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 ↓


Matter and Intermolecular Forces


The study of fluids and intermolecular forces plays an important role in understanding the properties of matter in chemistry. Intermolecular forces are forces of attraction or repulsion between neighbouring particles (atoms, molecules or ions). They are responsible for many properties of substances, including their states of matter - solid, liquid or gas.

Nature of liquids

Unlike solids, where particles are arranged in a fixed, repeating pattern, and unlike gases, where particles move around freely and occupy the entire available volume, liquids have intermediate properties. In a liquid, particles are much closer to each other than in a gas and have less kinetic energy. As a result, liquids have a fixed volume but can change shape depending on the container they are in.

Properties of liquids

  • Volume: A liquid has a definite volume, which means it occupies a fixed space irrespective of the shape of the container.
  • Shape: Unlike solids, liquids do not have a definite shape and they take the shape of their container.
  • Viscosity: This refers to the resistance of a fluid to flow. Fluids like honey have a high viscosity, while water has a low viscosity.
  • Surface tension: This is the tendency of liquid surfaces to shrink to reduce their surface area. This is why small insects can walk on water.

Visual example: Simple fluid model

liquid molecules in container

Intermolecular forces

Intermolecular forces are much weaker than the chemical bonds that hold atoms together within a molecule (covalent bonds, ionic bonds, etc.). However, they are important in determining the physical properties of substances at the macroscopic level. There are several types of intermolecular forces:

1. Dispersion force (London dispersion force)

These are the weakest intermolecular forces and occur between all atoms and molecules. They are the result of temporary fluctuations in electron density in atoms and nonpolar molecules, which create temporary dipoles that attract each other.

// Example: Dispersion forces Nonpolar molecules like CH4 (methane) still exhibit dispersion forces.

2. Dipole-dipole interaction

These occur between polar molecules, which have permanent dipoles. The positive end of one molecule is attracted to the negative end of the other molecule. This type of interaction is stronger than dispersion forces.

// Example: Dipole-Dipole interactions Polar molecules such as HCl (hydrochloric acid) exhibit dipole-dipole forces.

Visual example: Polar and nonpolar molecules

δ- δ+ Dipole–dipole interaction

3. Hydrogen bonding

Hydrogen bonds are stronger than dipole-dipole interactions and occur when hydrogen is bound to highly electronegative atoms such as nitrogen, oxygen or fluorine. This interaction plays an important role in the properties of biological molecules such as water and DNA.

// Example: Hydrogen bonding Water (H2O) molecules form hydrogen bonds with each other.

4. Ion-dipole force

These occur between ions and polar molecules. They are particularly important in solutions of ionic compounds in polar solvents such as water.

// Example: Ion-Dipole forces Sodium ions (Na+) interacting with water molecules represent ion-dipole forces.

Intermolecular forces and physical properties

The strength and type of intermolecular forces in a substance affect its boiling and melting point, vapor pressure, solubility, and viscosity. Here's how:

Boiling and melting point

Stronger intermolecular forces cause higher boiling and melting points, as more energy is required to overcome these forces to change to a different state.

// Concept: Boiling and melting points Water (H2O), with strong hydrogen bonding, has a high boiling point compared to ammonia (NH3).

Vapor pressure

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid. Substances with weaker intermolecular forces have higher vapor pressures because molecules escape into the vapor phase more easily.

// Concept: Vapor pressure Diethyl ether has a higher vapor pressure compared to water due to weaker intermolecular forces.

Stickiness

Viscosity is affected by intermolecular forces; stronger forces result in greater viscosity. Changes in temperature can also affect viscosity, as higher temperatures generally reduce the effects of intermolecular forces.

// Example: Viscosity Glycerin has a higher viscosity than water due to stronger intermolecular forces.

Applications of intermolecular forces

Understanding intermolecular forces is important for many scientific and industrial processes. Here are some examples:

1. Biological systems

Hydrogen bonding plays an important role in the structure and function of biological molecules. For example, it maintains the double helix structure of DNA.

2. Industrial processes

In the chemical industry, understanding vapor pressure and boiling point helps design processes such as distillation to separate the components of a mixture.

3. Materials science

The properties of polymers and other materials often depend on intermolecular forces. Engineers can design materials with specific properties by manipulating these forces.

Visual example: Water molecule with hydrogen bonds

H H Hey water molecule with hydrogen bond

In conclusion, fluid and intermolecular forces are a fascinating aspect of chemistry that provide information about the behavior and properties of matter. From determining the physical state of a substance to playing a vital role in biological processes, the study of these forces provides valuable knowledge for both scientific and practical applications.


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Matter and Intermolecular Forces

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