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.
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.
This document provides an overview of electrochemistry and galvanic cells. Some key points:
- Electron transfer reactions are oxidation-reduction (redox) reactions that can generate electric current or be driven by an applied current, making it the field of electrochemistry.
- Galvanic cells use spontaneous redox reactions to generate electricity, with oxidation occurring at the anode and reduction at the cathode. The potential difference between electrodes is called the cell voltage.
- Standard electrode potentials (E°) describe the tendency of half-reactions to occur and can be used to predict spontaneity of redox reactions in cells. Nernst equation relates cell potential to concentrations.
IB Chemistry on Lewis structure, ionic and covalent bondingLawrence kok
1. The document discusses different types of bonding including ionic bonding, covalent bonding, and dative bonding.
2. Ionic bonding occurs between metals and nonmetals where the metal loses electrons and forms cations and the nonmetal gains electrons and forms anions.
3. Covalent bonding occurs between nonmetals where electrons are shared between atoms to achieve stable octets. Single, double, and triple covalent bonds are formed by sharing one, two, or three electron pairs.
Electrochemistry deals with the interconversion of electrical and chemical energy through redox reactions. A galvanic or voltaic cell produces electricity through the spontaneous chemical change that occurs within it. A voltaic cell works because of the different abilities of materials like metals to donate or accept electrons, and the ability of electrons to flow through an external circuit. An example is a zinc-copper cell, where zinc oxidizes and copper reduces, driving electrons through the circuit.
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
This document discusses different types of diagrams used to represent the reactivity and stability of chemical species, including Latimer diagrams, Pourbaix diagrams, and Frost diagrams. Latimer diagrams show the standard reduction potential of oxidation states of an element using a horizontal line with values above it. Pourbaix diagrams represent the stability of a metal as a function of potential and pH, with different lines and boundaries indicating acid-base, redox, and solubility equilibria. Frost diagrams depict the free energy versus oxidation state of an element and can be constructed from Latimer diagrams. These diagramming methods provide a way to understand and predict the behavior of elements under various conditions.
This document discusses redox reactions and electrochemistry. It covers topics such as oxidation numbers, galvanic cells, cell notation, standard electrode potentials, and how cell potential relates to Gibbs free energy and equilibrium constants. It also discusses corrosion, batteries, fuel cells, and the differences between voltaic, electrolytic, and fuel cells. Redox reactions allow the interconversion of electrical and chemical energy.
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.
This document provides an overview of electrochemistry and galvanic cells. Some key points:
- Electron transfer reactions are oxidation-reduction (redox) reactions that can generate electric current or be driven by an applied current, making it the field of electrochemistry.
- Galvanic cells use spontaneous redox reactions to generate electricity, with oxidation occurring at the anode and reduction at the cathode. The potential difference between electrodes is called the cell voltage.
- Standard electrode potentials (E°) describe the tendency of half-reactions to occur and can be used to predict spontaneity of redox reactions in cells. Nernst equation relates cell potential to concentrations.
IB Chemistry on Lewis structure, ionic and covalent bondingLawrence kok
1. The document discusses different types of bonding including ionic bonding, covalent bonding, and dative bonding.
2. Ionic bonding occurs between metals and nonmetals where the metal loses electrons and forms cations and the nonmetal gains electrons and forms anions.
3. Covalent bonding occurs between nonmetals where electrons are shared between atoms to achieve stable octets. Single, double, and triple covalent bonds are formed by sharing one, two, or three electron pairs.
Electrochemistry deals with the interconversion of electrical and chemical energy through redox reactions. A galvanic or voltaic cell produces electricity through the spontaneous chemical change that occurs within it. A voltaic cell works because of the different abilities of materials like metals to donate or accept electrons, and the ability of electrons to flow through an external circuit. An example is a zinc-copper cell, where zinc oxidizes and copper reduces, driving electrons through the circuit.
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
This document discusses different types of diagrams used to represent the reactivity and stability of chemical species, including Latimer diagrams, Pourbaix diagrams, and Frost diagrams. Latimer diagrams show the standard reduction potential of oxidation states of an element using a horizontal line with values above it. Pourbaix diagrams represent the stability of a metal as a function of potential and pH, with different lines and boundaries indicating acid-base, redox, and solubility equilibria. Frost diagrams depict the free energy versus oxidation state of an element and can be constructed from Latimer diagrams. These diagramming methods provide a way to understand and predict the behavior of elements under various conditions.
This document discusses redox reactions and electrochemistry. It covers topics such as oxidation numbers, galvanic cells, cell notation, standard electrode potentials, and how cell potential relates to Gibbs free energy and equilibrium constants. It also discusses corrosion, batteries, fuel cells, and the differences between voltaic, electrolytic, and fuel cells. Redox reactions allow the interconversion of electrical and chemical energy.
This document outlines two methods for balancing redox reactions: the half-reaction method and the oxidation number method. The half-reaction method involves writing the oxidation and reduction as separate half-reactions, balancing atoms and charge separately, and combining. The oxidation number method assigns oxidation numbers and uses bracketing lines to connect oxidized and reduced atoms before balancing the overall change in oxidation. Both methods aim to balance the atoms and charges in redox reactions.
1. Electrochemistry involves electron transfer between chemical species in oxidation-reduction reactions.
2. Oxidation and reduction half-reactions can be balanced using the half-reaction method and combined to give the overall redox reaction.
3. Voltaic cells harness the energy of spontaneous redox reactions by allowing electrons to flow through an external circuit, and cell potential depends on the relative reduction potentials of the half-reactions.
B.tech. ii engineering chemistry unit 5 A electrochemistryRai University
Arrhenius proposed the theory of electrolytic dissociation to explain the properties of electrolytic solutions. The theory states that when an electrolyte dissolves in water, it breaks up into ions - positively charged cations and negatively charged anions. This process is called ionization. Ions are constantly recombining and dissociating, reaching a state of dynamic equilibrium. The extent of ionization depends on an equilibrium constant. Strong electrolytes have a high equilibrium constant and ionize completely, while weak electrolytes have a low constant and only partially ionize.
1. Electrochemistry examines phenomena resulting from combined chemical and electrical effects. It covers electrolytic and galvanic processes.
2. An electrochemical cell consists of two electrodes and an electrolyte. Charge is transported by electron motion in electrodes and ion motion in electrolytes.
3. At each electrode, an oxidation or reduction half-cell reaction occurs. The overall reaction is the sum of the half reactions. Thermodynamics predicts which reaction will occur as oxidation or reduction.
Option C Nernst Equation, Voltaic Cell and Concentration CellLawrence kok
This document provides a tutorial on voltaic cells, the Nernst equation, and concentration cells. It discusses the basic components and workings of voltaic cells, including the conversion of chemical energy to electrical energy through redox reactions. Equations for cell potential (Ecell), standard electrode potential (E°), and the Nernst equation are presented. Examples of specific voltaic cells like Daniell cells and their cell potentials are provided. The relationships between Gibbs free energy (ΔG), equilibrium constant (Kc), and cell potential are also summarized.
IB Chemistry on Redox Design and Nernst EquationLawrence kok
The document outlines research questions and procedures to investigate the effect of various factors on the emf and current of voltaic cells. Specifically, it will study how concentration, temperature, electrode size, salt bridge composition, and metal pairs affect measurements in zinc-copper and copper-copper cells. Tests will be conducted by varying one factor at a time while keeping others standard, and measuring the resulting emf and current.
This document contains tables of chemical data relevant to the Cambridge Pre-U Revised Syllabus in Chemistry. It includes tables with important physical constants, ionization energies of elements, bond energies, standard electrode potentials, atomic and ionic radii, and other reference data. The tables provide numerical data on topics like elemental properties, redox potentials, spectroscopic data, and more. This document serves as a compilation of reference material for students in an organized and easy-to-access manner.
