The difference between ionic and covalent bonds isn’t as complex as you’d think. Image by Decoded Science
Almost all the atoms found in nature, left alone to themselves, are stable structures. If they always remained such, there would be no need of chemists.
Fortunately, when in close contact, atoms can react in a number of ways.
Often they link to each other in various combinations through bonding, forming molecules called compounds.
Such interaction requires explanation, and so provides employment to humans educated in this field: The field called chemistry.
Chemical Bonds: Ionic and Covalent
There are a variety of ways atoms bond to one another. Some bonds are weaker, and some are stronger. Two of the strongest forms of chemical bond are the ionic and the covalent bonds. Chemical bonds form between two atoms, each with its own electron environment.
- If each of the two atoms shares an electron with the other atom nearly equally, the bond is called covalent.
- If one atom exerts considerable force over the other atom’s electron, while the other atom strives to give its electron over, the bond is largely ionic.
Which form of bond—covalent or ionic—is the stronger? The easy way to determine that is to measure the energy it takes to break the bond. That quantity is called bond dissociation energy. The greater the energy it takes to break the bond, the stronger that bond must be. It turns out that most ionic bonds are considerably more difficult to break than covalent bonds.
Ionic Bonds: Electronegativity
In the 1930s, Dr. Linus Pauling expounded on the quality he called electronegativity. Some atoms, if they can be humanized, desire to increase their electron density. Others desire the exact opposite. He developed a list of numbers that quantified that affinity. The presence of appreciable electronegativity favors ionic bond formation.
The simplest way to determine which of the ionic bonds are strongest is to examine the electronegativities of the anion (the negative portion of a compound) and its cation (the positive portion of the compound).
Linus Pauling developed a list of numbers that quantified the property of electronegativity. Image courtesy of the U.S. Library of Congress
Alkali Metals and Halogens
The alkali metals are the least electronegative elements found in the periodic table, whereas the halogens are the most electronegative elements. We list three combinations of these elements, lithium iodide, potassium chloride and rubidium fluoride.
Lithium iodide———-352 kilojoules per mole
Potassium chloride—-427 kilojoules per mole
Rubidium fluoride—–494 kilojoules per mole
Stronger than even these should be the ionic bond in the compound francium fluoride. It is the most electropositive of the elements. Hence, it follows that francium is the least electronegative element, as well.
Chemical Bonds and Electronegativity
In chemistry, we study the interactions between atoms. Elements stick together via chemical bonds; covalent and ionic, including Linus Pauling’s electronegative elements. Which bonds are the strongest? It depends on their properties.
University of Buenos Aires. Properties of Atoms, Radicals, and Bonds . Accessed October 15, 2013.
Linus Pauling. The Nature of the Chemical Bond . Electronegativity: Narrative. Oregon State University. Accessed October 15, 2013.
Which bond is more stable ionic or covalent and why??
I came to know that covalent is more stable than ionic bond but howw
Smriti… I can best illustrate what I believe you are thinking of using two compounds, methyl “cyanide” [CH3CN] and sodium cyanide [NaCN]. The first is covalent, the second, ionic. Covalent methyl cyanide does not ionize in water. That is, you do not get CH3 and CN ions by dissolving in water. NaCN, however, produces ions. That is, it breaks into two pieces, each piece of which is surrounded by water molecules. However, and this is the point, although it is less stable with regard to coming apart in water, that is only temporary. When the water evaporates, you still have NaCN. So this dissolution in a solvent does not mean the ionic bond is weaker. It is simply less stable as a single entity in an appropriate solvent such as water.Reply
Just trying to clarify for myself: The alkali metals are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). These are LEAST electronegative, therefore the most electropositive? Electropositivity increases as you go further up the periodic table?
Halogens fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). These are the most electronegative with the greatest electronegativity being in those at the TOP of the periodic table?
