molecule

Dictionary: mol·e·cule   (mŏl'ĭ-kyūl')

1. The smallest particle of a substance that retains the chemical and physical properties of the substance and is composed of two or more atoms; a group of like or different atoms held together by chemical forces.
2. A small particle; a tiny bit.

[French molécule, from New Latin mōlēcula, diminutive of Latin mōlēs, mass.]

Concept

Prior to the nineteenth century, chemists pursued science simply by taking measurements, before and after a chemical reaction, of the substances involved. This was an external approach, rather like a person reaching into a box and feeling of the contents without actually being able to see them. With the evolution of atomic theory, chemistry took on much greater definition: for the first time, chemists understood that the materials with which they worked were interacting on a level much too small to see. The effects, of course, could be witnessed, but the activities themselves involved the interactions of atoms in molecules. Just as an atom is the most basic particle of an element, a molecule is the basic particle of a compound. Whereas there are only about 90 elements that occur in nature, many millions of compounds are formed naturally or artificially. Hence the study of the molecule is at least as important to the pursuit of modern chemistry as the study of the atom. Among the most important subjects in chemistry are the ways in which atoms join to form molecules—not just the numbers and types of atoms involved, but the shape that they form together in the molecular structure.

How It Works

Introduction to the Molecule

Sucrose or common table sugar, of course, is grainy and sweet, yet it is made of three elements that share none of those characteristics. The formula for sugar is C₁₂H₂₂O₁₁, meaning that each molecule is formed by the joining of 12 carbon atoms, 22 hydrogens, and 11 atoms of oxygen. Coal is nothing like sugar—for one thing, it is as black as sugar is white, yet it is almost pure carbon. Carbon, at least, is a solid at room temperature, like sugar. The other two components of sugar, on the other hand, are gases, and highly flammable ones at that.

The question of how elements react to one another, producing compounds that are altogether unlike the constituent parts, is one of the most fascinating aspects of chemistry and, indeed, of science in general. Combined in other ways and in other proportions, the elements in sugar could become water (H₂O), carbon dioxide (CO₂), or even petroleum, which is formed by the joining of carbon and hydrogen.

Two different compounds of hydrogen and oxygen serve to further illustrate the curiosities involved in the study of molecules. As noted, hydrogen and oxygen are both flammable, yet when they form a molecule of water, they can be used to extinguish most fires. On the other hand, when two hydrogens join with two oxygens to form a molecule of hydrogen peroxide (H₂O₂), the resulting compound is quite different from water. In relatively high concentrations, hydrogen peroxide can burn the skin, and in still higher concentrations, it is used as rocket fuel. And whereas water is essential to life, pure hydrogen peroxide is highly toxic.

The Question of Molecular Structure

It is not enough, however, to know that a certain combination of atoms forms a certain molecule, because molecules may have identical formulas and yet be quite different substances. In English, for instance, there is the word "rose." Simply seeing the word, however, does not tell us whether it is a noun, referring to a flower, or a verb, as in "she rose through the ranks." Similarly, the formula of a compound does not necessarily tell what it is, and this can be crucial.

For instance, the formula C₂H₆O identifies two very different substances. One of these is ethyl alcohol, the type of alcohol found in beer and wine. Note that the elements involved are the same as those in sugar, though the proportions are different: in fact, some aspects of the body's reaction to ethyl alcohol are not so different from its response to sugar, since both lead to unhealthy weight gain. In reasonable small quantities, of course, ethyl alcohol is not toxic, or at least only mildly so; yet methyl ether—which has an identical formula—is a toxin.

But the distinction is not simply an external one, as simple as the difference between beer and a substance such as methyl ether, sometimes used as a refrigerant. To put it another way, the external difference reflects an internal disparity: though the formulas for ethyl alcohol and methyl ether are the same, the arrangements of the atoms within the molecules of each are not. The substances are therefore said to be isomers.

