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Stanley A Meyer Legacy Back up Secret Docs Save all Protect Spread print and give to schools NEVER STOP!!!!!!! Join Support here https://www.patreon.com/securesupplies/shop
Stanley A Meyer Legacy Back up Secret Docs Save all Protect Spread print and give to schools NEVER STOP!!!!!!! Join Support here https://www.patreon.com/securesupplies/shop
This document provides information on oxidation-reduction (redox) reactions and electrochemistry:
[1] Redox reactions involve the transfer of electrons between oxidizing and reducing agents. Common examples are corrosion reactions.
[2] Galvanic (voltaic) cells generate electricity through spontaneous redox reactions. The anode is where oxidation occurs and electrons are released. The cathode is where reduction occurs and electrons are gained.
[3] Cell potential depends on the relative tendencies of substances to be oxidized or reduced, as measured by standard reduction potentials. More negative potentials indicate greater reducing ability; more positive potentials indicate greater oxidizing ability.
This document discusses galvanic cells and cell potential. It begins by defining oxidation-reduction reactions and half-reactions. It then explains how galvanic cells use redox reactions to produce an electric current and discusses the components of galvanic cells including the salt bridge, electrodes, and direction of electron and ion flow. The document introduces standard reduction potentials and how to calculate cell potential from half-cell potentials. It explains how cell potential depends on concentration using Le Châtelier's principle and the Nernst equation. Examples are provided to demonstrate how to calculate cell potentials, determine reaction spontaneity, and predict changes in potential with changing concentrations.
The document provides information about lead-acid batteries, including:
1. It describes the basic construction of lead-acid batteries using lead and lead dioxide electrodes separated by a sulfuric acid electrolyte.
2. It explains the electrochemical reactions that occur during charging and discharging, where electrons are transferred between the electrodes through the external circuit.
3. It discusses factors that influence the battery voltage such as the electrolyte concentration, state of charge, and internal resistances, as well as models for predicting voltage.
This document discusses electrolysis and Faraday's law of electrolysis. It provides examples of predicting products of electrolysis for molten salts, aqueous salt solutions, and applying Faraday's law calculations. Key points include:
- During electrolysis, the cation is reduced at the cathode and the anion is oxidized at the anode
- In molten salts, the more easily oxidized/reduced species reacts at each electrode
- In aqueous solutions, overvoltage must be considered in addition to electrode potentials
- Faraday's law states the amount of substance reacted is directly proportional to the quantity of electricity passed through the cell
- Calculations can determine current, time, charge or mass from the other variables using Faraday's constant
The document discusses different types of batteries, including primary batteries that are not rechargeable and secondary batteries that are rechargeable. It provides examples of the lead acid battery, a secondary battery, and dry cells, a primary battery. It describes the basic components of cells, including electrodes, electrolyte, and separators. It also covers how batteries can be connected in series or parallel configurations and the chemical reactions that occur during charging and discharging of batteries. The document discusses new types of flexible paper batteries and applications of lithium batteries in wearable electronics and other innovative devices.
The document discusses different types of batteries, including primary batteries that are not rechargeable and secondary batteries that are rechargeable. It provides examples of the lead acid battery, a secondary battery, and dry cells, a primary battery. It describes the basic components of cells, including electrodes, electrolyte, and separators. It also covers how batteries can be connected in series or parallel configurations and the chemical reactions that occur during charging and discharging of batteries. The document discusses new types of flexible paper batteries and applications of lithium batteries in wearable electronics and other innovative devices.
Stanley A Meyer Legacy Back up Secret Docs Save all Protect Spread print and give to schools NEVER STOP!!!!!!! Join Support here https://www.patreon.com/securesupplies/shop
Stanley A Meyer Legacy Back up Secret Docs Save all Protect Spread print and give to schools NEVER STOP!!!!!!! Join Support here https://www.patreon.com/securesupplies/shop
This document provides information on oxidation-reduction (redox) reactions and electrochemistry:
[1] Redox reactions involve the transfer of electrons between oxidizing and reducing agents. Common examples are corrosion reactions.
[2] Galvanic (voltaic) cells generate electricity through spontaneous redox reactions. The anode is where oxidation occurs and electrons are released. The cathode is where reduction occurs and electrons are gained.
[3] Cell potential depends on the relative tendencies of substances to be oxidized or reduced, as measured by standard reduction potentials. More negative potentials indicate greater reducing ability; more positive potentials indicate greater oxidizing ability.
This document discusses galvanic cells and cell potential. It begins by defining oxidation-reduction reactions and half-reactions. It then explains how galvanic cells use redox reactions to produce an electric current and discusses the components of galvanic cells including the salt bridge, electrodes, and direction of electron and ion flow. The document introduces standard reduction potentials and how to calculate cell potential from half-cell potentials. It explains how cell potential depends on concentration using Le Châtelier's principle and the Nernst equation. Examples are provided to demonstrate how to calculate cell potentials, determine reaction spontaneity, and predict changes in potential with changing concentrations.
The document provides information about lead-acid batteries, including:
1. It describes the basic construction of lead-acid batteries using lead and lead dioxide electrodes separated by a sulfuric acid electrolyte.
2. It explains the electrochemical reactions that occur during charging and discharging, where electrons are transferred between the electrodes through the external circuit.
3. It discusses factors that influence the battery voltage such as the electrolyte concentration, state of charge, and internal resistances, as well as models for predicting voltage.
This document discusses electrolysis and Faraday's law of electrolysis. It provides examples of predicting products of electrolysis for molten salts, aqueous salt solutions, and applying Faraday's law calculations. Key points include:
- During electrolysis, the cation is reduced at the cathode and the anion is oxidized at the anode
- In molten salts, the more easily oxidized/reduced species reacts at each electrode
- In aqueous solutions, overvoltage must be considered in addition to electrode potentials
- Faraday's law states the amount of substance reacted is directly proportional to the quantity of electricity passed through the cell
- Calculations can determine current, time, charge or mass from the other variables using Faraday's constant
The document discusses different types of batteries, including primary batteries that are not rechargeable and secondary batteries that are rechargeable. It provides examples of the lead acid battery, a secondary battery, and dry cells, a primary battery. It describes the basic components of cells, including electrodes, electrolyte, and separators. It also covers how batteries can be connected in series or parallel configurations and the chemical reactions that occur during charging and discharging of batteries. The document discusses new types of flexible paper batteries and applications of lithium batteries in wearable electronics and other innovative devices.
The document discusses different types of batteries, including primary batteries that are not rechargeable and secondary batteries that are rechargeable. It provides examples of the lead acid battery, a secondary battery, and dry cells, a primary battery. It describes the basic components of cells, including electrodes, electrolyte, and separators. It also covers how batteries can be connected in series or parallel configurations and the chemical reactions that occur during charging and discharging of batteries. The document discusses new types of flexible paper batteries and applications of lithium batteries in wearable electronics and other innovative devices.
The document summarizes key concepts about lead-acid batteries, including:
1) Lead-acid batteries use lead and lead dioxide electrodes with a sulfuric acid electrolyte. Chemical reactions at the electrodes involve the transfer of electrons between the electrodes and ions in the electrolyte.
2) As the battery charges and discharges, the concentration of the sulfuric acid electrolyte changes. This affects the voltage according to the Nernst equation.
3) Factors like internal resistance and surface chemistry effects cause the terminal voltage to differ from the theoretical voltage. Battery models account for these factors.
This document provides an overview of key concepts in electrochemistry, including:
1) Electrochemical cells use spontaneous redox reactions to produce electrical energy through electron transfer along an external path between electrodes.
2) The standard cell potential (ΔE°cell) and free energy change (ΔG°) quantify a redox reaction's tendency to proceed.
3) Half-cell potentials determine ΔE°cell, with the more positive half having a greater tendency for oxidation.
4) The Nernst equation relates cell potential to non-standard state concentrations.
5) Corrosion occurs via redox at metal surfaces, and can be inhibited or protected against.
This document discusses oxidation-reduction (redox) reactions and electrochemistry.
1. It explains how to identify redox reactions by checking if the oxidation number (O.N.) of any species changes in the reaction. An example reaction between permanganate and oxalic acid is given.
