R2 Ca + 2 HCl 2 R-H + CaCl2 (Exhausted cation (Regenerated resin) exchange resin) 1.18 Engineering Chemistry Exhausted anion exchange resin is regenerated by passing a moderately concentrated (10-20%) sodium hydroxide solution through the chamber. R-Cl + NaOH R-OH + NaCl (Exhausted anion (Regenerated resin) exchange resin) Fig. 1.3. Demineralization process Advantages  Highly acidic and highly alkaline water can be treated.  Treated water has very low hardness (~ 2 ppm) and is suitable for high pressure boilers. Disadvantages  Ion exchange resins are expensive.  If water is turbid, efficiency of the process is reduced as the pores are blocked by turbid impurities.  Water containing Fe, Pb and Mn cannot be treated as they form stable complexes with resins and it becomes difficult to regenerate the resins.  For effective water treatment, ion exchanger should possess following properties. Water and its treatment 1.19  They should be non-toxic.  They should not discolour the water being treated.  They should possess high ion-exchange capacity.  They should be resistant to chemical attack.  They must be capable of being regenerated and back washed easily.  They should have large surface area since ion-exchange is a surface phenomenon (adsorption). 1.9. ZEOLITE PROCESS OR PERMUTIT PROCESS Zeolites are naturally occurring hydrated sodium aluminosilicate minerals. The term "zeolite" comes from the Greek for "boiling stone," as zeolites released water when heated. Zeolites do not have a fixed composition and are examples of nonstoichiometric compounds. Artificial zeolite known as permutit is widely used for water softening. It has the general formula Na2O. Al2O3.xSiO2.yH2O (x= 5-13; y = 3-4) but is generally represented by Na2Ze. Zeolites are highly porous and they have three dimensional tunnels and cavities of different sizes. Zeolites are insoluble in water and are capable of reversibly exchanging its sodium ions with the alkaline earth group cations generally present in water. Process A cylindrical column is initially packed with a layer of coarse gravel followed by a layer of fine gravel. Above this permutit or zeolites are packed as shown in the figure. A slow stream of hard water is passed through this material. As a result, calcium and magnesium ions present in the hard water are exchanged with sodium ions in the permutit. The water that flows out from the zeolite bed contains sodium ions which do not cause hardness. Na2Ze + CaCl2 → CaZe + 2 NaCl Na2Ze + MgSO4 → MgZe + Na2SO4 1.20 Engineering Chemistry Fig. 1.4. Zeolite process Regeneration Over a period of time, permutit bed loses its sodium exchange capacity as it gets converted into a mixture of calcium and magnesium aluminosilicates. The zeolite bed is said to be “exhausted”. Exhausted zeolite bed can be regenerated by treating with 5-10% brine. CaZe + 2NaCl → Na2Ze + CaCl2 Resulting calcium or magnesium chloride produced are washed out through the tap at the bottom. Regenerated permutit is reused for softening of water. Advantages  Water of very low hardness (~10 ppm hardness) is obtained.  This method is very cheap as exhausted zeolite can be regenerated.  No sludge is formed during this process.  The equipment used is compact and uses little space.  Easy operation.  The process can be made automatic and continuous. Water and its treatment 1.21 Disadvantages  Zeolite process cannot be used for acidic water as it may destroy the zeolite bed.  Turbid water or water with suspended impurities cannot be used as these may clog the fine pores in the zeolites.  This treatment only replaces only the cations leaving all anions like HCO3-, CO32- etc. in soft water. Presence of these impurities may result in boiler corrosion (caustic embrittlement).  Zeolite process cannot be used for brackish water because brackish water also contains Na+ ions. So the ion exchange reaction will not take place.  Water containing manganese, iron and lead ions cannot be used as they bind irreversibly with zeolite. Such zeolites cannot be regenerated back to sodium zeolites 1.10. DESALINATION OF BRACKISH WATER Water contains large amount of salts dissolved in it. The process of removal of these dissolved salts present in the water is known as desalination. Salinity of water is primarily due to dissolved sodium chloride and to smaller extent due to other inorganic salts (mostly of calcium and magnesium).The salinity of water is expressed in terms of mg/L of dissolved salts. On the basis of salinity, water is divided into three categories. (i) Fresh water: Water having salinity less than 1000 mg/L (ii) Brackish water: Water having salinity in the range 1000 – 35000 mg/L. Brackish water has a peculiar salty taste. (iii) Sea water: Water having salinity greater than 35000 mg/L. Desalination can be achieved by two methods. a) removal of water from salt (eg: distillation or evaporation, solvent extraction, reverse osmosis etc) b) removal of salt from water (eg: electrodialysis). Reverse osmosis and electrodialysis are most commonly used for the desalination on a large scale. 1.22 Engineering Chemistry Reverse Osmosis When two solutions of different concentrations are separated by a semi permeable membrane, the flow of solvent takes place from a region of low (solute) concentration to a high (solute) concentration until the concentration of the both solutions is equal. This process is known as osmosis. The driving force for this phenomenon is known as osmotic pressure. But when a greater pressure, in excess of osmotic pressure is applied on the higher concentration side, the solvent flow reverses, i.e. the solvent is forced to move from higher (solute) concentration to lower (solute) concentration. This is known as reverse osmosis. Principle: Brackish water contains large amounts of dissolved salts and is more concentrated as compared to fresh water. Thus, if sea water is kept in contact with fresh water through a semi permeable membrane, then a pressure of 15-40 kg/cm2 is applied on sea water, then water will pass thorough the semi permeable membrane leaving behind the salts. Thus, reverse osmosis separates water from dissolved salts. Pressure is usually applied by a pump. Charcoal pre filters are placed before the membrane unit to filter out the suspended matter which would otherwise clog the pores of the membrane. Membrane used for reverse osmosis Earlier semi permeable membranes used were made up of cellulose acetate. More recently superior membranes made up of polymethacrylate and polyamide polymers have been introduced. The membrane consists of a skin of about 0.25 μ and a support layer 100 μ in width. Fig. 1.5. Reverse osmosis Water and its treatment 1.23 Advantages:  Reverse osmosis is known as the finest hyper filtration technique known. This process will allow the removal of particles as small as ions.  