In fact C₂H₆O is just one of three types of formula for a compound: an empirical formula, or one that shows the smallest possible whole-number ratio of the atoms involved. By contrast, a molecular formula—a formula that indicates the types and numbers of atoms involved—shows the actual proportions of atoms. If the formula for glucose, a type of sugar (C₆H₁₂O₆), were rendered in empirical form, it would be CH₂O, which would reveal less about its actual structure. Most revealing of all, however, is a structural formula—a diagram that shows how the atoms are bonded together, complete with lines representing covalent bonds. (Structural formulas such as those that apply the Couper or Lewis systems are discussed in the Chemical Bonding essay, which also examines the subject of covalent bonds.)

Chemists involved in the area of stereo-chemistry, discussed below, attempt to develop three-dimensional models to show how atoms are arranged in a molecule. Such models for ethyl alcohol and methyl ether, for instance, would reveal that they are quite different, much as the two definitions of rose mentioned above illustrate the two distinctly different meanings. Because stereochemistry is a highly involved and complex subject, it can only be touched upon very briefly in this essay; nonetheless, an understanding of a molecule's actual shape is critical to the work of a professional chemist.

Molecules and Compounds

A molecule can be most properly defined as a group of atoms joined in a specific structure. A compound, on the other hand, is a substance made up of more than one type of atom—in other words, more than one type of element. Not all compounds are composed of discrete molecules, however. For instance, table salt (NaCl) is an ionic compound formed by endlessly repeating clusters of sodium and chlorine that are not, in the strictest sense of the word, molecules.

Salt is an example of a crystalline solid, or a solid in which the constituent parts are arranged in a simple, definite geometric pattern repeated in all directions. There are three kinds of crystalline solids, only one of which has a truly molecular structure. In an ionic solid such as table salt, ions (atoms, or groups of atoms, with an electric charge) bond a metal to a nonmetal—in this case, the metal sodium and the nonmetal chlorine. Another type of crystalline solid, an atomic solid, is formed by atoms of one element bonding to one another. A diamond, made of pure carbon, is an example. Only the third type of crystalline solid is truly molecular in structure: a molecular solid—sugar, for example—is one in which the molecules have a neutral electric charge.

Not all solids are crystalline; nor, of course, are all compounds solids: water, obviously, is a liquid at room temperature, while carbon dioxide is a gas. Nor is every molecule composed of more than one element. Oxygen, for instance, is ordinarily diatomic, meaning that even in its elemental form, it is composed of two atoms that join in an O₂ molecule. It is obvious, then, that the defining of molecules is more complex than it seems. One can safely say, however, that the vast majority of compounds are made up of molecules in which atoms are arranged in a definite structure.

In the essay that follows, we will discuss the ways atoms join to form molecules, a subject explored in more depth within the Chemical Bonding essay. (In addition, compounds themselves are examined in somewhat more detail within the Compounds essay.) We will also briefly examine how molecules bond to other molecules in the formation of solids and liquids. First, however, a little history is in order: as noted in the introduction to this essay, chemists did not always possess a clear understanding of the nature of a molecule.

A Brief History of the Molecule

In ancient and medieval times, early chemists—some of whom subscribed to an unscientific system known as alchemy—believed that one element could be transformed into another. Thus many an alchemist devoted an entire career to the vain pursuit of turning lead into gold. The alchemists were at least partially right, however: though one element cannot be transformed into another (except by nuclear fusion), it is possible to change the nature of a compound by altering the relations of the elements within it.

Modern understanding of the elements began to emerge in the seventeenth century, but the true turning point came late in the eighteenth century. It was then that French chemist Antoine Lavoisier (1743-1794) defined an element as a simple substance that could not be separated into simpler substances by chemical means. Around the same time, another French chemist, Joseph-Louis Proust (1754-1826) stated that a given compound always contained the same proportions of mass between elements. The ideas of Lavoisier and Proust were revolutionary at the time, and these concepts pointed to a substructure, invisible to the naked eye, underlying all matter.

In 1803, English chemist John Dalton (1766-1844) defined that substructure by introducing the idea that the material world is composed of tiny particles called atoms. Despite the enormous leap forward that his work afforded to chemists, Dalton failed to recognize that matter is not made simply of atoms. Water, for instance, is not just a collection of "water atoms": clearly, there is some sort of intermediary structure in which atoms are combined. This is the molecule, a concept introduced by Italian physicist Amedeo Avogadro (1776-1856).