So the greatest strength ionic bond is where the most electropositive (Fr) is matched with the most electronegative (F)? This would be FrF? Does this exist?Reply
You’ve got it. However, cesium is more electropositive than francium ‘because of relativistic effects’.Reply
Wow. That’s a lot of cool facts.Reply
Ionic bonds r stronger than covalent and it is well known fact but in reality diamond which has covalent bonding is hardest substance…and why is it so?i mean the hardest substance in universe should be ionic not covalentReply
Diamond is not the hardest substance known. And, even if it were… why should this depend upon being ionic or covalent?Reply
You are correct but the covalent bond is also present in the graphite why they are not hard because the covalent bond is depend on the bonding between the atom and the high melting point not on there physical behaviour.
On the other hand ionic bond is strong because the forces of attraction between the ion of opposite charge and repulsion between the ion of same charge.Reply
This is exactly the beauty of chemistry. The most important thing to consider in regards of a material strength is not its bonding type nor the elements, but rather on how those components are arrange to one another. Both diamond and graphite consisted of purely carbon atom, but on diamond those carbon atom bonded in a Three Dimension Vector, does making it have a very high bonding energy, thus difficult to break. While graphite only bond to the carbon existing on the same plane as your reference thus making it very brittle and easy to break.Reply
Yes diamond is the hardest substance known…
Its covalent bonds are arranged in tetrahedral manner which makes it one of the most stable compounds.Reply
Both ionic and covalent bonds are strong but unlike ionic structures, covalent once have weak intermolecular bonding and therefore they can break easily. When covalent bond form giant covalent lattices like in diamonds covalent bonds exist throughout the structure and it takes way more energy to break them and this makes them harder and an anomaly in the covalent structure. Hope it helpsReply
How did this conversation start? I never wrote about sexuality. Something strange is going on here.Reply
This has everything similar to the attraction between humans…think about it..we and everything in the multiverse is essentially energy constantly changing form. The pattern remain silent the same on a large scale or a quantum scale.Reply
The strength of Ionic and covalent bonds is similar the love between humans, where ionic represents woman and man (one has the electron the other needs), and covalent represent two people of the same sex (sharing electrons).
There’s possibility for two people of the same sex to possess a stronger love for one another than a man and woman, However, the strongest love possible for any one person, man or woman, is from that of the opposite sex; for from that bond, love can be made naturally in ways foreign to one of the same sex.Reply
An interesting allusion, illustration. Understandable to the reader.Reply
Maybe the ionic bond metaphor is correct but not as you said. Only because one atom is much more dominant in the relationship and they are not equalReply
Are the electrons a euphemism for the male genitalia?Reply
What an alienating thing to say. Why slander homosexuals to make a point about chemistry?Reply
Nice explanation sir
But the metallic bond is stronger than the ionic and covalent bond.
How to explain it sir ?Reply
Phosphates used in fertilizers contain no heavy metals.Reply
Nice Answer ..i love the way you explainReply
That’s kind of you. I appreciate it.Reply
Speed of the chemical reaction is affected by the type of the bond
When it is ionic it will be faster is that true?Reply
That’s a safe assumption.Reply
Ya its trueReply
I don’t belongs to chemistry but interested in it.
If any body knows than plZ tell me.
Which elements or chemicals can pull carbohydrates (rice) from one point to other.Reply
Let the reader use discernment: some of the comments above are not valid arguments.Reply
Yeah….this is absolutely right.Reply
Covalent bonds an ionic bonds are fairly equal in terms of strength, the reason covalent bonds are “weaker” (mp/bp) is actually due to most covalent compounds being simple molecular compounds. This means they are comprised of a maybe a dozen atoms or so…the molecule itself won’t decompose but it will melt due to intermolecular forces being weak, ionic compounds are all bonded together and every ion is bonded to 6 ions around it, meaning they have higher melting points. If you think about it giant covalent structures like diamond and silicon dioxide have very high melting points because even though they are covalent, every atom is bonded to the ones around it. So they aren’t weaker.Reply
what? It has to do with atomic mass. Electrons will automatically try to ionize (neutralize) with other atomic electrons in order to stabilize or (ionize). Think magnets vs. plastics.
