# Quantum Mechanics - Quantum Results, Theoretical Implications Of Quantum Mechanics

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Quantum mechanics is the theory used to provide an understanding of the behavior of microscopic particles such as electrons and **atoms**. More importantly, quantum mechanics describes the relationships between **energy** and **matter** on atomic and subatomic scale.

At the beginning of the twentieth century, German physicist Maxwell Planck (1858–1947) proposed that atoms absorb or emit electromagnetic **radiation** in bundles of energy termed quanta. This quantum concept seemed counter-intuitive to well-established Newtonian **physics**. Ultimately, advancements associated with quantum mechanics (e.g., the uncertainty principle) also had profound implications with regard to the philosophical scientific arguments regarding the limitations of human knowledge.

Planck proposed that atoms absorb or emit electro-magnetic radiation in defined and discrete units (quanta). Planck's quantum theory also asserted that the energy of **light** was directly proportional to its frequency, and this proved a powerful observation that accounted for a wide range of physical phenomena.

**Planck's constant** relates the energy of a **photon** with the frequency of light. Along with constant for the speed of light, Planck's constant (*h* = 6.626 10^{–34} Joule-second) is a fundamental constant of nature.

Prior to Planck's work, electromagnetic radiation (light) was thought to travel in waves with an infinite number of available frequencies and wavelengths. Planck's work focused on attempting to explain the limited **spectrum** of light emitted by hot objects and to explain the absence of what was termed the "violet catastrophe" predicted by 19th century theories developed by Prussian physicist Wilhelm Wien (1864–1928) and English physicist Baron (John William Strutt) Rayleigh (1842–1919).

Danish physicist Niels Bohr (1885–1962) studied Planck's quantum theory of radiation and worked in England with physicists J. J. Thomson (1856–1940), and Ernest Rutherford (1871–1937) improving their classical models of the atom by incorporating quantum theory. During this **time** Bohr developed his model of atomic structure. To account for the observed properties of **hydrogen**, Bohr proposed that electrons existed only in certain orbits and that, instead of traveling between orbits, electrons made instantaneous quantum leaps or jumps between allowed orbits. According to the **Bohr model**, when an **electron** is excited by energy it jumps from its ground state to an excited state (i.e., a higher energy orbital). The excited atom can then emit energy only in certain (quantized) amounts as its electrons jump back to lower energy orbits located closer to the nucleus. This excess energy is emitted in quanta of electromagnetic radiation (photons of light) that have exactly same energy as the difference in energy between the orbits jumped by the electron.

The electron quantum leaps between orbits proposed by the Bohr model accounted for Plank's observations that atoms emit or absorb electromagnetic radiation in quanta. Bohr's model also explained many important properties of the **photoelectric effect** described by Albert Einstein (1879–1955).

Using **probability theory**, and allowing for a wave-particle duality, quantum mechanics also replaced classical mechanics as the method by which to describe interactions between **subatomic particles**. Quantum mechanics replaced electron "orbitals" of classical **atomic models** with allowable values for angular **momentum** (angular **velocity** multiplied by **mass**) and depicted electrons position in terms of probability "clouds" and regions.

In the 1920s, the concept of quantization and its application to physical phenomena was further advanced by more mathematically complex models based on the work of the French physicist Louis Victor de Broglie (1892–1987) and Austrian physicist Erwin Schrödinger (1887–1961) that depicted the particle and wave nature of electrons. De Broglie showed that the electron was not merely a particle but a wave form. This proposal led Schrodinger to publish his wave equation in 1926. Schrödinger's work described electrons as "standing wave" surrounding the nucleus and his system of quantum mechanics is called wave mechanics. German physicist Max Born (1882–1970) and English physicist P.A.M Dirac (1902–1984) made further advances in defining the subatomic particles (principally the electron) as a wave rather than as a particle and in reconciling portions of quantum theory with relativity theory.

Working at about the same time, German physicist Werner Heisenberg (1901–1976) formulated the first complete and self-consistent theory of quantum mechanics. **Matrix mathematics** was well-established by the 1920s, and Heisenberg applied this powerful tool to quantum mechanics. In 1926, Heisenberg put forward his uncertainty principle that states that two complementary properties of a system, such as position and momentum, can never both be known exactly. This proposition helped cement the dual nature of particles (e.g., light can be described as having both wave and a particle characteristics). Electromagnetic radiation (one region of the spectrum of which comprises visible light) is now understood as having both particle and wave-like properties.

In 1925, Austrian-born physicist Wolfgang Pauli (1900–1958) published the **Pauli exclusion principle** that states that no two electrons in an atom can simultaneously occupy the same quantum state (i.e., energy state). Pauli's specification of spin (+1/2 or −1/2) on an electron gave the two electrons in any suborbital differing quantum numbers (a system used to describe the quantum state) and made completely understandable the structure of the **periodic table** in terms of electron configurations (i.e., the energy related arrangement of electrons in energy shells and suborbitals). In 1931, American chemist Linus Pauling published a paper that used quantum mechanics to explain how two electrons, from two different atoms, are shared to make a covalent bond between the two atoms. Pauling's work provided the connection needed in order to fully apply the new quantum theory to **chemical reactions**.

Quantum mechanics posed profound questions for scientists and philosophers. The concept that particles
such as electrons making quantum leaps from one
**orbit** to another, as opposed to simply moving between orbits, seems counter-intuitive, that is, outside the human experience with nature. Like much of quantum theory, the proofs of how nature works at the atomic level are mathematical. Bohr himself remarked, "Anyone who is not shocked by quantum theory has not understood it."

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over 1 year ago

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