Avogadro and the Idea of the Molecule

French chemist and physicist Joseph Gay-Lussac (1778-1850) had announced in 1809 that gases combine to form compounds in simple proportions by volume. As Gay-Lussac explained, the ratio, by weight, between hydrogen and oxygen in water is eight to one. The fact that this ratio was so "clean," involving whole numbers rather than decimals, intrigued Avogadro, who in 1811 proposed that equal volumes of gases have the same number of particles if measured at the same temperature and pressure. This, in turn, led him to the hypothesis that water is not composed simply of atoms, but of molecules in which hydrogen and oxygen combine.

For several decades, however, chemists largely ignored Avogadro's idea of the molecule. Only in 1860, four years after his death, was the concept resurrected by Italian chemist Stanislao Cannizzaro (1826-1910). Of course, the understanding of the molecule has progressed enormously in the years since then, and much of this progress is an outcome of advances in the study of subatomic structure. Only in the early twentieth century did physicists finally identify the electron, the negatively charged subatomic particle critical to the bonding of atoms.

Real-Life Applications

Molecular Mass

Just as the atoms of elements have a definite mass, so do molecules—a mass equal to that of the combined atoms in the molecule. The figures for the atomic mass of all elements are established, and can be found on the periodic table; therefore, when one knows the mass of a hydrogen atom and an oxygen atom, as well as the fact that there are two hydrogens and one oxygen in a molecule of water, it is easy to calculate the mass of a water molecule.

Individual molecules cannot easily be studied; therefore, the mass of molecules is compared by use of a unit known as the mole. The mole contains 6.022137 × 10²³ molecules, a figure known as Avogadro's number, in honor of the man who introduced the concept of the molecule. When necessary, it is possible today to study individual molecules, or even atoms and subatomic particles, using techniques such as mass spectrometry.

Bonding Within Molecules

Note that the mass of an atom in a molecule does not change; nor, indeed, do the identities of the individual atoms. An oxygen atom in water is the same oxygen atom in sugar, or in any number of other compounds. With regard to compounds, it should be noted that these are not the same thing as a mixture, or a solution. Sugar or salt can be dissolved in water at the appropriate temperatures, but the resulting solution is not a compound; the substances are joined physically, but they are not chemically bonded.

Chemical bonding is the joining, through electromagnetic force, of atoms representing different elements. Each atom possesses a certain valency, which determines its ability to bond with atoms of other elements. Valency, in turn, is governed by the configuration of valence electrons at the highest energy level (the shell) of the atom.

While studying noble gases, noted for their tendency not to bond, German chemist Richard Abegg (1869-1910) discovered that these gases always have eight valence electrons. This led to the formation of the octet rule: most elements (with the exception of hydrogen and a few others) are inclined to bond in such a way that they end up with eight valence electrons.

When a metal bonds to a nonmetal, this is known as ionic bonding, which results from attractions between ions with opposite electric charges. In ionic bonding, two ions start out with different charges and form a bond in which both have eight valence electrons. Nonmetals, however, tend to form covalent bonds. In a covalent bond, two atoms start out as most atoms do, with a net charge of zero. Each ends up possessing eight valence electrons, but neither atom "owns" them; rather, they share electrons.

Electronegativity

Not all elements bond covalently in the same way. Each has a certain value of electronegativity—the relative ability of an atom to attract valence electrons. Elements capable of bonding are assigned an electronegativity value ranging from a minimum of 0.7 for cesium to a maximum of 4.0 for fluorine. The greater the electronegativity value, the greater the tendency of an element to attract valence electrons.

When substances of differing electronegativity values form a covalent bond, this is described as polar covalent bonding. Water is an example of a molecule with a polar covalent bond. Because oxygen has a much higher electronegativity (3.5) than hydrogen (2.1), the electrons tend to gravitate toward the oxygen atom. By contrast, molecules of petroleum, a combination of carbon and hydrogen, tend to be nonpolar, because carbon (2.5) and hydrogen have very similar electronegativity values.

