Electron Cloud Model
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1) Atomic orbitals are mathematical descriptions of where electrons are likely to be found in an atom, obtained by solving the Schrödinger equation. Orbitals have different shapes depending on the electron's energy and angular momentum. 2) Electrons occupy orbitals from lowest to highest energy, with each orbital holding up to two electrons of opposite spin. Orbitals are grouped into shells corresponding to different energy levels. 3) The spatial distribution and probability of finding electrons in an atom can be described as an "electron cloud". Electron configurations contribute greatly to chemical bonding and reactivity patterns.
Electron Cloud Model
- 1. ELECTRON CLOUD MODEL Electron Density and Orbital Shapes Atomic orbitals are mathematical descriptions of where the electrons in an atom (or molecule) are most likely to be found. These descriptions are obtained by solving an equation known as the Schrödinger equation, which expresses our knowledge of the atomic world. As the angular momentum and energy of an electron increases, it tends to reside in differently shaped orbitals. The orbitals corresponding to the three lowest energy states are s, p, and d, respectively. The illustration shows the spatial distribution of electrons within these orbitals. The fundamental nature of electrons prevents more than two from ever being in the same orbital. The overall distribution of electrons in an atom is the sum of many such pictures. This description has been confirmed by many experiments in chemistry and physics, including an actual picture of a p-orbital made by a Scanning Tunneling Microscope. Electron Cloud Most of the physical and chemical properties of atoms, and hence of all matter, are determined by the nature of the electron cloud enclosing the nucleus. The nucleus of an atom, with its positive electric charge, attracts negatively charged electrons. This attraction is largely responsible for holding the atom together. The revolution of electrons about a nucleus is determined by the force with which they are attracted to the nucleus. The electrons move very rapidly, and determination of exactly where any particular one is at a given time is theoretically impossible (see Uncertainty Principle). If the atom were visible, the electrons might appear as a cloud, or fog, that is dense in some spots, thin in others. The shape of this cloud and the probability of finding an electron at any point in the cloud can be calculated from the equations of wave mechanics (see Quantum Theory). The solutions of these equations are called orbitals. Each orbital is associated with a definite energy, and
- 2. each may be occupied by no more than two electrons. If an orbital contains two electrons, the electrons must have opposite spins, a property related to the angular momentum of the electrons. The electrons occupy the orbitals of lowest energy first, then the orbitals next in energy, and so on, building out until the atom is complete (see Atom). The orbitals tend to form groups known as shells (so-called because they are analogous to the layers, or shells, around an onion). Each shell is associated with a different level of energy. Starting from the nucleus and counting outward, the shells, or principal energy levels, are numbered 1, 2, 3, … , n. The outer shells have more space than the inner ones and can accommodate more orbitals and therefore more electrons. The nth shell consists of 2n-1 orbitals, and each orbital can hold a maximum of 2 electrons. For example, the third shell contains five orbitals and holds a maximum of 10 electrons; the fourth shell contains seven orbitals and holds a maximum of 14 electrons. Among the known elements, only the first seven shells of an atom contain electrons, and only the first four shells are ever filled. Each shell (designated as n) contains different types of orbitals, numbered from 0 to n-1. The first four types of orbitals are known by their letter designations as s, p, d, and f. There is one s-orbital in each shell, and this orbital contains the most firmly bound electrons of the shell. The s-orbital is followed by the p-orbitals (which always occur in groups of three), the d-orbitals (which always occur in groups of five), and finally the f-orbitals (which always occur in groups of seven). The s-orbitals are always spherically shaped around the nucleus; each p-orbital has two lobes resembling two balls touching; each d-orbital has four lobes; and each f-orbital has eight lobes. The p-, d-, and f-orbitals have a directional orientation in space, but the spherical s-orbitals do not. The three p-orbitals are oriented perpendicular to one another along the axis of an imaginary three-dimensional Cartesian (x, y, z) coordinate system. The three p-orbitals are designated px, py, and pz, respectively. The d- and f-orbitals are similarly arranged about the nucleus at fixed angles to one another. When elements are listed in order of increasing atomic number, an atom of one element contains one more electron than an atom of the preceding element (see Chemical Elements). The added electrons fill orbitals in order of the increasing energy of the orbitals. The first shell contains the 1s orbital; the second shell contains the 2s orbital and the 2p orbitals; the third shell contains the 3s orbital, the 3p orbitals, and the 3d orbitals; the fourth shell contains the 4s orbital, the 4p orbitals, the 4d orbitals, and the 4f orbitals. After the two innermost shells, certain orbitals of outer shells have lower energies than the last orbitals of preceding shells. For this reason, some orbitals of the outer shells fill before the previous shells are complete. For example, the s-orbital of the
