Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → Physical chemistry → Chemical thermodynamics ↓


Thermodynamic efficiency


Chemical thermodynamics is a branch of physical chemistry that deals with the study of energy changes and conversions in chemical processes. One of the key concepts in chemical thermodynamics is thermodynamic potential. Thermodynamic potentials are functions used to describe the state and evolution of a thermodynamic system. They help us understand how energy is distributed within a system and how it can be used to do work. The aim of this article is to explore the importance of thermodynamic potentials in chemical thermodynamics.

Basic concepts

To fully understand what thermodynamic potential is, it is necessary to first understand some of the fundamental concepts of thermodynamics, such as systems, states, and processes.

  • Thermodynamic system: This is the part of the universe we are focusing on, which is separated from the surrounding environment by a boundary. For example, a gas in a cylinder is a thermodynamic system.
  • State of the system: The state of a system is described by its properties such as temperature (T), pressure (P), volume (V), and amount of matter (n). State functions are properties that depend only on the state of the system, not on how it reached that state.
  • Thermodynamic processes: These are the paths the system takes when moving from one state to another. Processes can be isothermal (constant temperature), isobaric (constant pressure), adiabatic (no heat exchange), and more.

What are thermodynamic potentials?

Thermodynamic potentials are scalar quantities derived from the internal energy of a system, which help to analyze and predict the behavior of thermodynamic systems. There are four main thermodynamic potentials:

  • Internal energy (U): It is the total energy possessed by a system, including kinetic and potential energy at the microscopic level.
  • Enthalpy (H): Defined as H = U + PV, where P is pressure and V is volume. Enthalpy is useful in processes that occur at constant pressure.
  • Helmholtz free energy (A): It is defined as A = U - TS where T is temperature and S is entropy. It is useful for processes at constant volume and temperature.
  • Gibbs free energy (G): Defined as G = H - TS. Gibbs free energy is particularly useful for processes at constant pressure and temperature, such as most chemical reactions.

Internal energy

Internal energy U is a fundamental concept in thermodynamics. It represents the total energy of a system, including kinetic and potential energy at the molecular level. Internal energy changes during heat transfer and work done by or on the system. For a process:

    ΔU = Q – W

Where:

  • ΔU is the change in internal energy.
  • Q is the heat added to the system.
  • W is the work done by the system.

The internal energy is a state function, meaning that it depends entirely on the state of the system, regardless of how it got there.

Enthalpy

Enthalpy H is another important thermodynamic potential, especially in processes occurring at constant pressure. It is defined by the equation:

    H = U + PV

The enthalpy change is often a convenient measure of the heat absorbed or released during a process at constant pressure, such as chemical reactions:

    ΔH = ΔU + PΔV

In exothermic reactions, enthalpy decreases when heat is released, while in endothermic reactions, enthalpy increases when heat is absorbed.

Helmholtz free energy

The Helmholtz free energy A is defined for systems located at constant volume and temperature. It is formulated as follows:

    A = U – TS

The Helmholtz free energy indicates the maximum work a system can do at constant volume and temperature, and is often used in statistical mechanics and quantum chemistry.

Its transformation is given by the equation:

    ΔA = ΔU – TΔS

This tells us that for a process at constant temperature, the change in the Helmholtz free energy is due to the change in internal energy and entropy.

Gibbs free energy

The Gibbs free energy G is the most important thermodynamic potential for chemists, because it applies to processes at constant pressure and temperature. It is defined as:

    G = H – TS

Gibbs free energy is important in predicting the spontaneity of chemical reactions:

  • If ΔG < 0, then the process is spontaneous.
  • If ΔG = 0, then the system is in equilibrium.
  • If ΔG > 0, then the process will occur spontaneously.

For a chemical reaction at constant pressure and temperature, the change in Gibbs free energy is given by:

    ΔG = ΔH – TΔS

Thermodynamic potential and equilibrium

Thermodynamic potentials also give information about the equilibrium properties of a system. When a system is at equilibrium, its thermodynamic potential reaches a minimum. For example:

  • In isothermal, isochoric processes, a system will achieve equilibrium by minimizing its Helmholtz free energy.
  • For isothermal, isobaric processes, systems achieve equilibrium by minimizing the Gibbs free energy.

Application of thermodynamic potential

Thermodynamic potentials are widely used in various fields from chemistry to physics. Here are some applications:

  • Chemical reactions: Gibbs free energy helps predict the spontaneity and equilibrium state of a reaction.
  • Phase transitions: Enthalpy is important in analyzing heat changes during phase transitions, such as melting or boiling.
  • Thermal systems: Internal energy considerations are important for the optimization of heat engines and refrigerators.
  • Statistical mechanics: Helmholtz free energy is essential to understanding the behavior of particles in a system.

Example with simple visualizations

Consider a simple reaction A + B ⇌ C occurring at constant pressure and temperature. We will investigate its spontaneity using Gibbs free energy changes:

    ΔG = ΔH – TΔS
Reactants Products ΔG

If ΔG < 0, then the reaction from reactants to products is spontaneous, as shown in the figure, where the products have a lower Gibbs free energy than the reactants.

Conclusion

Thermodynamic potentials are essential tools in chemical thermodynamics. They provide insight into the energy dynamics of various processes and help predict the direction and feasibility of reactions. Understanding these concepts enables us to effectively use and manipulate energy in a variety of chemical processes and applications. Mastery of thermodynamic potentials shapes the understanding of a wide range of chemical and physical systems, empowering both theoretical insights and practical applications.


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Thermodynamic efficiency

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