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The concepts of atoms and bonds in molecules which appeared in chemistry during the nineteenth century are unavoidable to explain the structure and the reactivity of the matter at a chemical level of understanding. Although they can be criticized from a strict reductionist point of view, because neither atoms nor bonds are observable in the sense of quantum mechanics, the topological and statistical interpretative approaches of quantum chemistry quantum theory of atoms in molecules, electron localization function and maximum probability domain provide consistent definitions which accommodate chemistry and quantum mechanics. This is a preview of subscription content, access via your institution.
A chemical bond is a lasting attraction between atoms , ions or molecules that enables the formation of chemical compounds. The bond may result from the electrostatic force of attraction between oppositely charged ions as in ionic bonds or through the sharing of electrons as in covalent bonds. The strength of chemical bonds varies considerably; there are "strong bonds" or "primary bonds" such as covalent, ionic and metallic bonds, and "weak bonds" or "secondary bonds" such as dipole—dipole interactions , the London dispersion force and hydrogen bonding.
Since opposite charges attract via a simple electromagnetic force , the negatively charged electrons that are orbiting the nucleus and the positively charged protons in the nucleus attract each other. An electron positioned between two nuclei will be attracted to both of them, and the nuclei will be attracted toward electrons in this position. This attraction constitutes the chemical bond. Due to the matter wave nature of electrons and their smaller mass, they must occupy a much larger amount of volume compared with the nuclei, and this volume occupied by the electrons keeps the atomic nuclei in a bond relatively far apart, as compared with the size of the nuclei themselves.
In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. The atoms in molecules , crystals , metals and diatomic gases—indeed most of the physical environment around us—are held together by chemical bonds, which dictate the structure and the bulk properties of matter.
All bonds can be explained by quantum theory , but, in practice, simplification rules allow chemists to predict the strength, directionality, and polarity of bonds. More sophisticated theories are valence bond theory , which includes orbital hybridization and resonance , and molecular orbital theory which includes linear combination of atomic orbitals and ligand field theory.
Electrostatics are used to describe bond polarities and the effects they have on chemical substances. A chemical bond is an attraction between atoms. This attraction may be seen as the result of different behaviors of the outermost or valence electrons of atoms.
These behaviors merge into each other seamlessly in various circumstances, so that there is no clear line to be drawn between them.
However it remains useful and customary to differentiate between different types of bond, which result in different properties of condensed matter. In the simplest view of a covalent bond , one or more electrons often a pair of electrons are drawn into the space between the two atomic nuclei.
Energy is released by bond formation. This is not as a result of reduction in potential energy, because the attraction of the two electrons to the two protons is offset by the electron-electron and proton-proton repulsions.
Instead, the release of energy and hence stability of the bond arises from the reduction in kinetic energy due to the electrons being in a more spatially distributed i. In a polar covalent bond , one or more electrons are unequally shared between two nuclei.
Covalent bonds often result in the formation of small collections of better-connected atoms called molecules , which in solids and liquids are bound to other molecules by forces that are often much weaker than the covalent bonds that hold the molecules internally together. Such weak intermolecular bonds give organic molecular substances, such as waxes and oils, their soft bulk character, and their low melting points in liquids, molecules must cease most structured or oriented contact with each other.
When covalent bonds link long chains of atoms in large molecules, however as in polymers such as nylon , or when covalent bonds extend in networks through solids that are not composed of discrete molecules such as diamond or quartz or the silicate minerals in many types of rock then the structures that result may be both strong and tough, at least in the direction oriented correctly with networks of covalent bonds. Also, the melting points of such covalent polymers and networks increase greatly.
In a simplified view of an ionic bond , the bonding electron is not shared at all, but transferred. In this type of bond, the outer atomic orbital of one atom has a vacancy which allows the addition of one or more electrons. These newly added electrons potentially occupy a lower energy-state effectively closer to more nuclear charge than they experience in a different atom. Thus, one nucleus offers a more tightly bound position to an electron than does another nucleus, with the result that one atom may transfer an electron to the other.
This transfer causes one atom to assume a net positive charge, and the other to assume a net negative charge. The bond then results from electrostatic attraction between the positive and negatively charged ions. Ionic bonds may be seen as extreme examples of polarization in covalent bonds.
Often, such bonds have no particular orientation in space, since they result from equal electrostatic attraction of each ion to all ions around them. Ionic bonds are strong and thus ionic substances require high temperatures to melt but also brittle, since the forces between ions are short-range and do not easily bridge cracks and fractures. This type of bond gives rise to the physical characteristics of crystals of classic mineral salts, such as table salt.
A less often mentioned type of bonding is metallic bonding. In this type of bonding, each atom in a metal donates one or more electrons to a "sea" of electrons that reside between many metal atoms. In this sea, each electron is free by virtue of its wave nature to be associated with a great many atoms at once. The bond results because the metal atoms become somewhat positively charged due to loss of their electrons while the electrons remain attracted to many atoms, without being part of any given atom.
