Dalton's Atomic Theory

• Atomic theory has evolved over centuries, thanks to numerous scientists.
• The concept of atoms traces back to ancient Greek philosophers, including Democritus.

Early Ideas About Matter

• Democritus (~460-370 BCE) posited that matter consists of indivisible particles termed "atomos."
• "Atomos" meant "uncuttable" or "indivisible" in ancient Greek.
• Aristotle refuted Democritus's atomic theory, proposing that matter was continuous, formed of earth, air, fire, and water.
• The views of Aristotle were accepted for nearly 2000 years, impeding atomic theory's progress.

Dalton's Atomic Theory

• John Dalton revived atomic theory with a scientific foundation the early 1800s.
• Dalton's atomic theory comprised several postulates:
  • All matter comprises extremely small particles called atoms.
  • Atoms of a given element share identical size, mass, and properties; atoms of different elements vary in these aspects.
  • Atoms are indivisible, indestructible, and cannot be created.
  • Atoms of different elements merge in simple, whole-number ratios, forming chemical compounds.
  • Atoms are combined, separated, or rearranged during chemical reactions.
• Dalton's theory successfully elucidated the laws of definite and multiple proportions.
• Dalton's model portrayed the atom as a solid, indivisible sphere.

Discovery of the Electron

• J.J. Thomson's 1897 cathode ray experiments led to discovering the electron.
• Cathode rays are streams of particles observed in vacuum tubes.
• Thomson found these particles to be negatively charged and much smaller than atoms.
• Thomson suggested the "plum pudding" model of the atom.
• The plum pudding model described the atom as a sphere of positive charge with negatively charged electrons embedded throughout.

Rutherford's Nuclear Model

• Ernest Rutherford, with Hans Geiger and Ernest Marsden, performed the gold foil experiment in 1909.
• In this experiment, alpha particles targeted a thin gold foil.
• Most alpha particles passed through, while some deflected at large angles, a few even bounced back.
• Rutherford determined that an atom's positive charge concentrates in a small, dense core called the nucleus.
• Rutherford proposed the nuclear model of the atom.
• The nuclear model depicts the atom with a small, positively charged nucleus, encircled by orbiting electrons.
• Rutherford's model did not explain why electrons didn't spiral into the nucleus.

Bohr's Model

• Niels Bohr refined Rutherford's model in 1913, using quantum theory.
• Bohr proposed that electrons orbit the nucleus in specific energy levels or shells.
• Electrons can only exist in specific energy levels, transitioning between levels by absorbing/emitting photons.
• Bohr's model explained hydrogen's line spectra.
• Bohr's model had limitations, struggling to predict the spectra of more complex atoms.

Quantum Mechanical Model

• The quantum mechanical model came about in the 1920s through contributions from scientists including Louis de Broglie, Werner Heisenberg, and Erwin Schrödinger.
• Louis de Broglie suggested particles, including electrons, possess wave-like traits.
• Werner Heisenberg formulated the uncertainty principle, stating the impossibility of knowing an electron's precise position and momentum simultaneously.
• Erwin Schrödinger created the Schrödinger equation, describing electron behavior in atoms as wave functions.
• The quantum mechanical model describes electrons via probabilities of finding them in space regions called orbitals.
• Orbitals are 3D regions where electrons are most likely found, not fixed paths, unlike in Bohr's model.
• The quantum mechanical model provides an accurate, complete description of atomic structure/behavior relative to previous models.

Subatomic Particles

• Atoms consist of protons, neutrons, and electrons.
• Protons are positively charged particles in the nucleus, determining the atomic number.
• Neutrons are electrically neutral particles located in the nucleus.
• Protons and neutrons share similar mass.
• Electrons are negatively charged particles surrounding the nucleus, with a much smaller mass than protons/neutrons.

Atomic Number and Mass Number

• Atomic number (Z) is the number of protons in an atom's nucleus.
• The atomic number defines the element.
• Mass number (A) is the total count of protons and neutrons in an atom's nucleus.
• Isotopes are atoms of the same element but differing neutron counts and mass numbers.
• The atomic mass unit (amu) expresses atom and subatomic particle masses.

Electron Configuration

• Electron configuration describes the arrangement of electrons within the atom.
• Electrons fill energy levels/sublevels in a specific order, following the Aufbau principle.
• Hund's rule dictates that electrons individually occupy each orbital within a sublevel before pairing.
• The Pauli exclusion principle specifies that no two electrons in the same atom can share the same four quantum numbers.

Ions

• An ion is an atom/molecule with a net electric charge from gaining/losing electrons.
• A cation is a positively charged ion, formed when an atom loses electrons.
• An anion is a negatively charged ion, formed when an atom gains electrons.