Chemical bond: Difference between revisions
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* Intermolecular (also known as Van der Waals) bonds. These are bonds between stable molecules, see the articles [[intermolecular forces]] and [[Van der Waals forces]] for more details. For many years it was believed that [[hydrogen bond]]ing should be classified as a separate type of bond, but modern theoretical chemistry recognizes it as a special type of intermolecular bond. | * Intermolecular (also known as Van der Waals) bonds. These are bonds between stable molecules, see the articles [[intermolecular forces]] and [[Van der Waals forces]] for more details. For many years it was believed that [[hydrogen bond]]ing should be classified as a separate type of bond, but modern theoretical chemistry recognizes it as a special type of intermolecular bond. | ||
It took several centuries before chemical bonding was fully understood, but at present it is generally accepted that non-relativistic quantum mechanical explanations based on Coulomb's electrostatic law<ref>That is, explanations derived from quantum mechanical energy operators containing electron-electron, electron-nucleus, and nucleus-nucleus Coulomb interactions plus electronic kinetic energies.</ref> give satisfactory accounts of all kinds of bonds.<ref>As late as 1916 the famous American chemist [[Gilbert N. Lewis]] disagreed strongly with this statement. He saw electrons as stationary without kinetic energy and not exerting Coulomb forces. In a lecture given at the December meeting of the Sections of Physics and Chemistry of the [[American Association for the Advancement of Science]], the [[American Physical Society]], and the [[American Chemical Society]] (see Science Magazine pp. 297-302 (1917); [http://dx.doi.org/10.1126/science.46.1187.297 DOI]), he declared the following: "Therefore, unless we are willing, under the onslaught of quantum theories, to throw overboard all of the basic principles of physical science, we must conclude that the electron in the Bohr atom not only ceases to obey Coulomb's law, but exerts no influence whatsoever upon another charged particle at any distance | It took several centuries before chemical bonding was fully understood, but at present it is generally accepted that non-relativistic quantum mechanical explanations based on Coulomb's electrostatic law<ref>That is, explanations derived from quantum mechanical energy operators containing electron-electron, electron-nucleus, and nucleus-nucleus Coulomb interactions plus electronic kinetic energies.</ref> give satisfactory accounts of all kinds of bonds.<ref>As late as 1916 the famous American chemist [[Gilbert N. Lewis]] disagreed strongly with this statement. He saw electrons as stationary without kinetic energy and not exerting Coulomb forces. In a lecture given at the December meeting of the Sections of Physics and Chemistry of the [[American Association for the Advancement of Science]], the [[American Physical Society]], and the [[American Chemical Society]] (see Science Magazine pp. 297-302 (1917); [http://dx.doi.org/10.1126/science.46.1187.297 DOI]), he declared the following: "Therefore, unless we are willing, under the onslaught of quantum theories, to throw overboard all of the basic principles of physical science, we must conclude that the electron in the Bohr atom not only ceases to obey Coulomb's law, but exerts no influence whatsoever upon another charged particle at any distance." </ref> | ||
Gravitational forces, strong nuclear forces, or even magnetic forces, do not play any significant role in chemical bonding. | Gravitational forces, strong nuclear forces, or even magnetic forces, do not play any significant role in chemical bonding. | ||
==Octet rule== | ==Octet rule== | ||
Revision as of 03:50, 16 April 2010
In chemistry, a chemical bond is a force between two atoms that is strong enough to see the two atoms that are exerting the force on each other as an entity. It can happen that the two atoms form a stable entity (a diatomic molecule), an example being two nitrogen atoms chemically bound to the stable molecule N—N (written as N2). It can also happen that the two atoms are part of a larger aggregate. For instance, the chemical bond C—O between a carbon (C) atom and an oxygen (O) atom in the molecule methanol H3C—OH is a strong and easily recognizable bond. Two atoms may also be bound in a crystal, such as sodium (Na) and chlorine (Cl) that appear in a 1-1 ratio in crystalline rock (kitchen) salt (NaCl). A diamond crystal can be seen as one huge molecule consisting of bonded C—C (carbon-carbon) pairs.