This process is very economical and convenient. It can be used at room temperature.  The membrane will last for more than 2 years. Replacement of membrane will take only few minutes.  It gives water, as and when required without any interruption.  RO process removes ionic and non-ionic impurities easily.  It is effective in removing colloidal impurities present in water.  Suitable for converting sea water into drinking water. Disadvantage:  The membrane must be strong enough to withstand high pressure without breaking. DID YOU KNOW? Chemicals can create monsters, foreign agents such as chemicals can enter a body of pregnant woman and can cause defects for the baby. The study of this subject is called teratology. Teratogens are substances that may cause birth defects. The word has been derived from the French word terat, meaning “monster”. Several environmental pollutants today have been proved to be the potential teratogens. CO, Cd, Pb, Hg are teratogenic to animals. 1.24 Engineering Chemistry QUESTIONS 1. What is meant by hardness of water? 2. Why Hardness is expressed in terms of Calcium carbonate? 3. What is temporary hardness and permanent hardness? 4. Identify the hardness causing salts among the following: NaHCO3, Ca(HCO3)2, Na2CO3, NaCl, MgCl2 and Al2(SO4)3. 5. What is Eriochrome Black-T? Write its Molecular structure 6. Identify the colours of the following chemical species Ca-EDTA, Mg-EDTA, Ca-EBT,EBT. 7. How is water softened by ion exchange method? Explain in detail with a neat diagram 8. What are zeolites? How do they function in removing the hardness of water? 9. What are cation and anion exchangers? Give one example? 10. A sample of water drawn by a jute mill near Calcutta has the following analytical data. Calculate the temporary and permanent hardness of the sample of water. Mg(HCO3)2 = 16.8 ppm, MgCl2 = 19ppm, Mg(NO3)2 = 29.6ppm , CaCO3 = 20 ppm, MgSO4 = 24 ppm, KOH = 0.9 ppm. 11. How hardness of water is determined by EDTA method? 12. What is reverse osmosis? How is it useful for desalination of brackish water? Explain with a diagram. 13. What is meant by desalination? 14. Enumerate the merits and demerits of reverse osmosis technique for desalination of water. 15. Define and give the importance of following: i) COD ii) Alkalinity iii) pH 16. Explain the significance of fluoride content in water 17. What is meant by exhaustion of zeolite? 18. What is semi-permeable membrane? 19. What is meant by exhaustion of cation and anion exchangers? How they are regenerated? Electrochemistry 2.1 UNIT - 2 ELECTROCHEMISTRY Conductance - Conductometric titration and its applications -estimation of strong acid, estimation of mixture of strong and weak acids (numerical based on conductance). Electrochemistry-redox reaction- types of Electrode-Ion selective electrodes – Glass electrodes-measurement of pH - potentiometry Green hydrogen-Energy storage systems for electric vehicles – Principle & Electrochemistry of a H2–O2 fuel cell, Li-ion battery, Na-ion battery. 2.1. INTRODUCTION Electrochemistry is a branch of chemistry that explores the relationship between electrical energy and chemical reactions. Electrochemical cell (also called Galvanic cell or Voltaic cell) deals with the conversion of chemical energy to electrical energy. While electrolytic cell deals with conversion of electrical energy into chemical energy. Fig. 2.1. Electrochemical cell Vs Electrolytic cell 2.2 Engineering Chemistry 2.2. APPLICATIONS OF ELECTROCHEMISTRY Electrochemistry has a vast array of applications that impact our daily lives and drive technological advancements: 1. Batteries and Fuel Cells: Electrochemical cells are at the heart of batteries and fuel cells. Lithium-ion batteries power a multitude of devices, from smartphones to electric vehicles. Fuel cells, which convert chemical energy from fuels like hydrogen into electricity, are being developed for clean energy applications. 2. Electroplating and Surface Coating: Electroplating involves coating a metal object with a thin layer of another metal using electrochemical processes. This technique is used to enhance the appearance, durability, and corrosion resistance of items ranging from jewellery to automotive parts. 3. Electrolysis: Electrolysis is used in various industrial processes, such as the extraction of metals from ores (e.g., aluminium production) and the electrolysis of water to produce hydrogen and oxygen gases, which have applications in energy storage and fuel. 4. Sensors and Analytical Devices: Electrochemical sensors detect and quantify chemical substances, playing a critical role in environmental monitoring, food safety, and medical diagnostics. For instance, glucose sensors are widely used by diabetic patients to monitor blood sugar levels. 5. Corrosion Prevention: Understanding electrochemical principles is essential for preventing and controlling corrosion, which is a major challenge in many industries, including construction, transportation, and marine engineering. Techniques such as cathodic protection and anodic inhibitors are employed to protect structures and materials from degradation. 6. Environmental Applications: Electrochemical methods are employed in water purification, waste treatment, and pollution control. Electrocoagulation and electrochemical oxidation are techniques used to remove contaminants from water and industrial effluents, contributing to environmental sustainability. Electrochemistry 2.3 Fig. 2.2. Applications of Electrochemistry DID YOU KNOW? Fig. 2.3. Electric Eel The electric eel (Electrophorus electricus), native to South America, can generate powerful electric shocks up to 600 volts. This ability comes from specialized organs that make up 80% of its body length, containing thousands of electrocytes (battery- like cells). 2.4 Engineering Chemistry How They Generate Electricity? Electric eels use electrocytes to create electricity. When signaled by the eel, these cells move ions across their membranes, generating a voltage difference. When many electrocytes discharge simultaneously, they produce a significant electric shock. 