2. Balancing redox reactions is important, and the document outlines the step-by-step process for balancing both acidic and basic redox reactions.
3. Electrochemical cells are described as either galvanic cells that generate potential or electrolytic cells that consume potential. The standard hydrogen electrode is used as a reference electrode with a standard potential of 0 V.
This document summarizes an electrochemistry chapter that covers:
- Types of electrochemical cells including galvanic and electrolytic cells
- Reversible electrodes like metal-metal ion, gas, and metal-insoluble electrodes
- Determining standard electrode potentials and using the Nernst equation
- Examples of calculating cell potentials and writing electrode half reactions
This document summarizes an electrochemistry chapter that covers:
- Types of electrochemical cells including galvanic and electrolytic cells
- Reversible electrodes like metal-metal ion, gas, and metal-insoluble electrodes
- Determining standard electrode potentials and using the Nernst equation
- Examples of calculating cell potentials and writing electrode half reactions
This document provides an overview of electrochemistry concepts including:
- Types of electrochemical processes including reversible and irreversible processes.
- Oxidation-reduction reactions and how they involve oxidation and reduction half-reactions.
- Galvanic/voltaic cells and how they generate electricity from spontaneous redox reactions.
- Components of electrochemical cells including electrodes, salt bridges, and how they allow indirect redox reactions.
- Standard electrode potentials and how they are used to determine if a reaction is spontaneous.
- The Nernst equation and how it describes the dependence of electrode potential on ion concentration.
Electrochemistry of some mono porphyrinssuchi ghosh
This document discusses various electrochemical techniques like polarography, linear sweep voltammetry, and cyclic voltammetry. It summarizes the results of applying these techniques to study the redox properties of porphyrins, metalloporphyrins, corroles, and manganese-containing corroles. Key findings include the effects of metal insertion, axial ligation, solvent, and linking groups like terephthalic acid on the redox potentials and mechanisms observed in cyclic voltammograms. Electrochemical techniques provide insight into electron transfer kinetics and mechanisms in these molecular systems.
Notes and Important Points on Electrochemistry - JEE Main 2015Ednexa
This document provides an introduction to electrochemistry and discusses key concepts such as electrolytes, electrolytic cells, and the preferential discharge theory of electrolysis. Some main points:
1) Electrolytes are substances that undergo decomposition into ions when an electric current is passed through them in solution. Electrolytic cells, also called voltameters, are devices used to carry out electrolysis where electrical energy is converted to chemical energy.
2) During electrolysis, oxidation occurs at the anode where anions are released and reduction occurs at the cathode where cations gain electrons. The preferential discharge theory states that ions with lower discharge potentials will be discharged first at the appropriate electrode.
3) Examples of products formed
The document provides an overview of fuel cell technology. It discusses the brief history of fuel cells and the basic principles of electrolysis and how fuel cells work by reversing the electrolysis process. It describes the main components of a fuel cell and the five most common types: alkaline, molten carbonate, phosphoric acid, proton exchange membrane, and solid oxide fuel cells. The benefits of fuel cells are highlighted such as efficiency, reliability and fuel flexibility. Challenges for different fuel cell types are also summarized, for example high operating temperatures of solid oxide fuel cells can limit applications.
Conductors and Non-Conductors
Substances can be classified as conductors and non-conductors based on their ability to conduct electricity.
Conductors: Substances that allow electric current to flow through them are called conductors. For example, Plastic, Wood, etc.
Non-Conductors: Non-conductors are insulators that do not allow electricity to pass through them. For example, Copper, Iron, etc.
Types of Conductors
Conductors are divided into two groups: Metallic conductors and Electrolytes.
Metallic Conductors: These conductors conduct electricity by the movement of electrons without any chemical change during the process. This type of conduction happens in solids and in the molten state.
Electrolytes: They conduct electricity by the movement of the ions in the solutions. It is present in the aqueous solution.
Distinguish between Metallic and Electrolytic Conduction
Metallic Conduction Electrolytic Conduction
The movement of electrons causes the electric current The movement of ions causes the electric current
There is no chemical reaction Ions get ionised or reduced at the electrodes
There is no transfer of matter It involves the transfer of matter in the form of ions
Follows Ohm’s law Follows Ohm’s law
Resistance increases with an increase in temperature Resistance decreases with an increase in temperature
Faraday’s law is not followed Follows Faraday’s law
Electrolytes
(a) Substances whose aqueous solutions allow the conductance of electric current and are chemically decomposed are called electrolytes.
(b) The positively charged ions furnished by the electrolyte are called cations, while the negatively charged ions furnished by the electrolyte are called anions.
Types of Electrolytes
(a) Weak electrolytes: Electrolytes that are decomposable to a very small extent in their dilute solutions are called weak electrolytes. For example, organic acids, inorganic acids and bases etc.
(b) Strong electrolytes: Electrolytes that are highly decomposable in aqueous solution and conduct electricity frequently are called electrolytes. For example, mineral acid and salts of strong acid.
Electrode
For the electric current to pass through an electrolytic conductor, the two rods or plates called electrodes are always needed. These plates are connected to the terminals of the battery to form a cell. The electrode through which the electric current flows into the electrolytic solution is called the anode, also called the positive electrode, and anions are oxidised here.
An electrode through which the electric current flows out of the electrolytic solution is called the cathode, also called the negative electrode, and cations are reduced there.
Electrolysis
Electrolysis is the process of chemical deposition of the electrolyte by passing an electric current. Electrolysis takes place in an electrolytic cell. This cell will convert the electrical energy to chemical energy.
New chm-152-unit-8-power-points-sp13-140227172047-phpapp01Cleophas Rwemera
The document provides information about electrochemistry including:
1) It discusses voltaic (galvanic) cells and electrolytic cells, how they are constructed using two electrodes in an electrolyte solution, and the definitions of anode and cathode.
2) It describes the zinc-copper cell as an example, showing the oxidation and reduction half-reactions, overall reaction, and cell notation. The initial voltage is given as 1.10 volts.
3) It explains how standard electrode potentials are measured relative to the standard hydrogen electrode, which has a defined potential of 0.00 V. Standard potentials allow comparison of an electrode's ability to be reduced or act as an oxidizing agent.
This document discusses electrochemical cells and how they produce electric current from redox reactions. It explains that a voltaic cell separates the oxidation and reduction half-reactions into different compartments connected by a salt bridge and wire. The document provides examples of how to construct half-cells for zinc-copper and aluminum-nickel voltaic cells and calculates cell voltage using reduction potentials from a table.
This document describes the characterization and density functional theory (DFT) calculations of a Fe, Cu dual-metal single atom catalyst on carbon black for oxygen reduction reaction. Transmission electron microscopy, X-ray diffraction, X-ray photoelectron spectroscopy and other techniques were used to characterize the catalyst's structure and composition. DFT calculations were performed to simulate the oxygen reduction reaction mechanism on the catalyst surface and calculate adsorption energies of reaction intermediates. The catalyst showed high activity and selectivity for oxygen reduction, outperforming monometallic analogues due to synergistic effects between isolated Fe and Cu sites.
- Electrochemical cells operate by allowing electrons to spontaneously flow from the reducing agent to the oxidizing agent, generating a current.
- A galvanic cell uses two half-cells with aqueous electrolyte solutions and electrodes to induce redox reactions. The half-reactions occur separately in each solution.
- The standard cell potential (ΔE°) can be calculated from the half-cell potentials and indicates whether the cell reaction is spontaneous.
This document provides an overview of basic electrochemistry concepts. It discusses the charge and current involved in electrochemical processes. It introduces Faraday's laws relating the amount of material transformed to the quantity of electricity passed. It also covers conductivity, Nernst equation, different types of electrodes, potentiometry, and various electrochemical techniques including cyclic voltammetry. The key concepts covered include electron transfer processes, Butler-Volmer equation, mass transport by diffusion and convection, and reversible cyclic voltammograms.
The document provides an overview of electrochemistry concepts including:
- Reference electrodes like the standard hydrogen electrode and calomel electrode are used to measure electrode potentials.
- The electrochemical series arranges metals based on their electrode potentials and can predict displacement reactions and reaction spontaneity.
- The Nernst equation relates cell potential to standard potential and activity of products and reactants, allowing prediction of cell potential under non-standard conditions.
