Covalent bond

A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, and the stable balance of attractive and repulsive forces between atoms, when they share electrons, is known as covalent bonding.[1] For many molecules, the sharing of electrons allows each atom to attain the equivalent of a full outer shell, corresponding to a stable electronic configuration.

Covalent bonding includes many kinds of interactions, including σ-bonding, π-bonding, metal-to-metal bonding, agostic interactions, bent bonds, and three-center two-electron bonds.[2][3] The term covalent bond dates from 1939.[4] The prefix co- means jointly, associated in action, partnered to a lesser degree, etc.; thus a "co-valent bond", in essence, means that the atoms share "valence", such as is discussed in valence bond theory.

In the molecule H
2
, the hydrogen atoms share the two electrons via covalent bonding.[5] Covalency is greatest between atoms of similar electronegativities. Thus, covalent bonding does not necessarily require that the two atoms be of the same elements, only that they be of comparable electronegativity. Covalent bonding that entails sharing of electrons over more than two atoms is said to be delocalized.

Covalent bond hydrogen
A covalent bond forming H2 (right) where two hydrogen atoms share the two electrons

History

Covalent
Early concepts in covalent bonding arose from this kind of image of the molecule of methane. Covalent bonding is implied in the Lewis structure by indicating electrons shared between atoms.

The term covalence in regard to bonding was first used in 1919 by Irving Langmuir in a Journal of the American Chemical Society article entitled "The Arrangement of Electrons in Atoms and Molecules". Langmuir wrote that "we shall denote by the term covalence the number of pairs of electrons that a given atom shares with its neighbors."[6]

The idea of covalent bonding can be traced several years before 1919 to Gilbert N. Lewis, who in 1916 described the sharing of electron pairs between atoms.[7] He introduced the Lewis notation or electron dot notation or Lewis dot structure, in which valence electrons (those in the outer shell) are represented as dots around the atomic symbols. Pairs of electrons located between atoms represent covalent bonds. Multiple pairs represent multiple bonds, such as double bonds and triple bonds. An alternative form of representation, not shown here, has bond-forming electron pairs represented as solid lines.

Lewis proposed that an atom forms enough covalent bonds to form a full (or closed) outer electron shell. In the diagram of methane shown here, the carbon atom has a valence of four and is, therefore, surrounded by eight electrons (the octet rule), four from the carbon itself and four from the hydrogens bonded to it. Each hydrogen has a valence of one and is surrounded by two electrons (a duet rule) – its own one electron plus one from the carbon. The numbers of electrons correspond to full shells in the quantum theory of the atom; the outer shell of a carbon atom is the n = 2 shell, which can hold eight electrons, whereas the outer (and only) shell of a hydrogen atom is the n = 1 shell, which can hold only two.

While the idea of shared electron pairs provides an effective qualitative picture of covalent bonding, quantum mechanics is needed to understand the nature of these bonds and predict the structures and properties of simple molecules. Walter Heitler and Fritz London are credited with the first successful quantum mechanical explanation of a chemical bond (molecular hydrogen) in 1927.[8] Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is good overlap between the atomic orbitals of participating atoms.

Types of covalent bonds

Atomic orbitals (except for s orbitals) have specific directional properties leading to different types of covalent bonds. Sigma (σ) bonds are the strongest covalent bonds and are due to head-on overlapping of orbitals on two different atoms. A single bond is usually a σ bond. Pi (π) bonds are weaker and are due to lateral overlap between p (or d) orbitals. A double bond between two given atoms consists of one σ and one π bond, and a triple bond is one σ and two π bonds.

Covalent bonds are also affected by the electronegativity of the connected atoms which determines the chemical polarity of the bond. Two atoms with equal electronegativity will make nonpolar covalent bonds such as H–H. An unequal relationship creates a polar covalent bond such as with H−Cl. However polarity also requires geometric asymmetry, or else dipoles may cancel out resulting in a non-polar molecule.