2.3. CONDUCTANCE Electrolytic conductance refers to the ability of an electrolyte solution to conduct electricity. It is a fundamental concept in electrochemistry and plays a crucial role in various industrial and scientific applications. What is Electrolytic Conductance? When an electrolyte (a substance that dissolves in water to produce ions) dissolves in water, it dissociates into its constituent ions. These ions are responsible for conducting electric current through the solution. Electrolytic conductance is measured by the ease with which these ions move under the influence of an electric field. The electrolytic conductance, G, of a medium is equal to the reciprocal of its electrical resistance (Resistance is the measure of how much the material opposes the flow of current) R in ohms: 1 1 R=G So, G = R (i) In a qualitative sense, the resistance tells us how difficult it is for an electrical current to pass. Since a solution is a three-dimensional conductor, the exact resistance will depend on the spacing l and area A of the electrodes (cell constant). The resistance of the solution in such situation is directly proportional to the distance between the electrodes and inversely proportional to the electrode surface area. If we consider the electrolytic cell with two electrodes having a cross sectional area of A [m2] and separated by l [m], then the resistance R of the electrolyte solution present between the electrodes is 𝑙 R=𝐴 (ii) Electrochemistry 2.5 where  is the proportionality constant called specific resistance and it is a characteristic property of material. Substituting the value of R from Eq. (ii ) in Eq. (i), the expression for the conductance, G , is 1𝐴 𝐴 G= 𝑙 = (iii) 𝑙 Where  is the reciprocal of specific resistance called specific conductance or conductivity. It is defined as the conductance of one cm3 of a solution of electrolyte. In the SI system, the unit of conductance is “Siemens”, (S) hence the unit for conductivity will be [S cm-1] (1S = 1  -1). Fig. 2.4. Specific conductance In order to compare quantitatively the conductivities of electrolytes, a quantity called molar conductance is frequently used. The molar conductance, m, is the conductivity per unit molar concentration of a dissolved electrolyte. It is connected with specific conductivity, , by the relation  m =𝑐 (iv) where c is the concentration in [mol m-3]. The molar conductivity is usually expressed in [S m2 mol-1] or [S cm2 mol-1]. Molarity = (moles/volume) ×1000 Since, we are working with 1 mole, Volume=1000/molarity Therefore, molar conductance =m = 1000  M Where, M is the number of moles of the electrolyte present in 1000 cc of electrolyte. 2.6 Engineering Chemistry Equivalent conductance is the conductance of a volume of solution containing one equivalent weight of dissolved substance. Normality = (gram equivalent/volume) × 1000 Since gram equivalent here is 1g Therefore, Volume = 1000/Normality So, equivalent conductance = e = 1000  N COMPARISON: Specific conductance Equivalent Molar conductance () conductance (e) (m) Conducting power of Conducting power of Conducting power of ions present in one cm3 ions produced by ions produced by of the solution dissolving one gram dissolving one gram equivalent of the mole of the electrolyte electrolyte Decreases with dilution Increases with dilution Increases with dilution Unit: Ohm-1cm-1 or Unit: Ohm-1cm2 equi-1 Unit: Ohm-1cm2 mol-1 or Sm-1 or S m2 equi-1 S m2 mol-1 2.4. EFFECT OF DILUTION ON CONDUCTANCE Specific conductance is the conductance of one cm3 of a solution of an electrolyte. If we dilute the solution the specific conductance decreases. On diluting the solution, the concentration of ions per cm3 decreases. Hence the specific conductance decreases. Equivalent or molar conductance and molar conductance increases with dilution. But it varies for strong electrolytes and weak electrolytes. For strong electrolytes (e.g., NaCl, HCl), which completely dissociate in solution, molar Electrochemistry 2.7 conductance increases with dilution due to reduced ion-ion interactions. Dilution decreases ion interactions, allowing ions to move more freely. For weak electrolytes (e.g., acetic acid), which partially dissociate, molar conductance increases more sharply with dilution due to increase in degree of dissociation. Graphical Representation  Strong Electrolytes: Linear increase in Λm Vs √c, showing moderate rise with dilution. This is because, with increased dilution, the number of ions per unit volume decreases, but the mobility of each ion increases due to less inter ionic attraction. This increase in mobility generally results in an increase in the overall conductance of the solution.  Weak Electrolytes: Steep, nonlinear increase, indicating significant rise in ionization. As the solution is diluted, the degree of dissociation of the weak electrolyte increases, leading to a more significant increase in the number of ions available for conduction. Fig. 2.5. Variation of molar conductance with dilution 2.8 Engineering Chemistry 2.5. FACTORS AFFECTING ELECTROLYTIC CONDUCTANCE (i) Nature of electrolyte The strong electrolytes undergo complete ionization and hence show higher conductivities since they furnish more number of ions. Example: HCl, H2SO4, HNO3, HBr, NaOH, KOH, NaCl, KCl etc. Whereas weak electrolytes undergo partial ionization and hence show comparatively low conductivities in their solutions. Examples: Acetic acid, Oxalic acid, Alkyl amines, Mercury(II) chloride, Lead acetate etc. (ii) Concentration of ions The sole reason for the conductivity of electrolytes is the ions present in them. The conductivity of electrolytes increases with increase in the concentration (number) of ions as there will be more charge carriers. As explained earlier for weak electrolytes and strong electrolytes, molar (or equivalent) conductance increases with dilution. (iii)Ionic size and mobility Ionic mobility decreases with increase in ionic size and extent of hydration. Hence, conductivity decreases. Example: In molten state, the conductivities of lithium salts are greater than those of cesium salts since the size of Li + ion is smaller than that of Cs+ ion. Fig. 2.6. Size and mobility of free ions However, in aqueous solutions the extent of hydration affects the mobility of the ion, which in turn affects the conductivity. Heavily hydrated ions show low conductance values due to larger size. Electrochemistry 2.9 Example: In aqueous solutions Li+ ion with high charge density is heavily hydrated than Cs+ ion with low charge density. Hence hydrated Li+ bigger than hydrated Cs+. As a result, lithium salts show lower conductivities compared to those of cesium salts in water. The ionic mobility is reduced in more viscous solvents. Hence, the conductivity decreases. Fig. 2.7. Size and mobility of hydrated ions (iv) Temperature The conductance of an electrolyte solution increases with increase in the temperature due to increase in the extent of ionization. Increasing the temperature of an electrolyte solution always increases the conductivity. The increase is significant, between 2-3 % per °C. With rise in temperature viscosity of the solvent decreases which makes ions to move freely to the electrode. 