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Electronegativity

Electronegativity, symbol χ (the Greek letter chi), is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself and thus the tendency to form negative ions. An atom's electronegativity is affected by both its atomic number and the distance that its valence electrons reside from the charged nucleus. The higher the associated electronegativity number, the more an element or compound attracts electrons towards it. First proposed by Linus Pauling in 1932 as a development of valence bond theory, it has been shown to correlate with a number of other chemical properties. Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties. Several methods of calculation have been proposed and, although there may be small differences in the numerical values of the electronegativity, all methods show the same periodic trends between elements.

The most commonly used method of calculation is that originally proposed by Pauling. This gives a dimensionless quantity, commonly referred to as the Pauling scale, on a relative scale running from around 0.7 to 3.98 (hydrogen = 2.20). When other methods of calculation are used, it is conventional (although not obligatory) to quote the results on a scale that covers the same range of numerical values: this is known as an electronegativity in Pauling units.

Electronegativity, as it is usually calculated, is not strictly an atomic property, but rather a property of an atom in a molecule: the equivalent property of a free atom is its electron affinity. It is to be expected that the electronegativity of an element will vary with its chemical environment, but it is usually considered to be a transferable property, that is to say that similar values will be valid in a variety of situations.
The opposite of electronegativity is electropositivity: a measure of an element's ability to donate electrons.

Electronegativities of the elements

Periodic table of electronegativity using the Pauling scale

See also Electronegativities of the elements (data page) and List of electronegativities

Methods of calculation

Pauling electronegativity

Pauling first proposed the concept of electronegativity in 1932 as an explanation of the fact that the covalent bond between two different atoms (A–B) is stronger than would be expected by taking the average of the strengths of the A–A and B–B bonds. According to valence bond theory, of which Pauling was a notable proponent, this "additional stabilization" of the heteronuclear bond is due to the contribution of ionic canonical forms to the bonding.

The difference in electronegativity between atoms A and B is given by:

\chi_{\rm A} - \chi_{\rm B} = ({\rm eV})^{-1/2} \sqrt{E_{\rm d}({\rm AB}) - [E_{\rm d}({\rm AA}) + E_{\rm d}({\rm BB})]/2}

where the dissociation energies, Ed, of the A–B, A–A and B–B bonds are expressed in electronvolts, the factor (eV)–½ being included to ensure a dimensionless result. Hence, the difference in Pauling electronegativity between hydrogen and bromine is 0.73 (dissociation energies: H–Br, 3.79 eV; H–H, 4.52 eV; Br–Br 2.00 eV)

As only differences in electronegativity are defined, it is necessary to choose an arbitrary reference point in order to construct a scale. Hydrogen was chosen as the reference, as it forms covalent bonds with a large variety of elements: its electronegativity was fixed first at 2.1, later revised to 2.20. It is also necessary to decide which of the two elements is the more electronegative (equivalent to choosing one of the two possible signs for the square root). This is done by "chemical intuition": in the above example, hydrogen bromide dissolves in water to form H+ and Br– ions, so it may be assumed that bromine is more electronegative than hydrogen.

To calculate Pauling electronegativity for an element, it is necessary to have data on the dissociation energies of at least two types of covalent bond formed by that element. Allred updated Pauling's original values in 1961 to take account of the greater availability of thermodynamic data, and it is these "revised Pauling" values of the electronegativity which are most usually used.

Mulliken electronegativity

Mulliken proposed that the arithmetic mean of the first ionization energy and the electron affinity should be a measure of the tendency of an atom to attract electrons. As this definition is not dependent on an arbitrary relative scale, it has also been termed absolute electronegativity, with the units of kilojoules per mole or electronvolts.

However, it is more usual to use a linear transformation to transform these absolute values into values which resemble the more familiar Pauling values. For ionization energies and electron affinities in electronvolts,

\chi = 0.187(E_{\rm i} + E_{\rm ea}) + 0.17 \,

and for energies in kilojoules per mole,

\chi = (1.97\times 10^{-3})(E_{\rm i} + E_{\rm ea}) + 0.19.

