Skip to Content
OpenStax Logo
Chemistry: Atoms First 2e

16.3 Electrode and Cell Potentials

Chemistry: Atoms First 2e16.3 Electrode and Cell Potentials
Buy book
  1. Preface
  2. 1 Essential Ideas
    1. Introduction
    2. 1.1 Chemistry in Context
    3. 1.2 Phases and Classification of Matter
    4. 1.3 Physical and Chemical Properties
    5. 1.4 Measurements
    6. 1.5 Measurement Uncertainty, Accuracy, and Precision
    7. 1.6 Mathematical Treatment of Measurement Results
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  3. 2 Atoms, Molecules, and Ions
    1. Introduction
    2. 2.1 Early Ideas in Atomic Theory
    3. 2.2 Evolution of Atomic Theory
    4. 2.3 Atomic Structure and Symbolism
    5. 2.4 Chemical Formulas
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  4. 3 Electronic Structure and Periodic Properties of Elements
    1. Introduction
    2. 3.1 Electromagnetic Energy
    3. 3.2 The Bohr Model
    4. 3.3 Development of Quantum Theory
    5. 3.4 Electronic Structure of Atoms (Electron Configurations)
    6. 3.5 Periodic Variations in Element Properties
    7. 3.6 The Periodic Table
    8. 3.7 Molecular and Ionic Compounds
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  5. 4 Chemical Bonding and Molecular Geometry
    1. Introduction
    2. 4.1 Ionic Bonding
    3. 4.2 Covalent Bonding
    4. 4.3 Chemical Nomenclature
    5. 4.4 Lewis Symbols and Structures
    6. 4.5 Formal Charges and Resonance
    7. 4.6 Molecular Structure and Polarity
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  6. 5 Advanced Theories of Bonding
    1. Introduction
    2. 5.1 Valence Bond Theory
    3. 5.2 Hybrid Atomic Orbitals
    4. 5.3 Multiple Bonds
    5. 5.4 Molecular Orbital Theory
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  7. 6 Composition of Substances and Solutions
    1. Introduction
    2. 6.1 Formula Mass
    3. 6.2 Determining Empirical and Molecular Formulas
    4. 6.3 Molarity
    5. 6.4 Other Units for Solution Concentrations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  8. 7 Stoichiometry of Chemical Reactions
    1. Introduction
    2. 7.1 Writing and Balancing Chemical Equations
    3. 7.2 Classifying Chemical Reactions
    4. 7.3 Reaction Stoichiometry
    5. 7.4 Reaction Yields
    6. 7.5 Quantitative Chemical Analysis
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  9. 8 Gases
    1. Introduction
    2. 8.1 Gas Pressure
    3. 8.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law
    4. 8.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions
    5. 8.4 Effusion and Diffusion of Gases
    6. 8.5 The Kinetic-Molecular Theory
    7. 8.6 Non-Ideal Gas Behavior
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  10. 9 Thermochemistry
    1. Introduction
    2. 9.1 Energy Basics
    3. 9.2 Calorimetry
    4. 9.3 Enthalpy
    5. 9.4 Strengths of Ionic and Covalent Bonds
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  11. 10 Liquids and Solids
    1. Introduction
    2. 10.1 Intermolecular Forces
    3. 10.2 Properties of Liquids
    4. 10.3 Phase Transitions
    5. 10.4 Phase Diagrams
    6. 10.5 The Solid State of Matter
    7. 10.6 Lattice Structures in Crystalline Solids
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  12. 11 Solutions and Colloids
    1. Introduction
    2. 11.1 The Dissolution Process
    3. 11.2 Electrolytes
    4. 11.3 Solubility
    5. 11.4 Colligative Properties
    6. 11.5 Colloids
    7. Key Terms
    8. Key Equations
    9. Summary
    10. Exercises
  13. 12 Thermodynamics
    1. Introduction
    2. 12.1 Spontaneity
    3. 12.2 Entropy
    4. 12.3 The Second and Third Laws of Thermodynamics
    5. 12.4 Free Energy
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  14. 13 Fundamental Equilibrium Concepts
    1. Introduction
    2. 13.1 Chemical Equilibria
    3. 13.2 Equilibrium Constants
    4. 13.3 Shifting Equilibria: Le Châtelier’s Principle
    5. 13.4 Equilibrium Calculations
    6. Key Terms
    7. Key Equations
    8. Summary
    9. Exercises
  15. 14 Acid-Base Equilibria
    1. Introduction
    2. 14.1 Brønsted-Lowry Acids and Bases
    3. 14.2 pH and pOH
