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1. The document provides a historical overview of water electrolysis from its discovery in 1789 to modern developments. Nicholson and Carlisle were the first to develop the technique in 1800, and by 1902 there were over 400 industrial units in operation.
2. It explains the theory behind water electrolysis, including the chemical reactions that produce hydrogen and oxygen, factors that determine minimum voltage requirements, and sources of inefficiency.
3. Various methods for producing hydrogen through water electrolysis are briefly described, including alkaline electrolysis, proton exchange membrane electrolysis, and producing hydrogen as a byproduct of chloralkali production. Advanced alkaline systems and high-pressure designs are highlighted.
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.
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.
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1. The document provides a historical overview of water electrolysis from its discovery in 1789 to modern developments. Nicholson and Carlisle were the first to develop the technique in 1800, and by 1902 there were over 400 industrial units in operation.
2. It explains the theory behind water electrolysis, including the chemical reactions that produce hydrogen and oxygen, factors that determine minimum voltage requirements, and sources of inefficiency.
3. Various methods for producing hydrogen through water electrolysis are briefly described, including alkaline electrolysis, proton exchange membrane electrolysis, and producing hydrogen as a byproduct of chloralkali production. Advanced alkaline systems and high-pressure designs are highlighted.
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.
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 information about oxidation-reduction (redox) reactions and electrochemical cells. It discusses how redox reactions can be separated into half-reactions that occur at different electrodes. As an example, it examines the redox reaction between copper and silver ions and how this reaction spontaneously occurs when the metals are placed in their respective ion solutions, but not in the reverse direction. It also explains how electrochemical cells like galvanic cells and electrolytic cells use separated half-reactions to generate a voltage or cause non-spontaneous reactions, respectively. Key concepts covered include the standard hydrogen electrode, conventions for assigning electrode potentials, and how voltages relate to the thermodynamics of redox processes.
Electrochemistry 1 the basic of the basicToru Hara
This document discusses key concepts in electrochemistry including the interface between electrode and electrolyte, thermodynamics and kinetics of electrode reactions, and overpotential. The interface contains an electric double layer consisting of an inner monomolecular layer, an outer diffuse region, and an intermediate layer. Overpotential arises from factors like activation energy needed for electrode reactions, concentration gradients that develop at the electrode surface, and resistance of the electrolyte. Overpotential is composed of ohmic drop, activation overpotential, and diffusion overpotential.
1. Electrolytes are substances that dissociate into ions when dissolved in water, allowing them to conduct electricity. They can be classified as strong, weak, or non-electrolytes based on their conductivity.
2. Electrodes are materials inserted into electrolytic cells, and are classified as anodes or cathodes depending on whether oxidation or reduction occurs. During electrolysis, ions migrate to electrodes and undergo chemical reactions.
3. Faraday's laws of electrolysis describe the relationships between electrical charge passed, mass of substance deposited, current over time, and equivalent weights of elements involved in electrolysis reactions.
Electrochemistry is the study of chemical reactions that involve the transfer of electrons between species. Key concepts include redox reactions, electrode potentials, and the Nernst equation. Electrochemical cells harness the energy of spontaneous redox reactions through the movement of electrons in an external circuit and the compensating flow of ions through an electrolyte. The standard cell potential (E°cell) is equal to the sum of the standard reduction potentials of the cathode and anode half-reactions.
This document provides information on electrochemistry and electrochemical cells. It defines electrochemistry as the study of electricity production from spontaneous chemical reactions and use of electrical energy for non-spontaneous reactions. It describes different types of electrochemical cells including galvanic cells that convert chemical to electrical energy and electrolytic cells that do the opposite. Key concepts discussed include electrode potentials, standard hydrogen electrode, Nernst equation, and factors affecting cell potential. Common electrochemical devices like batteries and the corrosion process are also summarized.
