History of Chemistry - Atomic Theory

History of Chemistry

• John Dalton proposed atomic theory, stating that elements are made up of tiny particles called atoms, and atoms of a given element are identical but different from those of other elements.
• He also proposed that chemical compounds form when atoms of different elements combine, with consistent ratios of atoms in a given compound.
• Joseph John Thomson discovered the electron by studying electrical discharges in partially evacuated tubes, and calculated the charge-to-mass ratio of the electron.
• Thomson's Plum Pudding model portrayed the atom as a sphere of positive charge with electrons embedded in it.
• Robert Millikan determined the mass of the electron using experiments with oil drops.
• Ernest Rutherford conducted the gold foil experiment to test Thomson's model, and found that most alpha particles passed through the foil without deflection, while some were deflected at large angles.
• Rutherford concluded that an atom has a small, dense, positively charged nucleus at its center, and most of the atom is empty space.
• He famously said, “It was almost as incredible as if your fired a 15-inch shell at a piece of tissue paper and it came back and hit you.”

Wave-Particle Duality

• The photoelectric effect, where light causes the emission of electrons from a metal surface, demonstrates the particle-like nature of light.
• Energy of a photon is quantized, E = hv, where h is Planck's constant and v is the frequency.
• The de Broglie hypothesis proposed that particles can exhibit wave-like behavior, with a wavelength inversely proportional to their momentum (λ = h/mv).

Line Spectra

• Line spectra are unique patterns of light emitted by atoms when excited.
• The line spectrum of hydrogen suggests that only specific energy levels are allowed for the electron in a hydrogen atom, leading to the concept of quantized energy levels.
• Energy emitted by the electron is quantized, and the energy change during an electronic transition between energy levels is: ΔE = Ef - Ei.

Heisenberg's Uncertainty Principle

• Heisenberg's Uncertainty Principle states that it is impossible to know both the momentum and position of a particle with certainty simultaneously.
• This principle relates to the wave nature of matter, where uncertainty in position is related to the number of waves combined to form a wave packet.

Schrödinger Equation

• Schrödinger formulated wave mechanics, which laid the foundation for modern quantum theory.
• The Schrödinger equation, ĤΨ = EΨ, describes the wave behavior of electrons in atoms and involves a wave function (Ψ), which is a mathematical function of the electron's position in three-dimensional space.
• The square of the wave function (Ψ²) represents the probability density of finding an electron in a specific region of space.

Quantum Numbers

• Quantum numbers are solutions to the Schrödinger equation that describe the properties of orbitals (wave functions).
• There are four main quantum numbers:
  • Principal quantum number, n: Determines the size and energy of an orbital, and the distance of the electron from the nucleus. Larger 'n' values indicate higher energy levels and larger orbital size.
  • Angular momentum quantum number, l: Determines the shape of an orbital and the number of subshells in a principal shell. It can take values from 0 to n-1, with 0 representing an s orbital (spherical), 1 representing a p orbital (dumbbell-shaped), 2 representing a d orbital (more complex shapes), and so on.
  • Magnetic quantum number, ml: Determines the orientation of an orbital in space relative to other orbitals. It can take integer values from -l to +l, including 0. For an s orbital, ml = 0, for a p orbital, ml = -1, 0, +1.
  • Spin quantum number, ms: Describes the intrinsic angular momentum of an electron, which is called spin. It can have two values: +1/2 or -1/2, corresponding to spin up or spin down.

Electronic Configuration

• Electronic configuration describes the distribution of electrons among different energy levels in an atom.
• Orbitals are filled in order of increasing energy, following the Aufbau principle.
• Hund's rule states that electrons fill degenerate orbitals singly before pairing up.
• The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers.
• For example, the electronic configuration of fluorine (Z=9) is 1s² 2s² 2p⁵, and neon (Z=10) is 1s² 2s² 2p⁶.
• Valence electrons are the electrons located in the outermost principal quantum level of an atom, and they are responsible for chemical bonding.

Magnetic Properties

• Paramagnetic substances have unpaired electrons and are attracted to a magnetic field.
• Diamagnetic substances have all paired electrons and are weakly repelled by a magnetic field.

Atomic Radius

• Atomic radius is defined as half the distance between the nuclei of two adjacent atoms in a molecule or solid.
• Atomic radius decreases across a period due to increased nuclear charge attracting electrons more strongly, resulting in a smaller atomic size.
• Atomic radius increases down a group due to the addition of electron shells, leading to larger atomic size.