Basic Concepts in Electrochemistry
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History of the Chlor-Alkali Industry
2 History of the Chlor-Alkali Industry During the last half of the 19th century, chlorine, used almost exclusively in the textile and paper industry, was made [1] by reacting manganese dioxide with hydrochloric acid 100–110◦C MnO2 + 4HCl −−−−−−→ MnCl2 + Cl2 + 2H2O (1) Recycling of manganese improved the overall process economics, and the process became known as the Weldon process [2]. In the 1860s, the Deacon process, which generated chlorine by direct catalytic oxidation of hydrochloric acid with air according to Eq. (2) was developed [3]. ◦ 450–460 C;CuCl2 cat. 4HCl + O2(air) −−−−−−−−−−−−−−→ 2Cl2 + 2H2O(2) The HCl required for reactions (1) and (2) was available from the manufacture of soda ash by the LeBlanc process [4,5]. H2SO4 + 2NaCl → Na2SO4 + 2HCl (3) Na2SO4 + CaCO3 + 2C → Na2CO3 + CaS + 2CO2 (4) Utilization of HCl from reaction (3) eliminated the major water and air pollution problems of the LeBlanc process and allowed the generation of chlorine. By 1900, the Weldon and Deacon processes generated enough chlorine for the production of about 150,000 tons per year of bleaching powder in England alone [6]. An important discovery during this period was the fact that steel is immune to attack by dry chlorine [7]. This permitted the first commercial production and distribu- tion of dry liquid chlorine by Badische Anilin-und-Soda Fabrik (BASF) of Germany in 1888 [8,9]. This technology, using H2SO4 for drying followed by compression of the gas and condensation by cooling, is much the same as is currently practiced. 17 “chap02” — 2005/5/2 — 09Brie:49 — page 17 — #1 18 CHAPTER 2 In the latter part of the 19th century, the Solvay process for caustic soda began to replace the LeBlanc process. -
Nernst Equation in Electrochemistry the Nernst Equation Gives the Reduction Potential of a Half‐Cell in Equilibrium
Nernst equation In electrochemistry the Nernst equation gives the reduction potential of a half‐cell in equilibrium. In further cases, it can also be used to determine the emf (electromotive force) for a full electrochemical cell. (half‐cell reduction potential) (total cell potential) where Ered is the half‐cell reduction potential at a certain T o E red is the standard half‐cell reduction potential Ecell is the cell potential (electromotive force) o E cell is the standard cell potential at a certain T R is the universal gas constant: R = 8.314472(15) JK−1mol−1 T is the absolute temperature in Kelvin a is the chemical activity for the relevant species, where aRed is the reductant and aOx is the oxidant F is the Faraday constant; F = 9.64853399(24)×104 Cmol−1 z is the number of electrons transferred in the cell reaction or half‐reaction Q is the reaction quotient (e.g. molar concentrations, partial pressures …) As the system is considered not to have reached equilibrium, the reaction quotient Q is used instead of the equilibrium constant k. The electrochemical series is used to determine the electrochemical potential or the electrode potential of an electrochemical cell. These electrode potentials are measured relatively to the standard hydrogen electrode. A reduced member of a couple has a thermodynamic tendency to reduce the oxidized member of any couple that lies above it in the series. The standard hydrogen electrode is a redox electrode which forms the basis of the thermodynamic scale of these oxidation‐ reduction potentials. For a comparison with all other electrode reactions, standard electrode potential E0 of hydrogen is defined to be zero at all temperatures. -
Electrolytic Cells
CHEMISTRY LEVEL 4C (CHM415115) ELECTROLYTIC CELLS THEORY SUMMARY & REVISION QUESTIONS (CRITERION 5) ©JAK DENNY Tasmanian TCE Chemistry Revision Guides by Jak Denny are licensed under a Creative Commons Attribution-NonCommercial-NoDerivatives 4.0 International License. INDEX: PAGES • INTRODUCTORY THEORY 3 • COMPARING ENERGY CONVERSIONS 4 • APPLICATIONS OF ELECTROLYSIS 4 • ELECTROPLATING 5 • COMPARING CELLS 6 • THE ELECTROCHEMICAL SERIES 7 • PREDICTING ELECTROLYSIS PRODUCTS 8-9 • ELECTROLYSIS PREDICTION FLOWCHART 1 0 • ELECTROLYSIS PREDICTION EXAMPLES 11 • ELECTROLYSIS PREDICTION QUESTIONS 12 • INDUSTRIAL ELECTROLYTIC PROCESSES 13 • IMPORTANT ELECTRICAL THEORY 14 • FARADAY’S ELECTROLYSIS LAWS 15 • QUANTITATIVE ELECTROLYSIS 16 • CELLS IN SERIES 17 • ELECTROLYSIS QUESTIONS 