Electrochemistry by Owen Borville 11.2.2025
A Redox reaction is a chemical reaction in which there are changes in oxidation states between the reactants and the products of the reaction.
Ex. Sn(s) + Cu2+(aq) => Cu(s) + Sn2+(aq), where the oxidation states of Sn changes from 0 to 2+ from solid to aqueous and Cu changes from 2+ to 0 from aqueous to solid.
Redox reactions can also be balanced by using the half-reaction method: (Fe2+) + (Cr2O7)2- => Fe3+ + Cr3+
(Step 1) Separate equation into half-reactions for oxidation and reduction. Oxidation: Fe2+ => Fe3+ Reduction: (Cr2O7)2- => Cr3+
(Step 2) Balance each half reaction: Fe2+ =>Fe3+ and => (Cr2O7)2- =>2Cr3+
(Step 3) Add H2O in the reduction to balance O = (Cr2O7)2- => Cr3+ + 7H2O
(Step 4) Add H+ in the reduction to balance H = (14H+) + (Cr2O7)2- => (2Cr3+) + 7H2O
(Step 5) Balance charges by Adding electrons to the Oxidation = Fe2+ => (Fe3+) + e- (each side has +2 electrons)
and Reduction = 6e- + (14H+) + (Cr2O7)2- =>(2Cr3+) + 7H2O (each side has +6 electrons)
(Step 6) If the number of electrons in the half-reactions are not equal, Multiply to make them equal. Therefore, multiply the oxidation reaction by: 6(Fe2+ =>Fe3+ + 6e-) = 6Fe2+=>(6Fe3+)+6e-
(Step 7) Add the half reactions together and cancel the electrons out on each side: 6e- +(6Fe2+) + (14H+)+ (Cr2O7)2- => (6Fe3+)+(2Cr3+)+(7H2O) + 6e-
=(6Fe2+) + (14H+)+ (Cr2O7)2- => (6Fe3+)+(2Cr3+)+(7H2O)
(Step 8) If the redox reaction occurs in basic solution: For each H+ ion, add one OH- to both sides of the equation. Then cancel out necessary reactants or products created by the new H2O molecules.
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Galvanic Cells (voltaic cell) are the experimental apparatus for generating electricity through the use of a spontaneous reaction. The galvanic cell produces electricity through spontaneous redox reactions. Electrons in the reaction are transferred between different chemical species, releasing energy that can be harnessed as electric power. A example of a galvanic cell is zinc metal placed in copper II solution. Zinc metal is oxidized at the anode (Zn => Zn2+ + 2e-), and copper ions are reduced at the cathode (Cu2++2e-=>Cu).
Electric current flows from anode to cathode because there is a difference in potential energy between the electrodes. The electrical potential is measured by a voltmeter and is called the cell potential (Ecell).
Galvanic cell notation: Zn(s)|Zn2+(1M)||Cu2+(1M)|Cu(s)
------------------------------|------anode-------|------cathode-----|
When the concentrations of Zn2+ and Cu2+ are 1 molar at 25 degrees C, the cell voltage is 1.10 V, as indicated by a voltmeter.
Zn(s) + Cu2+(aq) => Zn2+(aq) + Cu(s)
Standard Reduction Potentials: The hydrogen electrode is used as the arbitrary standard reference in measuring standard cell potentials of all other electrodes. H2 =>2H+ + 2e- E(SHE)= zero v. SHE = Standard Hydrogen Electrode at 0V, 1 M H+, 1 bar H2 pressure (almost 1 atm), 298K.
Zn(s)|Zn2+(1M)||H+(1M)|H2(g)|Pt(s)
Anode (oxidation) Zn(s) =>Zn2+(1M) + 2e-
Cathode (reduction) 2H+(1M) + 2e- =>H2 (1 atm)
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Zn(s) + 2H+(1M) => Zn2+(1M) + H2(1 atm) (electrons cancel)
Ecell = E(cathode) - E(anode)
Ecell = E(H+/H2) - E(Zn2+/Zn)
0.76 V = 0-E(Zn2+/Zn)
E(Zn2+/Zn) = -0.76 V
Pt(s)|H2(g)|H+(1M)||Cu2+(1M)|Cu(s)
Anode (oxidation) H2 (1 atm) => 2H+(1M) + 2e-
Cathode (reduction) Cu2+(1M) + 2e- =>Cu(s)
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H2(1 atm) + Cu2+(1M) => 2H+(1M) + Cu(s) (electrons cancel)
Ecell = E(cathode) - E(anode)
Ecell = E(Cu2+/Cu) - E(H+/H2)
0.34 V = 0-E(H+/H2)
E(H+/H2) = -0.34 V
Ecell = E(cathode)-E(anode)
Ecell = E(Cu2+/Cu) - E(Zn2+/Zn)
Ecell = 0.34 V - (-0.76 V)
Ecell = 1.10 V
Standard Reduction Potentials at 25 degrees C are listed in reference tables: Increasing strength as oxidizing agents or reducing agents.
