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 ↓


Phase transition


In the study of chemistry and particularly in the field of chemical thermodynamics, phase transitions are fascinating phenomena that involve the transformation of matter from one phase to another. This process is driven by changes in temperature and pressure, as well as other factors. In this detailed explanation, we will explore what phase transitions are, the different types of phase transitions, and discuss their importance in both natural processes and industrial applications.

Understanding phases of matter

Before diving into phase transitions, it is important to clearly understand what we mean by phases of matter. Matter exists in several different states or phases, and these include:

  • Solid: In this state the particles are tightly bound to each other and have a definite shape and volume. The intermolecular forces between the particles are strong, which maintain this definite structure.
  • Liquids: Here the particles are still packed close to each other but are free to move around each other. Liquids have a fixed volume but they take the shape of their container.
  • Gas: In the gaseous state the particles are far apart from each other and move around freely. Gases have neither a definite shape nor a definite volume.
  • Plasma: A less common stage of highly ionized gas that contains free electrons as a result of increased energy levels. This stage is found in stars, including the Sun.

What is a phase transition?

A phase transition, also known as a phase change, is a process in which matter changes from one phase to another. This transition occurs when energy, usually in the form of heat, is added to or removed from a system. The most common example of a phase transition is the change from solid ice to liquid water and eventually to water vapor (gas).

Example of a phase transition

Consider the transition of water from ice to liquid state. At 0°C (32°F), ice melts, meaning it changes from a solid state to a liquid state. This process is called melting. When water is heated to 100°C (212°F) at standard atmospheric pressure, it changes from a liquid state to a gaseous state, a process called evaporation or boiling.

In chemical terms, consider the notion of an equilibrium state during these transitions:

Ice (s) ↔ Water (l) Water (l) ↔ Vapor (g)

Types of phase transitions

Phase transitions are typically classified based on the nature of the transformation process. Here, we explore the most common types:

First-order phase transition

These transitions involve a latent heat, meaning they absorb or release a certain amount of energy during the transition. During first-order phase transitions, there are discontinuous changes in properties such as volume or entropy. Examples include:

  • Melting: Change from a solid to a liquid. For example, ice melting into water.
  • Evaporation: Change from a liquid to a gas. For example, water boiling and turning into steam.
  • Sublimation: Change from the solid state to the gas state without going through the liquid state. Dry ice (solid CO2) sublimation is an example.

Second-order phase transitions

In these transitions, there is no latent heat associated and there is no discontinuity in the first derivative of the free energy. Properties such as specific heat, magnetic susceptibility, and compressibility change during the transition. An example is the transition to the critical point of a liquid, where the liquid becomes indistinguishable from the gas phase.

Graphical representation of phase transition

To visualize phase transitions, it is helpful to look at phase diagrams and other diagrams that show pressure, temperature, and volume. Below is a simple line diagram that shows how phase transitions occur when the pressure is held constant:

PressureTemperatureMeltingBoil

Importance of phase transitions

Phase transitions are not just academic exercises; they have practical implications in a variety of fields. Let us explore some important applications:

Meteorology

In meteorology, the phase changes of water play an integral role in weather phenomena. The formation of clouds, rain, snow, hail and other forms of precipitation all depend on the phase changes of water from vapor to liquid and solid forms. Understanding these changes helps meteorologists predict weather patterns, leading to more accurate forecasts.

Industrial applications

Industrially, many processes take advantage of phase transitions for production and manufacturing. For example:

  • Distillation: This process separates components based on differences in boiling points, effectively utilizing the vaporization phase transition.
  • Cryogenics: Low-temperature technologies take advantage of phase transitions for applications such as liquefied natural gas (LNG). Cooling the gas into a liquid form makes storage and transport easier.

Mathematical description of the phase transition

To describe phase transitions mathematically, we refer to various thermodynamic potentials. A common approach is to use the Gibbs free energy function G = H - TS, where H is the enthalpy, T is the temperature, and S is the entropy.

Critical phenomena and phase transitions

Phase transitions are often accompanied by critical phenomena, which are dramatic changes in the physical properties of a substance near a critical point. At this turning point, properties such as density and heat capacity exhibit significant fluctuations.

Conclusion

Phase transitions play a vital role in understanding the properties and behavior of various substances under different environmental conditions. From explaining natural phenomena to expanding industrial capabilities, these transitions and their underlying principles remain a cornerstone of chemical science.

As we continue to explore and expand our understanding of matter and its transformations, phase transitions will undoubtedly play an enduring role in advancing scientific knowledge and technological innovations.


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Phase transition

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