Ionized particles are great conductors for this reason – stable..Reply
When the strength of ionic bond increases, its ionic nature decreases and covalent nature increases. But why covalent bonds are less stronger?Reply
“The greater the energy it takes to break the bond, the stronger that bond must be. It turns out that most ionic bonds are considerably more difficult to break than covalent bonds.”
That doesn’t really make sense. Covalent bonds are stronger than ionic and covalent bonds should have higher energy.Reply
Some do say covalent bonds are stronger; it is plausibly arguable both ways.Reply
Ya thats true…Reply
Think of it as if it is a magnet. Ionic bonds are between a metal and a non-metal. Metals generally need to lose electrons while non-metals need to gain them. this makes metals and non-metals attract, kinda like a magnet. The atoms exchange electrons and become ions, the metal has a positive charge because it has more protons than electrons, and the non-metal gains an electron and now the electrons are more numerous than the protons. The atoms then pretty much act like a magnet because one is positive and one is negative, so the stick together because opposites attract.
covalent bonds are shared electrons between two non-metals, they share the electron(s) to become stable, but both are negatively charged atoms, so the attraction is less. you don’t have the same magnetic styled attraction. I hope this kinda made things a bit clearer.Reply
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This article needs additional citations for verification . Please help improve this article by adding citations to reliable sources . Unsourced material may be challenged and removed. (March 2015) ( Learn how and when to remove this template message )
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.
Examples of Lewis dot -style representations of chemical bonds between carbon (C), hydrogen (H), and oxygen (O). Lewis dot diagrams were an early attempt to describe chemical bonding and are still widely used today.
All bonds can be explained by quantum theory , but, in practice, simplification rules allow chemists to predict the strength, directionality, and polarity of bonds. The octet rule and VSEPR theory are two examples. 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.
- 1 Overview of main types of chemical bonds
- 2 History
- 3 Bonds in chemical formulas
- 4 Strong chemical bonds
- 4.1 Ionic bond
- 4.2 Covalent bond
- 4.2.1 Single and multiple bonds
- 4.3 Coordinate covalent bond (Dipolar bond)
- 4.4 Metallic bonding
- 5 Intermolecular bonding
- 6 Theories of chemical bonding
- 7 References
- 8 External links
Overview of main types of chemical bonds[ edit ]
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 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.e. longer de Broglie wavelength ) orbital compared with each electron being confined closer to its respective nucleus.  These bonds exist between two particular identifiable atoms and have a direction in space, allowing them to be shown as single connecting lines between atoms in drawings, or modeled as sticks between spheres in models.
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 atoms and the atoms become positive or 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.
History[ edit ]
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 1704, 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.e. “hooked atoms”, “glued together by rest”, or “stuck together by conspiring motions”, Newton states that he would rather infer from their cohesion, that “particles attract one another by some force , which in immediate contact is exceedingly strong, at small distances performs the chemical operations, and reaches not far from the particles with any sensible effect.”
In 1819, on the heels of the invention of the voltaic pile , Jöns Jakob Berzelius developed a theory of chemical combination stressing the electronegative and electropositive characters of the combining atoms. By the mid 19th century, Edward Frankland , F.A. Kekulé , A.S. 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 1916, 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 (1904).
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 1927, the first mathematically complete quantum description of a simple chemical bond, i.e. that produced by one electron in the hydrogen molecular ion, H2+ , was derived by the Danish physicist Øyvind Burrau .  This work showed that the quantum approach to chemical bonds could be fundamentally and quantitatively correct, but the mathematical methods used could not be extended to molecules containing more than one electron. 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 1929, 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 F2 ( fluorine ) and O2 ( oxygen ) molecules, from basic quantum principles. This molecular orbital theory represented a covalent bond as an orbital formed by combining the quantum mechanical Schrödinger atomic orbitals which had been hypothesized for electrons in single atoms. The equations for bonding electrons in multi-electron atoms could not be solved to mathematical perfection (i.e., analytically), but approximations for them still gave many good qualitative predictions and results. 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 1933, H. H. James and A. S. 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.  With up to 13 adjustable parameters they obtained a result very close to the experimental result for the dissociation energy. 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.