Determining the overall cell reaction and Ecell at 25 degrees C of a galvanic cell composed of two different metallic ion electrode in two different 1 M solutions:
Step 1: Use the standard reduction potential table to determine which electrode is the cathode and which is the anode.
Step 2: Ecell = E(cathode)-E(anode)
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Spontaneity of Redox Reactions Under Standard-State Conditions
Ecell is related to the thermodynamic quantities ΔG and K.
Electric energy (J) = volts (V) x coulombs (C)
1(one) J = 1(one) C x 1(one) V
Total charge = mole of electrons x faraday (F)
1(one) F = 96,500 C/mol e-
1(one) F = 96,500 J/V*mol e-
ΔG° = -nFE°cell
E°cell = (RT/nF) lnK
n = number of moles
F = Faraday constant (96,500 C/mole e-)
R = gas constant (8.314 J/mol K)
By converting to the base-10 logarithm of K,
E°cell = [(0.0592 V)/n] log K (at 25 °C)
To calculate ΔG° and K for a reaction at 25 °C
(Step 1) Use E° to calculate Ecell. Look up values in the reference table.
E°cell = E(cathode) - E(anode)
(Step 2) ΔG° = -nFE°cell
Use E°cell = (RT/nF) lnK to calculate K
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Nernst Equation: Spontaneity of Redox Reactions Under Conditions Other Than the Standard State
Under nonstandard conditions, cell potential is calculated with the Nernst equation:
E =E° - (RT/nF)lnQ
When Q > 1, E < E°
When Q < 1, E > E°
To calculate whether a reaction will occur spontaneously between two reactants at a certain temperature K, and knowing the concentration M of each reactant:
(Step 1) Calculate E° using data from a standard reduction potential table. Determine which reactant is the cathode and which is the anode.
E°cell = E(cathode) - E(anode)
(Step 2)
E =E° - (RT/nF) lnQ
If E is positive, the reaction remains spontaneous. A negative value indicates that the reaction is not spontaneous.
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Concentration Cells are cells with two half-cells containing the same components, but differing ion concentrations. For example, a concentration cell has some solid Zn and also some Zn2+ ions. One half cell has 0.10 M (Zn2+) and the other half cell has 1.0 M (Zn2+). Use the Nernst Equation to solve for E. (E =E° - (RT/nF)lnQ)
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Batteries: Dry Cells = have no fluid component. For example, a zinc container as an anode and a graphite cathode of 1.5V. Alkaline batteries have a basic medium. Lead storage batteries have six cells, 2 V per cell, 12 V total and are rechargeable. Lithium ion batteries have a cell potential of 3.4 V and are rechargeable many times. Fuel cells have a cell potential of 1.23 V and are 70 percent efficient.
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Electrolysis: The use of electric energy to drive a nonspontaneous chemical reaction is called electrolysis. An electrolytic cell is used. While galvanic cells have spontaneous chemical reactions and electric energy is produces, electrolytic cells have non-spontaneous chemical reactions and energy is consumed.
An example of electrolysis is the splitting of water into hydrogen and oxygen gas using an electric current. Other examples include the production of metals like aluminum and sodium, and electroplating objects with a thin layer of another metal, such as gold or chromium. Electrolysis of molten chloride is used for the industrial production of highly reactive metals like sodium, magnesium, and aluminum, and for producing chlorine gas.
Quantitative application of electrolysis: Michael Faraday developed the quantitative treatment of electrolysis. Calculating the amount of mass produced when a constant current of known value is passed through an electrolytic cell for a certain amount of time:
(Step 1) Use the current value in Ampere (amp) and time to determine the charge in Coulombs.
(Step 2) Use this charge to determine the mass by converting Coulombs to moles to grams using stoichiometry.
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Corrosion is the deterioration of a metal by an electrochemical process. Corrosion occurs naturally as rusting of iron or tarnishing of silver. During corrosion, a galvanic cell is is produces where different parts of the metal function as anode and cathode, while the corrosion takes place at the anodic site. Methods of preventing corrosion include using a protective coating of zinc (by galvanization) and cathodic protection using connected sacrificial anodes.