This document discusses concentration cells, which generate electricity from differences in concentration between two solutions. There are two types of concentration cells: electrode concentration cells, which use electrodes of different concentrations, and electrolyte concentration cells, which use solutions of different concentrations. Electrolyte concentration cells can be further divided into those with and without transference, depending on whether the solutions are separated or in direct contact, allowing ion transfer between solutions. The electromotive force and liquid junction potential of concentration cells are also explained in terms of ion activities and transport numbers.
Half-reactions indicate the mole ratio of electrons to ions involved in redox reactions. The document discusses how current, time, and charge are related based on 1 mole of electrons equating to 96,500 Coulombs of charge. It provides examples calculating the mass of copper produced from electrolysis and the volume of chlorine gas produced from an industrial electrolysis process based on given values of current and time.
IB Chemistry on Voltaic Cell, Standard Electrode Potential and Standard Hydro...Lawrence kok
This document discusses voltaic cells and the potential differences between half-cells. It explains that connecting two half-cells with different electrode potentials through an external circuit and salt bridge allows electrons to flow spontaneously from the negative half-cell to the positive half-cell. Specifically, it gives the example of a Zn/Cu voltaic cell, where the Zn half-cell acts as the anode undergoing oxidation and the Cu half-cell acts as the cathode undergoing reduction. When connected, the potential difference between the half-cells can be measured as 1.10 volts using a high resistance voltmeter.
The document discusses different types of electrochemical cells including primary cells that produce electricity from non-reversible chemical reactions and secondary cells that can be recharged by passing electricity in the opposite direction of the spontaneous reaction. Examples of primary cells discussed include Daniel, mercury, dry, and alkaline cells, while examples of secondary cells include lead-acid, nickel-cadmium, nickel-metal hydride, and lithium-ion batteries. The working and reactions of common cells like lead-acid, alkaline, and dry cells are also explained.
IB Chemistry on Standard Reduction Potential, Standard Hydrogen Electrode and...Lawrence kok
The document discusses standard electrode potentials and how they are measured. It explains that the standard hydrogen electrode is used as a reference with a potential of 0 V. Other half-cell potentials are measured against this to determine their standard electrode potential. Common half-cells include metal/metal ion, gas/ion, and ion/ion systems. Standard conditions of 1 M concentrations, 1 atm pressure, and 298K temperature must be used. The potentials of zinc/zinc ion, iron III/iron II, and chlorine/chloride ion half-cells are given as examples.
Effect of Concentration Changes on Cell Potentialwewwchemistry
[ Visit http://www.wewwchemistry.com ] This example uses the Nernst equation to illustrate how changes in reactant or product concentration (effected by dilution) affect cell potentials.
This document discusses isotopes and atomic mass. It defines isotopes as atoms of the same element that have different mass numbers due to varying numbers of neutrons, while having the same number of protons. A nuclear symbol represents a particular atom and gives the mass number and atomic number. Examples are provided of writing nuclear symbols for atoms with given subatomic particle numbers. The document also discusses that the atomic mass of an element is based on the naturally occurring isotopes and their abundances, and is the weighted average mass compared to carbon-12. Examples of isotopes and atomic masses are provided for several elements.
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 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
The document discusses an electrochemistry unit covering chemical reactions that produce electrical currents or voltages. It provides information on voltaic cells, also known as galvanic cells, which harness spontaneous redox reactions to generate electricity. The document explains that voltaic cells use two half-reactions, an oxidation reaction at the anode and a reduction reaction at the cathode, to drive electrons from the anode to the cathode through an external circuit. Standard reduction potentials are used to predict if reactions will occur spontaneously.
This document provides an overview of redox reactions and electrochemistry applications. It discusses oxidation-reduction concepts like oxidation states and the OIL RIG mnemonic. Examples of redox reactions and electrochemistry applications are given, including galvanic cells, corrosion, electrolysis, and batteries. Key concepts covered include cell potential, the Nernst equation, and how concentration affects cell potential. Diagrams illustrate galvanic cells and how they function.
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.
This document outlines two methods for balancing redox reactions: the half-reaction method and the oxidation number method. The half-reaction method involves writing the oxidation and reduction as separate half-reactions, balancing atoms and charge separately, and combining. The oxidation number method assigns oxidation numbers and uses bracketing lines to connect oxidized and reduced atoms before balancing the overall change in oxidation. Both methods aim to balance the atoms and charges in redox reactions.
1. Electrochemistry involves electron transfer between chemical species in oxidation-reduction reactions.
2. Oxidation and reduction half-reactions can be balanced using the half-reaction method and combined to give the overall redox reaction.
3. Voltaic cells harness the energy of spontaneous redox reactions by allowing electrons to flow through an external circuit, and cell potential depends on the relative reduction potentials of the half-reactions.
B.tech. ii engineering chemistry unit 5 A electrochemistryRai University
Arrhenius proposed the theory of electrolytic dissociation to explain the properties of electrolytic solutions. The theory states that when an electrolyte dissolves in water, it breaks up into ions - positively charged cations and negatively charged anions. This process is called ionization. Ions are constantly recombining and dissociating, reaching a state of dynamic equilibrium. The extent of ionization depends on an equilibrium constant. Strong electrolytes have a high equilibrium constant and ionize completely, while weak electrolytes have a low constant and only partially ionize.
1. Electrochemistry examines phenomena resulting from combined chemical and electrical effects. It covers electrolytic and galvanic processes.
2. An electrochemical cell consists of two electrodes and an electrolyte. Charge is transported by electron motion in electrodes and ion motion in electrolytes.
3. At each electrode, an oxidation or reduction half-cell reaction occurs. The overall reaction is the sum of the half reactions. Thermodynamics predicts which reaction will occur as oxidation or reduction.
Option C Nernst Equation, Voltaic Cell and Concentration CellLawrence kok
This document provides a tutorial on voltaic cells, the Nernst equation, and concentration cells. It discusses the basic components and workings of voltaic cells, including the conversion of chemical energy to electrical energy through redox reactions. Equations for cell potential (Ecell), standard electrode potential (E°), and the Nernst equation are presented. Examples of specific voltaic cells like Daniell cells and their cell potentials are provided. The relationships between Gibbs free energy (ΔG), equilibrium constant (Kc), and cell potential are also summarized.
IB Chemistry on Redox Design and Nernst EquationLawrence kok
The document outlines research questions and procedures to investigate the effect of various factors on the emf and current of voltaic cells. Specifically, it will study how concentration, temperature, electrode size, salt bridge composition, and metal pairs affect measurements in zinc-copper and copper-copper cells. Tests will be conducted by varying one factor at a time while keeping others standard, and measuring the resulting emf and current.
This document contains tables of chemical data relevant to the Cambridge Pre-U Revised Syllabus in Chemistry. It includes tables with important physical constants, ionization energies of elements, bond energies, standard electrode potentials, atomic and ionic radii, and other reference data. The tables provide numerical data on topics like elemental properties, redox potentials, spectroscopic data, and more. This document serves as a compilation of reference material for students in an organized and easy-to-access manner.
This document discusses concentration cells, which generate electricity from differences in concentration between two solutions. There are two types of concentration cells: electrode concentration cells, which use electrodes of different concentrations, and electrolyte concentration cells, which use solutions of different concentrations. Electrolyte concentration cells can be further divided into those with and without transference, depending on whether the solutions are separated or in direct contact, allowing ion transfer between solutions. The electromotive force and liquid junction potential of concentration cells are also explained in terms of ion activities and transport numbers.
Half-reactions indicate the mole ratio of electrons to ions involved in redox reactions. The document discusses how current, time, and charge are related based on 1 mole of electrons equating to 96,500 Coulombs of charge. It provides examples calculating the mass of copper produced from electrolysis and the volume of chlorine gas produced from an industrial electrolysis process based on given values of current and time.