18-20 • ELECTROLYSIS TEST QUESTIONS 21-22 • TEST ANSWERS 23 2 ©JAK CHEMISTRY LEVEL 4C (CHM415115) ELECTROLYTIC CELLS (CRITERION 5) INTRODUCTION: In our recent investigation of electrochemical cells, we encountered spontaneous redox reactions that release electrical energy such as takes place in the familiar situations of “batteries”. When a car battery is ‘flat’ and needs to be recharged, a power supply (‘battery charger’) is connected to the flat battery and chemical changes take place and it is subsequently able to operate again as a power supply. The ‘recharging’ process is non-spontaneous and requires an input of energy to occur. When a cell is such that an energy input is required to make a non-spontaneous redox reaction take place, we describe the cell as an ELECTROLYTIC CELL. The redox process occurring in an ELECTROLYTIC CELL is referred to as ELECTROLYSIS. For example, consider the spontaneous redox reaction associated with an electrochemical (fuel) cell; i.e. 2H2(g) + O2(g) → 2H2O(l) + ELECTRICAL ENERGY This chemical reaction RELEASES energy which can be used to power an electric motor, to drive a machine, appliance or car. -
How Do We Learn Electrochemistry? by Jeffrey W
redcat_ad_IF_Sp2012_1.pdf 1 4/11/2012 1:36:17 PM research • news • events • resources | search • explore • connect • share • discover How Do We Learn Electrochemistry? by Jeffrey W. Fergus he importance of electrochemistry is undeniable—we on education. The fall 2006 issue was devoted to education literally cannot live without electrochemistry for proper and included articles discussing general needs for education in Tcell function and transmission of signals through the electrochemistry as well as some examples of approaches, and nervous system. Electrochemistry is also vital in a wide range even specific laboratory activities, to enhance electrochemical of important technological applications. For example, batteries education. More recently, in the summer 2010 issue, the ECS are important not only in storing energy for mobile devices Industrial Electrochemistry and Electrochemical Engineering and vehicles, but also for load leveling to enable the use of Division provided an additional evaluation of needs and some renewable energy conversion technologies. Electrochemistry activities for introducing electrochemistry into courses. The is involved in the production of materials by electrorefining current issue complements those earlier issues and provides or electrodeposition as well as the destruction of materials by insights on the status and needs in electrochemical education. TM corrosion. In spite of its ubiquity there are very few formal The first article provides a general backdrop on the state educational degree programs in electrochemistry. of higher education. Marye Ann Fox discusses the financial If electrochemistry is ubiquitous, but formal educational challenges being faced by academic institutions and how these programs are rare, how do the scientists and engineers working on challenges have an impact on the design and implementation of electrochemical products and processes learn the electrochemistry educational programs. -
Electrochemical Real-Time Mass Spectrometry: a Novel Tool for Time-Resolved Characterization of the Products of Electrochemical Reactions
Electrochemical real-time mass spectrometry: A novel tool for time-resolved characterization of the products of electrochemical reactions Elektrochemische Realzeit-Massenspektrometrie: Eine neuartige Methode zur zeitaufgelösten Charakterisierung der Produkte elektrochemischer Reaktionen Der Technischen Fakultät der Friedrich-Alexander-Universität Erlangen-Nürnberg zur Erlangung des Doktorgrades Dr.-Ingenieur vorgelegt von Peyman Khanipour Mehrin aus Shiraz, Iran Als Dissertation genehmigt von der Technischen Fakultät der Friedrich-Alexander-Universität Erlangen-Nürnberg Tag der mündlichen Prüfung: 17.11.2020 Vorsitzender des Promotionsorgans: Prof. Dr.