- 3. fourth shell (4s) fills before the d-orbitals of the third shell (3d). Orbitals generally fill in this order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s. In a notation frequently used to describe the electron configuration of an element, a superscript after the orbital letter gives the number of electrons in that orbital. Thus, 1s22s22p5 means that the atom has two electrons in the 1s orbital, two electrons in the 2s orbital, and five electrons in the 2p orbitals. Neutral atoms with exactly eight electrons in the outer shell (meaning the s- and p- orbitals of the outer shell are filled) are exceptionally stable. These neutral atoms are atoms of the noble gases, which are so stable that getting them to chemically react with other elements is very difficult. The unusual stability of the noble-gas electron structures is of great importance in chemical bonding and reactivity. All other elements tend to combine with each other in such a way as to imitate this stable structure. The structure of helium is 1s2; neon adds another stable shell, 2s22p6, to this; argon adds the orbitals 3s23p6; krypton adds the orbitals 4s23d104p6; and xenon adds the orbitals 5s24d105p6 (the s-orbital fills before the d-orbital of the previous shell). A.Electron Orbitals Electron Density and Orbital Shapes Atomic orbitals are mathematical descriptions of where the electrons in an atom (or molecule) are most likely to be found. These descriptions are obtained by solving an equation known as the Schrödinger equation, which expresses our knowledge of the atomic world. As the angular momentum and energy of an electron increases, it tends to reside in differently shaped orbitals. The orbitals corresponding to the three lowest energy states are s, p, and d, respectively. The illustration shows the spatial distribution of electrons within these orbitals. The fundamental nature of electrons prevents more than two from ever being in the same orbital. The overall distribution of electrons in an atom is the sum of many such pictures. This description has been confirmed by many experiments in chemistry and physics, including an actual picture of a p-orbital made by a Scanning Tunneling Microscope. Atomic Orbital Shapes Atomic orbitals are mathematical descriptions of where the electrons in an atom (or molecule) are most likely to be found. These descriptions are obtained by solving an equation known as the Schrödinger equation, which expresses our
- 4. knowledge of the atomic world. As the angular momentum and energy of an electron increases, it tends to reside in differently shaped orbitals. This description has been confirmed by many experiments in chemistry and physics, including an actual picture of a p-orbital made by a scanning tunneling microscope. Quantum Description of Electrons Scientists describe the properties of an electron in an atom with a set of numbers called quantum numbers. Electrons are a type of particle known as a fermion, and according to a rule of physics, no two fermions can be exactly alike. Each electron in an atom therefore has different properties and a different set of quantum numbers. Electrons that share the same principal quantum number form a shell in an atom. This chart shows the first three shells. The two electrons that share the principal quantum number 1 form the first shell. One of these electrons has the quantum numbers 1, s, 0, 1/2, and the other electron has the quantum numbers 1, s, 0, -1/2. Scientists cannot simultaneously measure both the exact location of an electron and its precise speed and direction, so they cannot measure the path a specific electron takes as it orbits the nucleus. The law of physics governing this phenomenon is called the uncertainty principle. Scientists can, however, determine the area an electron will probably occupy, and the probability of finding the electron at some place inside this area. A map of this area and its probabilities forms a cloudlike pattern known as an orbital. Each orbital can contain two electrons, but these electrons can not have identical properties, so they must spin in opposite directions. Orbitals are grouped into shells, like the layers of an onion, around the nucleus. Each shell can contain a limited number of orbitals, which means that each shell can contain a limited number of electrons. Each shell corresponds to a certain level of energy, and all the electrons in the shell have this same level of energy. As the shells get farther from the nucleus, they can contain more electrons, and the electrons in the shells have higher energy. See also Chemistry: Electron Cloud. Light Absorption and Emission When a photon, or packet of light energy, is absorbed by an atom, the atom gains the energy of the photon, and one of the atom’s electrons may jump to a higher energy level. The atom is then said to be excited. When an electron of an excited atom falls to a lower energy level, the atom may emit the electron’s excess energy in the form of a photon. The energy levels, or orbitals, of the
- 5. atoms shown here have been greatly simplified to illustrate these absorption and emission processes. For a more accurate depiction of electron orbitals, When an atom’s energy is at its minimum, it is said to be in a ground state. In this ground state, the atom’s electrons occupy the innermost available shells, those closest to the nucleus. When atoms are excited by heat, by an electric current, or by light or some other form of radiation, the atoms’ electrons can acquire energy and jump from an inner to an outer shell, leaving a vacancy in the inner shell. The atom seeks to shed this surplus energy, leading the electron in the outer orbit to fall back down to an inner vacancy. As it falls, the electron releases energy in the form of a photon, a tiny flash of light. The color of the light depends on the amount of energy emitted. Spectral Lines of Atomic Hydrogen When an electron makes a transition from one energy level to another, the electron emits a photon with a particular energy. These photons are then observed as emission lines using a spectroscope. The Lyman series involves transitions to the lowest or ground state energy level. Transitions to the second energy level are called the Balmer series. These transitions involve frequencies in the visible part of the spectrum. In this frequency range each transition is characterized by a different color. When an electron moves to a different shell, it does not gradually go from one shell to another, but instead jumps directly to the other shell. These jumps are like steps on a staircase (and are different from a smooth incline, or hill). The electron also absorbs or emits the energy to make jumps in steps. It cannot gradually build up or lose energy, but must instantly absorb the exact amount of energy needed to make a certain jump, or instantly emit the exact amount needed to fall to a lower shell. Each element has a different pattern of allowed jumps within its electronic structure, so the element’s atoms can only absorb or emit a distinct set of energies, or spectrum of colors. In this way, a scientist can tell which elements are present in a sample by looking at the colors absorbed or emitted when the sample is excited by heat, electricity, or light. See also Spectroscopy.