Covalent structures

There are several types of structures for covalent substances, including individual molecules, molecular structures, macromolecular structures and giant covalent structures. Individual molecules have strong bonds that hold the atoms together, but there are negligible forces of attraction between molecules. Such covalent substances are usually gases, for example, HCl, SO2, CO2, and CH4. In molecular structures, there are weak forces of attraction. Such covalent substances are low-boiling-temperature liquids (such as ethanol), and low-melting-temperature solids (such as iodine and solid CO2). Macromolecular structures have large numbers of atoms linked by covalent bonds in chains, including synthetic polymers such as polyethylene and nylon, and biopolymers such as proteins and starch. Network covalent structures (or giant covalent structures) contain large numbers of atoms linked in sheets (such as graphite), or 3-dimensional structures (such as diamond and quartz). These substances have high melting and boiling points, are frequently brittle, and tend to have high electrical resistivity. Elements that have high electronegativity, and the ability to form three or four electron pair bonds, often form such large macromolecular structures.[9]

One- and three-electron bonds

Graphical comparison of bonds
Lewis and MO diagrams of an individual 2e bond and 3e bond

Bonds with one or three electrons can be found in radical species, which have an odd number of electrons. The simplest example of a 1-electron bond is found in the dihydrogen cation, H+
2
. One-electron bonds often have about half the bond energy of a 2-electron bond, and are therefore called "half bonds". However, there are exceptions: in the case of dilithium, the bond is actually stronger for the 1-electron Li+
2
than for the 2-electron Li2. This exception can be explained in terms of hybridization and inner-shell effects.[10]

The simplest example of three-electron bonding can be found in the helium dimer cation, He+
2
. It is considered a "half bond" because it consists of only one shared electron (rather than two);[11] in molecular orbital terms, the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons. Another example of a molecule containing a 3-electron bond, in addition to two 2-electron bonds, is nitric oxide, NO. The oxygen molecule, O2 can also be regarded as having two 3-electron bonds and one 2-electron bond, which accounts for its paramagnetism and its formal bond order of 2.[12] Chlorine dioxide and its heavier analogues bromine dioxide and iodine dioxide also contain three-electron bonds.

Molecules with odd-electron bonds are usually highly reactive. These types of bond are only stable between atoms with similar electronegativities.[12]

Nitric oxide
Nitric oxide
Triplett-Sauerstoff
Dioxygen

Resonance

There are situations whereby a single Lewis structure is insufficient to explain the electron configuration in a molecule, hence a superposition of structures are needed. The same two atoms in such molecules can be bonded differently in different structures (a single bond in one, a double bond in another, or even none at all), resulting in a non-integer bond order. The nitrate ion is one such example with three equivalent structures. The bond between the nitrogen and each oxygen is a double bond in one structure and a single bond in the other two, so that the average bond order for each N–O interaction is 2 + 1 + 1/3 = 4/3.

Nitrate-ion-resonance-2D

Nitrate-ion-resonance-2D

Aromaticity

In organic chemistry, when a molecule with a planar ring obeys Hückel's rule, where the number of π electrons fit the formula 4n + 2 (where n is an integer), it attains extra stability and symmetry. In benzene, the prototypical aromatic compound, there are 6 π bonding electrons (n = 1, 4n + 2 = 6). These occupy three delocalized π molecular orbitals (molecular orbital theory) or form conjugate π bonds in two resonance structures that linearly combine (valence bond theory), creating a regular hexagon exhibiting a greater stabilization than the hypothetical 1,3,5-cyclohexatriene.

In the case of heterocyclic aromatics and substituted benzenes, the electronegativity differences between different parts of the ring may dominate the chemical behaviour of aromatic ring bonds, which otherwise are equivalent.

Hypervalence

Certain molecules such as xenon difluoride and sulfur hexafluoride have higher co-ordination numbers than would be possible due to strictly covalent bonding according to the octet rule. This is explained by the three-center four-electron bond ("3c–4e") model which interprets the molecular wavefunction in terms of non-bonding highest occupied molecular orbitals in molecular orbital theory and resonance of sigma bonds in valence bond theory.

Electron deficiency

In three-center two-electron bonds ("3c–2e") three atoms share two electrons in bonding. This type of bonding occurs in electron deficient compounds like diborane. Each such bond (2 per molecule in diborane) contains a pair of electrons which connect the boron atoms to each other in a banana shape, with a proton (nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron atoms. In certain cluster compounds, so-called four-center two-electron bonds also have been postulated.

Quantum mechanical description

After the development of quantum mechanics, two basic theories were proposed to provide a quantum description of chemical bonding: valence bond (VB) theory and molecular orbital (MO) theory. A more recent quantum description[13] is given in terms of atomic contributions to the electronic density of states.