This document discusses potentiometry, which is a method of analysis that determines concentration by measuring potential difference between two electrodes without current flow. It describes the principle, reference electrodes like standard hydrogen electrode and saturated calomel electrode, indicator electrodes like glass electrode, and how potentiometric titration can determine the endpoint using methods like the normal titration curve, first derivative curve, and second derivative curve. Potentiometry provides advantages over visual indicator methods by not requiring indicators and allowing the same instrument to be used for different titrations.
Potentiometry is the field of electro-analytical chemistry in which potential is measured without current flow.
It is a method of analysis in which we determine the concentration of solute in solution and the potential difference between two electrodes.
This document discusses electrogravimetry, which is a technique for analyzing substances electrolytically. It defines key terms used in electrogravimetry like cathode, anode, current density, and decomposition potential. It explains Faraday's two laws of electrolysis which relate the mass of a substance deposited to the quantity of electricity passed. The document also discusses concepts like back electromotive force, concentration polarization, and activation overpotential which influence the potential needed for electrolysis beyond the theoretical reversible potential.
This document discusses electrogravimetry, which is the quantitative analysis of substances by electrolysis. It defines key terms used in electrogravimetry like cathode, anode, current density, and overpotential. It explains Faraday's laws of electrolysis and how they relate to the amount of material deposited. It also describes how controlling variables like cathode potential can be used to selectively deposit metals and separate them from each other.
The document discusses several methods for producing hydrogen through water splitting, including:
- Steam reforming of methane, the most common current method.
- Electrolysis, where an electric current splits water into hydrogen and oxygen. More efficient variations include steam electrolysis and thermochemical electrolysis.
- Photochemical and photobiological systems use sunlight to drive the water splitting reaction.
- Thermal water splitting uses very high temperatures of around 1000°C.
- Gasification and biomass conversion also produce hydrogen from other feedstocks.
Low current electrolysis is discussed as a more efficient method, similar to the water splitting that occurs in photosynthesis. Producing hydrogen directly from water without electrolysis is also mentioned. Overall
The document discusses kinetics of electrochemical processes involved in corrosion. It defines corrosion rate (v) as the charge (electrons) transferred per unit time which can be measured as current. The kinetics are governed by the corrosion potential and polarization phenomena. Polarization occurs when the anodic and cathodic reactions are out of equilibrium. Factors like activation overpotential, concentration polarization and ohmic resistance affect the corrosion rate. Tafel and polarization curves are used to determine corrosion kinetics parameters like corrosion potential (Ecorr) and corrosion current density (icorr).
1) Electrolysis is the decomposition of electrolytes by the passage of an electric current through it. During electrolysis, ions move to the oppositely charged electrodes - cations to the cathode and anions to the anode.
2) At the anode, anions lose electrons in an oxidation reaction. At the cathode, cations gain electrons in a reduction reaction.
3) Which ion is discharged depends on factors like their position in the electrochemical series, concentration, and the electrode material. Ions higher in concentration or lower in the series will preferentially discharge.
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
This document discusses four types of electrode polarizations that occur in fuel cells: ohmic, concentration, and activation polarizations at both the anode and cathode. Ohmic polarization is due to the resistance of cell components to ion/electron flow. Concentration polarization occurs when the transport of reactants or products to/from the electrode-electrolyte interface is too slow. Activation polarization is associated with the reaction rate at the electrodes. Electrochemical impedance spectroscopy can be used to measure these polarizations by analyzing the different time dependencies and relaxation processes associated with each polarization type.
The three electrode electrochemical cell consists of a working electrode, reference electrode, and counter electrode immersed in an electrolyte. A potentiostat controls the voltage between the working and reference electrodes while measuring the current between the working and counter electrodes. This allows electrochemical experiments to be performed for applications like environmental monitoring, drug testing, and identifying contaminants. Electrochemical reactions involve the transfer of electrons to or from an electrode surface in oxidation or reduction processes. The rate of electrochemical reactions is affected by factors like the electrode material, mass transport method, solution conditions, and electrical variables.