    4. 14.3 Relative Strengths of Acids and Bases
    5. 14.4 Hydrolysis of Salts
    6. 14.5 Polyprotic Acids
    7. 14.6 Buffers
    8. 14.7 Acid-Base Titrations
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  16. 15 Equilibria of Other Reaction Classes
    1. Introduction
    2. 15.1 Precipitation and Dissolution
    3. 15.2 Lewis Acids and Bases
    4. 15.3 Coupled Equilibria
    5. Key Terms
    6. Key Equations
    7. Summary
    8. Exercises
  17. 16 Electrochemistry
    1. Introduction
    2. 16.1 Review of Redox Chemistry
    3. 16.2 Galvanic Cells
    4. 16.3 Electrode and Cell Potentials
    5. 16.4 Potential, Free Energy, and Equilibrium
    6. 16.5 Batteries and Fuel Cells
    7. 16.6 Corrosion
    8. 16.7 Electrolysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  18. 17 Kinetics
    1. Introduction
    2. 17.1 Chemical Reaction Rates
    3. 17.2 Factors Affecting Reaction Rates
    4. 17.3 Rate Laws
    5. 17.4 Integrated Rate Laws
    6. 17.5 Collision Theory
    7. 17.6 Reaction Mechanisms
    8. 17.7 Catalysis
    9. Key Terms
    10. Key Equations
    11. Summary
    12. Exercises
  19. 18 Representative Metals, Metalloids, and Nonmetals
    1. Introduction
    2. 18.1 Periodicity
    3. 18.2 Occurrence and Preparation of the Representative Metals
    4. 18.3 Structure and General Properties of the Metalloids
    5. 18.4 Structure and General Properties of the Nonmetals
    6. 18.5 Occurrence, Preparation, and Compounds of Hydrogen
    7. 18.6 Occurrence, Preparation, and Properties of Carbonates
    8. 18.7 Occurrence, Preparation, and Properties of Nitrogen
    9. 18.8 Occurrence, Preparation, and Properties of Phosphorus
    10. 18.9 Occurrence, Preparation, and Compounds of Oxygen
    11. 18.10 Occurrence, Preparation, and Properties of Sulfur
    12. 18.11 Occurrence, Preparation, and Properties of Halogens
    13. 18.12 Occurrence, Preparation, and Properties of the Noble Gases
    14. Key Terms
    15. Summary
    16. Exercises
  20. 19 Transition Metals and Coordination Chemistry
    1. Introduction
    2. 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
    3. 19.2 Coordination Chemistry of Transition Metals
    4. 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds
    5. Key Terms
    6. Summary
    7. Exercises
  21. 20 Nuclear Chemistry
    1. Introduction
    2. 20.1 Nuclear Structure and Stability
    3. 20.2 Nuclear Equations
    4. 20.3 Radioactive Decay
    5. 20.4 Transmutation and Nuclear Energy
    6. 20.5 Uses of Radioisotopes
    7. 20.6 Biological Effects of Radiation
    8. Key Terms
    9. Key Equations
    10. Summary
    11. Exercises
  22. 21 Organic Chemistry
    1. Introduction
    2. 21.1 Hydrocarbons
    3. 21.2 Alcohols and Ethers
    4. 21.3 Aldehydes, Ketones, Carboxylic Acids, and Esters
    5. 21.4 Amines and Amides
    6. Key Terms
    7. Summary
    8. Exercises
  23. A | The Periodic Table
  24. B | Essential Mathematics
  25. C | Units and Conversion Factors
  26. D | Fundamental Physical Constants
  27. E | Water Properties
  28. F | Composition of Commercial Acids and Bases
  29. G | Standard Thermodynamic Properties for Selected Substances
  30. H | Ionization Constants of Weak Acids
  31. I | Ionization Constants of Weak Bases
  32. J | Solubility Products
  33. K | Formation Constants for Complex Ions
  34. L | Standard Electrode (Half-Cell) Potentials
  35. M | Half-Lives for Several Radioactive Isotopes
  36. Answer Key
    1. Chapter 1
    2. Chapter 2
    3. Chapter 3
    4. Chapter 4
    5. Chapter 5
    6. Chapter 6
    7. Chapter 7
    8. Chapter 8
    9. Chapter 9
    10. Chapter 10
    11. Chapter 11
    12. Chapter 12
    13. Chapter 13
    14. Chapter 14
    15. Chapter 15
    16. Chapter 16
    17. Chapter 17
    18. Chapter 18
    19. Chapter 19
    20. Chapter 20
    21. Chapter 21
  37. Index
By the end of this section, you will be able to:
  • Describe and relate the definitions of electrode and cell potentials
  • Interpret electrode potentials in terms of relative oxidant and reductant strengths
  • Calculate cell potentials and predict redox spontaneity using standard electrode potentials