2.6. CONDUCTOMETRIC TITRATIONS Conductometric titrations are a type of titration where the conductance (or conductivity) of the solution is measured as a function of the volume of titrant added. PRINCIPLE Conductometric titration is based on the fact that the conductance of a solution at a constant temperature depends on the i) number of ions present in the solution ii) charge on the ions iii) mobility of ions. During the titration, one type of ions are replaced by the another kind of ions which differ in their mobilities. At the equivalence point, there is a sharp change in 2.10 Engineering Chemistry the conductance. The equivalence point can be located graphically by plotting the changes of conductance as a function of the volume of the added titrant. PROCEDURE A known volume of one of the reactant solutions is pipetted out into a beaker. Conductivity cell is dipped in this beaker and initial conductance is noted after stirring. Now, equal volumes of titrant (say 0.5 mL or 1 mL) is added to beaker, stirred well and conductance of the solution is measured after every addition of the titrant. Fig. 2.8. Conductometric titration set up Advantages:  Conductometric titration is simple and has wide selectivity.  Precise results with minimal errors  The conductometric titration is also suitable for weak acids and more dilute solutions.  Conductometric titration can be used to find out the endpoint of some samples such as coloured or turbid solutions  The temperature remains constant throughout the process. Disadvantages:  It is not possible to measure samples with high concentrations.  Changes in salt levels can increase the conductivity of the solution.  Only a limited number of redox titrations can be performed. Electrochemistry 2.11 DID YOU KNOW? H+ and OH- ions indeed have unusually high mobilities in aqueous solutions (water) compared to other ions. For H+ ions: Proton hopping: This hydronium ion can then transfer its proton very quickly to another water molecule due to the unique hydrogen bonding network in water. The proton essentially jumps between water molecules, effectively moving the H+ very fast. Fig. 2.9. H+ migration in water For OH- ions: OH⁻ ions can form when water dissociates into H⁺ and OH⁻. The hydroxide ion can then accept a proton from a neighbouring water molecule, creating a new OH⁻ ion and leaving behind a proton vacancy. This effectively allows the OH⁻ ion to move through the solution. Fig. 2.10. OH- migration in water 2.12 Engineering Chemistry 1. STRONG ACID WITH A STRONG BASE Let us consider a titration between strong acid (HCl) and strong base (NaOH). Here, a definite amount of HCl is pipetted into the beaker and conductivity cell is inserted. Initially high conductance is observed due to the complete ionization of strong acid (HCl) and high mobility of hydrogen ions. When the strong base (NaOH) is added as titrant, the conductance decreases due to the replacement of H+ ions by the added Na+ ions and replaced H+ ions react with OH- ions to form undissociated water. This decrease in the conductance given by line AB. The decrease continues till the equivalence point O is reached. After the equivalence point, addition of NaOH is continued, and the conductance increases due to the high conductivity of OH- ions given by line CD. H+ + Cl - + Na+ + OH - → Na+ + Cl- + H2O Fig. 2.11. Strong acid Vs Strong base – Conductometric graph Electrochemistry 2.13 2. MIXTURE OF A STRONG ACID (HNO3) AND A WEAK ACID (CH3COOH) Vs STRONG BASE In this type of titration, a mixture of strong acid and weak acid is pipetted out into the beaker. Conductivity cell is dipped, and conductance is measured. Now strong base is added from the burette as small aliquots (say 0.5 mL), stirred and conductance is noted. In the mixture, first strong acid reacts, after complete reaction of strong acid, weak acid begins to react. So, initial curve is similar to strong acid Vs strong base titration. First conductance decreases and reaches a minimum. The first break point (x) corresponds to the neutralization of strong acid. Once equivalence point for strong acid is reached, conductance increases slowly. The second break point (y) corresponds to the neutralization of weak acid and after that the conductance increases due to the excess of OH- ions in the case of a strong base as the titrant. H+ + Cl - + Na+ + OH - → Na+ + Cl- + H2O CH3COO-+ H+ + Na+ + OH - → CH3COO- +Na+ + H2O Fig. 2.12. Strong base Vs Strong acid+ Weak acid mixture – Conductometric graph 2.14 Engineering Chemistry ACTIVITY: Draw the conductance vs volume curve for (i) Strong acid Vs Weak base (ii) Weak acid Vs Strong base (iii) Weak acid Vs Weak base CHECK YOUR UNDERSTANDING: Consider conductometric titration between Na2SO4 (burette) and BaCl2. Precipitate BaSO4 is formed during the process. Predict the nature of the conductometric titration curve. 2.7. ELECTROCHEMICAL GLOSSARY Oxidation describes the loss of electrons. Zn(s)  Zn2+(aq) + 2 e- Reduction describes the gain of electrons Cu2+(aq)+ 2 e- Cu (s) Current is defined as flow of electricity which results from the ordered directional movement of electrically charged particles called electrons or ions. Electrode is a material which allows the electrons to pass through. It is used in the form of rod / bar / foil. Anode is the metal at which oxidation takes place. It is the electrode through which electrons leave the cell. It has a negative sign in electrochemical cell and positive sign in electrolytic cell. Cathode is the metal at which reduction takes place. It is the electrode through which electrons enter the cell. It has a positive sign in electrochemical cell and negative sign in electrolytic cell. Half-cell is a part of the cell containing electrode immersed in an electrolytic solution. If oxidation occurs at the electrode it is called oxidation half cell (anode Electrochemistry 2.15 compartment) and if reduction occurs it is called reduction half cell(cathode compartment). Redox reaction is the process in which oxidation and reduction occurs simultaneously. Oxidation-reduction reactions involve the