- Key applications of concepts like the electrochemical series, equilibrium constants, and Nernst equation include calculating standard cell potentials and determining reaction feasibility and direction.
This document discusses electrochemistry and galvanic cells. It defines oxidation and reduction, and explains how galvanic cells work by using half-reactions and a salt bridge or porous disk to allow ions to flow while preventing the electrons from mixing. It discusses how cell potential is calculated from standard reduction potentials of the half-reactions, and how the direction of electron flow determines the anode and cathode. Standard conditions and notation for describing complete galvanic cells are also covered.
A.) Comparison of Voltammetry to Other Electrochemical Methods
1.) Voltammetry: electrochemical method in which information about an analyte is
obtained by measuring current (i) as a function of applied potential
- only a small amount of sample (analyte) is used
Instrumentation – Three electrodes in solution containing analyte
Working electrode: microelectrode whose potential is varied with time
Reference electrode: potential remains constant (Ag/AgCl electrode or calomel)
Counter electrode: Hg or Pt that completes circuit, conducts e- from signal source through solution to the working electrode
Supporting electrolyte: excess of nonreactive electrolyte (alkali metal) to conduct current
B.) Theory of Voltammetry
1.) Excitation Source: potential set by instrument (working electrode)
- establishes concentration of Reduced and Oxidized Species at electrode based on Nernst Equation:
- reaction at the surface of the electrode
Analyte selectivity is provided by the applied potential on the working electrode.
Electroactive species in the sample solution are drawn towards the working electrode where a half-cell redox reaction takes place.
Another corresponding half-cell redox reaction will also take place at the counter electrode to complete the electron flow.
The resultant current flowing through the electrochemical cell reflects the activity (i.e. concentration) of the electroactive species involved
2.) Current generated at electrode by this process is proportional to concentration at
surface, which in turn is equal to the bulk concentration
For a planar electrode:
measured current (i) = nFADA( )
where:
n = number of electrons in ½ cell reaction
F = Faraday’s constant
A = electrode area (cm2)
D = diffusion coefficient (cm2/s) of A (oxidant)
= slope of curve between CMox,bulk and CMox,s
The document discusses electrochemistry and Daniel cells. It provides details on:
- How Daniel cells work by converting chemical energy from a redox reaction of zinc and copper into electrical energy.
- The components of a Daniel cell including zinc and copper electrodes, zinc sulfate and copper sulfate solutions, and a salt bridge to maintain electrical neutrality.
- How the cell produces a voltage through the oxidation of zinc and reduction of copper ions.
- How the voltage depends on the concentration of ions, as described by the Nernst equation.
The document provides information on phase diagrams including:
- Phase diagrams represent the phases present in materials at different conditions of temperature, pressure, and composition. They indicate solubility, solidification ranges, and melting points.
- Pure substances have solid, liquid, and vapor phases separated by phase boundaries and coexisting at triple points, as shown in pressure-temperature diagrams.
- Binary alloy phase diagrams show the phases present at different compositions and temperatures, including solid solutions, eutectic points where two solids form from liquid, and peritectic reactions where a solid and liquid form a new solid phase.
- The Gibbs phase rule and lever rule are used to analyze multi-phase regions. Cool
Stanley A Meyer Legacy Back up Secret Docs Save all Protect Spread print and give to schools NEVER STOP!!!!!!! Join Support here https://www.patreon.com/securesupplies/shop
Stanley A Meyer Legacy Back up Secret Docs Save all Protect Spread print and give to schools NEVER STOP!!!!!!! Join Support here https://www.patreon.com/securesupplies/shop
The document summarizes key concepts about lead-acid batteries, including:
1) Lead-acid batteries use lead and lead dioxide electrodes with a sulfuric acid electrolyte. Chemical reactions at the electrodes involve the transfer of electrons between the electrodes and ions in the electrolyte.
2) As the battery charges and discharges, the concentration of the sulfuric acid electrolyte changes. This affects the voltage according to the Nernst equation.
3) Factors like internal resistance and surface chemistry effects cause the terminal voltage to differ from the theoretical voltage. Battery models account for these factors.
This document provides an overview of key concepts in electrochemistry, including:
1) Electrochemical cells use spontaneous redox reactions to produce electrical energy through electron transfer along an external path between electrodes.
2) The standard cell potential (ΔE°cell) and free energy change (ΔG°) quantify a redox reaction's tendency to proceed.
3) Half-cell potentials determine ΔE°cell, with the more positive half having a greater tendency for oxidation.
4) The Nernst equation relates cell potential to non-standard state concentrations.
5) Corrosion occurs via redox at metal surfaces, and can be inhibited or protected against.
This document discusses oxidation-reduction (redox) reactions and electrochemistry.
1. It explains how to identify redox reactions by checking if the oxidation number (O.N.) of any species changes in the reaction. An example reaction between permanganate and oxalic acid is given.
2. Balancing redox reactions is important, and the document outlines the step-by-step process for balancing both acidic and basic redox reactions.
3. Electrochemical cells are described as either galvanic cells that generate potential or electrolytic cells that consume potential. The standard hydrogen electrode is used as a reference electrode with a standard potential of 0 V.
This document summarizes an electrochemistry chapter that covers:
- Types of electrochemical cells including galvanic and electrolytic cells
- Reversible electrodes like metal-metal ion, gas, and metal-insoluble electrodes
- Determining standard electrode potentials and using the Nernst equation
- Examples of calculating cell potentials and writing electrode half reactions
This document summarizes an electrochemistry chapter that covers:
- Types of electrochemical cells including galvanic and electrolytic cells
- Reversible electrodes like metal-metal ion, gas, and metal-insoluble electrodes
- Determining standard electrode potentials and using the Nernst equation
- Examples of calculating cell potentials and writing electrode half reactions
This document provides an overview of electrochemistry concepts including:
- Types of electrochemical processes including reversible and irreversible processes.
- Oxidation-reduction reactions and how they involve oxidation and reduction half-reactions.
- Galvanic/voltaic cells and how they generate electricity from spontaneous redox reactions.
- Components of electrochemical cells including electrodes, salt bridges, and how they allow indirect redox reactions.
- Standard electrode potentials and how they are used to determine if a reaction is spontaneous.
- The Nernst equation and how it describes the dependence of electrode potential on ion concentration.
Electrochemistry of some mono porphyrinssuchi ghosh
This document discusses various electrochemical techniques like polarography, linear sweep voltammetry, and cyclic voltammetry. It summarizes the results of applying these techniques to study the redox properties of porphyrins, metalloporphyrins, corroles, and manganese-containing corroles. Key findings include the effects of metal insertion, axial ligation, solvent, and linking groups like terephthalic acid on the redox potentials and mechanisms observed in cyclic voltammograms. Electrochemical techniques provide insight into electron transfer kinetics and mechanisms in these molecular systems.
Notes and Important Points on Electrochemistry - JEE Main 2015Ednexa
This document provides an introduction to electrochemistry and discusses key concepts such as electrolytes, electrolytic cells, and the preferential discharge theory of electrolysis. Some main points:
1) Electrolytes are substances that undergo decomposition into ions when an electric current is passed through them in solution. Electrolytic cells, also called voltameters, are devices used to carry out electrolysis where electrical energy is converted to chemical energy.
2) During electrolysis, oxidation occurs at the anode where anions are released and reduction occurs at the cathode where cations gain electrons. The preferential discharge theory states that ions with lower discharge potentials will be discharged first at the appropriate electrode.
3) Examples of products formed
The document provides an overview of fuel cell technology. It discusses the brief history of fuel cells and the basic principles of electrolysis and how fuel cells work by reversing the electrolysis process. It describes the main components of a fuel cell and the five most common types: alkaline, molten carbonate, phosphoric acid, proton exchange membrane, and solid oxide fuel cells. The benefits of fuel cells are highlighted such as efficiency, reliability and fuel flexibility. Challenges for different fuel cell types are also summarized, for example high operating temperatures of solid oxide fuel cells can limit applications.
Conductors and Non-Conductors
Substances can be classified as conductors and non-conductors based on their ability to conduct electricity.
Conductors: Substances that allow electric current to flow through them are called conductors. For example, Plastic, Wood, etc.