A knowledge of electronegativity values can be used to make predictions concerning bond polarities. Bonds that involve atoms whose electronegativities differ by more than 2 units are substantially ionic, whereas bonds between atoms whose electronegativities differ by less than 2 units are polar covalent. If the atoms have the same or similar electronegativity values, the bond is covalent.

Attractions Between Molecules

The energy required to pull apart a molecule is known as bond energy. Covalent bonds that involve hydrogen are among the weakest bonds between atoms, and hence it is relatively easy to separate water into its constituent parts, hydrogen and oxygen. (This is sometimes done by electrolysis, which involves the use of an electric current to disperse atoms.) Double and triple covalent bonds are stronger, but strongest of all is an ionic bond. The strength of the bond energy in salt, for instance, is reflected by its melting point of 1,472°F (800°C), much higher than that of water, at 32°F (0°C).

Bond energy relates to the attraction between atoms in a molecule, but in considering various substances, it is also important to recognize the varieties of bonds between molecules—that is, intermolecular bonding. For example, the polar quality of a water molecule gives it a great attraction for ions, and thus ionic substances such as salt and any number of minerals dissolve easily in water. On the other hand, we have seen that petroleum is essentially nonpolar, and therefore, an oil molecule offers no electric charge to bond it with a water molecule. For this reason, oil and water do not mix.

The bonding between water molecules is known as a dipole-dipole attraction. This type of intermolecular bond can be fairly strong in the liquid or solid state, though it is only about 1% as strong as a covalent bond within a molecule. When a substance containing molecules joined by dipole-dipole attraction is heated to become a gas, the molecules spread far apart, and these bonds become very weak. On the other hand, when hydrogen bonds to an atom with a high value of electronegativity (fluorine, for example), the dipole-dipole attraction between these molecules is particularly strong. This is known as hydrogen bonding.

Even a nonpolar molecule, however, must have some attraction to other nonpolar molecules. The same is true of helium and the other noble gases, which are highly nonattractive but can be turned into liquids or even solids at extremely low temperatures. The type of intermolecular attraction that exists in such a situation is described by the term London dispersion forces. The name has nothing to do with the capital of England: it is a reference to German-American physicist Fritz Wolfgang London (1900-1954), who in the 1920s studied the molecule from the standpoint of quantum mechanics.

Because electrons are not uniformly distributed around the nucleus of an atom at every possible moment, instantaneous dipoles are formed when most of the electrons happen to be on one side of an atom. Of course, this only happens for an infinitesimal fraction of time, but it serves to create a weak attraction. Only at very low temperatures do London dispersion forces become strong enough to result in the formation of a solid. (Thus, for instance, oil and rubbing alcohol freeze only at low temperatures.)

Molecular Structure

The Couper system and Lewis structures, discussed in the Chemical Bonding essay, provide a means of representing the atoms that make up a molecule. Though Lewis structures show the distribution of valence electrons, they do not represent the three-dimensional structure of the molecule. As noted earlier in this essay, the structure is highly important, because two compounds may be isomers, meaning that they have the same proportions of the same elements, yet are different substances.

Stereochemistry is the realm of chemistry devoted to the three-dimensional arrangement of atoms in a molecule. One of the most important methods used is known as the VSEPR model (valence shell electron pair repulsion). In bonding, elements always share at least one pair of electrons, and the VSEPR model begins with the assumption that the electron pairs must be as far apart as possible to minimize their repulsion, since like charges repel.

VSEPR structures can be very complex, and the rules governing them will not be discussed here, but a few examples can be given. If there are just two electron pairs in a bond between three atoms, the structure of a VSEPR model is like that of a stick speared through a ball, with two other balls attached at each end. The "ball" is an atom, and the "stick" represents the electron pairs. In water, there are four electron pairs, but still only three atoms and two bonds. In order to keep the electron pairs as far apart as possible, the angle between the two hydrogen atoms attached to the oxygen is 109.5°.

Where to Learn More

Basmajian, Ronald; Thomas Rodella; and Allen E. Breed. Through the Molecular Maze: A Helpful Guide to the Elements of Chemistry for Beginning Life Science Students. Merced, CA: Bioventure Associates, 1990.

Burnie, David. Microlife. New York: DK Publishing, 1997.

