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 ↓


Gibbs Free Energy


Chemical thermodynamics is a fascinating field of study that deals with the energy changes that accompany chemical reactions and physical transformations. One of the important concepts in this field is Gibbs free energy, named after American scientist Josiah Willard Gibbs. It is an important thermodynamic function that helps determine the favorability and spontaneity of processes. In this section, we will learn in more detail about Gibbs free energy, its importance, and its applications in undergraduate-level chemistry.

Defining Gibbs Free Energy

Gibbs free energy is represented by the symbol G. The change in Gibbs free energy (ΔG) during a process tells us whether it is spontaneous or non-spontaneous at constant temperature and pressure. The formula for Gibbs free energy is given as:

    ΔG = ΔH – TΔS

Where:

  • ΔG is the change in Gibbs free energy
  • ΔH is the change in enthalpy
  • T is the absolute temperature (in Kelvin)
  • ΔS is the change in entropy

This equation states that the Gibbs free energy change of a process is obtained from the balance between the enthalpy change (the absorbed or released energy) and the entropy change (the measure of disorder) at a given temperature. The nature of ΔG determines the direction of spontaneity:

  • If ΔG < 0, then the process is spontaneous.
  • If ΔG > 0, then the process is non-spontaneous.
  • If ΔG = 0, then the system is in equilibrium.

Understanding ease

The concept of spontaneity is important in chemical reactions. A spontaneous process occurs without external intervention. It can be compared to a ball rolling down a hill, as it will naturally move from a state of higher potential energy to a state of lower potential energy spontaneously.

Consider two scenarios:

  • Exothermic spontaneous reaction: When the reaction releases heat (ΔH is negative), and there is an increase in entropy (ΔS is positive), the process is spontaneous. Example: Combustion of glucose in oxygen.
  • Endothermic spontaneous reaction: Even if a reaction absorbs heat (ΔH is positive), if there is a sufficient increase in entropy (ΔS is positive), it can still be spontaneous. Example: Ice melting at room temperature.

Mathematical explanation

Let's take a closer look at the equations, where mathematicians and scientists often use these expressions to predict the feasibility of reactions.

    ΔG = ΔH – TΔS

This equation shows that spontaneity depends on the interaction between ΔH and ΔS. As the temperature changes, TΔS term can exceed the change in ΔH, changing the spontaneity of the reaction.

Examples to understand ΔG calculation:

Example 1: Consider the vaporization of water. At 100°C, calculate ΔG if ΔH = +40.65 kJ/mol and ΔS = +0.109 kJ/mol·K.

    ΔG = ΔH – TΔS
    ΔG = 40.65 kJ/mol – (373 K * 0.109 kJ/mol K)
    ΔG = 40.65 kJ/mol – 40.657 kJ/mol
    ΔG = -0.007 kJ/mol

Since ΔG is negative, evaporation is spontaneous at 100°C (373 K).

Example 2: For the given reaction at 25°C, calculate ΔG if ΔH = +24 kJ/mol and ΔS = -0.010 kJ/mol K.

    ΔG = ΔH – TΔS
    ΔG = 24 kJ/mol – (298 K * -0.010 kJ/mol K)
    ΔG = 24 kJ/mol - (-2.98 kJ/mol)
    ΔG = 24 kJ/mol + 2.98 kJ/mol
    ΔG = 26.98 kJ/mol

Since ΔG is positive, the reaction does not occur spontaneously at 25°C.

The concept of balance

In chemistry, equilibrium is the state where the rates of the forward and backward reactions are equal, resulting in no net change in the concentrations of reactants and products. In this state, ΔG is equal to zero. The system is stable, and no work is produced.

Example: Consider the formation of ammonia from a mixture of nitrogen and hydrogen gases at high pressure and temperature. A state of equilibrium is achieved when ammonia is produced as fast as it is broken down into nitrogen and hydrogen.

represented by the equation:

    N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g)

Graphical representation

To better understand Gibbs free energy consider the following visual example:

           ΔG ΔH
            ,
            ,
            ,
          Preliminary Final

In the graph, the energy of the initial state and the final state is plotted on the vertical axis. The change in Gibbs free energy is represented as the difference between the initial and final states.

Predicting the feasibility of a response

Chemists use Gibbs free energy calculations to estimate the feasibility and direction of chemical processes. Let's break it down into parts:

  • Negative ΔG: This result indicates that the reaction is favorable in terms of energy. For example, making water from hydrogen and oxygen releases energy, making it spontaneous.
  • Positive ΔG: In scenarios where ΔG is positive, the reaction requires external energy to proceed, such as the process of electrolysis of water to obtain hydrogen and oxygen gases.
  • Zero ΔG: This refers to equilibrium, allowing chemists to predict the concentration of a substance where the forward and reverse reaction rates are equal.

Applications in chemistry

Beyond theoretical discussions, Gibbs free energy is invaluable in practical chemical applications:

  • Biochemical processes: In biological systems, reactions such as ATP hydrolysis and cellular respiration involve cumulative free energy changes. This understanding is important for bioenergetics.
  • Industrial chemistry: Manufacturing reactions leverage knowledge of Gibbs free energy to optimize conditions. In ammonia synthesis or metal extraction, these insights facilitate cost-effective and sustainable practices.
  • Materials science: Predictive modeling of the stability and transformation of materials involves calculation of Gibbs free energy, which is important for the design of new advanced materials.

Additional considerations

  • Le Chatelier's principle: Changing conditions such as temperature can shift the equilibrium constants for reactions, affecting ΔG.
  • Standard conditions and state functions: The standard Gibbs free energy change (ΔG°) is also calculated under specific conditions (1 bar, 298.15 K), which provides a reference point for comparison among different reactions.
  • Non-standard conditions: In real-world applications, reactions do not always conform to standard conditions. The formula adjusts for non-standard Gibbs energy changes by including the reaction quotient.
    ΔG = ΔG° + RTlnQ

Where:

  • ΔG° is the standard Gibbs free energy
  • R is the universal gas constant
  • Q is the reaction quotient

Summary

In summary, Gibbs Free Energy is a key concept in chemical thermodynamics, providing information about the spontaneity and feasibility of reactions. From predicting the behavior of reactions under different conditions to practical applications in many fields, its importance cannot be underestimated. Armed with the equation ΔG = ΔH - TΔS, chemists and scientists predict, design and innovate chemical processes that are fundamental to technological progress and understanding the natural world.

Through Gibbs free energy emerges a rich fabric of knowledge and applications, reflecting the early explorations by scientists to better understand and harness the power of chemical reactions for the benefit of humanity.


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Gibbs Free Energy

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