Non-Conductors: Non-conductors are insulators that do not allow electricity to pass through them. For example, Copper, Iron, etc.
Types of Conductors
Conductors are divided into two groups: Metallic conductors and Electrolytes.
Metallic Conductors: These conductors conduct electricity by the movement of electrons without any chemical change during the process. This type of conduction happens in solids and in the molten state.
Electrolytes: They conduct electricity by the movement of the ions in the solutions. It is present in the aqueous solution.
Distinguish between Metallic and Electrolytic Conduction
Metallic Conduction Electrolytic Conduction
The movement of electrons causes the electric current The movement of ions causes the electric current
There is no chemical reaction Ions get ionised or reduced at the electrodes
There is no transfer of matter It involves the transfer of matter in the form of ions
Follows Ohm’s law Follows Ohm’s law
Resistance increases with an increase in temperature Resistance decreases with an increase in temperature
Faraday’s law is not followed Follows Faraday’s law
Electrolytes
(a) Substances whose aqueous solutions allow the conductance of electric current and are chemically decomposed are called electrolytes.
(b) The positively charged ions furnished by the electrolyte are called cations, while the negatively charged ions furnished by the electrolyte are called anions.
Types of Electrolytes
(a) Weak electrolytes: Electrolytes that are decomposable to a very small extent in their dilute solutions are called weak electrolytes. For example, organic acids, inorganic acids and bases etc.
(b) Strong electrolytes: Electrolytes that are highly decomposable in aqueous solution and conduct electricity frequently are called electrolytes. For example, mineral acid and salts of strong acid.
Electrode
For the electric current to pass through an electrolytic conductor, the two rods or plates called electrodes are always needed. These plates are connected to the terminals of the battery to form a cell. The electrode through which the electric current flows into the electrolytic solution is called the anode, also called the positive electrode, and anions are oxidised here.
An electrode through which the electric current flows out of the electrolytic solution is called the cathode, also called the negative electrode, and cations are reduced there.
Electrolysis
Electrolysis is the process of chemical deposition of the electrolyte by passing an electric current. Electrolysis takes place in an electrolytic cell. This cell will convert the electrical energy to chemical energy.
New chm-152-unit-8-power-points-sp13-140227172047-phpapp01Cleophas Rwemera
The document provides information about electrochemistry including:
1) It discusses voltaic (galvanic) cells and electrolytic cells, how they are constructed using two electrodes in an electrolyte solution, and the definitions of anode and cathode.
2) It describes the zinc-copper cell as an example, showing the oxidation and reduction half-reactions, overall reaction, and cell notation. The initial voltage is given as 1.10 volts.
3) It explains how standard electrode potentials are measured relative to the standard hydrogen electrode, which has a defined potential of 0.00 V. Standard potentials allow comparison of an electrode's ability to be reduced or act as an oxidizing agent.
This document discusses electrochemical cells and how they produce electric current from redox reactions. It explains that a voltaic cell separates the oxidation and reduction half-reactions into different compartments connected by a salt bridge and wire. The document provides examples of how to construct half-cells for zinc-copper and aluminum-nickel voltaic cells and calculates cell voltage using reduction potentials from a table.
This document describes the characterization and density functional theory (DFT) calculations of a Fe, Cu dual-metal single atom catalyst on carbon black for oxygen reduction reaction. Transmission electron microscopy, X-ray diffraction, X-ray photoelectron spectroscopy and other techniques were used to characterize the catalyst's structure and composition. DFT calculations were performed to simulate the oxygen reduction reaction mechanism on the catalyst surface and calculate adsorption energies of reaction intermediates. The catalyst showed high activity and selectivity for oxygen reduction, outperforming monometallic analogues due to synergistic effects between isolated Fe and Cu sites.
- Electrochemical cells operate by allowing electrons to spontaneously flow from the reducing agent to the oxidizing agent, generating a current.
- A galvanic cell uses two half-cells with aqueous electrolyte solutions and electrodes to induce redox reactions. The half-reactions occur separately in each solution.
- The standard cell potential (ΔE°) can be calculated from the half-cell potentials and indicates whether the cell reaction is spontaneous.
This document provides an overview of basic electrochemistry concepts. It discusses the charge and current involved in electrochemical processes. It introduces Faraday's laws relating the amount of material transformed to the quantity of electricity passed. It also covers conductivity, Nernst equation, different types of electrodes, potentiometry, and various electrochemical techniques including cyclic voltammetry. The key concepts covered include electron transfer processes, Butler-Volmer equation, mass transport by diffusion and convection, and reversible cyclic voltammograms.
The document provides an overview of electrochemistry concepts including:
- Reference electrodes like the standard hydrogen electrode and calomel electrode are used to measure electrode potentials.
- The electrochemical series arranges metals based on their electrode potentials and can predict displacement reactions and reaction spontaneity.
- The Nernst equation relates cell potential to standard potential and activity of products and reactants, allowing prediction of cell potential under non-standard conditions.
- Key applications of concepts like the electrochemical series, equilibrium constants, and Nernst equation include calculating standard cell potentials and determining reaction feasibility and direction.
This document discusses electrochemistry and galvanic cells. It defines oxidation and reduction, and explains how galvanic cells work by using half-reactions and a salt bridge or porous disk to allow ions to flow while preventing the electrons from mixing. It discusses how cell potential is calculated from standard reduction potentials of the half-reactions, and how the direction of electron flow determines the anode and cathode. Standard conditions and notation for describing complete galvanic cells are also covered.
A.) Comparison of Voltammetry to Other Electrochemical Methods
1.) Voltammetry: electrochemical method in which information about an analyte is
obtained by measuring current (i) as a function of applied potential
- only a small amount of sample (analyte) is used
Instrumentation – Three electrodes in solution containing analyte
Working electrode: microelectrode whose potential is varied with time
Reference electrode: potential remains constant (Ag/AgCl electrode or calomel)
Counter electrode: Hg or Pt that completes circuit, conducts e- from signal source through solution to the working electrode
Supporting electrolyte: excess of nonreactive electrolyte (alkali metal) to conduct current
B.) Theory of Voltammetry
1.) Excitation Source: potential set by instrument (working electrode)
- establishes concentration of Reduced and Oxidized Species at electrode based on Nernst Equation:
- reaction at the surface of the electrode
Analyte selectivity is provided by the applied potential on the working electrode.
Electroactive species in the sample solution are drawn towards the working electrode where a half-cell redox reaction takes place.
Another corresponding half-cell redox reaction will also take place at the counter electrode to complete the electron flow.
The resultant current flowing through the electrochemical cell reflects the activity (i.e. concentration) of the electroactive species involved
2.) Current generated at electrode by this process is proportional to concentration at
surface, which in turn is equal to the bulk concentration
For a planar electrode:
measured current (i) = nFADA( )
where:
n = number of electrons in ½ cell reaction
F = Faraday’s constant
A = electrode area (cm2)
D = diffusion coefficient (cm2/s) of A (oxidant)
= slope of curve between CMox,bulk and CMox,s
The document discusses electrochemistry and Daniel cells. It provides details on:
- How Daniel cells work by converting chemical energy from a redox reaction of zinc and copper into electrical energy.
- The components of a Daniel cell including zinc and copper electrodes, zinc sulfate and copper sulfate solutions, and a salt bridge to maintain electrical neutrality.
- How the cell produces a voltage through the oxidation of zinc and reduction of copper ions.
- How the voltage depends on the concentration of ions, as described by the Nernst equation.
The document provides information on phase diagrams including:
- Phase diagrams represent the phases present in materials at different conditions of temperature, pressure, and composition. They indicate solubility, solidification ranges, and melting points.
- Pure substances have solid, liquid, and vapor phases separated by phase boundaries and coexisting at triple points, as shown in pressure-temperature diagrams.
- Binary alloy phase diagrams show the phases present at different compositions and temperatures, including solid solutions, eutectic points where two solids form from liquid, and peritectic reactions where a solid and liquid form a new solid phase.
- The Gibbs phase rule and lever rule are used to analyze multi-phase regions. Cool
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Charging and Fueling Infrastructure Grant: Round 2 by Brandt HertensteinForth
Brandt Hertenstein, Program Manager of the Electrification Coalition gave this presentation at the Forth and Electrification Coalition CFI Grant Program - Overview and Technical Assistance webinar on June 12, 2024.
