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


Chemical thermodynamics is a branch of physical chemistry that deals with the study of the interrelation of heat and work with chemical reactions or physical changes of state within the limits of the laws of thermodynamics. It is basically concerned with the conservation and transformation of energy in matter. To understand chemical thermodynamics, it is necessary to understand several key concepts, such as systems, states, processes, energy, enthalpy, entropy, and Gibbs free energy.

Basic concepts of thermodynamics

Chemical thermodynamics is based on some fundamental concepts and laws that explain how energy is exchanged and transformed in chemical processes. We begin with some key terms:

System and environment

In thermodynamics, a system is the part of the universe we are studying, while the surroundings is everything outside the system. Systems can be classified into three types:

  • Open system: can exchange both energy and matter with its surrounding environment.
  • Closed system: can only exchange energy, not matter, with its surroundings.
  • Isolated system: cannot exchange energy or matter with its surroundings.

System status

The state of a system is described by its properties, which may change when changes occur in the system. State functions such as temperature, pressure, volume, and amount of matter describe these states. The total energy of the system is another important state function.

Processes

A process is a change from one state to another. Some processes are reversible, and others are irreversible. A reversible process is a hypothetical situation where the system changes state in such a way that the net energy or matter exchange with the surrounding environment can be reversed by infinitesimal adjustments. Most natural processes are irreversible.

First law of thermodynamics

The first law of thermodynamics is essentially the law of conservation of energy. It states that energy cannot be created or destroyed, only converted from one form to another. The internal energy U of a system changes when energy q is added by heating or when work w is done on or by the system. Mathematically, it is expressed as:

ΔU = q + w

Example: Consider heating a gas in a piston. The gas does work on the piston and its internal energy increases. The first law of thermodynamics helps us understand how much work is done due to the heat added to the system.

Gas Piston System

Second law of thermodynamics

The second law of thermodynamics introduces the concept of entropy, which is a measure of disorder or randomness in a system. It states that the total entropy of an isolated system can never decrease over time. Entropy can be transferred between the system and the surroundings, but for a process to be spontaneous, the total entropy must increase.

ΔS_total = ΔS_system + ΔS_surroundings ≥ 0

Example: The melting of ice is an example of an increase in entropy, as the ordered crystalline structure of ice becomes more random as it turns into liquid water.

Third law of thermodynamics

The third law of thermodynamics states that as the temperature of a pure, perfectly crystalline substance approaches absolute zero (0 Kelvin), its entropy approaches zero. This law provides a reference point for calculating the absolute entropy of substances.

Enthalpy

Enthalpy is a state function describing the total energy of a system, consisting of the internal energy and the product of its pressure and volume:

H = U + PV

Enthalpy is useful in chemical reactions occurring at constant pressure, where the change in enthalpy is equal to the heat exchanged with the surroundings.

Example: In a chemical reaction where reactants form products, the change in enthalpy tells us whether the reaction is endothermic (heat absorbing) or exothermic (heat releasing).

Gibbs free energy

Gibbs free energy combines enthalpy and entropy into a single value that describes the spontaneity of a system at constant temperature and pressure. The change in Gibbs free energy (ΔG) is given by:

ΔG = ΔH - TΔS

Negative ΔG indicates a spontaneous process, while positive ΔG indicates a non-spontaneous process.

Example: Combustion reactions have a negative ΔG value, indicating that they are spontaneous.

Conclusion

Chemical thermodynamics is a powerful tool that helps chemists understand how energy is transformed in chemical reactions and processes. By applying the laws of thermodynamics, chemists can predict whether a process will be spontaneous, how equilibrium will be established, and how energy will flow through systems.

In summary, the principles of chemical thermodynamics provide essential insights into the functioning of chemical systems, and guide both theoretical investigations and practical applications in chemistry.


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Chemical thermodynamics

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