Metallic bonding may be seen as an extreme example of delocalization of electrons over a large system of covalent bonds, in which every atom participates. This type of bonding is often very strong resulting in the tensile strength of metals. However, metallic bonding is more collective in nature than other types, and so they allow metal crystals to more easily deform, because they are composed of atoms attracted to each other, but not in any particularly-oriented ways.
This results in the malleability of metals. The cloud of electrons in metallic bonding causes the characteristically good electrical and thermal conductivity of metals, and also their shiny lustre that reflects most frequencies of white light.
Early speculations about the nature of the chemical bond , from as early as the 12th century, supposed that certain types of chemical species were joined by a type of chemical affinity. In , Sir Isaac Newton famously outlined his atomic bonding theory, in "Query 31" of his Opticks , whereby atoms attach to each other by some " force ".
Specifically, after acknowledging the various popular theories in vogue at the time, of how atoms were reasoned to attach to each other, i. By the mid 19th century, Edward Frankland , F. Couper, Alexander Butlerov , and Hermann Kolbe , building on the theory of radicals , developed the theory of valency , originally called "combining power", in which compounds were joined owing to an attraction of positive and negative poles.
In , chemist Gilbert N. Lewis developed the concept of the electron-pair bond , in which two atoms may share one to six electrons, thus forming the single electron bond , a single bond , a double bond , or a triple bond ; in Lewis's own words, "An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively. That same year, Walther Kossel put forward a theory similar to Lewis' only his model assumed complete transfers of electrons between atoms, and was thus a model of ionic bonding.
Both Lewis and Kossel structured their bonding models on that of Abegg's rule Niels Bohr proposed a model of the atom and a model of the chemical bond. According to his model for a diatomic molecule , the electrons of the atoms of the molecule form a rotating ring whose plane is perpendicular to the axis of the molecule and equidistant from the atomic nuclei.
The dynamic equilibrium of the molecular system is achieved through the balance of forces between the forces of attraction of nuclei to the plane of the ring of electrons and the forces of mutual repulsion of the nuclei. The Bohr model of the chemical bond took into account the Coulomb repulsion — the electrons in the ring are at the maximum distance from each other. In , the first mathematically complete quantum description of a simple chemical bond, i. A more practical, albeit less quantitative, approach was put forward in the same year by Walter Heitler and Fritz London.
The Heitler—London method forms the basis of what is now called valence bond theory. In , the linear combination of atomic orbitals molecular orbital method LCAO approximation was introduced by Sir John Lennard-Jones , who also suggested methods to derive electronic structures of molecules of F 2 fluorine and O 2 oxygen molecules, from basic quantum principles.
The equations for bonding electrons in multi-electron atoms could not be solved to mathematical perfection i. Most quantitative calculations in modern quantum chemistry use either valence bond or molecular orbital theory as a starting point, although a third approach, density functional theory , has become increasingly popular in recent years. In , H. James and A. Coolidge carried out a calculation on the dihydrogen molecule that, unlike all previous calculation which used functions only of the distance of the electron from the atomic nucleus, used functions which also explicitly added the distance between the two electrons.
Later extensions have used up to 54 parameters and gave excellent agreement with experiments. This calculation convinced the scientific community that quantum theory could give agreement with experiment. However this approach has none of the physical pictures of the valence bond and molecular orbital theories and is difficult to extend to larger molecules.
Because atoms and molecules are three-dimensional, it is difficult to use a single method to indicate orbitals and bonds. In molecular formulas the chemical bonds binding orbitals between atoms are indicated in different ways depending on the type of discussion. Sometimes, some details are neglected. For example, in organic chemistry one is sometimes concerned only with the functional group of the molecule. Thus, the molecular formula of ethanol may be written in conformational form, three-dimensional form, full two-dimensional form indicating every bond with no three-dimensional directions , compressed two-dimensional form CH 3 —CH 2 —OH , by separating the functional group from another part of the molecule C 2 H 5 OH , or by its atomic constituents C 2 H 6 O , according to what is discussed.
Sometimes, even the non-bonding valence shell electrons with the two-dimensional approximate directions are marked, e. Some chemists may also mark the respective orbitals, e. Strong chemical bonds are the intramolecular forces that hold atoms together in molecules. A strong chemical bond is formed from the transfer or sharing of electrons between atomic centers and relies on the electrostatic attraction between the protons in nuclei and the electrons in the orbitals.
The types of strong bond differ due to the difference in electronegativity of the constituent elements. A large difference in electronegativity leads to more polar ionic character in the bond. Ionic bonding is a type of electrostatic interaction between atoms that have a large electronegativity difference. There is no precise value that distinguishes ionic from covalent bonding, but an electronegativity difference of over 1.
Ionic bonding commonly occurs in metal salts such as sodium chloride table salt. A typical feature of ionic bonds is that the species form into ionic crystals, in which no ion is specifically paired with any single other ion in a specific directional bond.