Charging Fueling & Infrastructure (CFI) Program by Kevin MillerForth
Kevin Miller, Senior Advisor, Business Models of the Joint Office of Energy and Transportation gave this presentation at the Forth and Electrification Coalition CFI Grant Program - Overview and Technical Assistance webinar on June 12, 2024.
Dahua provides a comprehensive guide on how to install their security camera systems. Learn about the different types of cameras and system components, as well as the installation process.
Expanding Access to Affordable At-Home EV Charging by Vanessa WarheitForth
Vanessa Warheit, Co-Founder of EV Charging for All, gave this presentation at the Forth Addressing The Challenges of Charging at Multi-Family Housing webinar on June 11, 2024.
EV Charging at MFH Properties by Whitaker JamiesonForth
Whitaker Jamieson, Senior Specialist at Forth, gave this presentation at the Forth Addressing The Challenges of Charging at Multi-Family Housing webinar on June 11, 2024.
Charging Fueling & Infrastructure (CFI) Program Resources by Cat PleinForth
Cat Plein, Development & Communications Director of Forth, gave this presentation at the Forth and Electrification Coalition CFI Grant Program - Overview and Technical Assistance webinar on June 12, 2024.
Understanding Catalytic Converter Theft:
What is a Catalytic Converter?: Learn about the function of catalytic converters in vehicles and why they are targeted by thieves.
Why are They Stolen?: Discover the valuable metals inside catalytic converters (such as platinum, palladium, and rhodium) that make them attractive to criminals.
Steps to Prevent Catalytic Converter Theft:
Parking Strategies: Tips on where and how to park your vehicle to reduce the risk of theft, such as parking in well-lit areas or secure garages.
Protective Devices: Overview of various anti-theft devices available, including catalytic converter locks, shields, and alarms.
Etching and Marking: The benefits of etching your vehicle’s VIN on the catalytic converter or using a catalytic converter marking kit to make it traceable and less appealing to thieves.
Surveillance and Monitoring: Recommendations for using security cameras and motion-sensor lights to deter thieves.
Statistics and Insights:
Theft Rates by Borough: Analysis of data to determine which borough in NYC experiences the highest rate of catalytic converter thefts.
Recent Trends: Current trends and patterns in catalytic converter thefts to help you stay aware of emerging hotspots and tactics used by thieves.
Benefits of This Presentation:
Awareness: Increase your awareness about catalytic converter theft and its impact on vehicle owners.
Practical Tips: Gain actionable insights and tips to effectively prevent catalytic converter theft.
Local Insights: Understand the specific risks in different NYC boroughs, helping you take targeted preventive measures.
This presentation aims to equip you with the knowledge and tools needed to protect your vehicle from catalytic converter theft, ensuring you are prepared and proactive in safeguarding your property.
Implementing ELDs or Electronic Logging Devices is slowly but surely becoming the norm in fleet management. Why? Well, integrating ELDs and associated connected vehicle solutions like fleet tracking devices lets businesses and their in-house fleet managers reap several benefits. Check out the post below to learn more.
1. Different polarity firing sequences for electrodes in the steam generator as found in Stan
Meyer's literature
Trialling the multifunctional square wave generator to be used to test circuitry and
components of the VIC and steam resonator
Detailed scientific explanations of the electrolysis of water, magnetic effects on water,
electromagnetic effects on water and the water redox process
Different tube set configurations, thicknesses and gaps between the anode and cathode
tubes found in Stan Meyer's literature
ESP32 microcontroller programming principles for operation of steam resonator
ESP32 microcontroller programming principles for pulse train to VIC
Paul Butcher
2. Steam resonator: the different polarity firing sequences across the parallel plates (or
could be concentric tubes) are as follows:
Polarity firing sequence 1 for a pair of plates: WFC Hydrogen Gas Management System Memo
WFC 422DA page 3-25 and diagram 3-46 on page 3-50 and also Steam Resonator Memo WFC 430
figure 11-7 on page 11-13. For a pair of plates the sequence is: -
B - B-
OFF OFF
B+ B+
OFF OFF
B- B-
OFF OFF
B+ B+
The above firing sequence will cause the water molecules to oscillate between the plates, collide
and heat up.
Polarity firing sequence 2 for two pairs of plates: Steam Resonator Memo WFC 430 third
paragraph on page 11-4, figure 11-1 on page 11-7, figure 11-3 on page 11-9 and figure 11-4 on
page 11-10.
Different polarity firing
sequences for electrodes in
the steam generator as
found in Stan Meyer's
literature
3. Left-hand pair of plates
B+ OFF
OFF B-
B+ OFF
OFF B-
Right-hand pair of plates
B- OFF
OFF B+
B- OFF
OFF B+
The above firing sequence for either pair of plates will also cause the water molecules to oscillate,
collide and heat up. Now, as noted on page 11-4 (last paragraph), the voltage wave will travel up /
along the surfaces of the plates. So polarity firing sequence 2 is probably better suited for use in a
home heating unit because of the vertical pumping action that it also causes to happen which is
mentioned in the second paragraph of page 11-5 and also shown in figure 11-6 on page 11-12.
Polarity firing sequence 3 for two pairs of plates
Unfortunately, in the first and third paragraphs on page 11-5 of Steam Resonator Memo WFC 430
there is a different firing sequence which is: -
Left-hand pair of plates
B+ 0FF
0FF B+
B+ 0FF
0FF B+
Right-hand pair of plates
B- 0FF
0FF B-
B- 0FF
4. 0FF B-
I don’t understand this sequence and how it is supposed to work on the water molecule dipoles. Is
this sequence perhaps a mistake because it does not match firing sequence 2 ?
Polarity firing sequence 4 for a pair of plates: Steam Resonator WFC 427 DA Figure 1-2: Dual
switchover circuit. The firing sequence for pair of plates is the same as the left-hand pair of plates
in polarity firing sequence 3 above because it is shown as: -
B+ OFF
OFF B+
B+ OFF
OFF B+
Again, I don’t understand how this sequence is supposed to act on the water molecule dipoles. Is
this sequence also a mistake because it does not match the sequence for the left-hand pair of
plates in firing sequence 2 ?
Polarity firing sequence 5 (applicable to spherical water heater) as described on page K3 and
figure 32 on page K4 of the Water Fuel Cell Dealership Sales Manual 1986. This firing sequence
consists of only positive voltage pulses being applied to a spherical water bath which from all
angles causes repulsion of the positive side of the water molecule dipoles. This is supposed to
result in constant collision of the water molecules as they are repeatedly driven towards the middle
of the sphere ? I don’t understand why nothing is said on pages K3 or K4 about the interaction of
the positive voltage pulses and the negative side of the water molecule dipoles?
7. Detailed scientific explanations of the electrolysis of water is detailed below and the electric effects
on water, magnetic effects on water, electromagnetic effects on water and the water redox process
can all be accessed using the links immediately below: -
Electric effects on water
Magnetic effects on water
Electromagnetic effects on water
Water redox processes
Electrolysis of Water
Electrolysis of water is its decomposition to give hydrogen and oxygen gases due to an electric
current.