This document provides an overview of electrochemistry. It begins by defining electrochemistry as the study of chemical reactions at the interface of an electrode and electrolyte involving the interaction of electrical and chemical changes. The document then discusses the history and founders of electrochemistry, including Faraday's two laws of electrolysis. It explains key concepts such as oxidation-reduction reactions, balancing redox equations, and the Nernst equation. The document also covers applications including batteries, corrosion, electrolysis, and branches of electrochemistry like bioelectrochemistry and nanoelectrochemistry.
Depositacion electroforetica dentro de campos electricos moduladosMario ML
This document reviews electrophoretic deposition (EPD) under modulated electric fields such as pulsed direct current (PDC) and alternating current (AC). Classical EPD uses continuous direct current which can lead to issues depositing from aqueous suspensions due to water electrolysis. Modulated electric fields can reduce electrolysis and produce more uniform coatings. PDC and AC offer advantages over continuous DC like reducing bubble formation and particle aggregation. While deposition rates may decrease under modulated fields, they allow for depositing biochemical and biological materials in more active states. The document discusses EPD mechanisms and modulated field types, and their applications including in biotechnology.
Polarography uses a dropping mercury electrode (DME) to measure the current flowing through an electrochemical cell as a function of the applied potential. A polarogram plots this current versus potential and provides qualitative and quantitative information about species undergoing oxidation or reduction reactions. Jaroslav Heyrovsky invented the polarographic method in 1922 and won the Nobel Prize for his contributions to electroanalytical chemistry. All modern voltammetric methods originate from polarography. The DME provides advantages like a reproducible surface area and the ability to form amalgams with metal ions.
Hydrogen has the highest energy content by mass of any fuel and can be used as a substitute for hydrocarbons. It has a non-polluting burning process. There are several methods for producing hydrogen, including electrolysis of water, thermo-chemical processes, and from fossil fuels. Electrolysis uses electricity to split water into hydrogen and oxygen gases. Filter press electrolyzers are most widely used due to their ability to operate at high current densities and production rates. There are challenges to storing hydrogen including its low density and challenges maintaining it as a liquid. Storage methods include high pressure gas, liquid storage using cryogenics, underground storage, and chemically storing it in metal hydrides.
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More Related Content
Similar to detailed-scientific-explanations-of-the-electrolysis-of-water-magnetic-effects-on-water-electromagnetic-effects-on-water-and-the-water-redox-process (1).pdf
This document provides information about oxidation-reduction (redox) reactions and electrochemical cells. It discusses how redox reactions can be separated into half-reactions that occur at different electrodes. As an example, it examines the redox reaction between copper and silver ions and how this reaction spontaneously occurs when the metals are placed in their respective ion solutions, but not in the reverse direction. It also explains how electrochemical cells like galvanic cells and electrolytic cells use separated half-reactions to generate a voltage or cause non-spontaneous reactions, respectively. Key concepts covered include the standard hydrogen electrode, conventions for assigning electrode potentials, and how voltages relate to the thermodynamics of redox processes.
Electrochemistry 1 the basic of the basicToru Hara
This document discusses key concepts in electrochemistry including the interface between electrode and electrolyte, thermodynamics and kinetics of electrode reactions, and overpotential. The interface contains an electric double layer consisting of an inner monomolecular layer, an outer diffuse region, and an intermediate layer. Overpotential arises from factors like activation energy needed for electrode reactions, concentration gradients that develop at the electrode surface, and resistance of the electrolyte. Overpotential is composed of ohmic drop, activation overpotential, and diffusion overpotential.
1. Electrolytes are substances that dissociate into ions when dissolved in water, allowing them to conduct electricity. They can be classified as strong, weak, or non-electrolytes based on their conductivity.
2. Electrodes are materials inserted into electrolytic cells, and are classified as anodes or cathodes depending on whether oxidation or reduction occurs. During electrolysis, ions migrate to electrodes and undergo chemical reactions.