Rather, each species of ion is surrounded by ions of the opposite charge, and the spacing between it and each of the oppositely charged ions near it is the same for all surrounding atoms of the same type. It is thus no longer possible to associate an ion with any specific other single ionized atom near it. This is a situation unlike that in covalent crystals, where covalent bonds between specific atoms are still discernible from the shorter distances between them, as measured via such techniques as X-ray diffraction.
Ionic crystals may contain a mixture of covalent and ionic species, as for example salts of complex acids such as sodium cyanide, NaCN. However, the bonds between C and N atoms in cyanide are of the covalent type, so that each carbon is strongly bound to just one nitrogen, to which it is physically much closer than it is to other carbons or nitrogens in a sodium cyanide crystal.
When such crystals are melted into liquids, the ionic bonds are broken first because they are non-directional and allow the charged species to move freely. Similarly, when such salts dissolve into water, the ionic bonds are typically broken by the interaction with water but the covalent bonds continue to hold. In water, charged ions move apart because each of them are more strongly attracted to a number of water molecules than to each other.
The attraction between ions and water molecules in such solutions is due to a type of weak dipole-dipole type chemical bond. In melted ionic compounds, the ions continue to be attracted to each other, but not in any ordered or crystalline way. Covalent bonding is a common type of bonding in which two or more atoms share valence electrons more or less equally. The simplest and most common type is a single bond in which two atoms share two electrons.
Other types include the double bond , the triple bond , one- and three-electron bonds , the three-center two-electron bond and three-center four-electron bond. In non-polar covalent bonds, the electronegativity difference between the bonded atoms is small, typically 0 to 0.
Bonds within most organic compounds are described as covalent.
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A chemical bond is a lasting attraction between atoms , ions or molecules that enables the formation of chemical compounds. The bond may result from the electrostatic force of attraction between oppositely charged ions as in ionic bonds or through the sharing of electrons as in covalent bonds. The strength of chemical bonds varies considerably; there are "strong bonds" or "primary bonds" such as covalent, ionic and metallic bonds, and "weak bonds" or "secondary bonds" such as dipole—dipole interactions , the London dispersion force and hydrogen bonding. Since opposite charges attract via a simple electromagnetic force , the negatively charged electrons that are orbiting the nucleus and the positively charged protons in the nucleus attract each other. An electron positioned between two nuclei will be attracted to both of them, and the nuclei will be attracted toward electrons in this position.
The atoms of each element differ by their number of protons, neutrons, and electrons. For example, hydrogen has one proton, one electron, and no neutrons, while carbon has six protons, six neutrons, and six electrons. The number and arrangement of electrons of an atom determine the kinds of chemical bonds that it forms and how it reacts with other atoms to form molecules. There are three kinds of chemical bonds:. The atom that gains electrons has an overall negative charge, and the atom that donates electrons has an overall positive charge.
Skill 1. Organic chemistry studies the properties and reactions of organic compounds. Protons and neutrons are found in the nucleus of the atom, while electrons are found in the electron cloud around the nucleus. The mass number, A , is the sum of the number of protons and the number of neutrons in a nucleus. The type of element an atom represents is defined by the atomic number, Z in the atom. All atoms of one specific element have the same number of protons Z. Atoms that have the same atomic number Z , but different mass numbers A are called isotopes.
A. Structure of Atoms, Molecules and Chemical Bonds. 1. 1. A. III. Comparison Between Different Types Of Bonds And Their Biological Importance. 1.
At its most fundamental level, life is made up of matter. Matter is any substance that occupies space and has mass. Elements are unique forms of matter with specific chemical and physical properties that cannot break down into smaller substances by ordinary chemical reactions. There are elements, but only 98 occur naturally. The remaining elements are unstable and require scientists to synthesize them in laboratories.
Atoms separated by a great distance cannot link; rather, they must come close enough for the electrons in their valence shells to interact. But do atoms ever actually touch one another? Most physicists would say no, because the negatively charged electrons in their valence shells repel one another. No force within the human body—or anywhere in the natural world—is strong enough to overcome this electrical repulsion. So when you read about atoms linking together or colliding, bear in mind that the atoms are not merging in a physical sense.
Chemical bonding , any of the interactions that account for the association of atoms into molecules , ions , crystals , and other stable species that make up the familiar substances of the everyday world. When atoms approach one another, their nuclei and electrons interact and tend to distribute themselves in space in such a way that the total energy is lower than it would be in any alternative arrangement. If the total energy of a group of atoms is lower than the sum of the energies of the component atoms, they then bond together and the energy lowering is the bonding energy.
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Я не хотел тебя впутывать. - Я… понимаю, - тихо сказала она, все еще находясь под впечатлением его блистательного замысла. - Вы довольно искусный лжец. Стратмор засмеялся. - Годы тренировки.
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