Comparison of VB and MO theories

The two theories represent two ways to build up the electron configuration of the molecule.[14] For valence bond theory, the atomic hybrid orbitals are filled first according to the octet rule to produce a fully bonded valence configuration, followed by a linear combination of contributing structures (resonance) if there are several of them. In contrast, for molecular orbital theory a linear combination of atomic orbitals is performed first, followed by the filling of the resulting molecular orbitals according to the aufbau principle.

The two approaches are regarded as complementary, and each provides its own insights into the problem of chemical bonding. As valence bond theory builds the molecular wavefunction out of localized bonds, it is more suited for the calculation of bond energies and the understanding of reaction mechanisms. As molecular orbital theory builds the molecular wavefunction out of delocalized orbitals, it is more suited for the calculation of ionization energies and the understanding of spectral absorption bands.

At the qualitative level, both theories contain incorrect predictions. Simple (Heitler–London) valence bond theory correctly predicts the dissociation of homonuclear diatomic molecules into separate atoms, while simple (Hartree–Fock) molecular orbital theory incorrectly predicts dissociation into a mixture of atoms and ions. On the other hand, simple molecular orbital theory correctly predicts Hückel's rule of aromaticity, while simple valence bond theory incorrectly predicts that cyclobutadiene has a larger resonance energy than benzene.

Although the wavefunctions generated by both theories at the qualitative level do not agree and do not match the stabilization energy by experiment, they can be corrected by configuration interaction.[14] This is done by combining the valence bond covalent function with the functions describing all possible ionic configurations or by combining the molecular orbital ground state function with the functions describing all possible excited states using unoccupied orbitals. It can then be seen that the simple molecular orbital approach overestimates the weight of the ionic structures while the simple valence bond approach neglects them. This can also be described as saying that the molecular orbital approach neglects electron correlation while the valence bond approach overestimates it.[14]

Modern calculations in quantum chemistry usually start from (but ultimately go far beyond) a molecular orbital rather than a valence bond approach, not because of any intrinsic superiority in the former but rather because the MO approach is more readily adapted to numerical computations. Molecular orbitals are orthogonal, which significantly increases feasibility and speed of computer calculations compared to nonorthogonal valence bond orbitals. However, better valence bond programs are now available.

Covalency from atomic contribution to the electronic density of states

In COOP,[15] COHP[16] and BCOOP,[17] evaluation of bond covalency is dependent on the basis set. To overcome this issue, an alternative formulation of the bond covalency can be provided in this way.

The center mass cm(n,l,ml,ms) of an atomic orbital |n,l,ml,ms, with quantum numbers n, l, ml, ms, for atom A is defined as

where gA
|n,l,ml,ms
(E)
is the contribution of the atomic orbital |n,l,ml,ms of the atom A to the total electronic density of states g(E) of the solid

where the outer sum runs over all atoms A of the unit cell. The energy window [E0,E1] is chosen in such a way that it encompasses all relevant bands participating in the bond. If the range to select is unclear, it can be identified in practice by examining the molecular orbitals that describe the electron density along the considered bond.

The relative position CnAlA,nBlB of the center mass of |nA,lA levels of atom A with respect to the center mass of |nB,lB levels of atom B is given as

where the contributions of the magnetic and spin quantum numbers are summed. According to this definition, the relative position of the A levels with respect to the B levels is

where, for simplicity, we may omit the dependence from the principal quantum number n in the notation referring to CnAlA,nBlB.

In this formalism, the greater the value of CA,B, the higher the overlap of the selected atomic bands, and thus the electron density described by those orbitals gives a more covalent A–B bond. The quantity CA,B is denoted as the covalency of the A–B bond, which is specified in the same units of the energy E.

