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 → General chemistry ↓


Thermodynamics


Thermodynamics is a branch of physical chemistry that deals with the study of energy transformations in chemical processes. Understanding the laws of thermodynamics is important in predicting how chemical reactions occur and in determining the feasibility and ease of chemical processes.

Basic concepts and definitions

To understand thermodynamics it is necessary to understand some basic concepts:

  • System: This refers to the part of the universe we are interested in studying. This could be a reaction inside a beaker.
  • Environment: Everything outside the system.
  • Boundary: The separation between the system and the environment.
  • The state of a system: defined by properties such as pressure, volume, temperature, and composition.

Laws of thermodynamics

There are four laws of thermodynamics that help us understand the flow of energy. These laws are fundamental to physics and chemistry.

First law of thermodynamics

The first law is also called the law of conservation of energy. It states that energy can neither be created nor destroyed, it can only be converted from one form to another.

ΔU = Q - W

Here, ΔU is the change in internal energy of the system, Q is the heat added to the system, and W is the work done by the system.

For example, if you heat water in a pot, heat energy from the stove (surroundings) is transferred to the water (system), increasing its internal energy and making it hotter.

Second law of thermodynamics

The second law introduces the concept of entropy, which is a measure of disorder or randomness in a system. It states that for any spontaneous process, the total entropy of the system and its surroundings always increases.

heat sourceheat sink

In the diagram above, a heat source transfers energy to a heat sink. During this process, some energy always dissipates as unusable, increasing the total entropy.

Third law of thermodynamics

This law states that as the temperature of a system approaches absolute zero, the entropy of an ideal crystal approaches a stable minimum.

Think of it as trying to put all the pieces of a puzzle in the right order. As the temperature approaches absolute zero, molecular motion effectively stops, and the entropy of the system approaches a minimum.

Zeroth law of thermodynamics

The zeroth law of thermodynamics states that if two systems are in thermal equilibrium with a third system, then they are also in thermal equilibrium with each other. This is fundamental in defining temperature.

System ASystem C Thermal equilibriumSystem B

In the above figure, systems A and C are each in thermal equilibrium with system B, so A and C are also in equilibrium with each other.

Enthalpy and its importance

Enthalpy, denoted H, is a measure of the heat content in a system at constant pressure. It is a detailed state function used to determine the heat exchanged in chemical reactions.

ΔH = ΔU + PΔV

where ΔH is the change in enthalpy, ΔU is the change in internal energy, P is the pressure, and ΔV is the change in volume.

During an exothermic reaction (such as combustion) heat is released, resulting in a negative ΔH. For an endothermic reaction (such as melting ice) heat is absorbed, and ΔH is positive.

Gibbs free energy

Gibbs free energy, G, helps predict the spontaneity of reactions at constant pressure and temperature. It is defined as:

G = H - TS

where G is the Gibbs free energy, H is the enthalpy, T is the temperature, and S is the entropy.

A negative change in Gibbs free energy (ΔG) indicates a spontaneous process, while a positive change indicates a non-spontaneous process. If ΔG = 0, the system is in equilibrium.

Applications of thermodynamics

Thermodynamics plays an important role in various branches of chemistry and everyday applications, including:

  • Chemical reactions: Helps in predicting the feasibility of a reaction.
  • Phase transitions: Understanding melting point, boiling point.
  • Engineering: In designing engines and refrigerators.

For example, in engineering, thermodynamics aids in the design of engines and refrigeration systems by optimizing processes for energy efficiency. In biological systems, it helps explain how cells manage energy through processes such as respiration and photosynthesis.

Conclusion

Thermodynamics provides essential principles that are widely applicable in chemistry and engineering. By understanding how energy is transformed and conserved in chemical processes, scientists can predict and control the outcomes of chemical reactions. The second law concepts of entropy and the Gibbs free energy function, in particular, are invaluable in analyzing chemical processes and directing them toward desired outcomes.


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Thermodynamics

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