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


In the fascinating field of chemistry, the concept of "states of matter" is fundamental. Understanding states of matter is important because it forms the basis of how substances interact, transform, and exist. In general terms, matter is everything that has mass and occupies space. All matter is made up of atoms, and depending on conditions, it can exist in different forms or "states."

Types of states

There are three primary states of matter: solid, liquid, and gas. Although scientists have also identified other states of matter such as plasma and Bose-Einstein condensates, for the sake of simplicity, we will focus primarily on the three traditional states in this exploration.

Solids

Solids are characterized by their definite shape and volume. The particles in a solid are adjacent to each other, often in a regular pattern. The particles vibrate but do not move from their position, which is responsible for the definite shape and volume of a solid. This can be represented by the structure of a simple cube:

For example, consider a cube. The atoms in a solid are arranged in a structured pattern that enables it to maintain its shape unless broken or deformed by force.

Examples of solids include ice, wood, iron, and plastic. Each of these substances will retain its shape and volume unless an external force is applied to it, such as cutting or melting.

Liquids

Liquids have a fixed volume but they take the shape of their container. The particles in a liquid are still close to each other but are not in a fixed position, so that they flow around each other. This means that liquids can change shape but not volume. Take an example of water in a glass:

In this example, water conforms to the shape of the glass, highlighting its property of adjusting to its shape and remaining constant in volume.

Common examples of liquids include water, oil, and alcohol. These can be poured from one container to another, and their surfaces adapt to the shape of the container holding them.

Gases

Gases have neither a definite shape nor a definite volume. Instead, they expand to fill their container. This is due to the large amount of space between their particles, which move faster and are farther apart than those of solids and liquids. This can be illustrated in the following example:

Here, the gas molecules are dispersed within the container, thus showing the typical behavior of gases expanding to fill the available space.

Common examples include air, helium, and carbon dioxide. They can be easily compressed because of the large space between their particles.

Plasma

Plasma is an ionized gas, a gas in which enough energy is provided to release electrons from atoms or molecules and allow both species, ions and electrons, to coexist. Considered the fourth state of matter, plasma is not as commonly encountered in everyday life as the other states, but it is the most abundant state of matter in the universe:

Examples include lightning, stars, and neon signs. In plasmas, the energy is enough to break the bonds between electrons and nuclei, resulting in a soup of charged particles.

Although we don't regularly see plasma here on Earth, it has many applications, including electric arc welding, fluorescent lamps, and plasma televisions.

State changes

Matter can change from one state to another. These changes are known as "phase transitions" and usually result from changes in temperature or pressure.

Melting and freezing

Melting occurs when a solid substance turns into a liquid when heated. In contrast, freezing occurs when a liquid turns into a solid. For example, when ice is heated, it turns into water:

H2O (solid → liquid)

When the opposite process occurs, water becomes ice due to the removal of heat: H2O (liquid → solid).

Evaporation and condensation

Evaporation is when a liquid changes into a gas. There are two types of this: evaporation, where the process happens slowly, and boiling, where it happens quickly. In contrast, condensation is when a gas turns into a liquid. An example of this is steam from boiling water condensing back into water droplets on a cold surface:

H2O (liquid → gas) - vaporization
H2O (gas → liquid) - condensation

Sublimation and deposition

Sublimation is the process in which a solid substance changes directly into a gas, bypassing the liquid state. Dry ice (solid carbon dioxide) sublimes at room temperature:

CO2 (solid → gas)

Deposition is the opposite, where a gas becomes a solid without becoming a liquid. Frost formation from water vapor is an example of deposition.

Factors affecting the state of matter

Two main factors affect the states of matter: temperature and pressure.

Temperature

Temperature affects the movement of particles in matter. Increasing the temperature generally increases the energy in the system, causing the particles to move faster, possibly causing a change in state. For example:

  • By increasing the temperature of ice, it melts and turns into water, which later evaporates and becomes steam.
  • Conversely, lowering the temperature can slow the movement of particles, causing a gas to condense into a liquid, and a liquid to freeze into a solid.

Pressure

Pressure also affects the states of matter by bringing particles closer together. High pressure can turn a gas into a liquid, as seen in the case of carbonated beverages:

When you open a bottle of soda, the pressure decreases, causing the dissolved carbon dioxide to escape as bubbles. Similarly, lowering the pressure can cause a liquid to turn into vapor.

Practical applications

Understanding the states of matter has many practical applications in a variety of fields:

Everyday applications

In everyday life, recognizing the states of matter is helpful in cooking, preservation, and various forms of entertainment. Refrigeration, which relies on the condensation and evaporation of a coolant, is important for preserving food.

Industrial applications

Industries rely heavily on state changes, such as the use of chemical vapor deposition to produce thin films used in electronics, or the use of liquid metal for casting. The manipulation of these state changes allows precise control in manufacturing processes.

Conclusion

In short, the study of states of matter and the transitions between them remains a cornerstone of chemistry. Understanding how temperature and pressure affect these states allows scientists and engineers to invent and improve technologies that shape our world. From ice in drinks to nitrogen in the air, these concepts affect our daily lives in countless invisible ways.


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