3. Faraday's laws of electrolysis describe the relationships between electrical charge passed, mass of substance deposited, current over time, and equivalent weights of elements involved in electrolysis reactions.
Electrochemistry is the study of chemical reactions that involve the transfer of electrons between species. Key concepts include redox reactions, electrode potentials, and the Nernst equation. Electrochemical cells harness the energy of spontaneous redox reactions through the movement of electrons in an external circuit and the compensating flow of ions through an electrolyte. The standard cell potential (E°cell) is equal to the sum of the standard reduction potentials of the cathode and anode half-reactions.
This document provides information on electrochemistry and electrochemical cells. It defines electrochemistry as the study of electricity production from spontaneous chemical reactions and use of electrical energy for non-spontaneous reactions. It describes different types of electrochemical cells including galvanic cells that convert chemical to electrical energy and electrolytic cells that do the opposite. Key concepts discussed include electrode potentials, standard hydrogen electrode, Nernst equation, and factors affecting cell potential. Common electrochemical devices like batteries and the corrosion process are also summarized.
This document discusses potentiometry, which is a method of analysis that determines concentration by measuring potential difference between two electrodes without current flow. It describes the principle, reference electrodes like standard hydrogen electrode and saturated calomel electrode, indicator electrodes like glass electrode, and how potentiometric titration can determine the endpoint using methods like the normal titration curve, first derivative curve, and second derivative curve. Potentiometry provides advantages over visual indicator methods by not requiring indicators and allowing the same instrument to be used for different titrations.
Potentiometry is the field of electro-analytical chemistry in which potential is measured without current flow.
It is a method of analysis in which we determine the concentration of solute in solution and the potential difference between two electrodes.
This document discusses electrogravimetry, which is a technique for analyzing substances electrolytically. It defines key terms used in electrogravimetry like cathode, anode, current density, and decomposition potential. It explains Faraday's two laws of electrolysis which relate the mass of a substance deposited to the quantity of electricity passed. The document also discusses concepts like back electromotive force, concentration polarization, and activation overpotential which influence the potential needed for electrolysis beyond the theoretical reversible potential.
This document discusses electrogravimetry, which is the quantitative analysis of substances by electrolysis. It defines key terms used in electrogravimetry like cathode, anode, current density, and overpotential. It explains Faraday's laws of electrolysis and how they relate to the amount of material deposited. It also describes how controlling variables like cathode potential can be used to selectively deposit metals and separate them from each other.
The document discusses several methods for producing hydrogen through water splitting, including:
- Steam reforming of methane, the most common current method.
- Electrolysis, where an electric current splits water into hydrogen and oxygen. More efficient variations include steam electrolysis and thermochemical electrolysis.
- Photochemical and photobiological systems use sunlight to drive the water splitting reaction.
- Thermal water splitting uses very high temperatures of around 1000°C.
- Gasification and biomass conversion also produce hydrogen from other feedstocks.
Low current electrolysis is discussed as a more efficient method, similar to the water splitting that occurs in photosynthesis. Producing hydrogen directly from water without electrolysis is also mentioned. Overall
The document discusses kinetics of electrochemical processes involved in corrosion. It defines corrosion rate (v) as the charge (electrons) transferred per unit time which can be measured as current. The kinetics are governed by the corrosion potential and polarization phenomena. Polarization occurs when the anodic and cathodic reactions are out of equilibrium. Factors like activation overpotential, concentration polarization and ohmic resistance affect the corrosion rate. Tafel and polarization curves are used to determine corrosion kinetics parameters like corrosion potential (Ecorr) and corrosion current density (icorr).
1) Electrolysis is the decomposition of electrolytes by the passage of an electric current through it. During electrolysis, ions move to the oppositely charged electrodes - cations to the cathode and anions to the anode.
2) At the anode, anions lose electrons in an oxidation reaction. At the cathode, cations gain electrons in a reduction reaction.