2 H2
O + electrical energy (+ heat energy) O2
+ 2 H2
'I propose to distinguish these bodies by calling those anions which go to the anode ....and those
passing to the cathode, cations '
Detailed scientific
explanations of the
electrolysis of water,
magnetic effects on water,
electromagnetic effects on
water and the water redox
process
8. Michael Faraday 1834
Introduction
Creating an electric potential through water causes positive ions, including the inherent hydrogen
ions (H3
O+), to move towards the negative electrode (cathode) and negative ions, including the
inherent hydroxide ions (OH−), to move towards the positive electrode (anode). With a sufficient
potential difference, this may cause electrolysis with oxygen gas being produced at the anode and
hydrogen gas produced at the cathode (see [1878] for current reviews). f The electrolysis g of water
usually involves dilute, or moderately concentrated, salt solutions to reduce the power loss driving
the current through the solution and catalyze the reaction (see below). However, the presence of
salt is not a requirement for electrolysis. h Although often taught as an uncomplicated topic, the
electrolysis of water does not involve easy to understand concepts; particularly if including the
necessary mass transport and kinetics [4168],
Thus,
Anode +ve i 6H2
O(l) O2
(g) + 4H3
O+(aq) + 4e−(to
anode) b
E° = +1.229 V,
pH 0 d
E°' = +0.815 V
Cathode −ve
4e−(from
cathode) + 4H2
O(l) 2H2
(g) +
4OH−(aq)
E° = −0.828 V,
pH 14
E°' = −0.414 V
Overall
2H2
O(l) 2H2
(g)
+ O2
(g)
ΔG°' = +474.3 kJ ˣ mol−1
where (l), (g), and (aq) show the states of the material as being a liquid, a gas, or an aqueous
solution. The electrical circuit passes the electrons back from the anode to the cathode. The
reactions are heterogeneous, taking place at the boundary between the electrode and the
electrolyte with the aqueous boundary layer subject to concentration and electrical potential
gradients, and with the presence of the generated gaseous nanobubbles and microbubbles. When
salts are present, enabling greater electron flow, the primary reaction may differ; for example, on
electrolysis of an aqueous solution of copper chloride, a deposit of metallic copper and chlorine gas
are produced, with no production of oxygen or hydrogen gases. Even when oxygen and hydrogen
gases are produced, their production may not be the primary reactions [4167],
primary action 2 Na2
SO4
(aq) 4 Na° (at cathode) + 2 SO4
° (at anode)
secondary action at cathode 4 Na° + 4 H2
O (l) 4 NaOH (aq) + 2 H2
(g)
secondary action at anode 2 H2
O (l) + 2 SO4
° O2
(g) + 2 H2
SO4
(aq)
secondary action in bulk 2 H2
SO4
(aq) + 4 NaOH (aq) 2 Na2
SO4
(aq) + 4 H2
O (l)
overall 2H2
O(l) 2H2
(g) + O2
(g)
9. with the (regenerated) Na2
SO4
acting as a catalyst. Aqueous NaCl electrolysis, however, produces
mainly oxygen and hydrogen gases with only traces of chlorine gas or sodium metal unless the
NaCl is concentrated.
Water electrolysis electrode potentials with pH
The structural and thermodynamic properties for water surfaces in the vicinity of the electric field
exerted by the metal electrodes have been simulated [3829]. Generally, the water adjacent to the
electrodes c will change pH due to the ions produced or consumed. If a suitable porous membrane
separates the electrode compartments, then the concentration of H3
O+ next to the anode (anolyte)
and OH− next to the cathode (catholyte) are both expected to increase more than if there is free
mixing between the electrodes. There will also be an increase in their respective conductivities.
Without such a membrane, most of these ions will neutralize each other. Small but expected
differences in the anolyte and catholyte pHs cause only a slight change to the overall potential
difference required (1.229 V). Increasing the anolyte acid content due to the H3
O+ produced will
increase its electrode potential (for example: at pH 4, E = +0.992 V), and increasing the catholyte
alkaline content due to the OH− produced will make its electrode potential more negative (for
example: at pH 10, E = −0.592 V). If the anode reaction is forced to run at pH 14 and the cathode
reaction is run at pH 0.0, then the electrode potentials are +0.401 V and 0 V, respectively (see
above right). d
(a) Anode pH 0 2 H2
O O2
+ 4 H+ + 4 e− E° = +1.229 V
(b) Anode pH 14 4 OH− O2
+H2
O + 4 e− E° = +0.401 V
(c) Cathode pH 0 4 H+ + 4 e− 2 H2
E° = 0.0 V
(d) Cathode pH 14 4 H2
O + 4 e− 2 H2
+ 4 OH− E° = −0.828 V
Although electrolysis can be achieved with a (minimum) voltage of +0.403 V (see equations b and
c, above) [2515], it does not break the thermodynamic requirement of 1.229 V as further energy is
required to keep the electrode compartments at the required solute concentrations and pHs.
The layer next to the surface of the electrode determines the rate of the reaction [3831]. If it is
stagnant, molecules and ions have to diffuse to and from the electrode, restricting the rate of
reaction (mass transfer limitation) that is reported for current densities below 1.3 kA ˣ m−2 in chlor-
alkali electrolysis. e The mobility of the hydration layer nearest to the electrode (~5 Å) decreases
upon positive potentials while increasing upon negative potentials [3863]. This is because, at
positive potentials, the hydrogen bonding network gets ice-like structured parallel to the electrode,
while at negative interfaces, it is disrupted due to the hydrogen atoms pointing at the surface.
Additionally, there may be a high accumulation of hydroxide ions at positive electrodes that
significantly lower the oxygen solubility. Above 3.9 kA ˣ m−2, there is rapid convectional transport
and no mass transport limitations.
10. The current flowing indicates the rate of electrolysis. The amount of product formed can be
calculated directly from the duration and current flowing, as 96,485 coulombs (i.e., one faraday)
delivers one mole of electrons, with one faraday ideally producing 0.5 moles of H2
plus 0.25 moles
of O2
. Thus, one amp flowing for one second (one coulomb) produces 5.18 µmol H2
(10.455 µg,
0.1177 mL at STP) and 2.59 µmol O2
(82.888 µg, 0.0588 mL at STD; 4.9 kW h/m3 H2
at 60%
efficiency), if there are no side reactions at the electrodes;
that is
Number of moles = Coulombs/(unsigned numeric charge on the ion ˣ faraday)
Number of moles = (Current in amperes ˣ time in seconds)/(unsigned numeric charge on the ion ˣ
faraday)
The gases produced at the electrodes may dissolve, with their equilibrium solubility proportional to
their partial pressure as gases in the atmosphere above the electrolytic surface. Oxygen gas is
poorly soluble (≈ 44 mg ˣ kg−1, ≈ 1.4 mM at 0.1 MPa and 20 °C, but only ≈ 0.29 mM against its
normal atmospheric partial pressure). Hydrogen gas is less soluble (≈ 1.6 mg ˣ kg−1, ≈ 0.80 mM at
0.1 MPa and 20 °C but only ≈ 0.44 nM against its very low normal atmospheric partial pressure). It
may take a considerable time for the solubilities to drop from their initially-super-saturated state to
their equilibrium values after the electrolysis.
Although theoretically, as described above, the current passing should determine the amounts of
hydrogen and oxygen formed, several factors ensure that somewhat lower amounts of gas are
actually found;
(i) some electrons (and products) are used up in side-reactions,
(ii) some of the products are catalytically reconverted to water at the electrodes, particularly
if there is no membrane dividing the electrolysis compartments,
(iii) some hydrogen may absorb into the cathode (particularly if palladium is used),
(iv) some oxygen oxidizes the anode,
(v) some gas remains held up in the nanobubbles for a considerable time, and
(vi) some gas may escape measurement.
Current versus voltage in water electrolysis
The above description hides much important science and grossly over-simplifies the system. The
potential required at any position within the electrolytic cell is determined by the localized
concentration of the reactants and products, including the local pH of the solution, instantaneous
gas partial pressure, and effective electrode surface area loss due to attached gas bubbles.
11. The variation in potential across the cell is not uniform, and there is evidence of the formation of
somewhat kinetically stable large-scale charge zones [3557]. In addition, a greater potential
difference (called overpotential [3141]) is required at both electrodes to overcome the activation
energy barriers and insulating bubble coverage, and then to deliver a significant reaction rate.
Typically at suitable electrodes, such as those made of platinum, the overpotential adds about half
a volt to the potential difference between the electrodes. The use of different catalysts to reduce
the overpotential has been discussed [4213]. In addition, a further potential difference is required
to drive the current through the electrical resistance of the electrolytic cell and circuit. For a
(typical) one-ohm cell circuit resistance, a each amp current flow would require a further one volt
and waste one watt of power. This power (and consequent energy) loss (≈ 20%, [1978]) causes the
electrolyte to warm up during electrolysis.
To clarify:
The minimum necessary cell voltage to start water electrolysis is the potential 1.229 V.
