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 → Atomic Structure ↓


Atomic Model


The concept of the atom is central to the study of chemistry. Over time, various atomic models have been proposed to explain the structure of atoms. The atomic model describes how subatomic particles (protons, neutrons, and electrons) are arranged and interact within an atom. In this article, we will explore the various atomic models developed over history and their significance in understanding the nature of matter.

From ancient philosophy to early science

Before advanced scientific methods were developed, the idea of the atom began as early philosophical speculation. The word "atom" comes from the Greek word for "indivisible." Ancient Greek philosophers such as Democritus and Leucippus speculated that all matter was composed of tiny, indivisible particles. However, their ideas were not based on experimental evidence and were purely philosophical.

Dalton's atomic theory

In the early 19th century, John Dalton, an English chemist, developed the first scientific atomic theory. According to Dalton's theory:

  1. All matter is composed of atoms, which are indivisible and indestructible particles.
  2. All atoms of a given element are similar in mass and properties.
  3. Compounds are formed by the combination of two or more different types of atoms.
  4. A chemical reaction is a rearrangement of atoms.

Thomson's plum pudding model

In 1897, J.J. Thomson discovered the electron, a negatively charged particle found in atoms. This discovery resulted in the proposal of the "plum pudding model." In this model, the atom is viewed as a positively charged sphere, with the negatively charged electrons contained in it like raisins in a pudding or plums in a pie.

Thomson's model was important because it was the first model to incorporate the existence of subatomic particles, but it did not accurately describe how these particles were arranged within the atom.

Rutherford's atomic model

Ernest Rutherford discovered in 1909 that the atom contained a small, compact, positively charged nucleus through his famous gold foil experiment. In this experiment, alpha particles were passed through a thin sheet of gold foil. Most of the particles passed through, but some scattered, showing that a small portion of the atom was positively charged and contained most of the atom's mass.

The atomic model suggested that:

  • Most of the atom is empty space.
  • The nucleus, which contains protons and neutrons, is at the center of the atom.
  • Electrons orbit the nucleus like planets around the sun.

Although this model explained the existence of the nucleus, it left unanswered questions about the behavior of electrons and their stability in their orbits.

Bohr model

Niels Bohr improved Rutherford's model by introducing the concept of quantized electron orbits. In 1913, Bohr proposed that electrons orbit around the nucleus at fixed distances, called energy levels or shells, and that each electron in an atom is at a specific energy level.

Key principles of the Bohr model include:

  • Electrons revolve around the nucleus at a fixed distance without emitting energy.
  • Electrons can jump between orbitals by absorbing or emitting energy in the form of light.
  • Energy levels are quantized, that is, only specific energy levels are allowed.

Although the Bohr model was successful in explaining the spectrum of the hydrogen atom, it could not accurately predict the spectra of more complex atoms.

Quantum mechanical model

As researchers delved further into the mysteries of atomic structure, they developed a more complex and accurate model. The quantum mechanical model, often associated with Erwin Schrödinger and Werner Heisenberg, provides a more sophisticated understanding of electron behavior.

The quantum mechanical model differs from the Bohr model in that it does not define exact paths for the electrons. Instead, it uses probability distributions to describe where electrons are likely to be found, known as orbitals. This model is based on powerful mathematical descriptions of atomic behavior.

Features of the quantum mechanical model include:

  • Electrons lie in orbitals, regions of probability, rather than on fixed paths.
  • These orbitals are defined by quantum numbers, which describe the size, shape, and orientation of the regions where electrons are likely to be found.
  • Electrons have wave-particle duality, that is, they exhibit both particle-like and wave-like properties.

The quantum mechanical model is supported by Schrödinger's wave equation:

HΨ = EΨ

where H is the Hamiltonian operator, Ψ is the wave function, and E is the energy of the system.

Current understanding and applications

Today, the quantum mechanical model is the main framework used by scientists to understand atomic and molecular chemistry. It provides the basis for quantum chemistry and helps explain the behavior of atoms in chemical reactions, chemical bonding, and the properties of substances.

This model goes far beyond just explaining atoms; it predicts the structure and properties of molecules. An understanding of molecular orbital theory, derived from quantum mechanics, allows chemists to predict the bonding patterns and reactivity of molecules in various chemical reactions.

Conclusion

The evolution of atomic models throughout history highlights the development of scientific thought and technology. From the ancient philosophy of indivisible particles to the modern quantum mechanical model, each stage has contributed to our comprehensive understanding of atomic structure and properties. Atomic models continue to be a vital tool in advancing chemical research, which in turn drives further discoveries in the physical and biological sciences.


Undergraduate → 1.2.2


U
username
0%
completed in Undergraduate

← Prev (1.2.1)
Subatomic particles
Next (1.2.3) →
Quantum numbers

Comments 



Atomic Model

https://www.chemistryelite.com/ug/1.2.2?ceal=en