The Mulliken electronegativit

Electron acceptor

An electron acceptor is a chemical entity that accepts electrons transferred to it from another compound. It is an oxidizing agent that, by virtue of its accepting electrons, is itself reduced in the process.

Typical oxidizing agents undergo permanent chemical alteration through covalent or ionic reaction chemistry, resulting in the complete and irreversible transfer of one or more electrons. In many chemical circumstances, however, the transfer of electronic charge from an electron donor 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 accepting power of an acceptor molecule is measured by its electron affinity which is the energy released when filling the lowest unoccupied molecular orbital (LUMO).

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) of the electron donor:

{\Delta}E=A-I\,.

In chemistry, a class of electron acceptors that acquire not just one, but a set of two paired electrons that form a covalent bond with an electron donor molecule, is known as a Lewis acid. 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.

Examples

Examples of electron acceptors include oxygen, nitrate, iron (III), manganese (IV), sulfate, carbon dioxide, or in some microorganisms the chlorinated solvents such as tetrachloroethylene (PCE), trichloroethylene (TCE), dichloroethene (DCE), and vinyl chloride (VC). These reactions are of interest not only because they allow organisms to obtain energy, but also because they are involved in the natural biodegradation of organic contaminants. When clean-up professionals use monitored natural attenuation to clean up contaminated sites, biodegradation is one of the major contributing processes.

In biology, a terminal electron acceptor is a compound that receives or accepts an electron during cellular respiration or photosynthesis. All organisms obtain energy by transferring electrons from an electron donor to an electron acceptor. During this process (electron transport chain) the electron acceptor is reduced and the electron donor is oxidized.


Carbene

In chemistry, a carbene is a molecule containing a neutral carbon atom with a valence of two and two unshared valence electrons. The general formula is RR'C:, but the carbon can instead be double-bonded to one group. The term "carbene" may also merely refer to the compound H2C:, also called methylene, the parent hydride to which all other carbene compounds are related. Carbenes are classified as either singlets or triplets depending upon their electronic structure. Most carbenes are very short lived, although persistent carbenes are known.

One well studied carbene is Cl2C:, or dichlorocarbene, which can be generated in situfromchloroform and a strong base.

Structure and bonding

The two classes of carbenes are singlet and triplet carbenes. Singlet carbenes are spin-paired. In the language of valence bond theory, the molecule adopts an sp2hybrid structure. Triplet carbenes have two unpaired electrons. They may be either linear or bent, i.e. sp or sp2 hybridized, respectively. Most carbenes have a nonlinear triplet ground state, except for those with nitrogen, oxygen, or sulfur atoms, and halides directly bonded to the divalent carbon.

Carbenes are called singlet or triplet depending on the electronic spins they possess. Triplet carbenes are paramagnetic and may be observed by electron spin resonance spectroscopy if they persist long enough. The total spin of singlet carbenes is zero while that of triplet carbenes is one (in units of \hbar). Bond angles are 125-140° for triplet methylene and 102° for singlet methylene (as determined by EPR). Triplet carbenes are generally stable in the gaseous state, while singlet carbenes occur more often in aqueous media.

For simple hydrocarbons, triplet carbenes usually have energies 8 kcal/mol (33 kJ/mol) lower than singlet carbenes (see also Hund's rule of Maximum Multiplicity), thus, in general, triplet is the more stable state (the ground state) and singlet is the excited state species. Substituents that can donate electron pairs may stabilize the singlet state by delocalizing the pair into an empty p-orbital. If the energy of the singlet state is sufficiently reduced it will actually become the ground state. No viable strategies exist for triplet stabilization. The carbene called 9-fluorenylidene has been shown to be a rapidly equilibrating mixture of singlet and triplet states with an approximately 1.1 kcal/mol (4.6 kJ/mol) energy difference. It is however debatable whether diaryl carbenes such as the fluorene carbene are true carbenes because the electrons can delocalize to such an extent that they become in fact biradicals. In silico experiments suggest that triplet carbenes can be stabilized with electropositive groups such as trifluorosilyl groups.