3) Which ion is discharged depends on factors like their position in the electrochemical series, concentration, and the electrode material. Ions higher in concentration or lower in the series will preferentially discharge.
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
This document discusses four types of electrode polarizations that occur in fuel cells: ohmic, concentration, and activation polarizations at both the anode and cathode. Ohmic polarization is due to the resistance of cell components to ion/electron flow. Concentration polarization occurs when the transport of reactants or products to/from the electrode-electrolyte interface is too slow. Activation polarization is associated with the reaction rate at the electrodes. Electrochemical impedance spectroscopy can be used to measure these polarizations by analyzing the different time dependencies and relaxation processes associated with each polarization type.
The three electrode electrochemical cell consists of a working electrode, reference electrode, and counter electrode immersed in an electrolyte. A potentiostat controls the voltage between the working and reference electrodes while measuring the current between the working and counter electrodes. This allows electrochemical experiments to be performed for applications like environmental monitoring, drug testing, and identifying contaminants. Electrochemical reactions involve the transfer of electrons to or from an electrode surface in oxidation or reduction processes. The rate of electrochemical reactions is affected by factors like the electrode material, mass transport method, solution conditions, and electrical variables.
This document provides an overview of electrochemistry. It begins by defining electrochemistry as the study of chemical reactions at the interface of an electrode and electrolyte involving the interaction of electrical and chemical changes. The document then discusses the history and founders of electrochemistry, including Faraday's two laws of electrolysis. It explains key concepts such as oxidation-reduction reactions, balancing redox equations, and the Nernst equation. The document also covers applications including batteries, corrosion, electrolysis, and branches of electrochemistry like bioelectrochemistry and nanoelectrochemistry.
Depositacion electroforetica dentro de campos electricos moduladosMario ML
This document reviews electrophoretic deposition (EPD) under modulated electric fields such as pulsed direct current (PDC) and alternating current (AC). Classical EPD uses continuous direct current which can lead to issues depositing from aqueous suspensions due to water electrolysis. Modulated electric fields can reduce electrolysis and produce more uniform coatings. PDC and AC offer advantages over continuous DC like reducing bubble formation and particle aggregation. While deposition rates may decrease under modulated fields, they allow for depositing biochemical and biological materials in more active states. The document discusses EPD mechanisms and modulated field types, and their applications including in biotechnology.
Polarography uses a dropping mercury electrode (DME) to measure the current flowing through an electrochemical cell as a function of the applied potential. A polarogram plots this current versus potential and provides qualitative and quantitative information about species undergoing oxidation or reduction reactions. Jaroslav Heyrovsky invented the polarographic method in 1922 and won the Nobel Prize for his contributions to electroanalytical chemistry. All modern voltammetric methods originate from polarography. The DME provides advantages like a reproducible surface area and the ability to form amalgams with metal ions.
Hydrogen has the highest energy content by mass of any fuel and can be used as a substitute for hydrocarbons. It has a non-polluting burning process. There are several methods for producing hydrogen, including electrolysis of water, thermo-chemical processes, and from fossil fuels. Electrolysis uses electricity to split water into hydrogen and oxygen gases. Filter press electrolyzers are most widely used due to their ability to operate at high current densities and production rates. There are challenges to storing hydrogen including its low density and challenges maintaining it as a liquid. Storage methods include high pressure gas, liquid storage using cryogenics, underground storage, and chemically storing it in metal hydrides.
Similar to detailed-scientific-explanations-of-the-electrolysis-of-water-magnetic-effects-on-water-electromagnetic-effects-on-water-and-the-water-redox-process (1).pdf (20)
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
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
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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
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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
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1. 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
2. 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)
3. 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.
4. 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.
5. 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,
6. 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
7. 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,
8. 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,
9. 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]
10. 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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11. Revision #3
Created 14 June 2023 04:53:55 by Paul Butcher
Updated 14 June 2023 05:52:37 by Paul Butcher