-Ing. habil. Andreas Paul Fröba Gutachter: Prof. Dr. Karl J.J. Mayrhofer Prof. Dr. Frank-Michael Matysik I Acknowledgements This study is done in the electrosynthesis team of the electrocatalysis unit at Helmholtz- Institut Erlangen-Nürnberg (HI ERN) with the financial support of Forschungszentrum Jülich. I would like to express my deep gratitude to Prof. Dr. Karl J. J. Mayrhofer for accepting me as a Ph.D. student and also for all his encouragement, supports, and freedoms during my study. I’m grateful to Prof. Dr. Frank-Michael Matysik for kindly accepting to act as a second reviewer and also for the time he has invested in reading this thesis. This piece of work is enabled by collaboration with scientists from different expertise. I would like to express my appreciation to Dr. Sandra Haschke from FAU for providing shape-controlled high surface area platinum electrodes which I used for performing oxidation of primary alcohols and also the characterization of the provided material SEM, EDX, and XRD. Mr. Mario Löffler from HI ERN for obtaining the XPS data and his remarkable knowledge with the interpretation of the spectra on copper-based electrodes for the CO 2 electroreduction reaction. -
Electrochemical Cells
Electrochemical cells = electronic conductor If two different + surrounding electrolytes are used: electrolyte electrode compartment Galvanic cell: electrochemical cell in which electricity is produced as a result of a spontaneous reaction (e.g., batteries, fuel cells, electric fish!) Electrolytic cell: electrochemical cell in which a non-spontaneous reaction is driven by an external source of current Nils Walter: Chem 260 Reactions at electrodes: Half-reactions Redox reactions: Reactions in which electrons are transferred from one species to another +II -II 00+IV -II → E.g., CuS(s) + O2(g) Cu(s) + SO2(g) reduced oxidized Any redox reactions can be expressed as the difference between two reduction half-reactions in which e- are taken up Reduction of Cu2+: Cu2+(aq) + 2e- → Cu(s) Reduction of Zn2+: Zn2+(aq) + 2e- → Zn(s) Difference: Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) - + - → 2+ More complex: MnO4 (aq) + 8H + 5e Mn (aq) + 4H2O(l) Half-reactions are only a formal way of writing a redox reaction Nils Walter: Chem 260 Carrying the concept further Reduction of Cu2+: Cu2+(aq) + 2e- → Cu(s) In general: redox couple Ox/Red, half-reaction Ox + νe- → Red Any reaction can be expressed in redox half-reactions: + - → 2 H (aq) + 2e H2(g, pf) + - → 2 H (aq) + 2e H2(g, pi) → Expansion of gas: H2(g, pi) H2(g, pf) AgCl(s) + e- → Ag(s) + Cl-(aq) Ag+(aq) + e- → Ag(s) Dissolution of a sparingly soluble salt: AgCl(s) → Ag+(aq) + Cl-(aq) − 1 1 Reaction quotients: Q = a − ≈ [Cl ] Q = ≈ Cl + a + [Ag ] Ag Nils Walter: Chem 260 Reactions at electrodes Galvanic cell: -
Gas-Phase Electrochemistry: Measuring Absolute Potentials and Investigating Ion and Electron Hydration*
Pure Appl. Chem., Vol. 83, No. 12, pp. 2129–2151, 2011. doi:10.1351/PAC-CON-11-08-15 © 2011 IUPAC, Publication date (Web): 7 October 2011 Gas-phase electrochemistry: Measuring absolute potentials and investigating ion and electron hydration* William A. Donald1,‡ and Evan R. Williams2 1School of Chemistry, Bio21 Institute of Molecular Science and Biotechnology, and ARC Centre of Excellence for Free Radical Chemistry and Biotechnology, University of Melbourne, Melbourne, Victoria, Australia; 2Department of Chemistry, University of California, Berkeley, CA, USA Abstract: In solution, half-cell potentials and ion solvation energies (or enthalpies) are meas- ured relative to other values, thus establishing ladders of thermochemical values that are ref- erenced to the potential of the standard hydrogen electrode (SHE) and the proton hydration energy (or enthalpy), respectively, which are both arbitrarily assigned a value of 0. In this focused review article, we describe three routes for obtaining absolute solution-phase half- cell potentials using ion nanocalorimetry, in which the energy resulting from electron capture (EC) by large hydrated ions in the gas phase are obtained from the number of water mole- cules lost from the reduced precursor cluster, which was developed by the Williams group at the University of California, Berkeley. Recent ion nanocalorimetry methods for investigating ion and electron hydration and for obtaining the absolute hydration enthalpy of the electron are discussed. From these methods, an absolute electrochemical scale and ion solvation scale can be established from experimental measurements without any models. Keywords: absolute potentials; clusters; electrochemistry; electron capture dissociation; elec- tron transfer; gaseous state; hydrated ions; hydration; ion nanocalorimetry; solvation; thermo chemistry. -
Towards the Hydrogen Economy—A Review of the Parameters That Influence the Efficiency of Alkaline Water Electrolyzers