Reactivity

Singlet and triplet carbenes exhibit divergent reactivity. Singlet carbenes generally participate in cheletropic reactions as either electrophiles or nucleophiles. Singlet carbenes with unfilled p-orbital should be electrophilic. Triplet carbenes can be considered to be diradicals, and participate in stepwise radical additions. Triplet carbenes have to go through an intermediate with two unpaired electrons whereas singlet carbene can react in a single concerted step.

Due to these two modes of reactivity, reactions of singlet methylene are stereospecific whereas those of triplet methylene are stereoselective. This difference can be used to probe the nature of a carbene. For example, the reaction of methylene generated from photolysis of diazomethane with cis-2-butene or with trans-2-butene each give a single diastereomer of the 1,2-dimethylcyclopropane product: cis from cis and trans from trans, which proves that the methylene is a singlet. If the methylene were a triplet, one would not expect the product to depend upon the starting alkene geometry, but rather a nearly identical mixture in each case.

Reactivity of a particular carbene depends on the substituent groups. Their reactivity can be affected by metals. Some of the reactions carbenes can do are insertions into C-H bonds, skeletal rearrangements, and additions to double bonds. Carbenes can be classified as nucleophilic, electrophilic, or ambiphilic. For example, if a

Ionic radius

Ionic radius, rion, is a measure of the size of an atom's ion in a crystal lattice. It is measured in either picometres (pm) or Angstrom (Ã…), with 1 Ã… = 100 pm. Typical values range from 30 pm (0.3 Ã…) to over 200 pm (2 Ã…).

The concept of ionic radius was developed independently by Victor Goldschmidt and Linus Pauling in the 1920s to summarize the data being generated by the (at the time) new technique of X-ray crystallography: it is Pauling's approach which proved to be the more influential. X-ray crystallography can readily give the length of the side of the unit cell of a crystal, but it is much more difficult (in most cases impossible, even with more modern techniques) to distinguish a boundary between two ions. For example, it can be readily determined that each side of the unit cell of sodium chloride is 564.02 pm in length, and that this length is twice the distance between the centre of a sodium ion and the centre of a chloride ion:

2[rion(Na+) + rion(Cl−)] = 564.02 pm

However, it is not apparent what proportion of this distance is due to the size of the sodium ion and what proportion is due to the size of the chloride ion. By comparing many different compounds, and with a certain amount of chemical intuition, Pauling decided to assign a radius of 140 pm to the oxide ion O2−, at which point he was able to calculate the radii of the other ions by subtraction.

A major review of crystallographic data led to the publication of a revised set of ionic radii in 1976, and these are preferred to Pauling's original values. Some sources have retained Pauling's reference of rion(O2−) = 140 pm, while other sources prefer to list "effective" ionic radii based on rion(O2−) = 126 pm. The latter values are thought to be a more accurate approximation to the "true" relative sizes of anions and cations in ionic crystals.

The ionic radius is not a fixed property of a given ion, but varies with coordination number, spin state and other parameters. Nevertheless, ionic radius values are sufficiently transferable to allow periodic trends to be recognized. As with other types of atomic radius, ionic radii increase on descending a group. Ionic size (for the same ion) also increases with increasing coordination number, and an ion in a high-spin state will be larger than the same ion in a low-spin state. Anions (negatively charged) are almost invariably larger than cations (positively charged), although the fluorides of some alkali metals are rare exceptions. In general, ionic radius decreases with increasing positive charge and increases with increasing negative charge.

An "anomalous" ionic radius in a crystal is often a sign of significant covalent character in the bonding. No bond is completely ionic, and some supposedly "ionic" compounds, especially of the transition metals, are particularly covalent in character. This is illustrated by the unit cell parameters for sodium and silverhalides in the table. On the basis of the fluorides, one would say that Ag+ is larger than Na+, but on the basis of the chlorides and bromides the opposite appears to be true. This is because the greater covalent character of the bonds in AgCl and AgBr reduces the bond length and hence the apparent ionic radius of Ag+, an effect which is not present in the halides of the more electropositive sodium, nor in silver fluoride in which the fluoride ion is relatively unpolarizable.