IB Chemistry on Voltaic Cell, Standard Electrode Potential and Standard Hydro...Lawrence kok
This document discusses voltaic cells and the potential differences between half-cells. It explains that connecting two half-cells with different electrode potentials through an external circuit and salt bridge allows electrons to flow spontaneously from the negative half-cell to the positive half-cell. Specifically, it gives the example of a Zn/Cu voltaic cell, where the Zn half-cell acts as the anode undergoing oxidation and the Cu half-cell acts as the cathode undergoing reduction. When connected, the potential difference between the half-cells can be measured as 1.10 volts using a high resistance voltmeter.
The document discusses different types of electrochemical cells including primary cells that produce electricity from non-reversible chemical reactions and secondary cells that can be recharged by passing electricity in the opposite direction of the spontaneous reaction. Examples of primary cells discussed include Daniel, mercury, dry, and alkaline cells, while examples of secondary cells include lead-acid, nickel-cadmium, nickel-metal hydride, and lithium-ion batteries. The working and reactions of common cells like lead-acid, alkaline, and dry cells are also explained.
IB Chemistry on Standard Reduction Potential, Standard Hydrogen Electrode and...Lawrence kok
The document discusses standard electrode potentials and how they are measured. It explains that the standard hydrogen electrode is used as a reference with a potential of 0 V. Other half-cell potentials are measured against this to determine their standard electrode potential. Common half-cells include metal/metal ion, gas/ion, and ion/ion systems. Standard conditions of 1 M concentrations, 1 atm pressure, and 298K temperature must be used. The potentials of zinc/zinc ion, iron III/iron II, and chlorine/chloride ion half-cells are given as examples.
Effect of Concentration Changes on Cell Potentialwewwchemistry
[ Visit http://www.wewwchemistry.com ] This example uses the Nernst equation to illustrate how changes in reactant or product concentration (effected by dilution) affect cell potentials.
This document discusses isotopes and atomic mass. It defines isotopes as atoms of the same element that have different mass numbers due to varying numbers of neutrons, while having the same number of protons. A nuclear symbol represents a particular atom and gives the mass number and atomic number. Examples are provided of writing nuclear symbols for atoms with given subatomic particle numbers. The document also discusses that the atomic mass of an element is based on the naturally occurring isotopes and their abundances, and is the weighted average mass compared to carbon-12. Examples of isotopes and atomic masses are provided for several elements.
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 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
The document discusses an electrochemistry unit covering chemical reactions that produce electrical currents or voltages. It provides information on voltaic cells, also known as galvanic cells, which harness spontaneous redox reactions to generate electricity. The document explains that voltaic cells use two half-reactions, an oxidation reaction at the anode and a reduction reaction at the cathode, to drive electrons from the anode to the cathode through an external circuit. Standard reduction potentials are used to predict if reactions will occur spontaneously.
This document provides an overview of redox reactions and electrochemistry applications. It discusses oxidation-reduction concepts like oxidation states and the OIL RIG mnemonic. Examples of redox reactions and electrochemistry applications are given, including galvanic cells, corrosion, electrolysis, and batteries. Key concepts covered include cell potential, the Nernst equation, and how concentration affects cell potential. Diagrams illustrate galvanic cells and how they function.
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.
This document provides information about electrochemical cells. It begins by defining an electrochemical cell as consisting of two electrodes in contact with an electrolyte, with each electrode and electrolyte comprising an electrode compartment. It describes the two main types of electrochemical cells - electrolytic cells, where an external current causes non-spontaneous oxidation and reduction, and galvanic cells, where a spontaneous chemical reaction produces electricity. It then discusses standard reduction potentials, cell potentials, the Nernst equation, types of electrodes, and methods for determining standard electrode potentials, free energy changes, and equilibrium constants from cell potentials.
Redox titrations involve the reaction of an oxidizing titrant with a reducing analyte. The document discusses the redox reaction of permanganate (KMnO4) with iron in an ore sample. It also provides the theory behind redox titrations, including Nernst equations that describe the cell potential before, at, and after the equivalence point of the titration reaction. The equivalence point potential can be calculated from the standard reduction potentials of the oxidized and reduced forms of the analyte and titrant.
The document defines redox reactions as electron transfer reactions where oxidation is the loss of electrons and reduction is the gain of electrons. It provides examples of writing half reactions and ionic equations for redox reactions involving magnesium and hydrochloric acid. Practice problems are included for identifying oxidizing and reducing agents and writing half reactions and overall ionic equations.
The document defines redox reactions as reactions involving the transfer of electrons between reactants. It defines oxidation as the loss of electrons and reduction as the gain of electrons. It provides examples of writing half reactions and ionic equations for redox reactions involving magnesium and hydrochloric acid.
The document defines redox reactions as electron transfer reactions where oxidation is the loss of electrons and reduction is the gain of electrons. It provides examples of writing half reactions and ionic equations for redox reactions involving magnesium and hydrochloric acid. Practice problems are included for identifying oxidizing and reducing agents and writing half reactions and overall ionic equations.
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 an overview of electrochemistry and voltaic cells. It discusses redox reactions, how to balance redox reactions using the half-reaction method, and the components and operation of voltaic cells. Specifically, it explains that a voltaic cell uses a spontaneous redox reaction to generate electrical energy by separating the oxidation and reduction half-reactions into two half-cells connected by an external circuit and salt bridge. Electrons flow from the anode, where oxidation occurs, through the external circuit to the cathode, where reduction occurs.
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 oxidation-reduction (redox) reactions. It defines oxidation as the loss of electrons and reduction as the gain of electrons. Redox reactions always involve a transfer of electrons between reactants. The key principles are:
1. Oxidation and reduction always occur together in a redox reaction.
2. The total number of electrons lost must equal the total number of electrons gained to satisfy the conservation of charge.
3. Redox titrations can be used to determine the concentration of an unknown substance and rely on a redox reaction between a titrant and analyte with an indicator or potentiometer used to find the endpoint.
Electrochemistry involves the study of electricity produced from spontaneous chemical reactions in galvanic cells and the use of electricity to drive non-spontaneous reactions in electrolytic cells. Galvanic cells produce electricity through spontaneous redox reactions, with oxidation occurring at the anode and reduction at the cathode. Electrolytic cells use electricity to carry out non-spontaneous reactions. The potential difference between electrodes in a galvanic cell is called the cell potential, which can be calculated using standard electrode potentials and concentrations based on the Nernst equation.
Electrochemistry is the branch of chemistry that studies chemical reactions which involve charge transfer between electrodes and electrolytes. Key aspects include:
- The conversion of electrical energy to chemical energy in electrolytic cells where non-spontaneous redox reactions are driven by an external power source.
- The conversion of chemical energy to electrical energy in galvanic/voltaic cells where a spontaneous redox reaction generates an electric current.
2. What are some common types of electrochemical cells?
Some common types of electrochemical cells include:
- Galvanic/voltaic cells such as batteries, which harness the spontaneous redox reaction between two half-cells to generate a voltage. Examples include zinc-
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.
Similar to Ch17z5eelectrochem 110115232747-phpapp02 (20)
This document discusses suffixes and terminology used in medicine. It begins by listing common combining forms used to build medical terms and their meanings. It then defines several noun, adjective, and shorter suffixes and provides their meanings. Examples are given of medical terms built using combining forms and suffixes. The document also examines specific medical concepts in more depth, such as hernias, blood cells, acromegaly, splenomegaly, and laparoscopy.
The document is a chapter from a medical textbook that discusses anatomical terminology pertaining to the body as a whole. It defines the structural organization of the body from cells to tissues to organs to systems. It also describes the body cavities and identifies the major organs contained within each cavity, as well as anatomical divisions of the abdomen and back.
This document is from a textbook on medical terminology. It discusses the basic structure of medical words and how they are built from prefixes, suffixes, and combining forms. Some key points:
- Medical terms are made up of elements including roots, suffixes, prefixes, and combining vowels. Understanding these elements is important for analyzing terms.
- Common prefixes include hypo-, epi-, and cis-. Common suffixes include -itis, -algia, and -ectomy.
