Quantum Mechanics

Quantum mechanics is a physics branch exploring matters and energies' behavior at the smallest scales. It is also crucial in chemistry to understand electrons' behavior with atoms and molecules. 

Here are the major areas where quantum mechanics is present:

(Quantum mechanics = QM)

• Wave-Particle Duality: According to QM, particles like electrons exhibit wave and particle-like properties.

• Atomic Structure: QM provides a model for atomic structures. Electrons in atoms don't move in well-defined orbits, as in the classical model, but rather exist in regions of probability known as orbitals. These orbitals are described by mathematical functions called wavefunctions.

• Well-defined orbits: Orbits with high chances of having electrons.

• Quantization of Energy: QM introduces the concept of quantized energy levels. Atoms' electrons can only occupy certain discrete energy levels, and transitions between these levels result in the absorption or emission of photons, giving rise to the characteristic line spectra of elements.

• Heisenberg Uncertainty Principle: This principle, a fundamental aspect of quantum mechanics, states that there are inherent limits to the precision with which certain pairs of properties (like position and momentum) can be simultaneously known. This has implications for our understanding of electron behavior.

• Schrödinger Equation: The Schrödinger equation is a central equation in quantum mechanics that describes how the quantum state of a physical system changes over time. It is used to derive the wavefunctions and energy levels of electrons in atoms and molecules.

• Chemical Bonding: Quantum mechanics is crucial for understanding chemical bonding. It explains how atoms combine to form molecules and why certain arrangements of atoms are more stable than others. Molecular orbitals, which describe the distribution of electrons in molecules, are derived from quantum mechanical principles.

• Spectroscopy: Quantum mechanics is fundamental to the interpretation of spectroscopic data. It explains how different molecular and atomic transitions lead to the absorption or emission of specific wavelengths of light.

• Heisenberg Theory: Named after German physicist Werner Karl Heisenberg, the idea states the zero chance of prediction for an electron's location. 

The idea claims that it’s impossible to both know an object’s location and momentum. Either its position or its path is known, never at the same time. 

Unlike other equations, the Schrödinger equation can predict where electrons may be found, but couldn't predict exactly the exact location. But we can express the chance of finding electrons with the wavefunction, a mathematical function. 

So, orbitals went from being 2D circular orbits to 3D shapes.