"Common Molecules" (Web site). <http://www.recipnet.indiana.edu/common/common.html> (June 2, 2001).

Cooper, Christopher. Matter. New York: DK Publishing, 2000.

Mebane, Robert C. and Thomas R. Rybolt. Adventures with Atoms and Molecules, Book V: Chemistry Experiments for Young People. Springfield, NJ: Enslow Publishers, 1995.

"The Molecules of Life" (Web site). <http://biop.ox.ac.uk/www/mol_of_life/Molecules_of_Life.html> (June 2, 2001).

"Molecules of the Month." University of Oxford (Web site). <http://www.chem.ox.ac.uk/mom/> (June 2, 2001).

"Molecules with Silly or Unusual Names" (Web site). <http://www.bris.ac.uk/Depts/Chemistry/MOTM/silly/sillymols.htm> (June 2, 2001).

"Theory of Atoms in Molecules" (Web site). <http://www.chemistry.mcmaster.ca/faculty/bader/aim/> (June 2, 2001).

Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.

A molecule may be thought of either as a structure built of atoms bound together by chemical forces or as a structure in which two or more nuclei are maintained in some definite geometrical configuration by attractive forces from a surrounding swarm of negative electrons. Besides chemically stable molecules, short-lived molecular fragments called free radicals can be observed under special circumstances. See also Chemical bonding; Free radical; Molecular structure and spectra.

noun

A tiny amount: bit¹, crumb, dab¹, dash, dot, dram, drop, fragment, grain, iota, jot, minim, mite, modicum, ort, ounce, particle, scrap¹, scruple, shred, smidgen, speck, tittle, trifle, whit. Chiefly British spot. See big/small/amount.

A unit of matter that is the smallest particle of an element or chemical combination of atoms (as a compound) capable of retaining chemical identity with the substance in mass.

(click to enlarge)
Several methods of representing a molecule's structure. In Lewis structures, element symbols … (credit: © Merriam-Webster Inc.)

Smallest identifiable unit into which a pure substance can be divided and retain its composition and chemical properties. Division into still smaller parts, eventually atoms, involves destroying the bonding that holds the molecule together. For noble gases, the molecule is a single atom; all other substances have two (diatomic) or more (polyatomic) atoms in a molecule. The atoms are the same in elements, such as hydrogen (H₂), and different in compounds, such as glucose (C₆H₁₂O₆). Atoms always combine into molecules in fixed proportions. Molecules of different substances can have the same constituent atoms, either in different proportions, as in carbon monoxide (CO) and carbon dioxide (CO₂), or bonded in different ways (see isomer). The covalent bonds in molecules give them their shapes and most of their properties. (The concept of molecules has no significance in solids with ionic bonds.) Analysis with modern techniques and computers can determine and display the size, shape, and configuration of molecules, the positions of their nuclei and electron clouds, the lengths and angles of their bonds, and other details. Electron microscopy can even produce images of individual molecules and atoms. See also molecular weight.

For more information on molecule, visit Britannica.com.

Molecules are made of specific combinations of atoms. For example, carbon dioxide is made of one carbon atom and two oxygen atoms; water is made of two hydrogen atoms and one oxygen atom, and the atoms are joined by chemical bonds. Complex molecules such as starch may have hundreds of participating atoms linked in a specific pattern.

A combination of two or more atoms held together by a force between them. (See covalent bond and ionic bond.)

A group of atoms joined by chemical bonds; the smallest amount of a substance that possesses its characteristic properties.

Smallest particle of a compound that still retains its characteristics.

n.

The ultimate, indivisible unit of matter. It is distinguished from the corpuscle, also the ultimate, indivisible unit of matter, by a closer resemblance to the atom, also the ultimate, indivisible unit of matter. Three great scientific theories of the structure of the universe are the molecular, the corpuscular and the atomic. A fourth affirms, with Haeckel, the condensation of precipitation of matter from ether -- whose existence is proved by the condensation of precipitation. The present trend of scientific thought is toward the theory of ions. The ion differs from the molecule, the corpuscle and the atom in that it is an ion. A fifth theory is held by idiots, but it is doubtful if they know any more about the matter than the others.