See also

References

  1. ^ Campbell, Neil A.; Williamson, Brad; Heyden, Robin J. (2006). Biology: Exploring Life. Boston, MA: Pearson Prentice Hall. ISBN 0-13-250882-6. Retrieved 2012-02-05.
  2. ^ March, Jerry (1992). Advanced Organic Chemistry: Reactions, Mechanisms, and Structure. John Wiley & Sons. ISBN 0-471-60180-2.
  3. ^ Gary L. Miessler; Donald Arthur Tarr (2004). Inorganic Chemistry. Prentice Hall. ISBN 0-13-035471-6.
  4. ^ Merriam-Webster – Collegiate Dictionary (2000).
  5. ^ "Chemical Bonds". Hyperphysics.phy-astr.gsu.edu. Retrieved 2013-06-09.
  6. ^ Langmuir, Irving (1919-06-01). "The Arrangement of Electrons in Atoms and Molecules". Journal of the American Chemical Society. 41 (6): 868–934. doi:10.1021/ja02227a002.
  7. ^ Lewis, Gilbert N. (1916-04-01). "The atom and the molecule". Journal of the American Chemical Society. 38 (4): 762–785. doi:10.1021/ja02261a002.
  8. ^ Heitler, W.; London, F. (1927). "Wechselwirkung neutraler Atome und homöopolare Bindung nach der Quantenmechanik" [Interaction of neutral atoms and homeopolar bonds according to quantum mechanics]. Zeitschrift für Physik. 44 (6–7): 455–472. Bibcode:1927ZPhy...44..455H. doi:10.1007/bf01397394. English translation in Hettema, H. (2000). Quantum Chemistry: Classic Scientific Papers. World Scientific. p. 140. ISBN 978-981-02-2771-5. Retrieved 2012-02-05.
  9. ^ Stranks, D. R.; Heffernan, M. L.; Lee Dow, K. C.; McTigue, P. T.; Withers, G. R. A. (1970). Chemistry: A structural view. Carlton, Vic.: Melbourne University Press. p. 184. ISBN 0-522-83988-6.
  10. ^ Weinhold, F.; Landis, C. (2005). Valency and Bonding. Cambridge. pp. 96–100. ISBN 0-521-83128-8.
  11. ^ Harcourt, Richard D., ed. (2015). "Chapter 2: Pauling "3-Electron Bonds", 4-Electron 3-Centre Bonding, and the Need for an "Increased-Valence" Theory". Bonding in Electron-Rich Molecules: Qualitative Valence-Bond Approach via Increased-Valence Structures. Springer. ISBN 9783319166766.
  12. ^ a b Pauling, L. (1960). The Nature of the Chemical Bond. Cornell University Press. pp. 340–354.
  13. ^ Cammarata, Antonio; Rondinelli, James M. (21 September 2014). "Covalent dependence of octahedral rotations in orthorhombic perovskite oxides". Journal of Chemical Physics. 141 (11): 114704. Bibcode:2014JChPh.141k4704C. doi:10.1063/1.4895967. PMID 25240365.
  14. ^ a b c Atkins, P. W. (1974). Quanta: A Handbook of Concepts. Oxford University Press. pp. 147–148. ISBN 978-0-19-855493-6.
  15. ^ Hughbanks, Timothy; Hoffmann, Roald (2002-05-01). "Chains of trans-edge-sharing molybdenum octahedra: metal-metal bonding in extended systems". Journal of the American Chemical Society. 105 (11): 3528–3537. doi:10.1021/ja00349a027.
  16. ^ Dronskowski, Richard; Bloechl, Peter E. (2002-05-01). "Crystal orbital Hamilton populations (COHP): energy-resolved visualization of chemical bonding in solids based on density-functional calculations". The Journal of Physical Chemistry. 97 (33): 8617–8624. doi:10.1021/j100135a014.
  17. ^ Grechnev, Alexei; Ahuja, Rajeev; Eriksson, Olle (2003-01-01). "Balanced crystal orbital overlap population—a tool for analysing chemical bonds in solids". Journal of Physics: Condensed Matter. 15 (45): 7751. Bibcode:2003JPCM...15.7751G. doi:10.1088/0953-8984/15/45/014. ISSN 0953-8984.

Sources

External links

Acid

An acid is a molecule or ion capable of donating a hydron (proton or hydrogen ion H+), or, alternatively, capable of forming a covalent bond with an electron pair (a Lewis acid).The first category of acids is the proton donors or Brønsted acids. In the special case of aqueous solutions, proton donors form the hydronium ion H3O+ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted or Arrhenius acid usually contains a hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H+.

Aqueous Arrhenius acids have characteristic properties which provide a practical description of an acid. Acids form aqueous solutions with a sour taste, can turn blue litmus red, and react with bases and certain metals (like calcium) to form salts. The word acid is derived from the Latin acidus/acēre meaning sour. An aqueous solution of an acid has a pH less than 7 and is colloquially also referred to as 'acid' (as in 'dissolved in acid'), while the strict definition refers only to the solute. A lower pH means a higher acidity, and thus a higher concentration of positive hydrogen ions in the solution. Chemicals or substances having the property of an acid are said to be acidic.