energies Review Towards the Hydrogen Economy—A Review of the Parameters That Influence the Efficiency of Alkaline Water Electrolyzers Ana L. Santos 1,2, Maria-João Cebola 3,4,5 and Diogo M. F. Santos 2,* 1 TecnoVeritas—Serviços de Engenharia e Sistemas Tecnológicos, Lda, 2640-486 Mafra, Portugal; [email protected] 2 Center of Physics and Engineering of Advanced Materials (CeFEMA), Instituto Superior Técnico, Universidade de Lisboa, 1049-001 Lisbon, Portugal 3 CBIOS—Center for Research in Biosciences & Health Technologies, Universidade Lusófona de Humanidades e Tecnologias, Campo Grande 376, 1749-024 Lisbon, Portugal; [email protected] 4 CERENA—Centre for Natural Resources and the Environment, Instituto Superior Técnico, Universidade de Lisboa, Av. Rovisco Pais, 1049-001 Lisbon, Portugal 5 Escola Superior Náutica Infante D. Henrique, 2770-058 Paço de Arcos, Portugal * Correspondence: [email protected] Abstract: Environmental issues make the quest for better and cleaner energy sources a priority. Worldwide, researchers and companies are continuously working on this matter, taking one of two approaches: either finding new energy sources or improving the efficiency of existing ones. Hydrogen is a well-known energy carrier due to its high energy content, but a somewhat elusive one for being a gas with low molecular weight. This review examines the current electrolysis processes for obtaining hydrogen, with an emphasis on alkaline water electrolysis. This process is far from being new, but research shows that there is still plenty of room for improvement. The efficiency of an electrolyzer mainly relates to the overpotential and resistances in the cell. This work shows that the path to better Citation: Santos, A.L.; Cebola, M.-J.; electrolyzer efficiency is through the optimization of the cell components and operating conditions. -
Electrode Reactions and Kinetics
CHEM465/865, 2004-3, Lecture 14-16, 20 th Oct., 2004 Electrode Reactions and Kinetics From equilibrium to non-equilibrium : We have a cell, two individual electrode compartments, given composition, molar reaction Gibbs free ∆∆∆ energy as the driving force, rG , determining EMF Ecell (open circuit potential) between anode and cathode! In other words: everything is ready to let the current flow! – Make the external connection through wires. Current is NOT determined by equilibrium thermodynamics (i.e. by the composition of the reactant mixture, i.e. the reaction quotient Q), but it is controlled by resistances, activation barriers, etc Equilibrium is a limiting case – any kinetic model must give the correct equilibrium expressions). All the involved phenomena are generically termed ELECTROCHEMICAL KINETICS ! This involves: Bulk diffusion Ion migration (Ohmic resistance) Adsorption Charge transfer Desorption Let’s focus on charge transfer ! Also called: Faradaic process . The only reaction directly affected by potential! An electrode reaction differs from ordinary chemical reactions in that at least one partial reaction must be a charge transfer reaction – against potential-controlled activation energy, from one phase to another, across the electrical double layer. The reaction rate depends on the distributions of species (concentrations and pressures), temperature and electrode potential E. Assumption used in the following: the electrode material itself (metal) is inert, i.e. a catalyst. It is not undergoing any chemical transformation. It is only a container of electrons. The general question on which we will focus in the following is: How does reaction rate depend on potential? The electrode potential E of an electrode through which a current flows differs from the equilibrium potential Eeq established when no current flows. -
3 PRACTICAL APPLICATION BATTERIES and ELECTROLYSIS Dr
ELECTROCHEMISTRY – 3 PRACTICAL APPLICATION BATTERIES AND ELECTROLYSIS Dr. Sapna Gupta ELECTROCHEMICAL CELLS An electrochemical cell is a system consisting of electrodes that dip into an electrolyte and in which a chemical reaction either uses or generates an electric current. A voltaic or galvanic cell is an electrochemical cell in which a spontaneous reaction generates an electric current. An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction. Dr. Sapna Gupta/Electrochemistry - Applications 2 GALVANIC CELLS • Galvanic cell - the experimental apparatus for generating electricity through the use of a spontaneous reaction • Electrodes • Anode (oxidation) • Cathode (reduction) • Half-cell - combination of container, electrode and solution • Salt bridge - conducting medium through which the cations and anions can move from one half-cell to the other. • Ion migration • Cations – migrate toward the cathode • Anions – migrate toward the anode • Cell potential (Ecell) – difference in electrical potential between the anode and cathode • Concentration dependent • Temperature dependent • Determined by nature of reactants Dr. Sapna Gupta/Electrochemistry - Applications 3 BATTERIES • A battery is a galvanic cell, or a series of cells connected that can be used to deliver a self-contained source of direct electric current. • Dry Cells and Alkaline Batteries • no fluid components • Zn container in contact with MnO2 and an electrolyte Dr. Sapna Gupta/Electrochemistry - Applications 4 ALKALINE CELL • Common watch batteries − − Anode: Zn(s) + 2OH (aq) Zn(OH)2(s) + 2e − − Cathode: 2MnO2(s) + H2O(l) + 2e Mn2O3(s) + 2OH (aq) This cell performs better under current drain and in cold weather. It isn’t truly “dry” but rather uses an aqueous paste. -
Elements of Electrochemistry
Page 1 of 8 Chem 201 Winter 2006 ELEM ENTS OF ELEC TROCHEMIS TRY I. Introduction A. A number of analytical techniques are based upon oxidation-reduction reactions. B. Examples of these techniques would include: 1. Determinations of Keq and oxidation-reduction midpoint potentials. 2. Determination of analytes by oxidation-reductions titrations. 3. Ion-specific electrodes (e.g., pH electrodes, etc.) 4. Gas-sensing probes. 5. Electrogravimetric analysis: oxidizing or reducing analytes to a known product and weighing the amount produced 6. Coulometric analysis: measuring the quantity of electrons required to reduce/oxidize an analyte II. Terminology A. Reduction: the gaining of electrons B. Oxidation: the loss of electrons C. Reducing agent (reductant): species that donates electrons to reduce another reagent. (The reducing agent get oxidized.) D. Oxidizing agent (oxidant): species that accepts electrons to oxidize another species. (The oxidizing agent gets reduced.) E. Oxidation-reduction reaction (redox reaction): a reaction in which electrons are transferred from one reactant to another. 1. For example, the reduction of cerium(IV) by iron(II): Ce4+ + Fe2+ ! Ce3+ + Fe3+ a. The reduction half-reaction is given by: Ce4+ + e- ! Ce3+ b. The oxidation half-reaction is given by: Fe2+ ! e- + Fe3+ 2. The half-reactions are the overall reaction broken down into oxidation and reduction steps. 3. Half-reactions cannot occur independently, but are used conceptually to simplify understanding and balancing the equations. III. Rules for Balancing Oxidation-Reduction Reactions A. Write out half-reaction "skeletons." Page 2 of 8 Chem 201 Winter 2006 + - B. Balance the half-reactions by adding H , OH or H2O as needed, maintaining electrical neutrality. -
Chapter 10 – Chemical Reactions Notes
Chapter 8 – Chemical Reactions Notes Chemical Reactions: Chemical reactions are processes in which the atoms of one or more substances are rearranged to form different chemical compounds. How to tell if a chemical reaction has occurred (recap): Temperature changes that can’t be accounted for. o Exothermic reactions give off energy (as in fire). o Endothermic reactions absorb energy (as in a cold pack). Spontaneous color change. o This happens when things rust, when they rot, and when they burn. Appearance of a solid when two liquids are mixed. o This solid is called a precipitate. Formation of a gas / bubbling, as when vinegar and baking soda are mixed. Overall, the most important thing to remember is that a chemical reaction produces a whole new chemical compound. Just changing the way that something looks (breaking, melting, dissolving, etc) isn’t enough to qualify something as a chemical reaction! Balancing Equations Notes: Things to keep in mind when looking at the recipes for chemical reactions: 1) The stuff before the arrow is referred to as the “reactants” or “reagents”, and the stuff after the arrow is called the “products.” 2) The number of atoms of each element is the same on both sides of the arrow. Even though there may be different numbers of molecules, the number of atoms of each element needs to remain the same to obey the law of conservation of mass. 3) The numbers in front of the formulas tell you how many molecules or moles of each chemical are involved in the reaction. 4) Equations are nothing more than chemical recipes.