3D (left and center) and 2D (right) representations of the terpenoid molecule atisane.

In chemistry, a molecule is defined as a sufficiently stable, electrically neutral group of at least two atoms in a definite arrangement held together by very strong (covalent) chemical bonds.^[1][2] Molecules are distinguished from polyatomic ions in this strict sense. In organic chemistry and biochemistry, the term molecule is used less strictly and also is applied to charged organic molecules and biomolecules.

In the kinetic theory of gases the term molecule is often used for any gaseous particle regardless of its composition.^[3] According to this definition noble gas atoms are considered molecules despite the fact that they are composed of a single non-bonded atom.^[4]

A molecule may consist of atoms of a single chemical element, as with oxygen (O₂), or of different elements, as with water (H₂O). Atoms and complexes connected by non-covalent bonds such as hydrogen bonds or ionic bonds are generally not considered single molecules.

No typical molecule can be defined for ionic crystals (salts) and covalent crystals (network solids), although these are often composed of repeating unit cells that extend either in a plane (such as in graphene) or three-dimensionally (such as in diamond or sodium chloride). The theme of repeated unit-cellular-structure also holds for most condensed phases with metallic bonding. In glasses (solids that exist in a vitreous disordered state), atoms may also be held together by chemical bonds without any definable molecule, but also without any of the regularity of repeating units that characterises crystals.

Contents 1 Molecular science 2 History and Etymology 3 Molecular size 3.1 Radius 4 Molecular formula 5 Molecular geometry 6 Molecular spectroscopy 7 Theoretical aspects 8 See also 9 Notes 10 External links

Molecular science

The science of molecules is called molecular chemistry or molecular physics, depending on the focus. Molecular chemistry deals with the laws governing the interaction between molecules that results in the formation and breakage of chemical bonds, while molecular physics deals with the laws governing their structure and properties. In practice, however, this distinction is vague. In molecular sciences, a molecule consists of a stable system (bound state) comprising two or more atoms. Polyatomic ions may sometimes be usefully thought of as electrically charged molecules. The term unstable molecule is used for very reactive species, i.e., short-lived assemblies (resonances) of electrons and nuclei, such as radicals, molecular ions, Rydberg molecules, transition states, van der Waals complexes, or systems of colliding atoms as in Bose-Einstein condensate

History and Etymology

Main article: History of the molecule

According to Merriam-Webster and the Online Etymology Dictionary, the word "molecule" derives from the Latin "moles" or small unit of mass.

• Molecule (1794) - "extremely minute particle," from Fr. molécule (1678), from Mod.L. molecula, dim. of L. moles "mass, barrier". A vague meaning at first; the vogue for the word (used until late 18th century only in Latin form) can be traced to the philosophy of Descartes.

Although the existence of molecules has been accepted by many chemists since the early 19th century as a result of Dalton's laws of Definite and Multiple Proportions (1803-1808) and Avogadro's law (1811), there was some resistance among positivists and physicists such as Mach, Boltzmann, Maxwell, and Gibbs, who saw molecules merely as convenient mathematical constructs. The work of Perrin on Brownian motion (1911) is considered to be the final proof of the existence of molecules.

The definition of the molecule has evolved as knowledge of the structure of molecules has increased. Earlier definitions were less precise, defining molecules as the smallest particles of pure chemical substances that still retain their composition and chemical properties.^[5] This definition often breaks down since many substances in ordinary experience, such as rocks, salts, and metals, are composed of large networks of chemically bonded atoms or ions, but are not made of discrete molecules.

Molecular size

Most molecules are far too small to be seen with the naked eye, but there are exceptions. DNA, a macromolecule, can reach macroscopic sizes, as can molecules of many polymers. The smallest molecule is the diatomic hydrogen (H₂), with an overall length of roughly twice the 74 picometres (0.74 Å) bond length. Molecules commonly used as building blocks for organic synthesis have a dimension of a few Å to several dozen Å. Single molecules cannot usually be observed by light (as noted above), but small molecules and even the outlines of individual atoms may be traced in some circumstances by use of an atomic force microscope. Some of the largest molecules are macromolecules or supermolecules.