Unlike the spontaneous oxidation of copper by aqueous silver(I) ions described in section 17.2, immersing a copper wire in an aqueous solution of lead(II) ions yields no reaction. The two species, Ag+(aq) and Pb2+(aq), thus show a distinct difference in their redox activity towards copper: the silver ion spontaneously oxidized copper, but the lead ion did not. Electrochemical cells permit this relative redox activity to be quantified by an easily measured property, potential. This property is more commonly called voltage when referenced in regard to electrical applications, and it is a measure of energy accompanying the transfer of charge. Potentials are measured in the volt unit, defined as one joule of energy per one coulomb of charge, V = J/C.

When measured for purposes of electrochemistry, a potential reflects the driving force for a specific type of charge transfer process, namely, the transfer of electrons between redox reactants. Considering the nature of potential in this context, it is clear that the potential of a single half-cell or a single electrode can’t be measured; “transfer” of electrons requires both a donor and recipient, in this case a reductant and an oxidant, respectively. Instead, a half-cell potential may only be assessed relative to that of another half-cell. It is only the difference in potential between two half-cells that may be measured, and these measured potentials are called cell potentials, Ecell, defined as

Ecell=EcathodeEanodeEcell=EcathodeEanode

where Ecathode and Eanode are the potentials of two different half-cells functioning as specified in the subscripts. As for other thermodynamic quantities, the standard cell potential, E°cell, is a cell potential measured when both half-cells are under standard-state conditions (1 M concentrations, 1 bar pressures, 298 K):

E°cell=E°cathodeE°anodeE°cell=E°cathodeE°anode

To simplify the collection and sharing of potential data for half-reactions, the scientific community has designated one particular half-cell to serve as a universal reference for cell potential measurements, assigning it a potential of exactly 0 V. This half-cell is the standard hydrogen electrode (SHE) and it is based on half-reaction below:

2H+(aq)+2eH2(g)2H+(aq)+2eH2(g)

A typical SHE contains an inert platinum electrode immersed in precisely 1 M aqueous H+ and a stream of bubbling H2 gas at 1 bar pressure, all maintained at a temperature of 298 K (see Figure 16.5).

The figure shows a beaker just over half full of a blue liquid. A glass tube is partially submerged in the liquid. Bubbles, which are labeled “H subscript 2 ( g )” are rising from the dark grayquare, labeled “P t electrode” at the bottom of the tube. Below the bottom of the tube pointing to the solution in the beaker is the label “ 1 M H superscript plus ( a q).” A curved arrow points up to the right, indicating the direction of the bubbles. A black wire which is labeled “P t wire” extends from the dark grgrayare up the interior of the tube through a small port at the top. A second small port extends out the top of the tube to the left. An arrow points to the port opening from the left. The base of this arrow is labeled “H subscript 2 ( g ) at 1 a t m.” A light greygray points to a diagram in a circle at the right that illustrates the surface of the P t electrode in a magnified view. P t atoms are illustrated as a uniform cluster of grey sgray which are labeled “P t electrode atoms.” On the grey atograyace, the label “e superscript negative” is shown 4 times in a nearly even vertical distribution to show electrons on the P t surface. A curved arrow extends from a white sphere labeled “H superscript plus” at the right of the P t atoms to the uppermost electron shown. Just below, a straight arrow extends from the P t surface to the right to a pair of linked white spheres which are labeled “H subscript 2.” A curved arrow extends from a second white sphere labeled “H superscript plus” at the right of the P t atoms to the second electron shown. A curved arrow extends from the third electron on the P t surface to the right to a white sphere labeled “H superscript plus.” Just below, an arrow points left from a pair of linked white spheres which are labeled “H subscript 2” to the P t surface. A curved arrow extends from the fourth electron on the P t surface to the right to a white sphere labeled “H superscript plus.”
Figure 16.5 A standard hydrogen electrode (SHE).

The assigned potential of the SHE permits the definition of a conveniently measured potential for a single half-cell. The electrode potential (EX) for a half-cell X is defined as the potential measured for a cell comprised of X acting as cathode and the SHE acting as anode:

Ecell=EXESHE ESHE=0V(defined) Ecell=EXEcell=EXESHE ESHE=0V(defined) Ecell=EX

When the half-cell X is under standard-state conditions, its potential is the standard electrode potential, E°X. Since the definition of cell potential requires the half-cells function as cathodes, these potentials are sometimes called standard reduction potentials.

This approach to measuring electrode potentials is illustrated in Figure 16.6, which depicts a cell comprised of an SHE connected to a copper(II)/copper(0) half-cell under standard-state conditions. A voltmeter in the external circuit allows measurement of the potential difference between the two half-cells. Since the Cu half-cell is designated as the cathode in the definition of cell potential, it is connected to the red (positive) input of the voltmeter, while the designated SHE anode is connected to the black (negative) input. These connections insure that the sign of the measured potential will be consistent with the sign conventions of electrochemistry per the various definitions discussed above. A cell potential of +0.337 V is measured, and so

E°cell=E°Cu=+0.337VE°cell=E°Cu=+0.337V

Tabulations of E° values for other half-cells measured in a similar fashion are available as reference literature to permit calculations of cell potentials and the prediction of the spontaneity of redox processes.

This figure contains a diagram of an electrochemical cell. Two beakers are shown. Each is just over half full. The beaker on the left contains a clear, colorless solution and is labeled below as “1 M H superscript plus.” The beaker on the right contains a blue solution and is labeled below as “1 M C u superscript 2 plus.” A glass tube in the shape of an inverted U connects the two beakers at the center of the diagram. The tube contents are colorless. The ends of the tubes are beneath the surface of the solutions in the beakers and a small graylug is present at each end of the tube. The beaker on the left has a glass tube partially submersed in the liquid. Bubbles are rising from the gray square, labeled “Standard hydrogen electrode” at the bottom of the tube. A curved arrow points up to the right, indicating the direction of the bubbles. A black wire extends from the gray square up the interior of the tube through a small port at the top to a rectangle with a digital readout of “positive 0.337 V,” which is labeled “Voltmeter.” A second small port extends out the top of the tube to the left. An arrow points to the port opening from the left. The base of this arrow is labeled “1 a t m H subscript 2 ( g ).” The beaker on the right has an orange-brown strip that is labeled “C u electrode” at the top. A wire extends from the top of this strip to the voltmeter. An arrow points toward the voltmeter from the left which is labeled “e superscript negative flow.” Similarly, an arrow points away from the voltmeter to the right. A curved arrow extends from the surrounding solution to the standard hydrogen electrode in the beaker. The end of the arrow is labeled “H subscript 2” and tip of this arrow is labeled “2 H superscript plus.” A curved arrow extends from the “C u superscript 2 plus” label in the solution to a “C u” label at the lower edge of the C u electrode. Between the two beakers is the label “T equals 298 K.”
Figure 16.6 A cell permitting experimental measurement of the standard electrode potential for the half-reaction Cu2+(aq)+2eCu(s)Cu2+(aq)+2eCu(s)

Table 16.1 provides a listing of standard electrode potentials for a selection of half-reactions in numerical order, and a more extensive alphabetical listing is given in Appendix L.

Selected Standard Reduction Potentials at 25 °C
Half-Reaction E° (V)
F2(g)+2e2F(aq)F2(g)+2e2F(aq) +2.866
PbO2(s)+SO42−(aq)+4H+(aq)+2ePbSO4(s)+2H2O(l)PbO2(s)+SO42−(aq)+4H+(aq)+2ePbSO4(s)+2H2O(l) +1.69
MnO4(aq)+8H+(aq)+5eMn2+(aq)+4H2O(l)MnO4(aq)+8H+(aq)+5eMn2+(aq)+4H2O(l) +1.507
Au3+(aq)+3eAu(s)Au3+(aq)+3eAu(s) +1.498
Cl2(g)+2e2Cl(aq)Cl2(g)+2e2Cl(aq) +1.35827
O2(g)+4H+(aq)+4e2H2O(l)O2(g)+4H+(aq)+4e2H2O(l) +1.229
Pt2+(aq)+2ePt(s)Pt2+(aq)+2ePt(s) +1.20
Br2(aq)+2e2Br(aq)Br2(aq)+2e2Br(aq) +1.0873
Ag+(aq)+eAg(s)Ag+(aq)+eAg(s) +0.7996
Hg22+(aq)+2e2Hg(l)Hg22+(aq)+2e2Hg(l) +0.7973
Fe3+(aq)+eFe2+(aq)Fe3+(aq)+eFe2+(aq) +0.771
MnO4(aq)+2H2O(l)+3eMnO2(s)+4OH(aq)MnO4(aq)+2H2O(l)+3eMnO2(s)+4OH(aq) +0.558
I2(s)+2e2I(aq)I2(s)+2e2I(aq) +0.5355
NiO2(s)+2H2O(l)+2eNi(OH)2(s)+2OH(aq)NiO2(s)+2H2O(l)+2eNi(OH)2(s)+2OH(aq) +0.49
Cu2+(aq)+2eCu(s)Cu2+(aq)+2eCu(s) +0.34
Hg2Cl2(s)+2e2Hg(l)+2Cl(aq)Hg2Cl2(s)+2e2Hg(l)+2Cl(aq) +0.26808
AgCl(s)+eAg(s)+Cl(aq)AgCl(s)+eAg(s)+Cl(aq) +0.22233
Sn4+(aq)+2eSn2+(aq)Sn4+(aq)+2eSn2+(aq) +0.151
2H+(aq)+2eH2(g)2H+(aq)+2eH2(g) 0.00
Pb2+(aq)+2ePb(s)Pb2+(aq)+2ePb(s) −0.1262
Sn2+(aq)+2eSn(s)Sn2+(aq)+2eSn(s) −0.1375
Ni2+(aq)+2eNi(s)Ni2+(aq)+2eNi(s) −0.257
Co2+(aq)+2eCo(s)Co2+(aq)+2eCo(s) −0.28
PbSO4(s)+2ePb(s)+SO42−(aq)PbSO4(s)+2ePb(s)+SO42−(aq) −0.3505
Cd2+(aq)+2eCd(s)Cd2+(aq)+2eCd(s) −0.4030
Fe2+(aq)+2eFe(s)Fe2+(aq)+2eFe(s) −0.447
Cr3+(aq)+3eCr(s)Cr3+(aq)+3eCr(s) −0.744
Mn2+(aq)+2eMn(s)Mn2+(aq)+2eMn(s) −1.185
Zn(OH)2(s)+2eZn(s)+2OH(aq)Zn(OH)2(s)+2eZn(s)+2OH(aq) −1.245
Zn2+(aq)+2eZn(s)Zn2+(aq)+2eZn(s) −0.7618
Al3+(aq)+3eAl(s)Al3+(aq)+3eAl(s) −1.662
Mg2(aq)+2eMg(s)Mg2(aq)+2eMg(s) −2.372
Na+(aq)+eNa(s)Na+(aq)+eNa(s) −2.71
Ca2+(aq)+2eCa(s)Ca2+(aq)+2eCa(s) −2.868
Ba2+(aq)+2eBa(s)Ba2+(aq)+2eBa(s) −2.912
K+(aq)+eK(s)K+(aq)+eK(s) −2.931
Li+(aq)+eLi(s)Li+(aq)+eLi(s) −3.04
Table 16.1

Example 16.4

Calculating Standard Cell Potentials What is the standard potential of the galvanic cell shown in Figure 16.3?

Solution The cell in Figure 16.3 is galvanic, the spontaneous cell reaction involving oxidation of its copper anode and reduction of silver(I) ions at its silver cathode:

cell reaction:Cu(s)+2Ag+(aq)Cu2+(aq)+2Ag(s) anode half-reaction:Cu(s)Cu2+(aq)+2e cathode half-reaction:2Ag+(aq)+2e2Ag(s)cell reaction:Cu(s)+2Ag+(aq)Cu2+(aq)+2Ag(s) anode half-reaction:Cu(s)Cu2+(aq)+2e cathode half-reaction:2Ag+(aq)+2e2Ag(s)

The standard cell potential computed as

E°cell=E°cathodeE°anode =E°AgE°Cu =0.7996V0.34V =+0.46VE°cell=E°cathodeE°anode =E°AgE°Cu =0.7996V0.34V =+0.46V

Check Your Learning What is the standard cell potential expected if the silver cathode half-cell in Figure 16.3 is replaced with a lead half-cell: Pb2+(aq)+2ePb(s)Pb2+(aq)+2ePb(s)?

Answer:

−0. 47 V

Intrepreting Electrode and Cell Potentials

Thinking carefully about the definitions of cell and electrode potentials and the observations of spontaneous redox change presented thus far, a significant relation is noted. The previous section described the spontaneous oxidation of copper by aqueous silver(I) ions, but no observed reaction with aqueous lead(II) ions. Results of the calculations in Example 16.4 have just shown the spontaneous process is described by a positive cell potential while the nonspontaneous process exhibits a negative cell potential. And so, with regard to the relative effectiveness (“strength”) with which aqueous Ag+ and Pb2+ ions oxidize Cu under standard conditions, the stronger oxidant is the one exhibiting the greater standard electrode potential, E°. Since by convention electrode potentials are for reduction processes, an increased value of corresponds to an increased driving force behind the reduction of the species (hence increased effectiveness of its action as an oxidizing agent on some other species). Negative values for electrode potentials are simply a consequence of assigning a value of 0 V to the SHE, indicating the reactant of the half-reaction is a weaker oxidant than aqueous hydrogen ions.

Applying this logic to the numerically ordered listing of standard electrode potentials in Table 16.1 shows this listing to be likewise in order of the oxidizing strength of the half-reaction’s reactant species, decreasing from strongest oxidant (most positive E°) to weakest oxidant (most negative E°). Predictions regarding the spontaneity of redox reactions under standard state conditions can then be easily made by simply comparing the relative positions of their table entries. By definition, E°cell is positive when E°cathode > E°anode, and so any redox reaction in which the oxidant’s entry is above the reductant’s entry is predicted to be spontaneous.

Reconsideration of the two redox reactions in Example 16.4 provides support for this fact. The entry for the silver(I)/silver(0) half-reaction is above that for the copper(II)/copper(0) half-reaction, and so the oxidation of Cu by Ag+ is predicted to be spontaneous (E°cathode > E°anode and so E°cell > 0). Conversely, the entry for the lead(II)/lead(0) half-cell is beneath that for copper(II)/copper(0), and the oxidation of Cu by Pb2+ is nonspontaneous (E°cathode < E°anode and so E°cell < 0).

Recalling the chapter on thermodynamics, the spontaneities of the forward and reverse reactions of a reversible process show a reciprocal relationship: if a process is spontaneous in one direction, it is non-spontaneous in the opposite direction. As an indicator of spontaneity for redox reactions, the potential of a cell reaction shows a consequential relationship in its arithmetic sign. The spontaneous oxidation of copper by lead(II) ions is not observed,

Cu(s)+Pb2+(aq)Cu2+(aq)+Pb(s)E°forward=−0.47V(negative, non-spontaneous)Cu(s)+Pb2+(aq)Cu2+(aq)+Pb(s)E°forward=−0.47V(negative, non-spontaneous)

and so the reverse reaction, the oxidation of lead by copper(II) ions, is predicted to occur spontaneously:

Pb(s)+Cu2+(aq)Pb2+(aq)+Cu(s)E°forward=+0.47V(positive, spontaneous)Pb(s)+Cu2+(aq)Pb2+(aq)+Cu(s)E°forward=+0.47V(positive, spontaneous)

Note that reversing the direction of a redox reaction effectively interchanges the identities of the cathode and anode half-reactions, and so the cell potential is calculated from electrode potentials in the reverse subtraction order than that for the forward reaction. In practice, a voltmeter would report a potential of −0.47 V with its red and black inputs connected to the Pb and Cu electrodes, respectively. If the inputs were swapped, the reported voltage would be +0.47 V.

Example 16.5

Predicting Redox Spontaneity Are aqueous iron(II) ions predicted to spontaneously oxidize elemental chromium under standard state conditions? Assume the half-reactions to be those available in Table 16.1.

Solution Referring to the tabulated half-reactions, the redox reaction in question can be represented by the equations below:

Cr(s)+Fe2+(aq)Cr3+(aq)+Fe(s)Cr(s)+Fe2+(aq)Cr3+(aq)+Fe(s)

The entry for the putative oxidant, Fe2+, appears above the entry for the reductant, Cr, and so a spontaneous reaction is predicted per the quick approach described above. Supporting this predication by calculating the standard cell potential for this reaction gives

E°cell=E°cathodeE°anode =E°Fe(II)E°Cr =−0.447V−0.774V=+0.297VE°cell=E°cathodeE°anode =E°Fe(II)E°Cr =−0.447V−0.774V=+0.297V

The positive value for the standard cell potential indicates the process is spontaneous under standard state conditions.

Check Your Learning Use the data in Table 16.1 to predict the spontaneity of the oxidation of bromide ion by molecular iodine under standard state conditions, supporting the prediction by calculating the standard cell potential for the reaction. Repeat for the oxidation of iodide ion by molecular bromine.

Answer:

I2(s)+2Br(aq)2I(aq)+Br2(l)E°cell=+0.5518V(spontaneous) Br2(s)+2I(aq)2Br(aq)+I2(l)E°cell=−0.5518V(nonspontaneous)I2(s)+2Br(aq)2I(aq)+Br2(l)E°cell=+0.5518V(spontaneous) Br2(s)+2I(aq)2Br(aq)+I2(l)E°cell=−0.5518V(nonspontaneous)

Citation/Attribution

Want to cite, share, or modify this book? This book is Creative Commons Attribution License 4.0 and you must attribute OpenStax.

Attribution information
  • If you are redistributing all or part of this book in a print format, then you must include on every physical page the following attribution:
    Access for free at https://openstax.org/books/chemistry-atoms-first-2e/pages/1-introduction
  • If you are redistributing all or part of this book in a digital format, then you must include on every digital page view the following attribution:
    Access for free at https://openstax.org/books/chemistry-atoms-first-2e/pages/1-introduction
Citation information

© Feb 14, 2019 OpenStax. Textbook content produced by OpenStax is licensed under a Creative Commons Attribution License 4.0 license. The OpenStax name, OpenStax logo, OpenStax book covers, OpenStax CNX name, and OpenStax CNX logo are not subject to the Creative Commons license and may not be reproduced without the prior and express written consent of Rice University.