Traditionally one distinguishes in chemistry the following types of bonds:
- Covalent bonds. A covalent bond is an electron pair shared by two bonded atoms. These are the bonds most commonly found in organic chemistry. They take place mainly between hydrogen, carbon, nitrogen, and oxygen and lead to stable, recognizable, molecules that remain intact in the solid, liquid, and gaseous aggregation states. The phenomenon of electron pairing (covalent bonding) requires quantum mechanics for an explanation and deeper understanding.
- Ionic bonds. Here atom A loses an electron to its bonding partner B, so that A becomes the cation A+ and B the anion B−. Consecutively, the ions bind strongly through the Coulomb interaction. Systems of which the atoms are bound by ionic interactions are usually crystals, the example of rock salt (Na+—Cl−) was already mentioned. It requires advanced laboratory techniques to separate ionically bound molecules from crystals, because the crystals are very stable.
- Metallic bonds. A number of metal atoms can crystallize to form a metal, which is a solid recognized by high electric and thermal conductivity. The bonding is caused by delocalized electrons forming electronic bands. The mechanism is akin to the formation of molecular orbitals in molecules. An explanation of metallic bonding is offered by quantum mechanics. Metal molecules (M2, M3, etc.) are not easily prepared experimentally because the solids are usually very stable.
- Intermolecular (also known as Van der Waals) bonds. These are bonds between stable molecules, see the articles intermolecular forces and Van der Waals forces for more details. For many years it was believed that hydrogen bonding should be classified as a separate type of bond, but modern theoretical chemistry recognizes it as a special type of intermolecular bond.
It took several centuries before chemical bonding was fully understood, but at present it is generally accepted that non-relativistic quantum mechanical explanations based on Coulomb's electrostatic law[1] give satisfactory accounts of all kinds of bonds.[2] Gravitational forces, strong nuclear forces, or even magnetic forces, do not play any significant role in chemical bonding.
Octet rule
The octet rule is a simple rule that describes the valency of light atoms (atomic number Z ≤ 18, first and second row of the periodic system). The valency of an atom is the number of covalent bonds that the atom can make with a partner. The octet rule starts with the assumption that valency involves only the electrons in the outer shells of the (lowest-energy-state) atoms participating in the bonding. The following lists the number of valence electrons of the first 18 chemical elements
| H(1) | He(2) | ||||||
| Li(1) | Be(2) | B(3) | C(4) | N(5) | O(6) | F(7) | Ne(8) |
| Na(1) | Mg(2) | Al(3) | Si(4) | P(5) | S(6) | Cl(7) | Ar(8) |
The assumption is that the electron configuration of a noble gas (He, Ne, Ar) is particularly stable (since these gases are chemically inert) and that a bonded atom strives to a noble gas configuration. Except for hydrogen, which strives for for a helium configuration of two electrons (a duplet), it means that an atom tries to surround himself with an octet of eight electrons. Of course electrons must be shared with bonding partners to achieve a noble gas configuration.
(To be continued)
Note
- ↑ That is, explanations derived from quantum mechanical energy operators containing electron-electron, electron-nucleus, and nucleus-nucleus Coulomb interactions plus electronic kinetic energies.
- ↑ As late as 1916 the famous American chemist Gilbert N. Lewis disagreed strongly with this statement. He saw electrons as stationary without kinetic energy and not exerting Coulomb forces. In a lecture given at the December meeting of the Sections of Physics and Chemistry of the American Association for the Advancement of Science, the American Physical Society, and the American Chemical Society (see Science Magazine pp. 297-302 (1917); DOI), he declared the following: "Therefore, unless we are willing, under the onslaught of quantum theories, to throw overboard all of the basic principles of physical science, we must conclude that the electron in the Bohr atom not only ceases to obey Coulomb's law, but exerts no influence whatsoever upon another charged particle at any distance."