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Electropositive Radicals with their Valency

We know that elements and compounds are pure substances. An element cannot be decomposed into simpler substances. They are listed in the periodic table. Each element can represent by elemental symbols such as Ca for calcium, N for nitrogen, K for potassium etc. Similarly a compound is considered to be a pure substance which can be further broken into simpler substances as it is composed of two or more elements.

Elements cannot be decomposed into any further simpler substances.  In the elemental symbol, the 2nd letter in the symbol is a lower case letter such as ‘He’ for helium, ‘Ca’ for calcium, ‘Ne’ for neon. They can be in solid, liquid or gaseous state such as mercury (Hg), bromine (Br) and hydrogen (H). The charge on element forms electropositive or electronegative radicals. Radicals can be an atom or group of atom with some charge. A simple radical is composed of one atom while a compound radical is formed by the groups of atom. On the basis of charge on radical, they can be classified as electropositive and electronegative radicals.

An electropositive radical has positive charge while an electronegative radical has negative charge on it. The charge or valency of the radicals represents the combining capacity of the radicals. For example the valency of hydrogen atoms is one, it means that it can combine or displace one atom of the element and form a compound. Some common examples of it are hydrogen chloride [HCl], nitric acid [HNO3] and hydrofluoric acid [HF]. In sulphuric acid molecule, the valency of the sulphate radical is 2. Overall valency can be defined as the number of electrons which can donates or accepts by an atom to get the duplet state or octet state in its valence shell.

Valency of a radical is always a whole number. On the basis of valency, elements or radicals can be classified as monovalent (valency=1), divalent (valency=2), trivalent (valency=3) and so on. All metals form electropositive radicals while all non-metals form electronegative radicals

The charge on the radicals is due to lose or gain of electrons to get the stable valence shell configuration. The radicals with opposite charges attract each other to form electrovalent compounds which are also called as ionic compounds. The cation or electropositive radicals form ionic bond with electronegative radicals to form ionic compounds.

Since metals have tendency to lose electrons therefore they can easily form electropositive radicals such as Na+, Fe3+, Mn7+ etc. The charge on electropositive radicals depends on the valence shell configuration of elements. For example; alkali metals have one electron in their valence shell therefore they form 1+ ions while alkaline earth metals have 2 electrons in their valence shell and form M2+ radicals. The d-block elements can show variable valency due to incomplete d-orbitals.

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From Wikipedia


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


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.


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

Unpaired electron - Wikipedia, the free encyclopedia

Radicals are uncommon in s- and p-block chemistry, since the unpaired electron occupies a valence p orbital or an sp, sp2 or sp3 hybrid orbital. ...

Nitrogen group

The nitrogen group is a periodic table group consisting of nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), bismuth (Bi) and ununpentium (Uup) (unconfirmed).

In modern IUPAC notation, it is called Group 15. In the old IUPAC and CAS systems, it was called Group VB and Group VA, respectively (pronounced "group five B" and "group five A", because "V" is a Roman numeral). In the field of semiconductor physics, it is still universally called Group V. It is also collectively named the pnictogens. The "five" ("V") in the historical names comes from the fact that these elements have five valence electrons (see below).

Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:

This group has the defining characteristic that all the component elements have 5 electrons in their outermost shell, that is 2 electrons in the s subshell and 3 unpaired electrons in the p subshell. They are therefore 3 electrons short of filling their outermost electron shell in their non-ionized state. The most important element of this group is nitrogen(chemical symbolN), which in its diatomic form is the principal component of air.

Binary compounds of the group can be referred to collectively as pnictides. The spelling derives from the Greekπνίγειν (pnigein), to choke or stifle, which is a property of nitrogen; they are also mnemonic for the two most common members, P and N. The name pentels (from the Latin penta, five) was also used for this group at one time, stemming from the earlier group naming convention (Group VB).