- Dozens of combining forms are provided, such as gastro- meaning stomach, cardi- meaning heart, and aden- meaning gland.
- Rules are provided for analyzing terms, such as reading from the suffix backward and dropping combining vowels before suffixes starting with vowels
This document is the copyright information for Chapter 25 on Cancer from the 6th edition of the textbook Molecular Cell Biology published in 2008 by W. H. Freeman and Company. The chapter was authored by a team that includes Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This document is the copyright information for Chapter 24 on Immunology from the 6th edition of the textbook Molecular Cell Biology published in 2008 by W. H. Freeman and Company. The chapter was authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
Nerve cells, also known as neurons, are highly specialized cells that process and transmit information through electrical and chemical signals. This chapter discusses the structure and function of neurons, how they communicate with each other via synapses, and how signals are propagated along neurons through changes in their membrane potentials. Neurons play a vital role in the nervous system by allowing organisms to process information and coordinate their responses.
This document is the copyright information for Chapter 22 from the 6th edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "The Molecular Cell Biology of Development" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This document is the copyright information for Chapter 21 from the sixth edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Cell Birth, Lineage, and Death" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This document is the copyright page for Chapter 20 from the 6th edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Regulating the Eukaryotic Cell Cycle" and is authored by a group of scientists including Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This document is the copyright information for Chapter 19 from the 6th edition textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Integrating Cells into Tissues" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This chapter discusses microtubules and intermediate filaments, which are types of cytoskeletal filaments that help organize and move cellular components. Microtubules are involved in processes like cell division and intracellular transport, while intermediate filaments provide mechanical strength and help integrate the nucleus with the cytoplasm. Together, these filaments play important structural and functional roles in eukaryotic cells.
This chapter discusses microfilaments, which are one of the three main types of cytoskeletal filaments found in eukaryotic cells. Microfilaments are composed of actin filaments and play important roles in cell motility, structure, and intracellular transport. They allow cells to change shape and to move by contracting or extending parts of the cell surface.
This document is the copyright page for Chapter 16 from the 6th edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Signaling Pathways that Control Gene Activity" and is authored by a group of scientists including Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh and Matsudaira.
This document is the copyright page for Chapter 15 of the 6th edition textbook "Molecular Cell Biology" by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira. It provides the chapter title "Cell Signaling I: Signal Transduction and Short-Term Cellular Responses" and notes the copyright is held by W. H. Freeman and Company in 2008.
This document is the copyright page for Chapter 14 from the 6th edition textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Vesicular Traffic, Secretion, and Endocytosis" and is authored by a group of scientists including Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh and Matsudaira.
This chapter discusses how proteins are transported into membranes and organelles within cells. Proteins destined for membranes or organelles have targeting signals that are recognized by transport systems. The transport systems then direct the proteins to their proper destinations, such as inserting membrane proteins into membranes or delivering soluble proteins into organelles.
This document is the copyright information for Chapter 12 from the sixth edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Cellular Energetics" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
This chapter discusses the transmembrane transport of ions and small molecules across cell membranes. It covers topics such as passive transport through membrane channels and pumps, as well as active transport using ATP. The chapter is from the 6th edition of the textbook Molecular Cell Biology and is copyrighted by W. H. Freeman and Company in 2008.
This document is the copyright information for Chapter 10, titled "Biomembrane Structure", from the sixth edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter was written by a team of authors including Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh and Matsudaira.
This document is the copyright information for Chapter 9 from the 6th edition of the textbook "Molecular Cell Biology" published in 2008 by W. H. Freeman and Company. The chapter is titled "Visualizing, Fractionating, and Culturing Cells" and is authored by Lodish, Berk, Kaiser, Krieger, Scott, Bretscher, Ploegh, and Matsudaira.
ISO/IEC 27001, ISO/IEC 42001, and GDPR: Best Practices for Implementation and...PECB
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Leveraging Generative AI to Drive Nonprofit InnovationTechSoup
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How to Make a Field Mandatory in Odoo 17Celine George
In Odoo, making a field required can be done through both Python code and XML views. When you set the required attribute to True in Python code, it makes the field required across all views where it's used. Conversely, when you set the required attribute in XML views, it makes the field required only in the context of that particular view.
2. 2
17.1 Galvanic Cells
Oxidation reduction reactions involve a
transfer of electrons.
OIL- RIG
Oxidation Involves Loss
Reduction Involves Gain
LEO-GER
Lose Electrons Oxidation
Gain Electrons Reduction
3. 3
Applications
Moving electrons = electric current.
8H+
+MnO4
-
+ 5Fe+2
→ Mn+2
+ 5Fe+3
+4H2O
It helps to break the reactions into half
reactions.
8H+
+MnO4
-
+5e-
→ Mn+2
+4H2O
5(Fe+2
→ Fe+3
+ e-
)
In the same mixture this happens without
doing useful work, but if separate . . .
10. 10
Cell Potential
Oxidizing agent pulls the electron.
Reducing agent pushes the electron.
The push or pull (“driving force”) is called
the cell potential Ecell
Also called the electromotive force (emf).
Unit is the volt (V).
= 1 joule of work per coulomb of charge
transferred (1 V = 1 J/C).
Measured with a voltmeter.
11. 11
17.2 Standard Reduction Potentials
The reaction in a galvanic cell is a redox
reaction.
So, break it down into two half-reactions.
Assign a potential to each.
Sum the half-cell potentials to get the
overall cell potential.
“Active anodes” - the more active metal is
always the anode (good for multiple
choice questions).
13. 13
1 M HCl
H+
Cl-
H2 in
Standard Hydrogen Electrode
This is the reference
all other oxidations
are compared to.
Eº
= 0
º indicates standard
states of 25ºC,
1 atm, 1 M
solutions.
14. 14
Z5e 841 Figure 17.5: Zn/H Galvanic Cell.
Notice the electron flow also.
15. 15
Cell Potential
Zn(s) + Cu+2
(aq)→ Zn+2
(aq) + Cu(s)
The total cell potential is the sum of the potential
at each electrode.
Eº cell = EºZn→ Zn+2 + Eº Cu+2 →Cu
We can look up reduction potentials in a table
(see, p. 796).
For Br, s/b 1.07 (not 1.09) need for online HW!!)
Since one of the 1/2 reactions is oxidation its
table value must be reversed, so change its sign.
17. 17
Cell Potential pp
Determine the cell potential for a galvanic
cell based on the redox reaction . . .
Cu(s) + Fe+3
(aq)→ Cu+2
(aq) + Fe+2
(aq)
steps follow
Fe+3
(aq)+ e-
→ Fe+2
(aq) Eº = 0.77 V
Cu+2
(aq)+2e-
→ Cu(s) Eº = 0.34 V
Since one of these must be oxidation,
one of them needs to be reversed.
Which one?
18. 18
Cell Potential pp
Cu(s) + Fe+3
(aq)→ Cu+2
(aq) + Fe+2
(aq)
This is spontaneous only if Eºcell = (+), so reverse
the copper half-reaction.
Fe+3
(aq)+ e-
→ Fe+2
(aq) Eº = 0.77 V
Cu(s) → Cu+2
(aq)+2e-
Eº = -0.34 V
Must balance the e-
s, so multiply the Fe 1/2-
reaction by 2, BUT do not multiply Eº. Why?
Cell potential is an intensive property (doesn’t
depend on number of times the reaction occurs).
20. 20
Cell Potential
Be sure to use
the correct 1/2-
reactions!
For example,
table 17.1, p.
796 lists three
1/2-reactions for
MnO4
1-
Find them.
Answers next
slide.
21. 21
Cell Potential
Three 1/2-reactions for MnO4
1-
:
MnO4
1-
+ 4H1+
3e-
→ MnO2 + 2H2O Eº = 1.68
MnO4
1-
+ 8H1+
5e-
→ Mn2+
+ 4H2O Eº = 1.51
MnO4
1-
+ e-
→ MnO4
2-
Eº = 0.56
Pick the reaction that “works” with your overall
reaction (look at reactants and products).