transfer of electrons between substances. The gain of electrons is called reduction and the loss of electrons is called oxidation. The substance which loses electron is said to have undergone oxidation while the one which gains electron is said to have undergone reduction. Single Electrode Potential (E) The measure of tendency of an electrode to lose or gain electrons when it is contact with a solution of its own ion is called single electrode potential. Standard Electrode Potential (Eo) The measure of the tendency of the electrode to lose or gain electrons when it is dipped in its own salt solution of unit concentration (1M), at 25oC and 1atm is called standard single electrode potential. Oxidation potential (Eox) The measure of tendency of an electrode to lose electrons when it is contact with a solution of its own ion is called oxidation potential. Standard oxidation Potential (Eoox) The measure of the tendency of the electrodes loose electrons when it is dipped in its own salt solution of unit concentration (1M), at 25oC and 1atm is called standard oxidation potential. Reduction potential (Ered) The measure of tendency of an electrode to gain electrons when it is contact with a solution of its own ion is called reduction potential. Standard reduction Potential (Eored) The measure of the tendency of the electrode to gain electrons when it is dipped in its own salt solution of unit concentration (1M), at 25oC and 1atm is called standard reduction potential. 2.16 Engineering Chemistry Nernst equation for a single electrode For the reaction, M n+ (aq) + ne-  M (s) Nernst equation is given by Ered = Eored + 0.0591 log [M n+] n Nernst equation for a cell Nernst equation for a cell is given by Ecell= Eocell + 0.0591 log [R] n [P] 2.8. TYPES OF ELECTRODES Depending on the species involved in electron transfer reactions at the metal- electrolyte interface, the following types of electrodes are known i) Metal-metal ion electrode This process takes place within the very thin interfacial region at the electrode surface. It consists of a metal rod in contact with a solution of its own ion or ions (cations), with which the electrode is reversible. The potential of this electrode depends on the concentration of the metal ions in the solution and the temperature. Electrode representation (M/Mn+) Electrode reaction i) Zn(s) | Zn2+(aq) Zn → Zn2+ + 2e- ii) Cu(s) |Cu2+(aq) Cu → Cu2+ + 2e- ii) Redox or Ion-ion electrodes Many electrode reactions involve only ionic species, such as Fe2+ and Fe3. This type of electrodes consists of an inert metal like platinum in contact with an aqueous solution of the salt of an element in different oxidation state. The potential of this electrode depends on the concentration of the activities of the metal into in different oxidation state. Electrode representation Electrode reaction Fe3+aq| Fe2+ aq |Pt Fe3+(aq) + e- → Fe2+(aq) Electrochemistry 2.17 iii) Insoluble–salt or Metal- Metal sparingly salt electrodes This type of electrode consists of a metal in contact with one of the sparingly soluble salts and a solution containing a negative ion of the salt. The electrode involves a reversible reaction between the metal and the negative ion to form sparingly soluble salt with the liberation of electrons. A typical electrode of this kind consists of a silver wire covered with a thin coating of silver chloride, which is insoluble in water. The electrode reaction consists in the oxidation and reduction of the silver. Example: SCE and silver electrode Electrode representation (M|MX||X-) Electrode reaction Ag(s) | AgCl(s)| KCl AgCl(s) + e- → Ag(s) + Cl- Hg |Hg2Cl2(s)| KCl aq Hg2Cl2(s) + 2e- → 2Hg(l) + 2Cl- iv) Gas electrode: It consists of a gas bubbled around an inert metal like platinized platinum dipped in a solution containing ions to which the gas is reversible. The gas gets adsorbed on the inert metal and an equilibrium is established between the gas and its own ions in the solution. The potential of the gas electrode depends on the pressure of the gas, the concentration of the solution and the pressure Examples: Hydrogen gas electrode (SHE), Oxygen gas electrode, Chlorine gas electrode Electrode representation Electrode reaction + H │ H2(g) (1atm)│Pt H + + e- → ½ H2(g) Cl- │ Cl2(g)(1atm)│Pt ½ Cl2(g) + e- → Cl- 2.9. ION-SELECTIVE ELECTRODES An Ion selective electrode (ISE) is an indicator electrode that responds (produces a potential) when it is placed in a solution containing a certain ion. It responds to only one species of ions ignoring all other ions present in the solution. Potential developed at ion selective sensor is a measure of the concentration of the 2.18 Engineering Chemistry ionic species of interest. Now a large variety of ion-selective electrodes are available commercially which selectively respond to particular cations and anions, and certain gases. Most widely used among ion selective electrodes are pH electrodes. The main features of ISE are  Selectivity  Accuracy  Fast Response  Reproducibility of data.  Long Lifetime Construction: The three main components of making a measurement at an ISE are an inner reference, or standard solution an outer analyte, or sample, solution separated by a thin membrane. The potential developed at the membrane is the result of either an ion exchange process or an ion transport process occurring at each interface between the membrane and solution. 2.10. GLASS ELECTRODE Construction of glass electrode: 1. Glass Membrane: The core component of a glass electrode is a thin-walled glass bulb made of a special type of glass that is sensitive to hydrogen ions (H+). Special glass of low melting point and high electrical conductance is used for the purpose. 2. Internal Solution: Inside the glass bulb, there is an internal solution of known and constant pH, typically a dilute hydrochloric acid (HCl) solution. This internal solution ensures a constant internal potential. Electrochemistry 2.19 3. Internal Electrode: An internal reference electrode, usually a silver/silver chloride (Ag/AgCl) electrode, is immersed in the internal solution. The purpose of this electrode is to maintain a stable reference potential. 