quantum theory

(noun)

A theory developed in early 20th century, according to which nuclear and radiation phenomena can be explained by assuming that energy only occurs in discrete amounts called quanta.

Related Terms

Examples of quantum theory in the following topics:

  • Planck's Quantum Theory

    • As a result of these observations, physicists articulated a set of theories now known as quantum mechanics.
    • In some ways, quantum mechanics completely changed the way physicists viewed the universe, and it also marked the end of the idea of a clockwork universe (the idea that universe was predictable).
    • Max Planck named this minimum amount the "quantum," plural "quanta," meaning "how much."
    • One photon of light carries exactly one quantum of energy.
    • Planck is considered the father of the Quantum Theory.

  • The Bohr Model

    • The quantum theory from the period between Planck's discovery of the quantum (1900) and the advent of a full-blown quantum mechanics (1925) is often referred to as the old quantum theory.
    • where n = 1, 2, 3, ... is called the principal quantum number and ħ = h/2π.
    • Like Einstein's theory of the photoelectric effect, Bohr's formula assumes that during a quantum jump, a discrete amount of energy is radiated.
    • This marks the birth of the correspondence principle, requiring quantum theory to agree with the classical theory only in the limit of large quantum numbers.
    • The Bohr-Kramers-Slater theory (BKS theory) is a failed attempt to extend the Bohr model, which violates the conservation of energy and momentum in quantum jumps, with the conservation laws only holding on average.

  • Photochemistry

    • This "photoequivalence law" was derived by Albert Einstein during his development of the quantum (photon) theory of light.
    • The efficiency with which a given photochemical process occurs is given by its Quantum Yield (Φ).
    • Since many photochemical reactions are complex, and may compete with unproductive energy loss, the quantum yield is usually specified for a particular event.
    • The quantum yield of these products is less than 0.2, indicating there are radiative (fluorescence & phosphorescence) and non-radiative return pathways (green arrow).
    • Several secondary radical reactions then follow (shown in the gray box), making it difficult to assign a quantum yield to the primary reaction.

  • Indeterminacy and Probability Distribution Maps

    • An adequate account of quantum indeterminacy requires a theory of measurement.
    • Many theories have been proposed since the beginning of quantum mechanics, and quantum measurement continues to be an active research area in both theoretical and experimental physics.
    • Possibly the first systematic attempt at a mathematical theory for quantum measurement was developed by John von Neumann.
    • In quantum mechanical formalism, it is impossible that, for a given quantum state, each one of these measurable properties (observables) has a determinate (sharp) value.
    • In the world of quantum phenomena, this is not the case.

  • Description of the Hydrogen Atom

    • The hydrogen atom (consisting of one proton and one electron, not the diatomic form H2) has special significance in quantum mechanics and quantum field theory as a simple two-body problem physical system that has yielded many simple analytical solutions in closed-form.
    • This leads to a third quantum number, the principal quantum number n = 1, 2, 3, ....
    • The principal quantum number in hydrogen is related to the atom's total energy.
    • Note the maximum value of the angular momentum quantum number is limited by the principal quantum number: it can run only up to n − 1, i.e. ℓ = 0, 1, ..., n − 1.
    • Empirically, it is useful to group the fundamental constants into Rydbergs, which gives the much simpler equation below that turns out to be identical to that predicted by Bohr theory:

  • The Pauli Exclusion Principle

    • The Pauli exclusion principle, formulated by Austrian physicist Wolfgang Pauli in 1925, states that no two fermions of the same kind may simultaneously occupy the same quantum state.
    • In the theory of quantum mechanics, fermions are described by antisymmetric states.
    • In contrast, particles with integer spin (bosons) have symmetric wave functions; unlike fermions, bosons may share the same quantum states.
    • Electrons, being fermions, cannot occupy the same quantum state, so electrons have to "stack" within an atom—they have different spins while at the same place.
    • As spin is part of the quantum state of the electron, the two electrons are in different quantum states and do not violate the Pauli exclusion principle.

  • Explanation of Valence Bond Theory

    • In chemistry, valence bond (VB) theory is one of two basic theories—along with molecular orbital (MO) theory—that use quantum mechanics to explain chemical bonding.
    • VB theory complements molecular orbital (MO) theory, which does not adhere to the VB concept that electron pairs are localized between two specific atoms in a molecule.
    • MO theory can predict magnetic and ionization properties in a straightforward manner.
    • VB theory produces similar results, but is more complicated.
    • This theory is used to explain the covalent bond formation in many molecules.

  • Linear Combination of Atomic Orbitals (LCAO)

    • An LCAO approximation is a quantum superposition of atomic orbitals, used to calculate molecular orbitals in quantum chemistry.
    • These models provide a simple model of molecule bonding, understood through molecular orbital theory.
    • A linear combination of atomic orbitals, or LCAO, is a quantum superposition of atomic orbitals and a technique for calculating molecular orbitals in quantum chemistry.
    • In quantum mechanics, electron configurations of atoms are described as wave functions.

  • Bonding in Coordination Compounds: Valence Bond Theory

    • Valence bond theory is used to explain covalent bond formation in many molecules.
    • Valence bond theory is a synthesis of early understandings of how chemical bonds form.
    • In 1927, physicist Walter Heitler and collaborator Fritz London developed the Heitler-London theory, which enabled the calculation of bonding properties of the covalently bonded diatomic hydrogen molecule (H2) based on quantum mechanical considerations.
    • Finally, Linus Pauling integrated Lewis' proposal and the Heitler-London theory to give rise to two additional key concepts in valence bond theory: resonance and orbital hybridization.
    • It is in this aspect of valence bond theory that we see the concept of resonance.

  • Quantum Numbers

    • Quantum numbers provide a numerical description of the orbitals in which electrons reside.
    • Formally, the dynamics of any quantum system are described by a quantum Hamiltonian (H) applied to the wave equation.
    • The most prominent system of nomenclature spawned from the molecular orbital theory of Friedrich Hund and Robert S.
    • The average distance increases with n, thus quantum states with different principal quantum numbers are said to belong to different shells.
    • The second quantum number, known as the angular or orbital quantum number, describes the subshell and gives the magnitude of the orbital angular momentum through the relation.