Radius

Effective molecular radius is the size a molecule displays in solution.^[6][7] The table of permselectivity for different substances contains examples.

Molecular formula

A compound's empirical formula is the simplest integer ratio of the chemical elements that constitute it. For example, water is always composed of a 2:1 ratio of hydrogen to oxygen atoms, and ethyl alcohol or ethanol is always composed of carbon, hydrogen, and oxygen in a 2:6:1 ratio. However, this does not determine the kind of molecule uniquely - dimethyl ether has the same ratios as ethanol, for instance. Molecules with the same atoms in different arrangements are called isomers. Also carbohydrates, for example, haver the same ratio (carbon:hydrogen:oxygen= 1:2:1) (and thus the same empirical formula) but different total numbers of atoms in the molecule.

The molecular formula reflects the exact number of atoms that compose the molecule and so characterizes different isomers.

The empirical formula is often the same as the molecular formula but not always. For example the molecule acetylene has molecular formula C₂H₂, but the simplest integer ratio of elements is CH.

The molecular mass can be calculated from the chemical formula and is expressed in conventional atomic mass units equal to 1/12th of the mass of a neutral carbon-12 (¹²C isotope) atom. For network solids, the term formula unit is used in stoichiometric calculations.

Molecular geometry

Main article: Molecular geometry

Molecules have fixed equilibrium geometries—bond lengths and angles— about which they continuously oscillate through vibrational and rotational motions. A pure substance is composed of molecules with the same average geometrical structure. The chemical formula and the structure of a molecule are the two important factors that determine its properties, particularly its reactivity. Isomers share a chemical formula but normally have very different properties because of their different structures. Stereoisomers, a particular type of isomers, may have very similar physico-chemical properties and at the same time different biochemical activities.

Molecular spectroscopy

Main article: Spectroscopy

Molecular spectroscopy deals with the response (spectrum) of molecules interacting with probing signals of known energy (or frequency, according to Planck's formula). Molecules have quantized energy levels that can be analyzed by detecting the molecule's energy exchange through absorbance or emission.^[8] Spectroscopy does not generally refer to diffraction studies where particles such as neutrons, electrons, or high energy X-rays interact with a regular arrangement of molecules (as in a crystal).

Theoretical aspects

The study of molecules by molecular physics and theoretical chemistry is largely based on quantum mechanics and is essential for the understanding of the chemical bond. The simplest of molecules is the hydrogen molecule-ion, H₂⁺, and the simplest of all the chemical bonds is the one-electron bond. H₂⁺ is composed of two positively-charged protons and one negatively-charged electron, which means that the Schrödinger equation for the system can be solved more easily due to the lack of electron–electron repulsion. With the development of fast digital computers, approximate solutions for more complicated molecules became possible and are one of the main aspects of computational chemistry.

When trying to define rigorously whether an arrangement of atoms is "sufficiently stable" to be considered a molecule, IUPAC suggests that it "must correspond to a depression on the potential energy surface that is deep enough to confine at least one vibrational state".^[1] This definition does not depend on the nature of the interaction between the atoms, but only on the strength of the interaction. In fact, it includes weakly-bound species that would not traditionally be considered molecules, such as the helium dimer, He₂, which has one vibrational bound state^[9] and is so loosely bound that it is only likely to be observed at very low temperatures.

See also

• Atom
• Van der Waals molecule
• Diatomic molecule
• History of the molecule
• Chemical polarity
• Molecular geometry
• Covalent bond
• Noncovalent bonding
• list of compounds for a list of chemical compounds
• List of molecules in interstellar space
• Molecular Hamiltonian
• Molecular orbital
• Molecular modelling
• Molecular Design
• Small molecule