Generalization

The concept of ionic radii is based on the assumption of a spherical ion shape. However, from a group-theoretical point of view the assumption is only justified for ions that reside on high-symmetry crystal lattice sites like Na and Cl in halite or Zn and S in sphalerite. A clear distinction can be made, when the point symmetry group of the respective lattice site is considered, which are the cubic groups O6 and Td in NaCl and ZnS. For ions on lower-symmetry sites significant deviations of their electron density from a spherical shape may occur. This holds in particular for ions on lattice sites of polar symmetry, which are the crystallographic point groups C1, C1h, Cn or Cnv, n = 2, 3, 4 or 6. A thorough analysis of the bonding geometry was recently carried out for pyrite-type disulfides, where monovalent From Yahoo Answers

Question:Why is the bond between N and H in amines a polar bond? A) Hydrogen is more electronegative. B) Nitrogen is more electropositive. C) Nitrogen is more electronegative. D) Hydrogen is electrophilic. E) Nitrogen is electrophilic.

Answers:The answer is --- C) Nitrogen is more electronegative. because the electronegativity of N is 3.0 and H is 2.1 So nitrogen will have more influence on the electrons. N will have partial negative charge and H will have partial positive charge

Question:hey guyz please correct my concept as ceasium is most reactive metal in alkali metals due to its greatest electropositivity then why lithium is strong reducing agent(according to electrochemical series)?i mean if lithium cannot lose electron easily how it can reduce nonmetals strongly? plzz do help

Answers:I think lithium is greater electronegative. In one group, if the atom in periodic is upper than, it would have high electronegative. It's high electromagnetism is caused the electron is difficult to lose itself because the nucleon has pulled it. Reducing is reaction to catch electron. So easy for lithium to do it because it has high electronegative. Besides, it has small radius.

Question:A natural attraction (8 letters) Discoverer of the noble gases (6 letters Refers to elements of the sixth period ( 4 letters) Elements that are poor conductors of heat and electricity and are electronegative 27. To be in the form of gas (7 letters) 34. Rare earths of the seventh period ( 8 letters) hint: 6th letter of 34 is "i" 37. He used X rays to determine atomic numbers ( 6 letters) Hint for the discoverer of nobel gases...begings with and "r" 12. Those elements characterized by the belated filling of the next-to-the-outermost energy levels ( 10) 33. A series of rare earths of the sixth peroid. 37. A class of elements that show luster, conduct heat and electric current, and are electropositive 38. Elements that have certain metal properties but are classed as nonmetals. What is the chemical symbol for the most abundent element on the face of the earth

Answers:A natural attraction AFFINITY Discoverer of the noble gases (Sir William) RAMSAY Refers to elements of the sixth period RARE EARTHS (RARE??) Elements that are poor conductors of heat and electricity and are electronegative NON METALS http://www.tutorvista.com/content/science/science-ii/metals-non-metals/physical-properties-metals-non-metals.php 27. To be in the form of gas GASEOUS 34. Rare earths of the seventh period ACTINIDE 37. He used X rays to determine atomic numbers . I thought it was Henry MOSELEY (but that is 7 letters!) 33. A series of rare earths of the sixth peroid LANTHANOID. 37. A class of elements that show luster, conduct heat and electric current, and are electropositive METALS ? 38. Elements that have certain metal properties but are classed as nonmetals METALLOIDS. What is the chemical symbol for the most abundent element on the face of the earth. O for Oxygen (Silicon is second and Aluminium is third)

Question:Using what you know about the electronegativity of the main group elements, which of these would be a good oxidizing agent? A. H2 B. Cl2 C. C D. Cs Please Help!

Answers:B. Cl2 Remember that an oxidizing agent is reduced, it GAINS electrons. Naturally, an element with a HIGH electronegativity, a high tendency for an atom to attract a bonding pair of electrons, would GAIN electrons. Looking at the periodic table, Cl is right next to F, the most electronegative element! Cl has a high electronegativity. Cesium, on the other hand is the most electropositive element, so it would be a reducing agent. Cesium is really v cool. :)