Atomic Models - An introduction

  • In the late 19th century, atomic models were proposed to understand the structure of atoms.
  • These models helped in explaining various phenomena observed in nature.
  • The development of atomic models was a significant milestone in the field of physics.

Dalton’s Atomic Theory

  • Proposed by John Dalton in 1808.
  • According to Dalton’s theory, atoms are indivisible and indestructible.
  • Elements are made up of tiny particles called atoms.
  • Atoms of the same element are identical in mass and properties.
  • Compounds are formed by the combination of atoms of different elements.

Thomson’s Plum Pudding Model

  • Proposed by J.J. Thomson in 1904.
  • According to Thomson’s model, atoms are like a positively charged “pudding” with negatively charged electrons embedded in it.
  • This model explained the presence of electrons in atoms.
  • Thomson’s model lacked the concept of a nucleus.

Rutherford’s Nuclear Model

  • Proposed by Ernest Rutherford in 1911.
  • According to Rutherford’s model, atoms have a small, dense, and positively charged nucleus at the center.
  • Electrons revolve around the nucleus in circular orbits.
  • Most of the atom’s mass is concentrated in the nucleus.
  • The nucleus is positively charged due to the presence of protons.

Bohr’s Model of the Atom

  • Proposed by Niels Bohr in 1913.
  • Bohr’s model introduced the concept of energy levels or shells.
  • Electrons revolve around the nucleus in specific energy levels.
  • Electrons can jump from one energy level to another by absorbing or emitting energy.
  • Bohr’s model successfully explained the emission and absorption spectra of atoms.

Quantum Mechanical Model

  • Proposed by Schrödinger, Heisenberg, and others in the 1920s.
  • Based on the principles of quantum mechanics.
  • Describes the behavior of electrons in terms of probability distributions.
  • Electrons are found in regions called orbitals.
  • Orbitals are represented using mathematical functions called wave functions.

Subatomic Particles

  • Atoms are composed of subatomic particles.
  • The three main subatomic particles are protons, neutrons, and electrons.
  • Protons have a positive charge and are found in the nucleus.
  • Neutrons have no charge and are also found in the nucleus.
  • Electrons have a negative charge and revolve around the nucleus.

Atomic Number and Mass Number

  • Atomic number (Z) represents the number of protons in an atom’s nucleus.
  • The atomic number determines the element’s identity.
  • Mass number (A) represents the total number of protons and neutrons in the nucleus.
  • Isotopes are atoms of the same element with different mass numbers.

Electron Configuration

  • Electron configuration is the arrangement of electrons in an atom.
  • Electrons occupy specific energy levels or shells.
  • Each energy level can hold a certain number of electrons.
  • Electrons fill the lower energy levels before filling the higher ones.
  • The electron configuration determines the chemical properties of an element.