22. 22
Line Notation pp
solidAqueousAqueoussolid
Anode on the leftCathode on the right
Single line to show different phases.
Double line → porous disk or salt bridge.
If all the substances on one side are
aqueous, a platinum electrode is used.
For: Cu(s) + Fe+3
(aq)→ Cu+2
(aq) + Fe+2
(aq)
Cu(s)Cu+2
(aq)Fe+3
(aq),Fe+2
(aq)Pt(s)
Remember to show the electrodes!!
23. 23
Complete Galvanic Cell Description (AP Test) pp
The reaction always runs
spontaneously in the direction that
produced a positive cell potential.
Four things for a complete description:
1. Cell Potential and balanced reaction
2. Direction of flow
3. Designation of anode and cathode
4. Nature of all components -- electrodes
& ions (plus inert conductor like Pt if
needed). Use line notation.
24. 24
Practice pp
Completely describe the galvanic cell
based on the following half-reactions
under standard conditions.
MnO4
-
+ 8 H+
+5e-
→ Mn+2
+ 4H2O
Eº = 1.51 V
Fe+2
+2e-
→ Fe(s) Eº = -0.44 V
25. 25
Practice - Item 1 pp
Determine cell potential & balanced reaction
MnO4
-
+ 8 H+
+5e-
→ Mn+2
+ 4H2O Eº = 1.51 V
Fe+2
+2e-
→ Fe(s) Eº = -0.44 V
Since (+) E required, reverse the 2nd reaction
Eºcell =1.51 + (+0.44) = 1.95 V
Complete, balanced reaction is 2MnO4
-
+ 5Fe(s) + 16 H+
→ 2Mn+2
+ 5Fe2+
(aq) + 8H2O(l)
Note: multiplying the half reactions to balance
the reaction does NOT multiply the Eº values!!!
(intensive property).
26. 26
Practice - Item 2 pp
Determine e-
flow by inspecting 1/2 rxns &
using the direction that gives a (+) Eºcell
MnO4
-
+ 8 H+
+5e-
→ Mn+2
+ 4H2O Eº = 1.51 V
Fe(s) → Fe+2
+2e-
Eº =+0.44 V Eºcell
=1.51 + (+0.44) = 1.95 V
So, electrons flow from Fe(s) to MnO4
-
(aq)
27. 27
Practice - Item 3 pp
Designate the anode and cathode
MnO4
-
+ 8 H+
+5e-
→ Mn+2
+ 4H2O Eº = 1.51 V
Fe(s) → Fe+2
+2e-
Eº =+0.44 V Eºcell
=1.51 - (0.44) = 1.95 V
Elections flow from Fe(s) to MnO4
-
So, oxidation occurs in the compartment
containing Fe(s) -- the anode
Reduction occurs in the compartment
containing MnO4
-
-- Use Pt for the cathode
Note: e-
always flow from anode to cathode
”Red cat ate an ox". Red/cat = reduct/cathode
28. 28
Practice - Item 4 pp
Describe nature of each electrode & ions
present (use line notation)
Complete, balanced reaction is 2MnO4
-
+
5Fe(s) + 16 H+
→ 2Mn+2
+ 5Fe2+
(aq) + 8H2O(l)
Electrode in Fe/Fe2+
compartment is iron metal
An inert conductor like Pt must be used in MnO4
-
1
/ Mn+2
compartment (don’t forget).
Line notation is: Fe(s)Fe+2
(aq)MnO4
-
1
(aq),Mn+2
(aq)Pt(s)
29. 29
pp Figure 17.7:
A Schematic of
the previous
Galvanic Cell
Eº = 1.95 v
Be able to draw
this as well as
write the line
notation for the
AP exam.
30. 30
17.3 Cell Potential, Work & ∆G
emf = potential (V) = work (J) / Charge(C)
E = work done by system / charge
E = -w/q (emf & work have opposite signs)
Use (-)w because it is flowing out from system.
Charge is measured in coulombs.
-w = qE (where q = the charge)
Faraday = 96 485 C/mol e-
q = nF = moles of e-
x charge per mole e-
w = -qE = -nFE = ∆G
31. 31
Potential, Work and ∆G pp
∆Gº = -nFE º (at standard conditions)
if E º > 0, then ∆Gº < 0 spontaneous
if E º < 0, then ∆Gº > 0 nonspontaneous
In fact, reverse is spontaneous.
32. 32
Potential, Work and ∆G
Calculate ∆Gº for the following reaction:
Cu+2
(aq)+ Fe(s) → Cu(s)+ Fe+2
(aq)
Fe+2
(aq)+ 2e-
→ Fe(s) Eº = -0.44 V
Cu+2
(aq)+2e-
→ Cu(s) Eº = 0.34 V
∆Gº = -nFE º Answer?
-1.5 x 105
J Calculation with units is . . .
-(2 mol e-
)(96 485 C/mol e-
)(0.78 J/C)
33. 33
Putting It Together pp
Using Table 17.1, predict if 1 M HNO3 will
dissolve gold to form a 1 M Au3+
solution?
What are the half reactions? . . .
Gold needs to be oxidized so HNO3 must
be reduced. Look for a half-reaction with
HNO3 where NO3
1-
is being reduced . . .
NO3
1-
+ 4H1+
+ 3e-
→ NO + 2H2O Eº = 0.96 v
Au → Au3+
+ 3e-
Eº = -1.50 v
34. 34
Putting It Together pp
Using Table 17.1, predict if 1 M HNO3 will
dissolve gold to form a 1 M Au3+
solution?
NO3
1-
+ 4H1+
+ 3e-
→ NO + 2H2O Eº = 0.96 v
Au → Au3+
+ 3e-
Eº = -1.50 v
Au + NO3
1-
+ 4H1+
→ Au3+
+ NO + 2H2O Eºcell = -0.54 v
Since Eºcell = negative, this cannot be
spontaneous because . . .
∆Gº = -nFE º = (-)(3)(96 485)(-0.54) = +156kJ
Since ∆Gº is (+), not spontaneous.
35. 35
17.4 Cell Potential and Concentration pp
Notes for online HW
1 lb = 453.6 g
1 Faraday = 96 485 c/s (not 96 500 c/s)
Use 0.0592 in Nernst equation (not
0.0591)
36. 36
17.4 Cell Potential and Concentration17.4 Cell Potential and Concentration pppp
Qualitatively - Can predict direction of
change in E from LeChâtelier.
2Al(s) + 3Mn+2
(aq) → 2Al+3
(aq) + 3Mn(s)
Predict if Ecell will be greater or less than
Eºcell of 0.48 v if
[Al+3
] = 1.5 M and [Mn+2
] = 1.0 M if
[Al+3
] = 1.0 M and [Mn+2
] = 1.5 M if
[Al+3
] = 1.5 M and [Mn+2
] = 1.5 M
Steps . . .
37. 37
Cell Potential pp
2Al(s) + 3Mn2Al(s) + 3Mn+2+2
(aq)(aq) →→ 2Al2Al+3+3
(aq)(aq) + 3Mn(s)+ 3Mn(s)
Predict ifPredict if EEcellcell will be greater or less thanwill be greater or less than
EEººcellcell of 0.48 vof 0.48 v
ifif [Al[Al+3+3
] = 1.5] = 1.5 MM andand [Mn[Mn+2+2
] = 1.0] = 1.0 MM
Answer . . .Answer . . .