Common aqueous acids include hydrochloric acid (a solution of hydrogen chloride which is found in gastric acid in the stomach and activates digestive enzymes), acetic acid (vinegar is a dilute aqueous solution of this liquid), sulfuric acid (used in car batteries), and citric acid (found in citrus fruits). As these examples show, acids (in the colloquial sense) can be solutions or pure substances, and can be derived from acids (in the strict sense) that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid.

The second category of acids are Lewis acids, which form a covalent bond with an electron pair. An example is boron trifluoride (BF3), whose boron atom has a vacant orbital which can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia (NH3). Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly or by releasing protons (H+) into the solution, which then accept electron pairs. However, hydrogen chloride, acetic acid, and most other Brønsted-Lowry acids cannot form a covalent bond with an electron pair and are therefore not Lewis acids. Conversely, many Lewis acids are not Arrhenius or Brønsted-Lowry acids. In modern terminology, an acid is implicitly a Brønsted acid and not a Lewis acid, since chemists almost always refer to a Lewis acid explicitly as a Lewis acid.

Chemical bond

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. 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.

Chemical polarity

In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end.

Polar molecules must contain polar bonds due to a difference in electronegativity between the bonded atoms. A polar molecule with two or more polar bonds must have a geometry which is asymmetric in at least one direction, so that the bond dipoles do not cancel each other.

Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Polarity underlies a number of physical properties including surface tension, solubility, and melting and boiling points.

Coordinate covalent bond

A coordinate covalent bond, also known as a dative bond or coordinate bond is a kind of 2-center, 2-electron covalent bond in which the two electrons derive from the same atom. The bonding of metal ions to ligands involves this kind of interaction. This type of interaction is central to Lewis acid-base theory.

Covalent bond classification method

The covalent bond classification (CBC) method is also referred to as the LXZ notation. It was published by M. L. H. Green in the mid-1990s as a solution for the need to describe covalent compounds such as organometallic complexes in a way that is not prone to limitations resulting from the definition of oxidation state. Instead of simply assigning a charge to an atom in the molecule (i.e. the oxidation state), the covalent bond classification method analyzes the nature of the ligands surrounding the atom of interest, which is often a transition metal. According to this method, there are three basic types of interactions that allow for coordination of the ligand. The three types of interaction are classified according to whether the ligating group donates two, one, or zero electrons. These three classes of ligands are respectively given the symbols L, X and Z.

Covalent radius

The covalent radius, rcov, is a measure of the size of an atom that forms part of one covalent bond. It is usually measured either in picometres (pm) or angstroms (Å), with 1 Å = 100 pm.

In principle, the sum of the two co equal the covalent bond length between two atoms, R(AB) = r(A) + r(B). Moreover, different radii can be introduced for single, double and triple bonds (r1, r2 and r3 below), in a purely operational sense. These relationships are certainly not exact because the size of an atom is not constant but depends on its chemical environment. For heteroatomic A–B bonds, ionic terms may enter. Often the polar covalent bonds are shorter than would be expected on the basis of the sum of covalent radii. Tabulated values of covalent radii are either average or idealized values, which nevertheless show a certain transferability between different situations, which makes them useful.

The bond lengths R(AB) are measured by X-ray diffraction (more rarely, neutron diffraction on molecular crystals). Rotational spectroscopy can also give extremely accurate values of bond lengths. For homonuclear A–A bonds, Linus Pauling took the covalent radius to be half the single-bond length in the element, e.g. R(H–H, in H2) = 74.14 pm so rcov(H) = 37.07 pm: in practice, it is usual to obtain an average value from a variety of covalent compounds, although the difference is usually small. Sanderson has published a recent set of non-polar covalent radii for the main-group elements, but the availability of large collections of bond lengths, which are more transferable, from the Cambridge Crystallographic Database has rendered covalent radii obsolete in many situations.

Delocalized electron

In chemistry, delocalized electrons are electrons in a molecule, ion or solid metal that are not associated with a single atom or a covalent bond.

The term is general and can have slightly different meanings in different fields. In organic chemistry, this refers to resonance in conjugated systems and aromatic compounds. In solid-state physics, this refers to free electrons that facilitate electrical conduction. In quantum chemistry, this refers to molecular orbital electrons that extend over several adjacent atoms.