Bonds in chemical formulas[ edit ]
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 (CH3–CH2–OH), by separating the functional group from another part of the molecule (C2H5OH), or by its atomic constituents (C2H6O), according to what is discussed. Sometimes, even the non-bonding valence shell electrons (with the two-dimensional approximate directions) are marked, e.g. for elemental carbon .‘C‘. Some chemists may also mark the respective orbitals, e.g. the hypothetical ethene−4 anion (\/C=C/\ −4) indicating the possibility of bond formation.
Strong chemical bonds[ edit ]
|Typical bond lengths in pm|
and bond energies in kJ/mol.
Bond lengths can be converted to Å
by division by 100 (1 Å = 100 pm).
Data taken from University of Waterloo. 
|H — Hydrogen|
|C — Carbon|
|N — Nitrogen|
|O — Oxygen|
|F, Cl, Br, I — Halogens|
Strong chemical bonds are the intramolecular forces which 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 bond[ edit ]
Ionic bonding is a type of electrostatic interaction between atoms which have a large electronegativity difference. There is no precise value that distinguishes ionic from covalent bonding, but a difference of electronegativity of over 1.7 is likely to be ionic, and a difference of less than 1.7 is likely to be covalent.  Ionic bonding leads to separate positive and negative ions . Ionic charges are commonly between −3 e to +3 e .
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. X-ray diffraction shows that in NaCN, for example, the bonds between sodium cations (Na+) and the cyanide anions (CN−) are ionic, with no sodium ion associated with any particular cyanide. 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. For example, in solution, the cyanide ions, still bound together as single CN− ions, move independently through the solution, as do sodium ions, as Na+. 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 bond[ edit ]
Nonpolar covalent bonds in methane (CH4). The Lewis structure shows electrons shared between C and H atoms.
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 nonpolar covalent bonds, the electronegativity difference between the bonded atoms is small, typically 0 to 0.3. Bonds within most organic compounds are described as covalent. The figure shows methane (CH4), in which each hydrogen forms a covalent bond with the carbon. See sigma bonds and pi bonds for LCAO-description of such bonding.
Molecules which are formed primarily from non-polar covalent bonds are often immiscible in water or other polar solvents , but much more soluble in non-polar solvents such as hexane .
A polar covalent bond is a covalent bond with a significant ionic character . This means that the two shared electrons are closer to one of the atoms than the other, creating an imbalance of charge. Such bonds occur between two atoms with moderately different electronegativities and give rise to dipole–dipole interactions . The electronegativity difference between the two atoms in these bonds is 0.3 to 1.7.
Single and multiple bonds[ edit ]
A single bond between two atoms corresponds to the sharing of one pair of electrons. The electron density of these two bonding electrons is concentrated in the region between the two atoms, which is the defining quality of a sigma bond .
Two p-orbitals forming a pi-bond.
A double bond between two atoms is formed by the sharing of two pairs of electrons, one in a sigma bond and one in a pi bond , with electron density concentrated on two opposite sides of the internuclear axis. A triple bond consists of three shared electron pairs, forming one sigma and two pi bonds.
Quadruple and higher bonds are very rare and occur only between certain transition metal atoms.
Coordinate covalent bond (Dipolar bond)[ edit ]
Adduct of ammonia and boron trifluoride
A coordinate covalent bond is a covalent bond in which the two shared
bonding electrons are from the same one of the atoms involved in the bond.
For example, boron trifluoride (BF3) and ammonia (NH3) from an adduct or coordination complex F3B←NH3 with a B–N bond in which a lone pair of electrons on N is shared with an empty atomic orbital on B. BF3 with an empty orbital is described as an electron pair acceptor or Lewis acid , while NH3 with a lone pair which can be shared is described as an electron-pair donor or Lewis base . The electrons are shared roughly equally between the atoms in contrast to ionic bonding. Such bonding is shown by an arrow pointing to the Lewis acid.
