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 → States of matter ↓


Phase diagram


Phase diagrams in chemistry are essential tools used to understand the conditions under which different phases of matter exist and coexist in equilibrium. Typically, these diagrams display temperature and pressure as axes and show where the different phases of a substance are stable. By studying phase diagrams, one can determine the state of a substance at a given temperature and pressure, and predict how this state will change under different conditions.

Importance of phase diagrams

Phase diagrams help represent the state of matter. They help chemists predict how a substance such as water will behave when temperature or pressure changes. This information is important for many applications, from industrial processes to understanding natural phenomena.

Basic components of a phase diagram

The phase diagram usually includes the following features:

  • Axis: The horizontal axis usually represents temperature, while the vertical axis represents pressure.
  • Lines or curves: These separate different phases (solid, liquid, gas). Each line represents conditions where two phases can coexist in equilibrium.
  • Triple point: A unique point on the phase diagram where all three phases exist in equilibrium.
  • Critical point: The point beyond which a supercritical fluid exists, a state where the liquid and gas phases are indistinguishable.

Visual representation of a basic phase diagram

Pressure Temperature Solid Liquid Gas Three Points Critical point

Understanding areas and lines

Solid area

In the solid region the substance remains in a rigid form with a definite shape and volume. This region lies below the melting line on the phase diagram.

Fluid area

The liquid state lies between the melting line and the evaporation line. In this state the substance has a definite volume but no definite shape, conforming to the shape of its container.

Gas field

Above the vaporization line, the substance exists as a gas. Here, it fills the entire volume of its container, having neither a defined volume nor a shape.

Lines of balance

  • Melting line: This line represents the conditions at which the solid and liquid states exist simultaneously. Along this line, a solid can change into a liquid and vice versa.
    • Solid ⇌ Liquid
  • Evaporation line: It represents the conditions where the liquid and gas phases are in equilibrium.
    • Liquid ⇌ Gas
  • Sublimation line: This line describes the direct transition between the solid and gas state without changing into the liquid state.
    • Solid ⇌ Gas

Explanation of triple point

The triple point is a singular point on the diagram where solid, liquid, and gas phases all coexist in perfect equilibrium. For example, the triple point of water is at exactly 0.01°C and 611.657 Pascals. Interestingly, at the triple point, all three phases are indistinguishable, creating a unique situation where the state of matter cannot be described simply as solid, liquid, or gas.

Critical point and supercritical fluid

The critical point marks the end of the liquid-gas boundary on the phase diagram. Beyond this point, the substance exists as a supercritical fluid, where distinct liquid and gas phases do not exist. Supercritical fluids have unique properties that combine aspects of both liquids and gases, making them highly useful in applications such as supercritical fluid extraction.

Water: An exceptional case

The phase diagram of water has unique characteristics due to hydrogen bonding. One area where water differs is the slope of the melting line, which is negative. This means that increasing pressure causes ice to melt, a situation that is the opposite in many other substances.

Pressure(P) Temperature(T) Snow Water Steam Three Points Critical point

Specific phase changes

Phase changes occur when a substance transitions between different states of matter. Here are some common phase changes:

  • Melting: The transition from a solid to a liquid. This occurs when a solid gains enough energy to overcome a stable lattice structure while maintaining its solid state.
  • Freezing: The process of a liquid freezing. The opposite of melting, this is the process when a liquid loses enough energy to become a solid.
  • Vaporization: The change from a liquid to a gas. This can occur through boiling (throughout the liquid) or evaporation (surface level).
  • Condensation: The conversion of a gas into a liquid. This occurs when a gas loses energy and turns back into a liquid.
  • Sublimation: Direct transition from solid to gas without going through the liquid state.
  • Deposition: The opposite of sublimation; a gas changes into a solid without first becoming a liquid.

Real-world applications of phase diagrams

Industrial processes

Many industries rely on phase diagrams to understand and control the conditions needed for different phases to occur in order to efficiently manufacture products. This includes metallurgy, where phase diagrams guide alloying processes and ensure structural integrity.

Food industry

In the food industry, phase diagrams help in processes such as freeze-drying, which combines sublimation and freezing techniques to preserve foods by removing moisture through a transition from a solid to a gaseous state.

Natural phenomena

Understanding phase diagrams explains natural phenomena such as ice floating on water, which affect ecosystems and the global climate. The unique properties of water are important in these phenomena.

Advanced concepts: Phase rule

Phase diagrams follow the Gibbs phase rule, which is a formula that describes the number of degrees of freedom in a closed system. The formula is:

F = C − P + 2

Where:

  • F: degrees of freedom
  • C: number of components in the system
  • P: number of steps present

This rule gives information about how many variables can be changed independently without changing the number of steps present.

Conclusion

Phase diagrams serve as invaluable tools in understanding the behavior of substances under different conditions. By depicting distinct phases and the transitions between them, they provide a graphical perspective of phase stability. This knowledge is important not only in academic and research contexts, but also has widespread industrial and real-world applications.


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Phase diagram

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