Bohr’s Formula for the Energy of Electrons

  • According to Bohr’s formula, the energy of an electron in an energy level can be calculated using the formula:
    • E = -13.6 Z^2 / n^2 eV
    • E: Energy of the electron
    • Z: Atomic number (number of protons)
    • n: Principal quantum number (energy level)
  1. Quantum Numbers
  • Quantum numbers describe the properties and behavior of electrons in an atom.
  • Principal quantum number (n) determines the energy level of an electron.
  • It can have integer values starting from 1.
  • The maximum number of electrons that can be accommodated in a shell is given by 2n².
  1. Orbital Shapes
  • Orbitals are regions in an atom where electrons are likely to be found.
  • Different types of orbitals have different shapes and orientations.
  • s orbitals are spherical in shape.
  • p orbitals have a dumbbell shape and are oriented along three mutually perpendicular axes.
  • d orbitals have complex shapes and are oriented in different directions.
  1. Electron Spin
  • The spin quantum number (s) represents the spin of an electron.
  • It determines the orientation of the electron’s spin.
  • An electron can have two possible spin states: +1/2 (spin-up) or -1/2 (spin-down).
  • This property is fundamental in determining the magnetic properties of atoms.
  1. Pauli Exclusion Principle
  • The Pauli exclusion principle states that no two electrons in an atom can have the same set of quantum numbers.
  • This principle helps explain the arrangement of electrons in different orbitals.
  • It ensures that only a maximum of two electrons can occupy an orbital, and they must have opposite spins.
  1. Aufbau Principle
  • The Aufbau principle states that electrons fill the lowest energy levels or orbitals first before filling higher ones.
  • This principle determines the order in which electrons occupy different energy levels and sublevels.
  • By following this principle, we can determine the electron configuration of any given atom.
  1. Hund’s Rule
  • Hund’s rule states that electrons occupy orbitals of the same energy level one by one with parallel spins before pairing up.
  • This rule helps explain the distribution of electrons in orbitals in a way that minimizes electron-electron repulsion.
  • Electron configuration can be represented using orbital diagrams or electron configuration notation.
  1. Electron Configuration Notation
  • Electron configuration notation represents the distribution of electrons in energy levels and orbitals.
  • It uses the noble gas shorthand to simplify writing electron configurations.
  • For example, the electron configuration of oxygen (8 electrons) can be written as 1s² 2s² 2p⁴ or [He] 2s² 2p⁴.
  1. Valence Electrons
  • Valence electrons are the electrons in the outermost shell of an atom.
  • They determine the chemical properties of an element.
  • Elements in the same group or column of the periodic table have the same number of valence electrons.
  • Valence electrons are involved in chemical bonding and the formation of compounds.
  1. Lewis Dot Diagrams
  • Lewis dot diagrams (Lewis structures) represent the valence electrons of an atom.
  • The symbol of the element is surrounded by dots representing the valence electrons.
  • These diagrams are useful for understanding bonding and predicting chemical reactions.
  • For example, the Lewis dot diagram for oxygen (valence electron configuration 2s² 2p⁴) is shown as O:⋅.
  1. Summary
  • Atomic models have evolved over time, from Dalton’s atomic theory to the quantum mechanical model.
  • Each model contributed to our understanding of atoms and their properties.
  • Quantum numbers describe the properties of electrons, including their energy levels, orbital shapes, and spin.
  • Electron configurations determine the arrangement of electrons in an atom.
  • Valence electrons play a crucial role in chemical bonding and the behavior of elements.
  1. Quantum Mechanical Model
  • Proposed by Schrödinger, Heisenberg, and others in the 1920s.
  • Based on the principles of quantum mechanics.
  • Describes the behavior of electrons in terms of probability distributions.
  • Electrons are found in regions called orbitals.
  • Orbitals are represented using mathematical functions called wave functions.
  1. Subatomic Particles
  • Atoms are composed of subatomic particles.
  • The three main subatomic particles are protons, neutrons, and electrons.
  • Protons have a positive charge and are found in the nucleus.
  • Neutrons have no charge and are also found in the nucleus.
  • Electrons have a negative charge and revolve around the nucleus.
  1. Atomic Number and Mass Number
  • Atomic number (Z) represents the number of protons in an atom’s nucleus.
  • The atomic number determines the element’s identity.
  • Mass number (A) represents the total number of protons and neutrons in the nucleus.
  • Isotopes are atoms of the same element with different mass numbers.
  • Isotopes have the same atomic number but different mass numbers.
  1. Electron Configuration
  • Electron configuration is the arrangement of electrons in an atom.
  • Electrons occupy specific energy levels or shells.
  • Each energy level can hold a certain number of electrons.
  • Electrons fill the lower energy levels before filling the higher ones.
  • The electron configuration determines the chemical properties of an element.
  1. Bohr’s Formula for the Energy of Electrons
  • According to Bohr’s formula, the energy of an electron in an energy level can be calculated using the formula:
    • E = -13.6 Z^2 / n^2 eV
    • E: Energy of the electron
    • Z: Atomic number (number of protons)
    • n: Principal quantum number (energy level)
  1. Quantum Numbers
  • Quantum numbers describe the properties and behavior of electrons in an atom.
  • Principal quantum number (n) determines the energy level of an electron.
  • It can have integer values starting from 1.
  • The maximum number of electrons that can be accommodated in a shell is given by 2n².
  • Other quantum numbers include the azimuthal quantum number (l), magnetic quantum number (ml), and spin quantum number (ms).
  1. Orbital Shapes
  • Orbitals are regions in an atom where electrons are likely to be found.
  • Different types of orbitals have different shapes and orientations.
  • s orbitals are spherical in shape.
  • p orbitals have a dumbbell shape and are oriented along three mutually perpendicular axes (px, py, pz).
  • d orbitals have complex shapes and are oriented in different directions.
  1. Electron Spin
  • The spin quantum number (s) represents the spin of an electron.
  • It determines the orientation of the electron’s spin.
  • An electron can have two possible spin states: +1/2 (spin-up) or -1/2 (spin-down).
  • Spin is an intrinsic property of electrons and contributes to their magnetic behavior.
  1. Pauli Exclusion Principle
  • The Pauli exclusion principle states that no two electrons in an atom can have the same set of quantum numbers.
  • This principle helps explain the arrangement of electrons in different orbitals.
  • It ensures that only a maximum of two electrons can occupy an orbital, and they must have opposite spins.
  • The Pauli exclusion principle prevents electron-electron repulsion and provides stability to atomic structures.
  1. Aufbau Principle
  • The Aufbau principle states that electrons fill the lowest energy levels or orbitals first before filling higher ones.
  • This principle determines the order in which electrons occupy different energy levels and sublevels.
  • By following this principle, we can determine the electron configuration of any given atom.
  • The Aufbau principle helps in understanding the stability and reactivity of elements.
  • It provides the basis for the periodic table’s organization and the prediction of chemical behavior.