Transition metal complexes are generally bound by coordinate covalent bonds. For example, the ion Ag+ reacts as a Lewis acid with two molecules of the Lewis base NH3 to form the complex ion Ag(NH3)2+, which has two Ag←N coordinate covalent bonds.
Metallic bonding[ edit ]
In metallic bonding, bonding electrons are delocalized over a lattice of atoms. By contrast, in ionic compounds, the locations of the binding electrons and their charges are static. The freely-moving or delocalization of bonding electrons leads to classical metallic properties such as luster (surface light reflectivity ), electrical and thermal conductivity , ductility , and high tensile strength .
Intermolecular bonding[ edit ]
There are four basic types of bonds that can be formed between two or more (otherwise non-associated) molecules, ions or atoms. Intermolecular forces cause molecules to be attracted or repulsed by each other. Often, these define some of the physical characteristics (such as the melting point ) of a substance.
- A large difference in electronegativity between two bonded atoms will cause a permanent charge separation, or dipole, in a molecule or ion. Two or more molecules or ions with permanent dipoles can interact within dipole-dipole interactions . The bonding electrons in a molecule or ion will, on average, be closer to the more electronegative atom more frequently than the less electronegative one, giving rise to partial charges on each atom, and causing electrostatic forces between molecules or ions.
- A hydrogen bond is effectively a strong example of an interaction between two permanent dipoles. The large difference in electronegativities between hydrogen and any of fluorine , nitrogen and oxygen , coupled with their lone pairs of electrons cause strong electrostatic forces between molecules. Hydrogen bonds are responsible for the high boiling points of water and ammonia with respect to their heavier analogues.
- The London dispersion force arises due to instantaneous dipoles in neighbouring atoms. As the negative charge of the electron is not uniform around the whole atom, there is always a charge imbalance. This small charge will induce a corresponding dipole in a nearby molecule; causing an attraction between the two. The electron then moves to another part of the electron cloud and the attraction is broken.
- A cation–pi interaction occurs between a pi bond and a cation.
Theories of chemical bonding[ edit ]
In the (unrealistic) limit of “pure” ionic bonding , electrons are perfectly localized on one of the two atoms in the bond. Such bonds can be understood by classical physics . The forces between the atoms are characterized by isotropic continuum electrostatic potentials. Their magnitude is in simple proportion to the charge difference.
Covalent bonds are better understood by valence bond theory or molecular orbital theory . The properties of the atoms involved can be understood using concepts such as oxidation number . The electron density within a bond is not assigned to individual atoms, but is instead delocalized between atoms. In valence bond theory, the two electrons on the two atoms are coupled together with the bond strength depending on the overlap between them. In molecular orbital theory, the linear combination of atomic orbitals (LCAO) helps describe the delocalized molecular orbital structures and energies based on the atomic orbitals of the atoms they came from. Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties. These may have their own names, such as sigma bond and pi bond .
In the general case, atoms form bonds that are intermediate between ionic and covalent, depending on the relative electronegativity of the atoms involved. This type of bond is sometimes called polar covalent .
References[ edit ]
- ^ Rioux, F. (2001). “The Covalent Bond in H2“. The Chemical Educator. 6 (5): 288. doi : 10.1007/s00897010509a .
- ^ Lewis, Gilbert N. (1916). “The Atom and the Molecule” . Journal of the American Chemical Society . 38 (4): 772. doi : 10.1021/ja02261a002 . a copy
- ^ Бор Н. (1970). Избранные научные труды (статьи 1909-1925). 1. М.: «Наука». p. 133.
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External links[ edit ]
|Wikiquote has quotations related to: Chemical bond|
|Wikimedia Commons has media related to Chemical bonding .|
- W. Locke (1997). Introduction to Molecular Orbital Theory . Retrieved May 18, 2005.
- Carl R. Nave (2005). HyperPhysics . Retrieved May 18, 2005.
- Linus Pauling and the Nature of the Chemical Bond: A Documentary History . Retrieved February 29, 2008.
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