Since aSince a productproduct [ ][ ] has been raised abovehas been raised above
1.01.0 MM, Le Chatelier predicts a shift, Le Chatelier predicts a shift leftleft
and Eand Ecellcell < Eº< Eºcellcell
38. 38
Cell Potential pp
2Al(s) + 3Mn2Al(s) + 3Mn+2+2
(aq)(aq) →→ 2Al2Al+3+3
(aq)(aq) + 3Mn(s)+ 3Mn(s)
Predict ifPredict if EEcellcell will be greater or less thanwill be greater or less than
EEººcellcell of 0.48 vof 0.48 v ifif
[Al[Al+3+3
] = 1.0] = 1.0 MM andand [Mn[Mn+2+2
] = 1.5] = 1.5 MM
Answer . . .Answer . . .
Since aSince a reactantreactant [ ][ ] has been raisedhas been raised
above 1.0above 1.0 MM, Le Chatelier predicts a shift, Le Chatelier predicts a shift
rightright and Eand Ecellcell > Eº> Eºcellcell
39. 39
Cell Potential pp
2Al(s) + 3Mn+2
(aq) → 2Al+3
(aq) + 3Mn(s)
Predict if Ecell will be greater or less than
Eºcell of 0.48 v if
[Al+3
] = 1.5 M and [Mn+2
] = 1.5 M
Answer . . .
Since both [ ]s have been raised above
1.0 M, cannot use Le Chatelier!
Must use the Nernst Equation (coming to
your community soon!)
40. 40
Le Chatelier, ∆G, & Concentration Cells
Figure 17.9
A Concentration Cell
That Contains a Sliver
Electrode and Aqueous
Silver Nitrate in Both
Compartments
Since the right
compartment has higher
[Ag1+
] there is a shift right
of e-
So, Ag metal is
deposited on the right
side while [Ag1+
]
decreases on right side
and increases on the left.
41. 41
Le Chatelier, ∆G, & Concentration Cells pp
So, Ag metal is
deposited on the right
side while [Ag1+
]
decreases on right
side and increases
on the left.
You will need to
recognize this
concept to solve for
∆G on the AP exam.
Hint: the electrode
with the largest [ ] will
always be the
cathode (where
reduction occurs).
42. 42
The Nernst Equation
∆G = ∆Gº +RTln(Q), since ∆G = -nFE . . .
-nFE = -nFEº + RTln(Q)
E= Eº - RTln(Q)
nF
What is n in Al(s) + Mn2+
→ Al3+
+ Mn(s)?
Always have to figure out “n” by balancing.
2Al(s) + 3Mn+2
(aq) → 2Al+3
(aq) + 3Mn(s) Eº = 0.48 V
n = 6. Why? . . .
n = mole of e-
, not mole of compound.
43. 43
The Nernst Equation continued
∆G = ∆Gº +RTln(Q)
-nFE = -nFEº + RTln(Q)
E= Eº - RTln(Q) = Eº - 2.303RT log (Q) nF
nF
Since we know R and F, at 25o
C, above is aka
E= Eº - 0.0592log(Q)
n
Textbook has typo: 0.0591 should be 0.0592
Use 0.0592 for all your calculations!
44. 44
The Nernst Equation pp
E= Eº - 0.0592log(Q)
n
For concentration cells (i.e., not at 1 M) this
equation must be done separately for each 1/2
cell, then subtract the results (or flip one and
add) to get the cell potential.
See the following problem . . .
46. 46
Ecell when [Ag1+
] on the right = 1.0 M pp
Since [Ag1+
] is same on
both sides Ecell = Eºcell which
is 0 because . . .
Ag1+
+ e-
→ Ag Eº = .80v
Ag → Ag1+
+ e-
Eº = -.80v
Eºcell = 0
E= Eº - 0.0592log(Q)
n
Since log(1) = 0, Ecell = Eºcell and
Eº = 0 from above, so Ecell also
= 0.
47. 47
Ecell when [Ag1+
] on the right = 2.0 M pp
Cathode always has the
higher [ ] and e-
always
flow from the anode to
the cathode.
Q = [ ]anode ÷ [ ]cathode
Here, [cathode] is on the
right (2.0 M), & in the
denominator for Q
E= Eº - 0.0592log(Q)
n
E= 0 - 0.0592log 1.0= .018 v
1 2.0
48. 48
Ecell when [Ag1+
] on the right = 0.10 M pp
Cathode always has the
higher [ ] and e-
always
flow from the anode to
the cathode.
So, [cathode] is on the
left & in the denominator
for Q
E= Eº - 0.0592log(Q)
n
E= 0 - 0.0592log 0.1 = .059 v
1 1.0
You will have a test
question on this and it is
NOT on the pre-test, so . . .
Do p. 832 #53!!
49. 49
Ecell when [Ag1+
] on the right = 0.10 M pp
Notes for all of these:
Eº is always 0 for
concentration cells
because the 1/2 reactions
cancel.
E= Eº - 0.0592log(Q)
n
E= 0 - 0.0592log anode .
n cathode
Memorize this equation!
You will have a test
question on this and it is
NOT on the pre-test, so . . .
Do p. 832 #53!!
50. 50
The Nernst Equation
As reactions proceed concentrations of
products increase and reactants decrease.
Reaches equilibrium where Q = K and Ecell = 0
Since at equilibrium Ecell = 0 = Eº - RTln(K)
nF
Eº = RTln(K)
nF
nFEº = ln(K) at 25º C aka log(K) = nEº RT
0.0592
51. 51
Nernst Equation & K pp
Calculate KCalculate Kspsp of silver iodide at 298 Kof silver iodide at 298 K
AgI(s)AgI(s) →→ AgAg++
+ I+ I--
where Ewhere Eoo
AgI(s) + eAgI(s) + e--
→→ Ag(s) + IAg(s) + I--
-0.15 v-0.15 v
II22(s) + 2e(s) + 2e--
→→ 2I2I--
+0.54 v+0.54 v AgAg++
+ e+ e--
→→ Ag(s)Ag(s) +0.80 v+0.80 v
logK = nEº/0.0592logK = nEº/0.0592
1st, get Eº1st, get Eºcellcell for overall reaction. Steps.for overall reaction. Steps.
Find overall reaction, then EºFind overall reaction, then Eºcellcell
Only need 1st & 3rd equations to get . . .Only need 1st & 3rd equations to get . . .
AgI(s)AgI(s) →→ AgAg++
+ I+ I--
where Ewhere Eoo
cellcell = -0.95 v= -0.95 v
52. 52
Nernst Equation & K pp
Calculate Ksp of silver iodide at 298 K
AgI(s) → Ag+
+ I-
where Eo
cell = -0.95 v
at 25ºC log(K) = nEº = (1)(-0.95) = -16.05
0.0592 0.0592
Ksp = 10-16.05
= 9.0 x 10-17
53. 53
17.5 Batteries are Galvanic Cells
Car batteries are lead storage batteries.
Pb +PbO2 +H2SO4 →PbSO4(s) +H2O
Be able to recognize the anode &
cathode from the half reactions.
57. 57
17.6 Corrosion
RustingRusting - spontaneous- spontaneous oxidationoxidation..
Most structural metals haveMost structural metals have reductionreduction
potentials that arepotentials that are lessless positive than Opositive than O22 ..
So, theySo, they oxidizeoxidize while Owhile O22 isis reduced.reduced.
FeFe+2+2
+2e+2e--
→→ FeFe EEº=º= --0.44 V0.44 V
OO22 + 2H+ 2H22O + 4eO + 4e--
→→ 4OH4OH--
EEº=º= ++0.40 V0.40 V
Reverse top half reaction, then add bothReverse top half reaction, then add both
Fe + OFe + O22 + H+ H22OO →→ FeFe22 OO33(rust)(rust)+ H+ H++
EºEºcellcell = 0.84 v= 0.84 v
Reaction happens in two places . . .Reaction happens in two places . . .
60. 60
Preventing Corrosion
Coating - to keep out air and water.
Galvanizing - Putting on a zinc coat
Zinc has a lower reduction potential than
iron, so it is more easily oxidized.
So, zinc is a more active metal than iron.
Alloying with metals that form oxide coats.
Cathodic Protection - Attaching large
pieces of a more active metal like
magnesium that get oxidized instead of
iron (iron stays reduced).