4. External Electrode and Junction: The glass electrode is usually paired with a reference electrode (Ag/AgCl electrode) that is placed in the solution being tested. This reference electrode provides a stable reference potential. Fig. 2.13. Measurement of pH using glass electrode Principle When a glass surface is kept in contact with a solution, a potential is established between the glass and the solution. The value of the potential is a function of H+ ion concentration of solution and the nature of the glass electrode. The magnitude of the potential difference at 25 C is given by 0 Eg = E - 0.0591 pH g 0 E determined using solution of known pH g 2.20 Engineering Chemistry When glass electrode is dipped into an analyte solution, there is a change in composition of the glass membrane due to ion exchange process involving solution and the membrane. [H+] + Na+Gl- → Na+ + H+Gl- (membrane) (membrane) A corresponding change in membrane potential, proportional to pH is measured. All other potentials are constant. In short, the membrane potential (variable) is measured against two standard external and internal Ag/AgCl reference electrodes. 2.11. ACID-BASE TITRATION-DETERMINATION OF pH During a chemical reaction, if a change in the concentration of anion (H+ in case of neutralization) causes a change in potential of a suitable electrode, then the progress of the reaction can be monitored by change in potential of electrodes. For an acid-base titration, hydrogen electrode or a pH-sensitive glass electrode may be used as the indicator electrode.  Reference electrode: Saturated Calomel Electrode (SCE)  Indicator electrode : Glass electrode Combined glass electrode which is commercially available has the combination of these two electrodes. Cell representation is Pt, H2 (1 atm), H+ (unknown concentration C) II KCl solution, Hg2Cl2(s), Hg Emf is measured by potentiometer which is connected to both electrodes while base is added slowly from the burette. Ecell = Ecathode – Eanode = ESCE –E glass = 0.2422 – (E0G - 0.0591 log [H+]) Ecell = 0.2422 + 0.0591 pH pH = Ecell – 0.2422 + E0G 0.0591V By knowing emf of the cell, we can find the pH of the solution. Electrochemistry 2.21 Procedure Pipette solution: unknown acid solution Burette solution: standard base Base is slowly added from burette. During the addition of base from the burette concentration of H+ keeps on decreasing (pH increasing). From the earlier equation, it is clear that Ecell increases with pH. A graph is plotted with electrode potential against volume of base added. The curve is steep near endpoint. More accurate end point can be determined by plotting pH/V Vs volume of base. Fig. 2.14. pH metry titration curves Advantages of glass electrode  It is very easy to construct and simple to operate.  The potential developed remains constant for long time.  It can be used with very small amount of the test solution.  It can be used even in the presence of oxidizing impurities, reducing impurities etc.  Results are accurate.  Equilibrium is easily achieved. 2.22 Engineering Chemistry Limitations of glass electrode  Very fragile and to be used with care  Can measure pH 0-10 only 2.12. POTENTIOMETRY Potentiometry is an electroanalytical technique in which amount of the substance in solution is determined by placing two electrodes in a test solution. A potential difference is set up between two electrodes by the addition of the titrant. Since this is an equilibrium measurement, therefore, the Nernst equation is applicable. A potentiometer is used to measure the emf of the cell and emf is measured no current passes through the circuit under equilibrium condition (i.e.), when neither the current is drawn from the cell nor it passes through the cell. Applications of potentiometry  Determination of concentration of analyte (redox active species)  Determination of concentration of H+ (pH)  Determination of concentration of one particular ion in presence of other ions Principle Potentiometric measurement system consists of two electrodes called reference electrode and indicator electrode, potentiometer and a solution of analyte.  Reference electrode is an electrode whose potential is accurately known and constant. Its potential does not change to whichever solution it is dipped. It should maintain a constant potential on passing small current and follows Nernst equation. Most widely used reference electrodes are Saturated Calomel Electrode (SCE) and Ag-AgCl electrode.  Indicator electrode is an electrode whose potential depends upon the concentration of the analyte ions. Ideally it must respond to changes in concentration of an analyte or group of analytes rapidly and reproducibly. During the course of the titration, the electrode potential of indicator electrode changes with the concentration of active ion in the analyte according to Nernst equation. A Electrochemistry 2.23 digital potentiometer is used to measure the potential difference between indicator and reference electrodes. Ecell= Eindicator − Ereference Fig. 2.15. Potentiometry electrodes Potentiometric titrations can be used for redox titrations, acid-base titrations, precipitation titrations etc. Redox titration Any redox reaction (reduction- oxidation reaction) can be studied using a potentiometer. The most common system is the ferrous- ferric system. Ferrous ion can be oxidized to Ferric ion by an oxidizing agent viz., KMnO4, K2Cr2O7 etc. The reaction can be represented as 2KMnO4+10FeSO4+8H2SO4  2MnSO4 + 5Fe2(SO4)3 + 8H2O+ K2SO4 K2Cr2O7 + 7H2SO4 + 6FeSO4 Cr2(SO4)3 + 3Fe2(SO4)3 +7H2O + K2SO4  Reference electrode: Saturated Calomel Electrode (SCE)  Indicator electrode: Pt Oxidation of Fe2+ to Fe3+ happens in the test solution and indicator electrode measures the change in potential The cell is represented as SCE // Pt / Fe2+ - Fe3+ Since the potential of the reference SCE remains fixed throughout the experiment, the variation of potential of the indicator electrode for a particular redox system could be described as, 2.24 Engineering Chemistry For this system, On reaching the end point, entire Fe2+ is converted to Fe3+, thus the denominator term decreases drastically, the potential increases drastically which is observed as a jump. Procedure Pipette solution : Ferrous salt solution Additional solution: dil.H2SO4 Burette solution : standard KMnO4 or K2Cr2O7 Fig. 2.16. Potentiometry set up 20 ml Fe2+ salt solution is pipetted out into a beaker. Equal volume of dil.H2SO4 is added and a working Pt electrode and a reference calomel electrode are Electrochemistry 2.25 placed into it in such a way that the Pt foil is completely immersed. The other end of the electrodes is connected to appropriate terminals of the potentiometer (Pt electrode to +ve terminal). Solution is stirred uniformly by a magnetic stirrer bit. The emf of the made up solution is measured before adding oxidizing agent. Oxidizing agent is added from a burette, aliquots of 1ml at a time. The solution is stirred and emf is noted. The titration is carried out till the end point is reached which is indicated by a sudden increase in the emf. Titration is continued for another 4-5 ml after the sudden rise. First graph is plotted volume of oxidizing agent vs emf. Second graph is drawn with volume of oxidizing agent vs. E/V. Volume corresponding to the peak of the second graph gives the end point. Fig. 2.17. Potentiometry curves Advantages of potentiometric titrations  Even coloured solutions can be used.  