Double bond

A double bond in chemistry is a chemical bond between two chemical elements involving four bonding electrons instead of the usual two. The most common double bond, that is between two carbon atoms, can be found in alkenes. Many types of double bonds exist between two different elements. For example, in a carbonyl group with a carbon atom and an oxygen atom. Other common double bonds are found in azo compounds (N=N), imines (C=N) and sulfoxides (S=O). In skeletal formula the double bond is drawn as two parallel lines (=) between the two connected atoms; typographically, the equals sign is used for this. Double bonds were first introduced in chemical notation by Russian chemist Alexander Butlerov.Double bonds involving carbon are stronger than single bonds and are also shorter. The bond order is two. Double bonds are also electron-rich, which makes them potentially more reactive in the presence of a strong electron acceptor (as in addition reactions of the halogens).

Electron donor

An electron donor is a chemical entity that donates electrons to another compound. It is a reducing agent that, by virtue of its donating electrons, is itself oxidized in the process.

Typical reducing agents undergo permanent chemical alteration through covalent or ionic reaction chemistry. This results in the complete and irreversible transfer of one or more electrons. In many chemical circumstances, however, the transfer of electronic charge to an electron acceptor may be only fractional, meaning an electron is not completely transferred, but results in an electron resonance between the donor and acceptor. This leads to the formation of charge transfer complexes in which the components largely retain their chemical identities.

The electron donating power of a donor molecule is measured by its ionization potential which is the energy required to remove an electron from the highest occupied molecular orbital.

The overall energy balance (ΔE), i.e., energy gained or lost, in an electron donor-acceptor transfer is determined by the difference between the acceptor's electron affinity (A) and the ionization potential (I):

In chemistry, the class of electron donors that donate not just one, but a set of two paired electrons that form a covalent bond with an electron acceptor molecule, is known as a Lewis base. This phenomenon gives rise to the wide field of Lewis acid-base chemistry. The driving forces for electron donor and acceptor behavior in chemistry is based on the concepts of electropositivity (for donors) and electronegativity (for acceptors) of atomic or molecular entities.

Electrophilic addition

In organic chemistry, an electrophilic addition reaction is an addition reaction where, in a chemical compound, a π bond is broken and two new σ bonds are formed. The substrate of an electrophilic addition reaction must have a double bond or triple bond.The driving force for this reaction is the formation of an electrophile X+ that forms a covalent bond with an electron-rich unsaturated C=C bond. The positive charge on X is transferred to the carbon-carbon bond, forming a carbocation during the formation of the C-X bond.

In step 2 of an electrophilic addition, the positively charged intermediate combines with (Y) that is electron-rich and usually an anion to form the second covalent bond.

Step 2 is the same nucleophilic attack process found in an SN1 reaction. The exact nature of the electrophile and the nature of the positively charged intermediate are not always clear and depend on reactants and reaction conditions.

In all asymmetric addition reactions to carbon, regioselectivity is important and often determined by Markovnikov's rule. Organoborane compounds give anti-Markovnikov additions. Electrophilic attack to an aromatic system results in electrophilic aromatic substitution rather than an addition reaction.

In Electrophilic addition the electrophile with the positive charge effects the formation of the total structure, which thus bears a positive charge as well, to make up for the new addition, which then results in the intermediate, bearing that positive charge. This intermediate is the key to understanding the electrophilic addition, which is due to the positive nature of the particles involved. If this is done, then the reactions can be understood by these additions as positively charged reactions. The positive charge bears result as the intermediate form otherwise known as the total structure of such an intermediate. The end product thus contains the complete structure, with the addition of Y, a nucleophile.

Heterolysis (chemistry)

In chemistry, heterolysis or heterolytic fission (from Greek ἕτερος, heteros, "different", and λύσις, lusis, "loosening") is the process of cleaving a covalent bond where one previously bonded species takes both original bonding electrons from the other species. During heterolytic bond cleavage of a neutral molecule, a cation and an anion will be generated. Most commonly the more electronegative atom keeps the pair of electrons becoming anionic while the more electropositive atom becomes cationic.

Heterolytic fission almost always happens to single bonds, the process usually produces two fragment species.

The energy required to break the bond is called the heterolytic bond dissociation energy, which is not equivalent to homolytic bond dissociation energy commonly used to represent the energy value of a bond.