These elements are also noted for their stability in compounds due to their tendency for forming double and triple covalent bonds. This is the property of these elements which leads to their potential toxicity, most evident in phosphorus, arsenic and antimony. When these substances react with various chemicals of the body, they create strong free radicals not easily processed by the liver, where they accumulate. Paradoxically it is this strong bonding which causes nitrogen and bismuth's reduced toxicity (when in molecules), as these form strong bonds with other atoms which are difficult to split, creating very unreactive molecules. For example N2, the diatomic form of nitrogen, is used for inert atmosphere in situations where argon or another noble gas would be prohibitively expensive.

The nitrogen group consists of two non-metals, two metalloids, one metal, and one synthetic (presumably metallic) element. All the elements in the group are a solid at room temperature except for nitrogen which is a gas at room temperature. Nitrogen and bismuth, despite both being part of the nitrogen group, are very different in their physical properties. For example, at STP nitrogen is a transparent nonmetallic gas, while bismuth is a brittle pinkish metallic solid.

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Answers:Here are couple of answers from internet but any good Chemistry text book will have all this information. Look up one from a library if you do not have one: http://home.att.net/~cat6a/class_elem-VI.htm http://www.syvum.com/cgi/online/serve.cgi/squizzes/chem/valency2a.html

Question:and i dont want it in the periodic table form i want it as a list of elements and radicals with their valencies alongwith their atomic no. , mass no., electronic configuration, and of course, their symbols.

Answers:here its inter active or color coded and expanded wow for all that i need a local

Question:practical inorganic chemistry

Answers:Acid radical: Acid radical is an anion left after removal of hydrogen atoms from an acid. Anions : Anions are atoms or groups of atoms that have gained electrons. Having more negatively charged electrons than positively charged protons, they are negatively charged. The atoms that form ions most easily are the Group 17 (or VII) atoms, also called the halides: F, Cl, Br and I. All these form anions with a -1 charge. O, S, N and P also form anions, carrying charges of -2 (oxygen and sulfur) or -3 (N and P). Most anions are composed from multiple atoms, and are called polyatomic ions (polyatomic = many atoms). Polyatomic ions are usually built around a core atom which is more often than not a non-metal, but some metals, notably manganese and chromium, form polyatomic ions as well. In most polyatomic ions, these atoms combine with oxygen and sometimes with hydrogen as well. As with every other generalisation, there are exceptions. For example, SCN-, the thiocyanate ion, is polyatomic, but has neither oxygen nor hydrogen. (Worse, NH4+ is polyatomic, but is a cation!) But back to the story. The negative charge (the extra electron) in the polyatomic ion is shared around the entire ion. It is not associated with a particular nucleus in the ion, specifically not with the nucleus to which oxygen and/or hydrogen are attached. This is true whether the charge is single or multiple. COMMON NEGATIVE IONS (ANIONS) acetate CH3COO- nitride N3- bromide Br- nitrite NO2- carbonate CO32- oxalate C2O42- hydrogen carbonate HCO3- oxide O2- chlorate ClO3- permanganate MnO4- perchlorate ClO4- phosphide P3- chloride Cl- phosphate PO43- chlorite ClO2- (mono)hydrogen phosphate HPO42- hypochlorite ClO- dihydrogen phosphate H2PO4- chromate CrO42- sulphate SO42- dichromate Cr2O72- hydrogen sulphate HSO4- cyanide CN- sulphide S2- fluoride F- hydrogen sulphide HS- hydride H- sulphite SO32- hydroxide OH- hydrogen sulphite HSO3- iodide I- thiocyanate SCN- nitrate NO3- thiosulphate S2O32- Note that, with the exceptions of hydroxide (OH) and cyanide (CN-), all the names ending in -ide are monatomic. The rest are -ates or -ites. In the days of yore, chemists gathered samples of anions and then tried to give them some order. The rules they came up with went like this: The most frequently occurring version of a polyatomic ion got the name -ate. The most frequently occurring anion of chlorine and oxygen is ClO3-. It was given the name chlorate. One more oxygen? Put a per- on the front. ClO4- is perchlorate. (Per is from the Greek hyper for too much.) One less oxygen? