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Introductory Chemistry - 1st Canadian Edition: Free Energy under Nonstandard Conditions

Introductory Chemistry - 1st Canadian Edition
Free Energy under Nonstandard Conditions
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table of contents
  1. Cover
  2. Title Page
  3. Copyright
  4. Table Of Contents
  5. Acknowledgments
  6. Dedication
  7. About BCcampus Open Education
  8. Chapter 1. What is Chemistry
    1. Some Basic Definitions
    2. Chemistry as a Science
  9. Chapter 2. Measurements
    1. Expressing Numbers
    2. Significant Figures
    3. Converting Units
    4. Other Units: Temperature and Density
    5. Expressing Units
    6. End-of-Chapter Material
  10. Chapter 3. Atoms, Molecules, and Ions
    1. Acids
    2. Ions and Ionic Compounds
    3. Masses of Atoms and Molecules
    4. Molecules and Chemical Nomenclature
    5. Atomic Theory
    6. End-of-Chapter Material
  11. Chapter 4. Chemical Reactions and Equations
    1. The Chemical Equation
    2. Types of Chemical Reactions: Single- and Double-Displacement Reactions
    3. Ionic Equations: A Closer Look
    4. Composition, Decomposition, and Combustion Reactions
    5. Oxidation-Reduction Reactions
    6. Neutralization Reactions
    7. End-of-Chapter Material
  12. Chapter 5. Stoichiometry and the Mole
    1. Stoichiometry
    2. The Mole
    3. Mole-Mass and Mass-Mass Calculations
    4. Limiting Reagents
    5. The Mole in Chemical Reactions
    6. Yields
    7. End-of-Chapter Material
  13. Chapter 6. Gases
    1. Pressure
    2. Gas Laws
    3. Other Gas Laws
    4. The Ideal Gas Law and Some Applications
    5. Gas Mixtures
    6. Kinetic Molecular Theory of Gases
    7. Molecular Effusion and Diffusion
    8. Real Gases
    9. End-of-Chapter Material
  14. Chapter 7. Energy and Chemistry
    1. Formation Reactions
    2. Energy
    3. Stoichiometry Calculations Using Enthalpy
    4. Enthalpy and Chemical Reactions
    5. Work and Heat
    6. Hess’s Law
    7. End-of-Chapter Material
  15. Chapter 8. Electronic Structure
    1. Light
    2. Quantum Numbers for Electrons
    3. Organization of Electrons in Atoms
    4. Electronic Structure and the Periodic Table
    5. Periodic Trends
    6. End-of-Chapter Material
  16. Chapter 9. Chemical Bonds
    1. Lewis Electron Dot Diagrams
    2. Electron Transfer: Ionic Bonds
    3. Covalent Bonds
    4. Other Aspects of Covalent Bonds
    5. Violations of the Octet Rule
    6. Molecular Shapes and Polarity
    7. Valence Bond Theory and Hybrid Orbitals
    8. Molecular Orbitals
    9. End-of-Chapter Material
  17. Chapter 10. Solids and Liquids
    1. Properties of Liquids
    2. Solids
    3. Phase Transitions: Melting, Boiling, and Subliming
    4. Intermolecular Forces
    5. End-of-Chapter Material
  18. Chapter 11. Solutions
    1. Colligative Properties of Solutions
    2. Concentrations as Conversion Factors
    3. Quantitative Units of Concentration
    4. Colligative Properties of Ionic Solutes
    5. Some Definitions
    6. Dilutions and Concentrations
    7. End-of-Chapter Material
  19. Chapter 12. Acids and Bases
    1. Acid-Base Titrations
    2. Strong and Weak Acids and Bases and Their Salts
    3. Brønsted-Lowry Acids and Bases
    4. Arrhenius Acids and Bases
    5. Autoionization of Water
    6. Buffers
    7. The pH Scale
    8. End-of-Chapter Material
  20. Chapter 13. Chemical Equilibrium
    1. Chemical Equilibrium
    2. The Equilibrium Constant
    3. Shifting Equilibria: Le Chatelier’s Principle
    4. Calculating Equilibrium Constant Values
    5. Some Special Types of Equilibria
    6. End-of-Chapter Material
  21. Chapter 14. Oxidation and Reduction
    1. Oxidation-Reduction Reactions
    2. Balancing Redox Reactions
    3. Applications of Redox Reactions: Voltaic Cells
    4. Electrolysis
    5. End-of-Chapter Material
  22. Chapter 15. Nuclear Chemistry
    1. Units of Radioactivity
    2. Uses of Radioactive Isotopes
    3. Half-Life
    4. Radioactivity
    5. Nuclear Energy
    6. End-of-Chapter Material
  23. Chapter 16. Organic Chemistry
    1. Hydrocarbons
    2. Branched Hydrocarbons
    3. Alkyl Halides and Alcohols
    4. Other Oxygen-Containing Functional Groups
    5. Other Functional Groups
    6. Polymers
    7. End-of-Chapter Material
  24. Chapter 17. Kinetics
    1. Factors that Affect the Rate of Reactions
    2. Reaction Rates
    3. Rate Laws
    4. Concentration–Time Relationships: Integrated Rate Laws
    5. Activation Energy and the Arrhenius Equation
    6. Reaction Mechanisms
    7. Catalysis
    8. End-of-Chapter Material
  25. Chapter 18. Chemical Thermodynamics
    1. Spontaneous Change
    2. Entropy and the Second Law of Thermodynamics
    3. Measuring Entropy and Entropy Changes
    4. Gibbs Free Energy
    5. Spontaneity: Free Energy and Temperature
    6. Free Energy under Nonstandard Conditions
    7. End-of-Chapter Material
  26. Appendix A: Periodic Table of the Elements
  27. Appendix B: Selected Acid Dissociation Constants at 25°C
  28. Appendix C: Solubility Constants for Compounds at 25°C
  29. Appendix D: Standard Thermodynamic Quantities for Chemical Substances at 25°C
  30. Appendix E: Standard Reduction Potentials by Value
  31. Glossary
  32. About the Authors
  33. Versioning History

Free Energy under Nonstandard Conditions

Jessie A. Key

Learning Objectives

  • To be able to determine free energy at nonstandard conditions using the standard change in Gibbs free energy, ΔG°.
  • To understand and use the relationship between ΔG° and the equilibrium constant K.

Many reactions do not occur under standard conditions, and therefore we need some ways of determining the free energy under nonstandard conditions.

Using Standard Change in Gibbs Free Energy, ΔG°

The change in Gibbs free energy under nonstandard conditions, ΔG, can be determined from the standard change in Gibbs free energy, ΔG°:

\Delta G=\Delta G^{\circ}+RT \ln Q

where R is the ideal gas constant 8.314 J/mol K, Q is the reaction quotient, and T is the temperature in Kelvin.

Under standard conditions, the reactant and product solution concentrations are 1 M, or the pressure of gases is 1 bar, and Q is equal to 1. Taking the natural logarithm simplifies the equation to:

\Delta G=\Delta G^{\circ}\text{ (under standard conditions)}

Under nonstandard conditions, Q must be calculated (in a manner similar to the calculation for an equilibrium constant). For gases, the concentrations are expressed as partial pressures in the units of either atmospheres or bars, and solutes in the units of molarity.

For the reaction aA + bB\leftrightharpoons cC+dD:

Q_{\text{gases}}=\dfrac{\left(P_C\right)^c\left(P_D\right)^d}{\left(P_A\right)^a\left(P_B\right)^b}\hspace{2em}\text{or}\hspace{2em}Q_{\text{solutes}}=\dfrac{[C]^c[D]^d}{[A]^a[B]^b}

Example 18.7

Consider the following reaction:

4NH3(g) + 5O2(g) ⇌ 6H2O(g) + 4NO(g)

  1. Use the thermodynamic data in the appendix to calculate ΔG° at 298 K.
  2. Calculate ΔG at 298 K for a mixture of 2.0 bar NH3(g), 1.0 bar O2(g), 1.5 bar H2O(g), and 1.2 bar NO(g).

Solution

  1. \phantom{start}

        \begin{align*} \Delta G^{\circ}&=\underset{\text{products}}{\sum n\Delta G^{\circ}{}_{\text{f}}}-\underset{\text{reactants}}{\sum m\Delta G^{\circ}{}_{\text{f}}} \\ &=[(6\times -228.6\text{ kJ/mol})+(4\times 87.6\text{ kJ/mol})] \\ &\phantom{=}-[(4\times -16.4\text{ kJ/mol})+(5\times 0.0\text{ kJ/mol})] \\ &=(-1021.2\text{ kJ/mol})-(-65.6\text{ kJ/mol}) \\ &=-955.6\text{ kJ/mol}=-9.56\times 10^5\text{ J/mol} \end{align*}

  2. \phantom{start}

        \begin{align*} \Delta G&=\Delta G^{\circ}+RT \ln Q \\ Q&=\dfrac{(1.5\text{ bar})^6(1.2\text{ bar})^4}{(2.0\text{ bar})^4(1.0\text{ bar})^5}=1.5 \\ \\ \Delta G&=(-9.56\times 10^5\text{ J/mol})+(8.314\text{ J/mol K})(298\text{ K})\ln(1.5) \\ &=(-9.56\times 10^5\text{ J/mol})+(1.0\times 10^3\text{ J/mol}) \\ &=-9.6\times 10^5\text{ J/mol} \end{align*}

The Relationship between ΔG° and K

There is a direct relationship between ΔG⁰ and the equilibrium constant K. We can establish this relationship by substituting the equilibrium values (ΔG = 0, and K = Q) into the equation for determining free energy change under nonstandard conditions:

    \begin{align*} \Delta G&=\Delta G^{\circ}+RT\ln Q \\ 0&=\Delta G^{\circ}+RT\ln K \\ \Delta G^{\circ}&=-RT\ln K \end{align*}

We now have a way of  relating the equilibrium constant directly to changes in enthalpy and entropy. As well, we can now determine the equilibrium constant from thermochemical data tables or determine the standard free energy change from equilibrium constants.

Example 18.8

The Ksp for CuI(s) at 25°C is 1.27 × 10−12. Determine ΔG° for the following:

Cu+(aq) + I−(aq) → CuI(s)

Solution

\Delta G^{\circ}=-RT\ln K

The equation given is in the opposite direction to the definition of Ksp:

    \begin{align*} K&=\dfrac{1}{K_{\text{sp}}}=\left(\dfrac{1}{1.27}\times 10^{-12}\right)=7.87\times 10^{11} \\ \\ \Delta G^{\circ}&=-(8.314\text{ J/mol K})(298\text{ K})\ln(7.87\times 10^{11}) \\ &=-(8.314\text{ J/mol K})(298\text{ K})(27.4) \\ &=-6.79\times 10^4\text{ J/mol}=-67.9\text{ kJ/mol} \\ \end{align*}

Key Takeaways

  • The free energy at nonstandard conditions can be determined using ΔG = ΔG° + RT ln Q.
  • There is a direct relationship between ΔG° and the equilibrium constant K: ΔG° = −RT ln K.

Icon for the Creative Commons Attribution 4.0 International License

Free Energy under Nonstandard Conditions by Jessie A. Key is licensed under a Creative Commons Attribution 4.0 International License, except where otherwise noted.

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Copyright © 2014

                                by Jessie A. Key

            Introductory Chemistry - 1st Canadian Edition by Jessie A. Key is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.
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