Accurate and fast results.  Can be conducted even in microscale. 2.26 Engineering Chemistry 2.13. GREEN HYDROGEN Green hydrogen refers to hydrogen produced through the process of electrolysis using renewable energy sources, which results in no greenhouse gas emissions. It is a key component in the transition to a sustainable and low-carbon energy system. Methods of Producing Green Hydrogen Green hydrogen is produced exclusively through electrolysis, a process that splits water into hydrogen and oxygen using electricity. However, the source of electricity is what differentiates green hydrogen from other types. Electrolysis Powered by Renewable Energy The core principle of green hydrogen production is the use of renewable energy sources to power the electrolysis process. Here are the primary methods: 1. Electrolysis with Solar Power: o Solar panels generate electricity from sunlight. This electricity is then used to power the electrolyzer, splitting water into hydrogen and oxygen. Fig. 2.18. Solar green hydrogen production Electrochemistry 2.27 2. Electrolysis with Wind Power: o Wind turbines convert wind energy into electricity. This electricity is fed into the electrolyzer to produce hydrogen. Fig. 2.19. Wind green hydrogen production 3. Electrolysis with Hydropower: o Hydropower plants harness the energy of flowing water to generate electricity. This electricity is used for electrolysis. Applications of Green Hydrogen as energy source 1. Transportation:  Hydrogen Fuel Cell Vehicles (FCVs): Companies like Toyota, Honda, and Hyundai have developed hydrogen fuel cell vehicles (e.g., Toyota Mirai, Honda Clarity, Hyundai Nexo) that emit only water vapor. These vehicles are being used in various regions with growing hydrogen infrastructure, like California, Japan, and parts of Europe.  Buses and Trucks: Green hydrogen is being used in public transport. For instance, the European Clean Hydrogen Alliance has been promoting hydrogen-powered buses across the EU. Similarly, companies like Nikola Motors are developing hydrogen-powered trucks. 2.28 Engineering Chemistry 2. Industrial Applications:  Steel Production: Companies like SSAB in Sweden are using green hydrogen in a project called HYBRIT to replace coal in steel production, aiming to produce fossil-free steel.  Chemical Industry: Green hydrogen is used to produce ammonia for fertilizers through a cleaner process. Projects like Yara in Norway are developing green ammonia plants using renewable energy sources. 3. Energy Storage and Grid Balancing:  Power-to-Gas: Green hydrogen can be produced from surplus renewable electricity (from wind or solar) and stored for later use. For example, Germany's "Energiepark Mainz" uses excess wind power to produce hydrogen, which can be injected into the natural gas grid or used as fuel.  Hydrogen-Powered Backup Generators: Data centers and remote facilities are exploring hydrogen fuel cells as backup power sources. For instance, Microsoft has experimented with hydrogen fuel cells as a backup power option for its data centers. 4. Maritime Applications:  Hydrogen-Powered Ships: The shipping industry is exploring green hydrogen as a fuel alternative to reduce emissions. Projects like the “HySeas III” in Scotland aim to develop the world's first sea-going vehicle and passenger ferry powered by hydrogen fuel cells. 5. Residential Heating:  Hydrogen Blending in Natural Gas Grids: Some countries are testing the blending of green hydrogen into existing natural gas grids to reduce carbon emissions from residential heating. The UK has projects like the HyDeploy project, which blends hydrogen into the natural gas supply for homes. Electrochemistry 2.29 DID YOU KNOW? Fig. 2.20. Schematics of green hydrogen mission Indian government took major initiative on the production of green hydrogen to the capacity of 5 million metric tons per annum by 2030, thanks to its favorable geographic location with the abundance of sunlight and wind energy. The mission leads to the reduction to the worth of 1 lakh crore of the fossil fuel imports and an abatement of annual greenhouse gas emissions (~ 50 MT). Advantages of Green Hydrogen  Zero Emissions: Produces no greenhouse gases during production if renewable energy is used, making it a clean energy source.  Energy Storage: Acts as a means to store excess renewable energy for use when renewable sources are not available.  Versatility: Can be used across various sectors, including transportation, industry, and power generation.  Energy Carrier: Can transport energy over long distances, useful for regions rich in renewable resources but far from demand centres. Challenges 1. Cost: Currently more expensive than hydrogen produced from fossil fuels (grey or blue hydrogen). 2. Infrastructure: Requires development of production, storage, transportation, and distribution infrastructure. 2.30 Engineering Chemistry 3. Efficiency: Overall efficiency of production, storage, and conversion processes needs improvement. 4. Materials: Dependence on expensive and rare materials for certain electrolysis technologies. 2.14. ENERGY STORAGE SYSTEMS FOR ELECTRIC VEHICLES Energy storage systems are critical for the performance, range, and overall viability of electric vehicles (EVs). Here’s an overview of the primary types of energy storage systems and key considerations: Types of Energy Storage Systems 1. Hydrogen Fuel Cells  Description: Convert hydrogen gas into electricity through a chemical reaction with oxygen.  Advantages: High energy density, fast refuelling, and no direct emissions (only water vapor).  Limitations: Infrastructure challenges, high production and storage costs for hydrogen, and lower efficiency compared to battery-electric systems. 2. Lithium-Ion Batteries:  Description: The most widely used battery type in EVs due to its high energy density and long cycle life.  Advantages: High efficiency, good power-to-weight ratio, relatively fast charging, and decreasing costs.  Limitations: Safety concerns such as thermal runaway, environmental impact of mining lithium and cobalt, and recycling challenges. 