The potential necessary to start water electrolysis without withdrawing heat from the surroundings
is
−ΔH°'/nF = 1.481 V
This results in at least a 21% unavoidable loss of efficiency. Usually, further heat is generated, and
efficiency lost, from the overpotentials applied. Additionally, energy is wasted due to the
evaporation of water from within the wet gases evolved.
The efficiency of electrolysis [1876] increases with the temperature as the hydrogen-bonding
reduces. However, due to the endergonic process, the heat demand increases as the electrical
demand decreases, mostly balancing overall energy demand. If the pressure over the electrolysis is
increased, then more current passes for the same applied voltage. However, the output of gas per
coulomb and the heating effect are both decreased. This is due to the increased solubility of the
gases and smaller bubbles, reducing cell resistance and increasing recombination reactions.
Although reducing the distance between electrodes reduces the resistance of the electrolysis
medium, the process may suffer if the closeness allows a build-up of gas between these electrodes
[1876]. Low to higher pulsed potential increases the reaction (current) and accelerates both the
movement of bubbles from the electrode surface and the mass transfer rate in the electrolyte,
which lowers the electrochemical polarization in the diffusion layer and further increases hydrogen
production efficiency [2075]. The rate of change of the current density (and hence efficiency) can
be increased using a magnetic field [2075, 3041] with or without optical enhancement. [2941]. The
investment costs of electrolysis have been reviewed [3255].
Pure water conducts an electric current very poorly and, for this reason, is difficult (slow) to
electrolyze, except if using deep-sub-Debye-length nanogap electrochemical cells [4304]. Usually,
12. however, some salts will be added or present in tap and ground waters which will be sufficient to
allow electrolysis to proceed significantly. The gases produced may be due to secondary reactions
(see above) [4167]. Such salts, and particularly chloride ions, may then undergo redox reactions at
an electrode. These side reactions both reduce the efficiency of the electrolysis reactions (above)
and produce new solutes. Other electrolytic reactions may occur at the electrodes so producing
further solutes and gases. In addition, these solutes may react together to produce other materials.
Together the side reactions are complex, and this complexity increases somewhat when the
voltage applied to the cell is greater than that required by the above reactions and processes. The
likely reactions within the electrode compartments are described below. Some of these may only
occur to a minimal extent, and other reactions may also be occurring that are not included.
Standard electrode potentials are shown below.
Electrode compartment contents in water (NaCl) electrolysis
Electrolysis compartments.
The effects of current, salt concentration, and time on the pH and alkalinity of the electrolytic
solutions has been investigated [4039]. A representation of the compartments in the electrolytic
cell is shown right, with some of their constituent molecules, ions, and radicals. Other materials
may be present, and some of the materials given may be at very low concentrations or have short
half-lives.
Ozone, O3
Noteworthy amongst the side products is ozone (O3
, see left). The relative amount of O3
produced
(relative to molecular oxygen) depends on the overpotential, pH, radicals present, and anode
material. O3
evolution is much lower than that for O2
due to the higher potential required. Very
little O3
may be produced at low overpotentials, but at high current densities and overpotential, up
to a sixth (or more) of the oxidized molecules may be O3
. As O3
is more soluble than O2
, there may
be twice the dissolved O3
than O2
, but the bubble gas will contain about 20 times the O2
than O3
[
2358]. Tin oxide anodes have proved helpful for the production of O3
, particularly if doped with Sb
and Ni, as they bind both oxygen molecules and hydroxyl radicals to facilitate the O3
production [
2359]. Ozone decomposes in water in a few minutes. Decomposition of ozone (particularly in
alkaline solution) gives rise to several strong oxidants, including hydroxyl radicals (·OH), that form
a powerful oxidizing agent capable of killing viruses, amoebae, algae, and dangerous bacteria,
such as MRSA and Legionella.
2 O3
3 O2
O3
+ OH− HO2
− + O2
O3
+ HO2
− ·OH + O2
·− + O2
Although charged ions are attracted into the compartments under the applied potential, oppositely
charged ions are created in both compartments due to the electrolytic reactions. Thus, for
13. example, Na+ ions enter the catholyte from the anode compartment, but excess OH− is produced
simultaneously at the cathode. The concentration of the OH− ions will be generally expected to be
greater than the increase in cations in the catholyte, and the concentration of the H3
O+ ions will be
generally expected to be greater than any increase in anions in the anolyte. Often a conductive but
semi-permeable membrane (for example, Nafion, a highly hydrated sulfonated tetrafluoroethylene
based copolymer [1880]) is used to separate the two compartments and reduce the movement of
the products between the electrode compartments; a process that improves the yield by reducing
back and side reactions [1978]. Due to the easier electrolysis of water containing 1H rather than 2H
(D) or 3H (T), electrolysis can produce water with reduced or enriched isotopic composition.
Local inhomogeneities of surface tension in the produced gas bubbles may be caused by
temperature or altered material concentration gradients at the interface. The resulting solute
currents enhance the mass transfer and bubble growth [3264].
When electrolysis uses short voltage pulses of alternating polarity at above 100 kHz, the
nanobubbles produced contain both H2
and O2
gases that can spontaneously react (combust) to
form water while producing pressure jumps [2900].
Proposed mechanism for electrolysis on platinum
What is less well understood?
Although much time has been spent on investigating and modeling the electrolytic system [1877],
it is still not entirely clear how water is arranged on the surface of the electrodes. Alignment of the
water dipoles with the field is expected, together with the consequential breakage of a proportion
of the water molecules’ hydrogen bonds. Whether the water at the electrode surface is “free” or
coordinated to strong electrolytes (such as Li+ and Na +) affects the ease of electrolysis, with
coordinating water more reactive than “free” water [3516].
When the electrode processes occur, singly-linked hydrogen atoms and singly-linked oxygen atoms
are bound to the platinum atoms at the cathode and anode. The binding energies of these
hydrophilic intermediates are strongly influenced by hydrogen-bonding (HB) to surface water
molecules and the electrode composition [3082]. These bound atoms can diffuse around in two
dimensions on the surface of their respective electrodes until they take part in their further
reaction. Peroxide (···O-O-H) may also be bound to the electrode as part of the O2
dissociation
process [3913]. Other atoms and polyatomic groups may also bind similarly to the electrode
surfaces and subsequently undergo reactions [2899]. Molecules such as O2
and H2
produced at the
surfaces may enter nanoscopic cavities in the liquid water (nanobubbles) as gases, or become
solvated by the water.
Gas-containing cavities in liquid solution (often called bubbles) grow or shrink by diffusion
according to whether the solution is over-saturated or under-saturated with the dissolved gas.
Given suitable electrodes, the size of the cathodic hydrogen bubbles depends on the overvoltage,
14. with nanobubbles being formed at low overvoltages and larger bubbles being formed at higher
overvoltages [2068]. Larger micron-plus sized bubbles have sufficient buoyancy to rise through the
solution and release contained gas at the surface before all the gas dissolves. With smaller bubbles
a pressure is exerted by the surface tension in inverse proportion to their diameter, and bubbles
may be expected to collapse. However, as the nanobubble gas/liquid interface is charged, an
opposing force to the surface tension is introduced, slowing or preventing their dissipation.
Electrolytic solutions have been proven to contain vast numbers of gaseous nanobubbles [974].
The ‘natural’ state of such interfaces appears to be negative [1266]. Other ions with low surface
charge density (such as Cl−, ClO−, HO2
− and O2
·−) will also favor the gas/liquid interfaces [928a]
as probably do hydrated electrons [1841, 1874]. Aqueous radicals also prefer to reside at such
interfaces [939]. From this known information, it seems clear that the nanobubbles present in the
catholyte will be negatively charged. However, those in the anolyte [1881] will probably possess
little charge (with the produced excess positive H3
O+ ions canceling out the natural negative
charge). Therefore, catholyte nanobubbles are not likely to lose their charge on mixing with the
anolyte stream and are otherwise known to be stable for many minutes [974]. Additionally, gas
molecules may become charged within the nanobubbles (such as the superoxide radical ion, O2
·−),
due to the decay of ozone present. Also, the excess potential on the cathode increasies the overall
charge of the nanobubbles and the stability of that charge. The raised temperature at the electrode
surface, due to the excess power loss over that required for the electrolysis, may also increase
nanobubble formation by reducing local gas solubility. Raising the pressure on solutions containing
nanobubbles will also slow down their dissipation if this pressure increases the dissolved gas
content.