One example of the differences in the energies is the energy required to break a Hydrogen-Hydrogen bond

Hydrazide

Hydrazides in organic chemistry are a class of organic compounds sharing a common functional group characterized by a nitrogen to nitrogen covalent bond with 4 substituents with at least one of them being an acyl group. The general structure for a hydrazide is E(=O)-NR-NR2, where the R's are frequently hydrogens. Hydrazides can be further classified by atom attached to the oxygen: carbohydrazides (R-C(=O)-NH-NH2), sulfonohydrazides (R-S(=O)2-NH-NH2) and phosphonic dihydrazides (R-P(=O)(-NH-NH2)2. The related hydrazines do not carry an acyl group. Some important members of this class are sulfonylhydrazides such as p-toluenesulfonylhydrazide which are useful reagents in organic chemistry such as in the Shapiro reaction and the Eschenmoser–Tanabe fragmentation. This reagent can be prepared by reaction of tosyl chloride with hydrazine.

Intramolecular force

An intramolecular force is any force that binds together the atoms making up a molecule or compound, not to be confused with intermolecular forces, which are the forces present between molecules. The subtle difference in the name comes from the Latin roots of English with inter meaning between or among and intra meaning inside. Chemical bonds are considered to be intramolecular forces, for example. These forces are often stronger than intermolecular forces, which are present between atoms or molecules that are not bonded.

Irreversible antagonist

An irreversible antagonist is a type of antagonist that binds permanently to a receptor, either by forming a covalent bond to the active site, or alternatively just by binding so tightly that the rate of dissociation is effectively zero at relevant time scales. This permanently deactivates the receptor and is usually followed by rapid internalisation and recycling of the non-functional receptor protein. Irreversible enzyme inhibitors that act similarly are clinically used and include drugs such as aspirin, omeprazole and monoamine oxidase inhibitors.

Non-covalent interactions

A non-covalent interaction differs from a covalent bond in that it does not involve the sharing of electrons, but rather involves more dispersed variations of electromagnetic interactions between molecules or within a molecule. The chemical energy released in the formation of non-covalent interactions is typically on the order of 1-5 kcal/mol (1000–5000 calories per 6.02 x 10^23 molecules). Non-covalent interactions can be classified into different categories, such as electrostatic, π-effects, van der Waals forces, and hydrophobic effects.Non-covalent interactions are critical in maintaining the three-dimensional structure of large molecules, such as proteins and nucleic acids. In addition, they are also involved in many biological processes in which large molecules bind specifically but transiently to one another (see the properties section of the DNA page). These interactions also heavily influence drug design, crystallinity and design of materials, particularly for self-assembly, and, in general, the synthesis of many organic molecules.Intermolecular forces are non-covalent interactions that occur between different molecules, rather than between different atoms of the same molecule

Single bond

In chemistry, a single bond is a chemical bond between two atoms involving two valence electrons. That is, the atoms share one pair of electrons where the bond forms. Therefore, a single bond is a type of covalent bond. When shared, each of the two electrons involved is no longer in the sole possession of the orbital in which it originated. Rather, both of the two electrons spend time in either of the orbitals which overlap in the bonding process. As a Lewis structure, a single bond is denoted as AːA or A-A, for which A represents an element (Moore, Stanitski, and Jurs 329). In the first rendition, each dot represents a shared electron, and in the second rendition, the bar represents both of the electrons shared in the single bond.

A covalent bond can also be a double bond or a triple bond. A single bond is weaker than either a double bond or a triple bond. This difference in strength can be explained by examining the component bonds of which each of these types of covalent bonds consists (Moore, Stanitski, and Jurs 393).

Usually, a single bond is a sigma bond. An exception is the bond in diboron, which is a pi bond. In contrast, the double bond consists of one sigma bond and one pi bond, and a triple bond consists of one sigma bond and two pi bonds (Moore, Stanitski, and Jurs 396). The number of component bonds is what determines the strength disparity. It stands to reason that the single bond is the weakest of the three because it consists of only a sigma bond, and the double bond or triple bond consist not only of this type of component bond but also at least one additional bond.

The single bond has the capacity for rotation, a property not possessed by the double bond or the triple bond. The structure of pi bonds does not allow for rotation (at least not at 298 K), so the double bond and the triple bond which contain pi bonds are held due to this property. The sigma bond is not so restrictive, and the single bond is able to rotate using the sigma bond as the axis of rotation (Moore, Stanitski, and Jurs 396-397).

Another property comparison can be made in bond length. Single bonds are the longest of the three types of covalent bonds as interatomic attraction is greater in the two other types, double and triple. The increase in component bonds is the reason for this attraction increase as more electrons are shared between the bonded atoms (Moore, Stanitski, and Jurs 343).

Single bonds are often seen in diatomic molecules. Examples of this use of single bonds include H2, F2, and HCl.

Single bonds are also seen in molecules made up of more than two atoms. Examples of this use of single bonds include:

Both bonds in H2O

All 4 bonds in CH4Single bonding even appears in molecules as complex as hydrocarbons larger than methane. The type of covalent bonding in hydrocarbons is extremely important in the nomenclature of these molecules. Hydrocarbons containing only single bonds are referred to as alkanes (Moore, Stanitski, and Jurs 334). The names of specific molecules which belong to this group end with the suffix –ane' Examples include ethane, 2-methylbutane, and cyclopentane (Moore, Stanitski, and Jurs 335).

Substitution reaction

Substitution reaction (also known as single displacement reaction or single substitution reaction) is a chemical reaction during which one functional group in a chemical compound is replaced by another functional group. Substitution reactions are of prime importance in organic chemistry. Substitution reactions in organic chemistry are classified either as electrophilic or nucleophilic depending upon the reagent involved. There are other classifications as well that are mentioned below.

Organic substitution reactions are classified in several main organic reaction types depending on whether the reagent that brings about the substitution is considered an electrophile or a nucleophile, whether a reactive intermediate involved in the reaction is a carbocation, a carbanion or a free radical or whether the substrate is aliphatic or aromatic. Detailed understanding of a reaction type helps to predict the product outcome in a reaction. It also is helpful for optimizing a reaction with regard to variables such as temperature and choice of solvent.

A good example of a substitution reaction is halogenation. When chlorine gas (Cl-Cl) is irradiated, some of the molecules are split into two chlorine radicals (Cl.) whose free electrons are strongly nucleophilic. One of them breaks a weak C-H covalent bond and grabs the liberated proton to form the electrically neutral H-Cl. The other radical reforms a covalent bond with the CH3. to form CH3Cl (methyl chloride).

Triple bond

A triple bond in chemistry is a chemical bond between two atoms involving six bonding electrons instead of the usual two in a covalent single bond. The most common triple bond, that between two carbon atoms, can be found in alkynes.

Other functional groups containing a triple bond are cyanides and isocyanides. Some diatomic molecules, such as dinitrogen and carbon monoxide, are also triple bonded. In skeletal formula the triple bond is drawn as three parallel lines (≡) between the two connected atoms.Triple bonds are stronger than the equivalent single bonds or double bonds, with a bond order of three.

Valence electron

In chemistry, a valence electron is an outer shell electron that is associated with an atom, and that can participate in the formation of a chemical bond if the outer shell is not closed; in a single covalent bond, both atoms in the bond contribute one valence electron in order to form a shared pair. The presence of valence electrons can determine the element's chemical properties, such as its valence—whether it may bond with other elements and, if so, how readily and with how many. For a main group element, a valence electron can exist only in the outermost electron shell; in a transition metal, a valence electron can also be in an inner shell.

An atom with a closed shell of valence electrons (corresponding to an electron configuration s2p6) tends to be chemically inert. Atoms with one or two more valence electrons than are needed for a "closed" shell are highly reactive due to the following reasons:

1) It requires relatively low energy (compared to the lattice enthalpy) to remove the extra valence electrons to form a positive ion.

2) Because of their tendency either to gain the missing valence electrons (thereby forming a negative ion), or to share valence electrons (thereby forming a covalent bond).

Similar to an electron in an inner shell, a valence electron has the ability to absorb or release energy in the form of a photon. An energy gain can trigger an electron to move (jump) to an outer shell; this is known as atomic excitation. Or the electron can even break free from its associated atom's valence shell; this is ionization to form a positive ion. When an electron loses energy (thereby causing a photon to be emitted), then it can move to an inner shell which is not fully occupied.

Valence energy levels correspond to the principal quantum numbers (n = 1, 2, 3, 4, 5 ...) or are labeled alphabetically with letters used in the X-ray notation (K, L, M, …).

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