A Redox reaction is a chemical reaction in which there are changes in oxidation states between the reactants and the products of the reaction.
Ex. Sn(s) + Cu2+(aq) => Cu(s) + Sn2+(aq), where the oxidation states of Sn changes from 0 to 2+ from solid to aqueous and Cu changes from 2+ to 0 from aqueous to solid.
Redox reactions can also be balanced by using the half-reaction method: (Fe2+) + (Cr2O7)2- => Fe3+ + Cr3+
(Step 1) Separate equation into half-reactions for oxidation and reduction. Oxidation: Fe2+ => Fe3+ Reduction: (Cr2O7)2- => Cr3+
(Step 2) Balance each half reaction: Fe2+ =>Fe3+ and => (Cr2O7)2- =>2Cr3+
(Step 3) Add H2O in the reduction to balance O = (Cr2O7)2- => Cr3+ + 7H2O
(Step 4) Add H+ in the reduction to balance H = (14H+) + (Cr2O7)2- => (2Cr3+) + 7H2O
(Step 5) Balance charges by Adding electrons to the Oxidation = Fe2+ => (Fe3+) + e- (each side has +2 electrons)
and Reduction = 6e- + (14H+) + (Cr2O7)2- =>(2Cr3+) + 7H2O (each side has +6 electrons)
(Step 6) If the number of electrons in the half-reactions are not equal, Multiply to make them equal. Therefore, multiply the oxidation reaction by: 6(Fe2+ =>Fe3+ + 6e-) = 6Fe2+=>(6Fe3+)+6e-
(Step 7) Add the half reactions together and cancel the electrons out on each side: 6e- +(6Fe2+) + (14H+)+ (Cr2O7)2- => (6Fe3+)+(2Cr3+)+(7H2O) + 6e-
=(6Fe2+) + (14H+)+ (Cr2O7)2- => (6Fe3+)+(2Cr3+)+(7H2O)
(Step 8) If the redox reaction occurs in basic solution: For each H+ ion, add one OH- to both sides of the equation. Then cancel out necessary reactants or products created by the new H2O molecules.
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Galvanic Cells (voltaic cell) are the experimental apparatus for generating electricity through the use of a spontaneous reaction. The galvanic cell produces electricity through spontaneous redox reactions. Electrons in the reaction are transferred between different chemical species, releasing energy that can be harnessed as electric power. A example of a galvanic cell is zinc metal placed in copper II solution. Zinc metal is oxidized at the anode (Zn => Zn2+ + 2e-), and copper ions are reduced at the cathode (Cu2++2e-=>Cu).
Electric current flows from anode to cathode because there is a difference in potential energy between the electrodes. The electrical potential is measured by a voltmeter and is called the cell potential (Ecell).
Galvanic cell notation: Zn(s)|Zn2+(1M)||Cu2+(1M)|Cu(s)
------------------------------|------anode-------|------cathode-----|
When the concentrations of Zn2+ and Cu2+ are 1 molar at 25 degrees C, the cell voltage is 1.10 V, as indicated by a voltmeter.
Zn(s) + Cu2+(aq) => Zn2+(aq) + Cu(s)
Standard Reduction Potentials: The hydrogen electrode is used as the arbitrary standard reference in measuring standard cell potentials of all other electrodes. H2 =>2H+ + 2e- E(SHE)= zero v. SHE = Standard Hydrogen Electrode at 0V, 1 M H+, 1 bar H2 pressure (almost 1 atm), 298K.
Zn(s)|Zn2+(1M)||H+(1M)|H2(g)|Pt(s)
Anode (oxidation) Zn(s) =>Zn2+(1M) + 2e-
Cathode (reduction) 2H+(1M) + 2e- =>H2 (1 atm)
----------------------------------------------------------------
Zn(s) + 2H+(1M) => Zn2+(1M) + H2(1 atm) (electrons cancel)
Ecell = E(cathode) - E(anode)
Ecell = E(H+/H2) - E(Zn2+/Zn)
0.76 V = 0-E(Zn2+/Zn)
E(Zn2+/Zn) = -0.76 V
Pt(s)|H2(g)|H+(1M)||Cu2+(1M)|Cu(s)
Anode (oxidation) H2 (1 atm) => 2H+(1M) + 2e-
Cathode (reduction) Cu2+(1M) + 2e- =>Cu(s)
----------------------------------------------------------------
H2(1 atm) + Cu2+(1M) => 2H+(1M) + Cu(s) (electrons cancel)
Ecell = E(cathode) - E(anode)
Ecell = E(Cu2+/Cu) - E(H+/H2)
0.34 V = 0-E(H+/H2)
E(H+/H2) = -0.34 V
Ecell = E(cathode)-E(anode)
Ecell = E(Cu2+/Cu) - E(Zn2+/Zn)
Ecell = 0.34 V - (-0.76 V)
Ecell = 1.10 V
Standard Reduction Potentials at 25 degrees C are listed in reference tables: Increasing strength as oxidizing agents or reducing agents.