61. 61
Preventing Corrosion
Cathodic Protection - Attaching largeCathodic Protection - Attaching large
pieces of an active metal like magnesiumpieces of an active metal like magnesium
that get oxidized instead of iron.that get oxidized instead of iron.
Attach Mg wireAttach Mg wire to iron pipe (and replaceto iron pipe (and replace
periodically).periodically).
Attach titanium barsAttach titanium bars to ships’ hulls. Into ships’ hulls. In
salt water the Ti acts as the anode and issalt water the Ti acts as the anode and is
oxidized instead of the steel hull, whichoxidized instead of the steel hull, which
now acts as the cathode.now acts as the cathode.
63. 63
Running a galvanic cell backwards.
Put a voltage whose magnitude is bigger than
the potential which reverses the direction of the
redox reaction.
Produces a chemical change which would not
normally happen because the potential is
negative.
Remember: 1 A = 1 C/s and 1 F = 96 485 C
Used for electroplating -- depositing the neutral
metal onto the electrode by reducing the metal
ions in solution.
17.7 Electrolysis
67. 67
Calculating plating
Have to include the charge.Have to include the charge.
Measure currentMeasure current II (in amperes)(in amperes)
1 amp = 1 coulomb of charge per second1 amp = 1 coulomb of charge per second
1 A = 1 C/s1 A = 1 C/s
q =q = II x t = the chargex t = the charge
q/nF = moles of metalq/nF = moles of metal
Mass of plated metalMass of plated metal
68. 68
Calculating plating pp
How many minutes must a 5.00 amp
current be applied to produce 10.5 g of
Ag from Ag+
Steps follow . . .
Set up a picket fence that includes
– Current & time
– Quantity of charge (in coulombs)
– Moles of electrons
– Moles of metal (may be different)
– Grams of metal
Arrange the picket fence so the units give
you what you’re looking for.
69. 69
Calculating plating pp
How many minutes must a 5.00 amp
current be applied to produce 10.5 g of
Ag from Ag+
Steps follow . . .
The pieces of the picket fence are:
– 5.00 amp (rewrite as 5.00 C/s)
– 10.5 g Ag
– 107.868 g Ag/1mol Ag
– 1 mol e-
/mol Ag (Ag → Ag1+
+ 1e-
)
– 96 485 C/mol e-
– 60 sec/min
Your answer? . . .
70. 70
Calculating platingCalculating plating pppp
How many minutes must a 5.00 amp
current be applied to produce 10.5 g of Ag
from Ag+
(amp = C/s)
31.3 minutes31.3 minutes
(10.5 g Ag)(1 mol Ag/107.868 g Ag)(1 mol e-
/1 mol
Ag)(96 485 C/1 mol e-
)(1 s/5.00C)(1 min1 min/60 s)
71. 71
Calculating plating pp
An antique automobile bumper is to be
chrome plated by dipping into an acidic
Cr2O7
2-
solution serving as the cathode of
an electrolytic cell. MCr = 51.996
Using 10.0 amperes, how long to deposit
1.00 x 102
grams of Cr(s)? Steps . . .
Find overall reaction from 1/2-reactions
(to get moles of e-). Steps follow.
The only substances you start with are
Cr2O7
2-
and H2O. Which is oxidized?
Reduced? . . .
72. 72
Calculating plating pp
Using 10.0 amperes, how long to depositUsing 10.0 amperes, how long to deposit
1.00 x 101.00 x 1022
grams of Cr(s) by dipping intograms of Cr(s) by dipping into
acidic Cracidic Cr22OO77
2-2-
(aq) ,(aq) , MMCrCr = 51.996?= 51.996? TheThe
only substances you start with are Cronly substances you start with are Cr22OO77
2-2-
and Hand H22O.O. Which is oxidized? Which isWhich is oxidized? Which is
reduced?reduced? . . .. . .
CrCr22OO77
2-2-
must be reducedmust be reduced to get to Crto get to Cr(s)(s) soso
HH22O must be oxidizedO must be oxidized. Use Table 17.1 p.. Use Table 17.1 p.
796 to find the 1/2-reaction of H796 to find the 1/2-reaction of H22O . . .O . . .
73. 73
Calculating plating pp
Use Table 17.1 p. 843 to find the 1/2-Use Table 17.1 p. 843 to find the 1/2-
reaction of Hreaction of H22O. Which is it?O. Which is it?
HH22OO22 + 2H+ 2H1+1+
+ 2e+ 2e--
→→ 2H2H2200
2H2H2200 →→ OO22 + 4H+ 4H1+1+
+ 4e+ 4e--
OO22
+ 2H+ 2H220 + 4e0 + 4e--
→→ 4OH4OH1-1-
2H2H220 +0 +
2e2e--
→→ HH22 + 2OH+ 2OH1-1-
*****
#2 above is the only one that starts#2 above is the only one that starts
with water and is oxidized.with water and is oxidized.
74. 74
Calculating plating pp
Using 10.0 amperes, how long to deposit
1.00 x 102
grams of Cr(s) by dipping into
acidic Cr2O7
2-
(aq) ? , MCr = 51.996
So, the 1/2-reactions needed are . . . 6e-
+ 14H1+
+ Cr2O7
2-
→ 2Cr3+
+ 7H2O 2H20 → O2
+ 4H1+
+ 4e-
and
one more! (Why?)
You only got to Cr3+
, but you need Cr(s)
Also need Cr3+
+ 3e-
→ Cr(s)
Now, write the balanced equation . . .
75. 75
Calculating plating pp
Using 10.0 amperes, how long to deposit
1.00 x 102
grams of Cr(s) by dipping into
acidic Cr2O7
2-
(aq) ?, MCr = 51.996
Now, write the balanced equation . . .
2H1+
+ Cr2O7
2-
→ H2O + 3O2 + 2Cr(s)
How many mol e-
changed in this reaction
(comprised of three 1/2-reactions)?
12 mol e-
Now, you can start to do the problem.
Your answer in days? . . .
76. 76
Calculating plating pp
Using 10.0 amperes, how long to deposit
1.00 x 102
grams of Cr(s) by dipping into
acidic Cr2O7
2-
(aq)? , MCr = 51.996 with 2H1+
+
Cr2O7
2-
→ H2O + 3O2 + 2Cr(s) and 12 mol
e-
moving.
Your answer is . . .
1.29 days . . .
100.g Cr(s) • 1mol Cr/51.996g Cr(s) • 12 mol e-
/2 mol
Cr(s) • 96 486 C/1 mol e-
• 1s/10.0C • 1 h/3600 s • 1
d/24 h = 1.29 days
77. 77
Calculating plating
Electrolysis of a molten salt , MCl, usingElectrolysis of a molten salt , MCl, using
3.86 amps for 16.2 min deposits 1.52 g of3.86 amps for 16.2 min deposits 1.52 g of
metal. What is the metal?metal. What is the metal?
Use picket fence to get moles of metal.Use picket fence to get moles of metal.
Since g/mol = M, use calculated moles ofSince g/mol = M, use calculated moles of
metal and 1.52 g to get M. Answer . . .metal and 1.52 g to get M. Answer . . .
Potassium.Potassium.
The solution is . . .The solution is . . .
78. 78
Calculating plating
Electrolysis of a molten salt , MCl, usingElectrolysis of a molten salt , MCl, using
3.86 amps for 16.2 min deposits 1.52 g of3.86 amps for 16.2 min deposits 1.52 g of
metal. What is the metal?metal. What is the metal?
3.86 C/s • 16.2 min • 60s/1m • 1mol e-/96 485 C • 1 mol3.86 C/s • 16.2 min • 60s/1m • 1mol e-/96 485 C • 1 mol
MM1+1+
/1mol e- = 0.039 = moles of metal./1mol e- = 0.039 = moles of metal.