3. Sodium-Ion Batteries  Description: Sodium-ion (Na-ion) batteries are an emerging alternative to lithium-ion batteries, operating similarly but using sodium-ions as charge carriers. They are gaining attention due to the abundance and low cost of sodium. Electrochemistry 2.31  Advantages: Use of abundant and inexpensive materials, lower environmental impact compared to lithium-ion batteries, improved safety with a lower risk of thermal runaway, and potential for lower overall costs.  Limitations: Currently lower energy density than lithium-ion batteries, resulting in shorter driving ranges for EVs, technology is less mature with ongoing research needed to improve performance and cycle life, and potentially heavier batteries due to the larger size of sodium-ions. Features: 1. Energy Density: Determines the driving range of the EV. Higher energy density means longer range. 2. Power Density: Affects the vehicle's acceleration and the ability to handle regenerative braking. 3. Cycle Life: The number of charge and discharge cycles a battery can undergo before significant degradation. Longer cycle life means longer battery lifespan. 4. Charging Time: The time required to recharge the battery. Faster charging is crucial for the convenience and adoption of EVs. 5. Safety: Concerns such as the risk of thermal runaway, fire, or explosion. Safety features and stable chemistries are essential. 6. Cost: Includes manufacturing, raw materials, and recycling. Reducing costs is critical for making EVs affordable. 7. Environmental Impact: The impact of raw material extraction, battery manufacturing, and disposal. Sustainable practices and recycling are increasingly important. 2.15. FUEL CELLS Fuel cells are electrochemical cells in which chemical energy present in the fuel is converted into electrical energy without burning. Fuel is oxidized at the anode by electrochemical process. Fuel + oxygen  Oxidized products + electricity 2.32 Engineering Chemistry Like any other electrochemical cell, the fuel cell has two electrodes and electrolyte. Fuel and oxidant are continuously and separately supplied to the two electrodes of the fuel cell, where they undergo reactions. Fuel cell can supply current as long as they are supplied with reactants. A fuel cell essentially consists of the following arrangement. Fuel/electrode//electrolyte/electrode oxidant At anode electrochemical oxidation of the fuel happens Fuel  Oxidant product + ne- At cathode, the oxidant gets reduced Oxidant + ne-  Reduction products The electrons produced by the oxidation process at anode can perform useful work when they pass through the external circuit to the cathode. HYDROGEN-OXYGEN FUEL CELL Hydrogen-oxygen fuel cell is the simplest type of fuel cell in which hydrogen gas is used as the fuel and oxygen gas as oxidant. A schematic diagram of H2-O2 fuel cell is shown in the figure below. Fig. 2.21. H2-O2 fuel cell Electrochemistry 2.33 Construction: The cell consists of a porous carbon (graphite) electrodes (both anode and cathode) impregnated with catalysts such as finely divided platinum or palladium. Electrolyte is an aqueous solution of KOH (25-30%). Working: Hydrogen fuel is continuously supplied at anode and oxygen gas is supplied at the cathode. As hydrogen gas diffuses through the anode, it is adsorbed on the electrode surface, gets oxidized to H+ ions. It reacts with hydroxyl ions to form water. The electrons produced moves through the external circuit, reaches cathode where oxygen gas is passed. The adsorbed oxygen gas is reduced to hydroxyl ions. Thus, electrolyte is regenerated. The electrode reactions are summarized as below: Anode: 2H2 → 4H++ 4e- 4H++ 4OH- → 4 H2O ------------------------------------ 2H2 + 4 OH- → 4 H2O + 4e- Cathode: O2 + 4e- → 2 O2- 2O2- + 2 H2O → 4 OH- ----------------------------------- O2 + 2 H2O +4e- → 4 OH- Overall reaction: 2H2 + O2 → 2 H2O A typical fuel cell produces a voltage from 0.6 V to 0.7 V 2.34 Engineering Chemistry Advantages of fuel cell:  High efficiency (50-80% efficiency)  Pollution free  High power density  Quiet operation  No moving parts, so elimination of wear and tear Disadvantages of fuel cell:  Fuel cells were large and extremely expensive to manufacture  Energy cannot be stored  Power output is moderate  Fuels in the form of gases and oxygen need to be stored in tanks under high pressure. 2.16. LITHIUM-ION BATTERIES Lithium-ion batteries are a type of rechargeable battery powering everything from smartphones and laptops to electric vehicles and grid-scale energy storage systems. Their high energy density, long lifespan, and lightweight design have made them the preferred choice for a vast array of applications. How Do They Work? A lithium-ion battery consists of several components:  Anode: Typically made of graphite, it stores lithium-ions during charging.  Cathode: Composed of materials like lithium cobalt oxide or lithium iron phosphate, it receives lithium-ions during charging.  Separator: A porous membrane that prevents the anode and cathode from touching but allows lithium-ions to pass through.  Electrolyte: A liquid or gel-like substance that conducts ions between the anode and cathode. Electrochemistry 2.35 Fig. 2.22. Lithium ion battery  Charging: LiCoO2+C6→LiC6+CoO2 During charging, lithium-ions move from the cathode (LiCoO₂) to the anode (graphite,C6) through the electrolyte, while electrons flow through the external circuit to balance the charge.  Discharging: : LiC6+CoO2→C6+LiCoO2 During discharging, lithium-ions move back from the anode to the cathode, releasing energy that can be used to power electronic devices. Advantages  High Energy Density: Lithium-ion batteries can store a large amount of energy in a compact and lightweight package, making them ideal for portable applications.  Long Cycle Life: They can be recharged many times without significant loss of capacity, providing a long operational lifespan. 2.36 Engineering Chemistry  Low Self-Discharge: Lithium-ion batteries have a low rate of self-discharge compared to other rechargeable batteries, retaining their charge for longer periods.  Fast Charging: Advanced lithium-ion batteries support rapid charging, reducing downtime for users. Applications  Consumer Electronics: Widely used in smartphones, laptops, tablets, and wearable devices due to their high energy density and compact size.  Electric Vehicles (EVs): Powering electric cars, buses, and bikes, contributing to the reduction of greenhouse gas

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