Determining the overall cell reaction and Ecell at 25 degrees C of a galvanic cell composed of two different metallic ion electrode in two different 1 M solutions:
Step 1: Use the standard reduction potential table to determine which electrode is the cathode and which is the anode.
Step 2: Ecell = E(cathode)-E(anode)
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Spontaneity of Redox Reactions Under Standard-State Conditions
Ecell is related to the thermodynamic quantities ΔG and K.
Electric energy (J) = volts (V) x coulombs (C)
1(one) J = 1(one) C x 1(one) V
Total charge = mole of electrons x faraday (F)
1(one) F = 96,500 C/mol e-
1(one) F = 96,500 J/V*mol e-
ΔG° = -nFE°cell
E°cell = (RT/nF) lnK
n = number of moles
F = Faraday constant (96,500 C/mole e-)
R = gas constant (8.314 J/mol K)
By converting to the base-10 logarithm of K,
E°cell = [(0.0592 V)/n] log K (at 25 °C)
To calculate ΔG° and K for a reaction at 25 °C
(Step 1) Use E° to calculate Ecell. Look up values in the reference table.
E°cell = E(cathode) - E(anode)
(Step 2) ΔG° = -nFE°cell
Use E°cell = (RT/nF) lnK to calculate K
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Nernst Equation: Spontaneity of Redox Reactions Under Conditions Other Than the Standard State
Under nonstandard conditions, cell potential is calculated with the Nernst equation:
E =E° - (RT/nF)lnQ
When Q > 1, E < E°
When Q < 1, E > E°
To calculate whether a reaction will occur spontaneously between two reactants at a certain temperature K, and knowing the concentration M of each reactant:
(Step 1) Calculate E° using data from a standard reduction potential table. Determine which reactant is the cathode and which is the anode.
E°cell = E(cathode) - E(anode)
(Step 2)
E =E° - (RT/nF) lnQ
If E is positive, the reaction remains spontaneous. A negative value indicates that the reaction is not spontaneous.
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Concentration Cells are cells with two half-cells containing the same components, but differing ion concentrations. For example, a concentration cell has some solid Zn and also some Zn2+ ions. One half cell has 0.10 M (Zn2+) and the other half cell has 1.0 M (Zn2+). Use the Nernst Equation to solve for E. (E =E° - (RT/nF)lnQ)
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Batteries: Dry Cells = have no fluid component. For example, a zinc container as an anode and a graphite cathode of 1.5V. Alkaline batteries have a basic medium. Lead storage batteries have six cells, 2 V per cell, 12 V total and are rechargeable. Lithium ion batteries have a cell potential of 3.4 V and are rechargeable many times. Fuel cells have a cell potential of 1.23 V and are 70 percent efficient.
----------------------------------------------------------------------------------
Electrolysis: The use of electric energy to drive a nonspontaneous chemical reaction is called electrolysis. An electrolytic cell is used. While galvanic cells have spontaneous chemical reactions and electric energy is produces, electrolytic cells have non-spontaneous chemical reactions and energy is consumed.
An example of electrolysis is the splitting of water into hydrogen and oxygen gas using an electric current. Other examples include the production of metals like aluminum and sodium, and electroplating objects with a thin layer of another metal, such as gold or chromium. Electrolysis of molten chloride is used for the industrial production of highly reactive metals like sodium, magnesium, and aluminum, and for producing chlorine gas.
Quantitative application of electrolysis: Michael Faraday developed the quantitative treatment of electrolysis. Calculating the amount of mass produced when a constant current of known value is passed through an electrolytic cell for a certain amount of time:
(Step 1) Use the current value in Ampere (amp) and time to determine the charge in Coulombs.
(Step 2) Use this charge to determine the mass by converting Coulombs to moles to grams using stoichiometry.
-------------------------------------------------------------------------------------------------------------------------------------------
Corrosion is the deterioration of a metal by an electrochemical process. Corrosion occurs naturally as rusting of iron or tarnishing of silver. During corrosion, a galvanic cell is is produces where different parts of the metal function as anode and cathode, while the corrosion takes place at the anodic site. Methods of preventing corrosion include using a protective coating of zinc (by galvanization) and cathodic protection using connected sacrificial anodes.