Orbital motion (quantum)
Orbital motion in quantum mechanics refers to the behavior of electrons around the nucleus of an atom, fundamentally different from classical physics predictions. Traditional models suggested that electrons orbit the nucleus like planets around the sun, but this posed a problem, as classical physics would imply that they should spiral into the nucleus. The introduction of quantum mechanics by scientists such as Niels Bohr resolved this issue by proposing that electrons occupy discrete energy levels, known as orbitals, which are defined by quantum numbers.
In this framework, electrons exhibit quantum motion, meaning they "jump" between these energy levels without occupying the space in between. This behavior can be observed in phenomena like the emission spectrum of hydrogen, where specific wavelengths of light are emitted as electrons transition between orbitals, showcasing quantized energy levels. Additionally, electrons can be described as both particles and waves, leading to probabilistic models for their locations.
The development of quantum numbers allows for the categorization of orbitals by size, shape, and orientation. Each orbital can accommodate a limited number of electrons, with specific rules governing their arrangement based on their spins. Overall, quantum orbital motion is crucial for understanding atomic structure and behavior, influencing fields such as chemistry and physics.
Orbital motion (quantum)
The atom is the basic building block of matter. As scientists began to study the structure of the atom, they realized that it consists of a dense nucleus surrounded by a cloud of one or more electrons. However, Newton’s laws of motion predict that orbiting electrons will soon fall into the nucleus. Albert Einstein, Niels Bohr, and other scientists of the early twentieth century devised quantum mechanics to explain their observations.
Historical Background
In 1897, J.J. Thomson discovered that atoms contain electrons. However, Thomson’s model described a uniform atom with electrons scattered throughout. This changed in 1909, when Ernest Rutherford experimented by shooting particle rays at an extremely thin sheet of gold. Rutherford’s results pointed to an atom with a dense center, or nucleus. He reasoned that electrons orbit the nucleus, much like planets orbit the sun.
However, Newton’s laws of motion predict that orbiting electrons will soon fall into the nucleus. This posed a problem for Rutherford’s model. In 1913, Rutherford’s student Niels Bohr published the idea that electron motion is quantum: That is, electrons do not move according to the laws of classical physics. The term quantum, which refers to a discrete amount of energy, originated with Max Planck in 1899. Bohr’s model of quantum orbitals was controversial when first published. In 1905, however, Albert Einstein used the idea of quantum energy in his theories.
Quantum Motion of Electrons
Imagine walking up or down a staircase. At any point in your journey, your foot may rest on a step, or it may be suspended in the air between steps. This is ordinary motion. Now imagine, instead, that as you moved along the staircase your feet were always firmly planted on one step or another. How would you move? You would need to disappear from one step while, at the same instant, materializing on a different step. This bizarre scenario describes how electrons move among different quantum orbitals.
According to the Bohr atomic model, electron orbitals are simply different energy levels. Usually, an electron is found in the lowest energy level, closest to the nucleus. When an electron gains energy (such as light or heat), it can briefly move away from the nucleus to a higher-energy orbital and then drop back down. Evidence tells us that electrons never occupy the spaces between orbitals. Instead, they seemingly jump from one orbital to another. This movement is known as quantum motion. (Note that orbitals are not physical objects, like staircase steps, but simply locations around the nucleus where electrons may be found.)
Evidence from the Emission Spectrum of Hydrogen
Sunlight passing through a prism forms a rainbow of colors from red to violet. This is called a continuous spectrum, because it includes light of every wavelength between red and violet. Each wavelength of light carries a particular amount of energy. To learn about elements, early scientists would heat gases until they glowed (produced light) and then pass this light through a prism. Instead of a continuous spectrum, this produced a few distinct lines of particular colors. These spectral lines are unique for each element. Hydrogen, for example, produces one red line, a blue line, and a violet-blue line, as well as other lines outside the visible spectrum.
Recall that both light and heat are forms of energy. When an electron absorbs energy, it moves (briefly) to a higher-energy orbital. When it drops back down to a lower energy level, it releases this energy as light. The appearance of spectral lines means that only specific amounts of energy (corresponding to the wavelength of each line) are released as electrons move from one orbital to another. This energy, and therefore the energy levels occupied by the electrons, is quantized, meaning that it occurs only in specific quantities. This is how the terms quantum physics and quantum leap came about.
Electrons Are Also Waves
During the time when Bohr developed his model, scientists were realizing that light can be described as both a wave and a particle. Similarly, scientists began to describe electrons and other subatomic particles as waves. Erwin Schrodinger further developed Bohr’s model by describing electrons, and their energy levels, as waves. Instead of indicating locations where electrons can be found, Schrodinger’s equations give probabilities of electrons being found in particular locations. An electron is more likely to be found in some parts of an orbital than in others. Fortunately, most physics and chemistry depends on electrons’ energy levels, rather than their specific locations.
Quantum Numbers and Orbitals
Bohr’s model was improved by the mathematical work of Schrodinger, which describes orbitals in terms of three quantum numbers: n, l, and m. The principle quantum number n describes the orbital’s size and energy. Larger orbitals have higher energy levels. The angular quantum number l describes the orbital’s shape. Orbitals can be shaped like spheres, barbells, cloverleafs, or other shapes. Some of these can be aligned in different directions. The magnetic quantum number m specifies how an orbital is oriented in space. Importantly, all of these quantum numbers are integers, or whole numbers. There can be no in-between energy levels, shapes, or orientations for the orbitals. Across the periodic table, orbitals repeat in specific patterns. Elements in a family or group (column) will have similar types of orbitals.
Each orbital can hold one or, at most, two electrons. Electrons have a property called spin (which is unrelated to their actual spinning). Spin can be either positive or negative. Because no two electrons in an atom can have the exact same combination of quantum numbers and spin, two electrons sharing an orbital must have opposite signs for spin.
Bibliography
Goldman, Martin V. “Quantum Atom.” Physics 2000. University of Colorado, Boulder. Web. 30 Dec. 2014. http://www.colorado.edu/physics/2000/index.pl
Nave, Carl R. “Quantum Physics.” Hyperphysics Concepts. Georgia State University. 2012. Web. 30 Dec. 2014. http://hyperphysics.phy-astr.gsu.edu/hbase/quacon.html#quacon
“Physics and Astronomy.” PBS, A Science Odyssey. WGBH. Web. 30 Dec. 2014. http://www.pbs.org/wgbh/aso/databank/physastro.html
“Rutherford-Bohr Model.” Understanding Radiation. Environmental Protection Agency. Web. 30 Dec. 2014. http://www.epa.gov/radiation/understand/rutherford.html