Sunlight, as an external electric field in water electrolysis, has proven to increase hydrogen
production. This increase has been associated with an effect on the surface tension [3291].
Oxygen may be reduced (hydrogenated) in acid solution at the cathode. On platinum, Pt(111),
there are two possibilities for the reduction route, that shift with the electrode potential [4092],
O2 (adsorbed)
2 O(adsorbed)
O(adsorbed)
+ H+ + e− OH(adsorbed)
possibly followed by,
OH(adsorbed)
+ H+ + e− H2
O(adsorbed)
or,
O2 (adsorbed)
+ H+ + e− OOH(adsorbed)
possibly followed by,
15. OOH (adsorbed)
+ H+ + e− O(adsorbed)
+ H2
O(adsorbed)
[Back to Top ]
Commercial systems
Commercial systems are more complex relative to the above descriptions. They must be safe,
efficient, and cheap to run. The electrodes must reduce the overpotentials required while keeping
their capital costs low. The electrolytes must be clear of impurities that may poison the electrode
surfaces and usually consist of concentrated alkali or acid. As the efficiency of an electrolyzer
improves with the increased temperature, industrial electrolyzers run warm to hot. The best
electrolyzers operate at 70-80% electricity-to-hydrogen efficiency and produce high-purity (about
99.9%) hydrogen at about one MPa pressure while providing intrinsically safe operation at all times
[3367]. A close collaboration between chemists and engineers is required to develop industrially
relevant catalysts for the hydrogen and oxygen evolution reactions [4314]. There are 2021 reviews
of hydrogen production from solar powered water electrolysis [4350] and from seawater [4360].
Although electrolysis of seawater to make hydrogen gas might seem attractive, it has many
drawbacks. Impurities, such as ions, bacteria, plastics and small particulates, limit the membrane
lifetime, and pH changes cause precipitation and electrode degradation [4437]. Because of these
drawbacks, seawater must be somewhat purified by reverse osmosis before electrolysis, with the
extra cost of this stage being marginal compared with the cost of electrolysis [4438].
[Back to Top ]
Footnotes
a The approximate resistivities of pure water, tap water, and seawater are 18 MΩ ˣ cm, 5 kΩ ˣ cm,
and 20 Ω ˣ cm, respectively. Thus the electrolysis rate is speeded up by factors of about 1000 ˣ or
1,000,000 ˣ using tap water or seawater respectively, rather than pure water. The overpotential is
increased in deuterated water (up to about twice ) and is affected by the ion species present and
electrolyte concentrations [3840]. [Back]
b Traditionally, such equations are written with the electrons on the left-hand side and (however
written) the redox potential refers to so directed equations. Here it is written reversed to show how
the cell reaction is balanced, as this is how the reaction occurs.
O2
(g) + 4 H3
O+(aq) + 4 e− 6 H2
O(l) E° = +1.229 V, pH 0
[Back]
16. c The electrodes should preferably be made from a material with high conductivity, resistance to
corrosion and erosion during the electrolysis, and catalyzing the electrode reactions. Also, for
industrial use, they should be relatively inexpensive. Platinum is an excellent but expensive
electrode material. Industrial cathodes may be made from steel or nickel, and those used as
anodes are metals such as titanium coated with the oxides and mixed oxides of metals such as
nickel and cobalt. Water next to the surface will organize dependent on the surface material [2521]
and considerably reduce their refractive index; the Pockels electro-optical effect [2874]. [Back]
d At the anode, E° = +1.229 − 0.059 pH V. At the cathode E° = −0.059 pH V. The value '0.059' is
derived from the Nernst constant = Loge
(10) ˣ RT/F = 0.059 V (25 °C). [Back]
e Chlor-alkali electrolyzers convert oxygen to hydroxide ions.
O2
(g) + 2H2
O(aq) + 4e− 4OH−
The industrial operating conditions have NaOH concentrations exceeding 10 M, temperatures in the
range of 80°C to 90°C, and current densities of 4 - 6 kA ˣ m−2. [Back]
f Electrolysis was first discovered by Alessandro Volta (1745-1827) with his invention of the battery
in 1799 (The Voltaic pile). A. Volta, On the electricity excited by the mere contact of conducting
substances of different kinds, Philosophical Transactions of the Royal Society, 90 (1800) 403-431.[
Back]
g 'Electrolysis' is the process of being decomposed by the direct action of electricity. [Back]
h Many high-school textbooks give the reactions,
Cathode 2 H+ + 2 e− → H2
Anode 4 OH− − 4e−→ 2 H2
O + O2
to simplify the progressive learning of this complex subject [4167]. [Back]
i The anode is negatively charged for galvanic cells (part of batteries generating voltages) but
positive for electrolytic cells (where the potential is applied to the electrodes). [Back]
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17.
18. There seems to be a BIG PROBLEM with the assumptions being made about the dimensions of the
tube sets and everyone needs to understand this because the absolutely critical gap between
cathode and anode is different in each of the following 3 information sources. The Patents are not
consistent, as one Patent states the INSIDE diameter of the cathode and the other states the
OUTSIDE diameter of the cathode. Also, one Patent has a solid rod anode and the other a tubular
anode. A solid rod anode inside a cathode tube may have different performance characteristics
from an anode tube inside a cathode tube.
The following dimensions are all in INCHES.
1.Don Gable’s sketches for a solid rod anode and tubular cathode:-
a) Cathode outer tube OUTSIDE diameter MEASURED as 0.75,
b) Cathode outer tube wall thickness MEASURED as 0.03,
c) Cathode outer tube INSIDE diameter BY CALCULATION = 0.69 (i.e. 0.75 – (2 x 0.03))
d) Anode SOLID ROD diameter MEASURED as 0.5,
e) Gap between INSIDE surface of cathode tube and the surface of the
SOLID anode rod BY CALCULATION = 0.095 (i.e. (0.69 – 0.5)/2)
2. World Patent WO92/07861 dated 14 May 1992 for a solid rod anode and tubular cathode:-
a) Cathode outer tube OUTSIDE diameter NOT STATED,
b) Cathode outer tube wall thickness NOT STATED and CANNOT BE CALCULATED,
c) Cathode outer tube INSIDE diameter STATED in Patent as 0.75
d) Anode SOLID ROD diameter STATED in Patent as 0.5,
e) Gap between INSIDE surface of cathode tube and the surface of the
SOLID anode rod NOT STATED in Patent but BY CALCULATION = 0.125 (i.e. (0.75 – 0.5) / 2)
3. US Patent US4936961 dated 26 June 1990 for a TUBULAR anode and a tubular cathode:-
a) Cathode outer tube OUTSIDE diameter STATED in Patent as 0.75,
b) Cathode outer tube wall thickness NOT STATED but CALCULATED
Different tube set
configurations, thicknesses
and gaps between the anode
and cathode tubes found in
Stan Meyer's literature
19. as 0.0625 i.e. ((outside diameter of 0.75 – (2 x 0.0625 gap stated in Patent) – 0.5 anode outside
diameter)) / 2 = 0.0625 thickness of cathode tube wall,
c) Cathode outer tube INSIDE diameter NOT STATED but CALCULATED as (0.5 anode diameter +
(2x 0.0625 gap) = 0.625,
d) Anode TUBE outside diameter STATED in Patent as 0.5,
e) Gap between INSIDE surface of cathode tube and the OUTER surface of the
Anode TUBE STATED in Patent as 0.0625
f) Anode TUBE thickness not STATED and cannot be CALCULATED
g) Anode TUBE inside diameter not STATED and cannot be CALCULATED
So we have three different gaps between the electrodes:
Don Gable = 0.095,
World Patent = 0.125 and
US Patent = 0.0625
The World Patent was nearly 2 years after the US Patent.
Obviously the capacitance of these cathode/anode combinations is very different.
Which combination was actually used with the 10 VIC Circuits and the 10 bobbin/5 coil packs that
pulsed the tube sets inside the resonant cavity chamber that provided the HHO for Stan Meyer’s
Dune Buggy engine?
Why did Don Gable measure different dimensions from the World Patent ?
Which combination gives the highest HHO yield at a given voltage with minimum amps?