1.52g/0.039 mol = 39.1 g/mol = potassium1.52g/0.039 mol = 39.1 g/mol = potassium
79. 79
Other uses pp
Electrolysis of water.Electrolysis of water.
Separating mixtures of ionsSeparating mixtures of ions
A more positive reduction potentialA more positive reduction potential
means that reaction proceeds forward.means that reaction proceeds forward.
The metal with the mostThe metal with the most positivepositive
reductionreduction potential is easiest to plate outpotential is easiest to plate out
of solution (and is the best oxidizer).of solution (and is the best oxidizer).
80. 80
Relative Oxidizing Abilities pp
An acidic solution has CeAn acidic solution has Ce4+4+
, VO, VO22
1+1+
, & Fe, & Fe3+3+
. Use. Use
Table 17.1 p. 796 to predict the order ofTable 17.1 p. 796 to predict the order of
oxidizingoxidizing ability. Answer . . .ability. Answer . . .
Order of oxidizing ability is the same as theOrder of oxidizing ability is the same as the
order fororder for beingbeing reduced, So. . .reduced, So. . . CeCe4+4+
+ e+ e--
→→ CeCe3+3+
E = 1.70 vE = 1.70 v VOVO22
1+1+
+ 2H+ 2H1+1+
+ e+ e--
→→ VOVO2+2+
+ H+ H22OO E = 1.00 vE = 1.00 v FeFe3+3+
+ e+ e--
→→ FeFe2+2+
E = 0.77 vE = 0.77 v
Also, predict which one will beAlso, predict which one will be reducedreduced at theat the
cathode of ancathode of an electrolyticelectrolytic cell at thecell at the lowestlowest
voltage. Answer . . .voltage. Answer . . .
SinceSince CeCe4+4+
is the greatest oxidizer it is the mostis the greatest oxidizer it is the most
easily reduced (needs least voltage).easily reduced (needs least voltage).
Do section 17.8 on your own.Do section 17.8 on your own.
Editor's Notes
Z5e 837 Ch 17 Intro
Z5e Section 17.1 Galvanic Cells
Z5e 838 Fig. 17.1
Z5e 839 Fig. 17.1
Electrons flow in the wire, ions flow through salt bridge
Z5e 839 Fig. 17.1
Electrons flow in the wire, ions flow through salt bridge
Z5e 839 Fig. 17.3
Oxidation occurs at the anode because the species in solution acting as the reducing agent (so it is oxidized) supplies electrons to the anode.
The species in the solution acting as the oxidizing agent (so it is reduced) receives electrons from the cathode (where reduction is occurs)
Z5e 851 Section 17.4 Dependence of Cell Potential on Concentration
Rf. SE 17.5
Lesser since product [ ] is raised to greater than 1.0 M
Greater
Need Nernst equation!
Z5e 851 Section 17.4 Dependence of Cell Potential on Concentration
Rf. SE 17.5
Lesser since product [ ] is raised to greater than 1.0 M
Greater
Need Nernst equation!
Z5e 851 Section 17.4 Dependence of Cell Potential on Concentration
Rf. SE 17.5
Lesser since product [ ] is raised to greater than 1.0 M
Greater
Need Nernst equation!
Z5e 851 Section 17.4 Dependence of Cell Potential on Concentration
Rf. SE 17.5
Lesser since product [ ] is raised to greater than 1.0 M
Greater
Need Nernst equation!
Z5e 851 Section 17.4 Dependence of Cell Potential on Concentration
Rf. SE 17.5
Lesser since product [ ] is raised to greater than 1.0 M
Greater
Need Nernst equation!
Z5e 851 Section 17.4 Dependence of Cell Potential on Concentration
Rf. SE 17.5
Lesser since product [ ] is raised to greater than 1.0 M
Greater
Need Nernst equation!
Z5e 851 Section 17.4 Dependence of Cell Potential on Concentration
Rf. SE 17.5
Lesser since product [ ] is raised to greater than 1.0 M
Greater
Need Nernst equation!
Z5e 853
n = mole of electrons not mole of the compounds
Z7e text p. 832 #53
Z5e 882 357
Z5e 882 #59
Z5e 882 #59
Z5e 882 #59
See SE 17.8 page 856
Taken from test bank 5e Ch 17 #76
Eo = -0.95 (don’t use 2nd equation, above; its extraneous)
Log Ksp = (1)(-0.95)/0.0592 so Ksp = 9.0 x 10-17
Taken from test bank 5e Ch 17 #76
Eo = -0.95 (don’t use 2nd equation, above; its extraneous)
Log Ksp = (1)(-0.95)/0.0592 so Ksp = 9.0 x 10-17
Z5e 857 Section 17.5 Batteries
Z5e 858 Figure 17.13
Z5e 857 Section 17.5 Batteries
Z5e 858 Figure 17.14
Z5e 861 Section 17.6 Corrosion
Z5e 862 Figure 17.17
Z5e 865
Z5e 865 Figure 17.18(a)
Z5e 866 Section 17.7 Electrolysis
Z5e 867 Figure 17.19 (a) and (b)
Z5e 866
Z5e 868 SE 17.9 electroplating
Answer: 31.3 minutes
Z5e 868 SE 17.9 electroplating
Answer: 31.3 minutes
Z5e 868 SE 17.9 electroplating
Answer: 31.3 minutes
Z5e Test Bank #78
6e- + 14H+ + Cr2O72- --&gt; 2Cr3+ + 7H2O
(2H20 --&gt; O2 + 4H+ + 4e-) x 3
[since water is oxidized, write out and balance this half reaction)
(Cr3+ + 3e- --&gt; Cr(s)) x 2
Balanced equation is:
2H+ + Cr2O7 --&gt; H2O + 3O2 + 2Cr(s) with 12 e- changing
Z5e Test Bank #78
6e- + 14H+ + Cr2O72- --&gt; 2Cr3+ + 7H2O
(2H20 --&gt; O2 + 4H+ + 4e-) x 3
[since water is oxidized, write out and balance this half reaction)
(Cr3+ + 3e- --&gt; Cr(s)) x 2
Balanced equation is:
2H+ + Cr2O7 --&gt; H2O + 3O2 + 2Cr(s) with 12 e- changing
Z5e Test Bank #78
6e- + 14H+ + Cr2O72- --&gt; 2Cr3+ + 7H2O
(2H20 --&gt; O2 + 4H+ + 4e-) x 3
[since water is oxidized, write out and balance this half reaction)
(Cr3+ + 3e- --&gt; Cr(s)) x 2
Balanced equation is:
2H+ + Cr2O7 --&gt; H2O + 3O2 + 2Cr(s) with 12 e- changing
Z5e Test Bank #78
6e- + 14H+ + Cr2O72- --&gt; 2Cr3+ + 7H2O
(2H20 --&gt; O2 + 4H+ + 4e-) x 3
[since water is oxidized, write out and balance this half reaction)
(Cr3+ + 3e- --&gt; Cr(s)) x 2
Balanced equation is:
2H+ + Cr2O7 --&gt; H2O + 3O2 + 2Cr(s) with 12 e- changing
Z5e Test Bank #78
6e- + 14H+ + Cr2O72- --&gt; 2Cr3+ + 7H2O
(2H20 --&gt; O2 + 4H+ + 4e-) x 3
[since water is oxidized, write out and balance this half reaction)
(Cr3+ + 3e- --&gt; Cr(s)) x 2
Balanced equation is:
2H+ + Cr2O7 --&gt; H2O + 3O2 + 2Cr(s) with 12 e- changing
Z5e Test Bank #78
6e- + 14H+ + Cr2O72- --&gt; 2Cr3+ + 7H2O
(2H20 --&gt; O2 + 4H+ + 4e-) x 3
[since water is oxidized, write out and balance this half reaction)
(Cr3+ + 3e- --&gt; Cr(s)) x 2
Balanced equation is:
2H+ + Cr2O7 --&gt; H2O + 3O2 + 2Cr(s) with 12 e- changing