Chemistry Revamped E01 | Quantum Mechanics: The Mechanics of Reality

Democritus, an ancient Greek philosopher, proposed an unpopular school of thought that challenged Aristotle’s four-element theory. His ideas remained dormant for centuries until modern scientists revived them.

John Dalton, an English chemist, revived and refined Democritus’s ancient atomism. He transformed a philosophical concept into a scientific theory supported by experimental observations in chemistry.

J.J. Thomson conducted experiments revealing atoms contain internal structure. Robert Millikan later confirmed the charge of these particles. Together they demonstrated atoms possess divisible components.

Ancient Greek philosopher Democritus proposed atomism, later refined by chemist John Dalton in the 19th century. Both challenged prevailing theories about matter’s composition.

Ernest Rutherford designed experiments to test the plum pudding model. He fired alpha particles (helium atoms without electrons) at gold foil to probe atomic structure.

Rutherford’s nuclear model, though revolutionary, contained a catastrophic flaw recognized by physicists applying classical electromagnetic theory to atomic structure.

Louis de Broglie proposed that particles can behave as waves and vice versa. This revolutionary idea broadened physicists’ perspective on subatomic particles, particularly electrons.

Physicists conducted double-slit experiments firing electrons through two closely spaced slits. This experimental setup provides concrete evidence for wave-particle duality.

Niels Bohr combined his ideas with Louis de Broglie’s wave theory to create a solid atomic model. Their collaboration integrated quantum concepts with atomic structure.

Electrons in atoms change orbits by absorbing or emitting electromagnetic energy. Scientists control these transitions by adjusting the light’s frequency.

High-frequency electromagnetic waves like UV, X-rays, and gamma rays interact dangerously with atoms. Lower-frequency radiation like infrared and radio waves pose minimal danger.

Werner Heisenberg, the German physicist after whom Walter White’s alias was named, proposed a fundamental limit to measurement precision in quantum mechanics.

Erwin Schrödinger, a cat-lover and physicist, took an alternate route from Bohr’s model. He developed a new mathematical framework using wave functions to describe quantum systems.

Louis de Broglie proposed that particles can behave as waves and vice versa, fundamentally challenging classical physics’ distinction between particles and waves. His hypothesis resolved contradictions in atomic stability and electron behavior.

Physicists use the Hamiltonian, a mathematical tool named after William Rowan Hamilton, to construct wave functions and solve quantum systems. Schrödinger incorporated it into his equation.

The Schrödinger equation produces atomic orbitals as solutions. These orbitals give the same energy levels as Bohr’s model but with additional structural details and accuracy.

Electrons possess an intrinsic property called spin. This quantum property determines how many electrons can occupy each orbital.

Atomic wave functions from neighboring atoms overlap and interfere, creating new molecular orbitals that extend across multiple nuclei. This process explains chemical bond formation.

Niels Bohr combined with Louis de Broglie to establish that electrons occupy only specific, quantized energy levels. Electrons never exist between these discrete states, inhabiting only allowed orbitals with whole-number quantum characteristics.

Atoms forming bonds create hybrid orbitals by mixing their atomic orbitals. This mixing occurs within individual atoms before bonding to neighboring atoms.

Computational chemists apply density functional theory (DFT) when systems contain too many electrons for exact Schrödinger equation solutions. DFT enables practical quantum chemistry calculations.

The 1s orbital represents the first and lowest energy orbital in hydrogen atoms and all other elements. S orbitals exist at every principal energy level, maintaining spherical symmetry regardless of quantum number.

P orbitals appear in the second energy level and beyond, providing three perpendicular orbitals per energy level. Each p orbital accommodates two electrons with opposite spins, contributing to atoms’ bonding capability and geometry.

Electrons filling atomic orbitals at the second energy level encounter mutual repulsion that creates sub-energy levels. Negatively charged electrons repel each other, forcing s and p orbitals to differ in energy despite sharing the same principal quantum number.

Atomic wave functions from adjacent atoms interfere constructively when orbitals overlap, combining crest with crest and trough with trough. This interference produces bonding molecular orbitals with enhanced electron density between nuclei.

Atomic wave functions combine destructively when positive and negative phases align oppositely, creating antibonding molecular orbitals. Electrons in these orbitals destabilize molecules rather than stabilizing them like bonding orbitals do.

Atomic wave functions create antibonding orbitals through destructive interference when atoms approach. The star symbol denotes these orbitals, as in sigma-star or two-pi-star. Electrons occupying antibonding orbitals destabilize molecules.

Helium atoms demonstrate noble gas behavior by refusing to bond with each other. Their electron configuration forces occupation of both bonding and antibonding orbitals, preventing stable molecule formation.

P orbitals from adjacent atoms form pi bonding orbitals when they overlap side-by-side. Electrons occupying these orbitals create pi bonds, distinct from sigma bonds formed by head-on overlap.

Carbon atoms form four equivalent sp3 hybrid orbitals by combining one 2s orbital with three 2p orbitals. This hybridization produces methane’s characteristic tetrahedral geometry with 109.5-degree bond angles.

Carbon atoms combine one 2s orbital with two 2p orbitals to create three sp2 hybrid orbitals, leaving one unhybridized p orbital. This configuration produces benzene’s flat hexagonal structure and enables pi bonding.

Carbon atoms mix one 2s orbital with one 2p orbital to form two sp hybrid orbitals, leaving two p orbitals unhybridized. This minimal hybridization creates linear molecules with triple bonds, as in acetylene.

Glowing organisms like certain mushrooms emit light through bioluminescence, where the enzyme luciferase catalyzes reactions that release energy as visible photons. Chemical reactions excite electrons, which then return to ground states by emitting light.

Plants exploit quantum mechanical electron transitions to capture light energy for photosynthesis. Chlorophyll molecules absorb photons, exciting electrons to higher energy states that drive chemical reactions producing glucose and oxygen.

Modern electronic devices depend fundamentally on quantum mechanical principles governing electron behavior. Engineers design transistors and LEDs by exploiting how electrons absorb and release energy through discrete orbital transitions.