Notes

1. ^ ^{a b} International Union of Pure and Applied Chemistry (1994). "molecule". Compendium of Chemical Terminology Internet edition.
2. ^ Pauling, Linus (1970). General Chemistry. New York: Dover Publications, Inc.. ISBN 0-486-65622-5. 
   Ebbin, Darrell, D. (1990). General Chemistry, 3th Ed.. Boston: Houghton Mifflin Co.. ISBN 0-395-43302-9. 
   Brown, T.L. (2003). Chemistry – the Central Science, 9th Ed.. New Jersey: Prentice Hall. ISBN 0-13-066997-0. 
   Chang, Raymond (1998). Chemistry, 6th Ed.. New York: McGraw Hill. ISBN 0-07-115221-0. 
   Zumdahl, Steven S. (1997). Chemistry, 4th ed.. Boston: Houghton Mifflin. ISBN 0-669-41794-7. 
3. ^ E.g. see [1]
4. ^ Chandra, Sulekh. Comprehensive Inorganic Chemistry. New Age Publishers. ISBN 8122415121. 
5. ^ Molecule Definition (Frostburg State University)
6. ^ Chang RL, Deen WM, Robertson CR, Brenner BM. (Oct 1975). "Permselectivity of the glomerular capillary wall: III. Restricted transport of polyanions". Kidney Int. 8 (4): 212-218. PMID 1202253. 
7. ^ Chang RL, Ueki IF, Troy JL, Deen WM, Robertson CR, Brenner BM. (Sept 1975). "Permselectivity of the glomerular capillary wall to macromolecules. II. Experimental studies in rats using neutral dextran". Biophys J. 15 (9): 887-906. PMID 1182263. 
8. ^ International Union of Pure and Applied Chemistry (1997,2006). "spectroscopy". Compendium of Chemical Terminology Internet edition.
9. ^ Anderson JB (May 2004). "Comment on "An exact quantum Monte Carlo calculation of the helium-helium intermolecular potential" [J. Chem. Phys. 115, 4546 (2001)]". J Chem Phys 120 (20): 9886–7. doi:10.1063/1.1704638. PMID 15268005. 

External links

• Molecule of the Month - School of Chemistry, University of Bristol

v • d • e Particles in physics Elementary particles Fermions:  Quarks: u · d · c · s · t · b • Leptons: e− · e+ · μ− · μ+ · τ− · τ+ · νe · νμ · ντ · νe · νμ · ντ Bosons:  Gauge bosons: γ · g · W± · Z  Other:  Ghosts Composite particles Hadrons:  Baryons/Hyperons/Nucleons: p · n · Δ · Λ · Σ · Ξ · Ω • Mesons/Quarkonia: π · K · ρ · J/ψ · ϒ · B · D · η · η′ · φ Other:  Atomic nuclei • Atoms • Exotic atoms: Positronium · Muonium · Onia • Molecules Hypothetical elementary particles Superpartners: Axino · Chargino · Gaugino · Gluino · Gravitino · Higgsino · Neutralino · Sfermion Other: A0 · Dilaton · G · H0 · J · Tachyon · X · Y · W' · Z' · Sterile neutrino Hypothetical composite particles Exotic hadrons:  Exotic baryons: Dibaryon · Pentaquark • Exotic mesons: Glueball · Tetraquark Other:  Mesonic molecule · Pomeron Quasiparticles Davydov soliton · Exciton · Magnon · Phonon · Plasmon · Polariton · Polaron · Roton Lists List of particles · List of quasiparticles · List of baryons · List of mesons · Timeline of particle discoveries

v • d • e Into matter Macroscopic object > Molecule > Atom > Subatomic particle (Composite particle · Elementary particle)

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Dansk (Danish)
n. - molekyle

Nederlands (Dutch)
molecule

Français (French)
n. - molécule

Deutsch (German)
n. - Molekül, winziges Teilchen

Ελληνική (Greek)
n. - (φυσ., μτφ.) μόριο, κομματάκι

Italiano (Italian)
molecola

Português (Portuguese)
n. - molécula (f) (Quím.)

Русский (Russian)
молекула

Español (Spanish)
n. - molécula

Svenska (Swedish)
n. - molekyl

中文（简体）(Chinese (Simplified))
分子, 些微

中文（繁體）(Chinese (Traditional))
n. - 分子, 些微

한국어 (Korean)
n. - 분자, 소량

日本語 (Japanese)
n. - 分子, 微粒子, 微量, 微分子

العربيه (Arabic)
‏(الاسم) جزيء‏

עברית (Hebrew)
n. - ‮פרודה, מולקולה‬

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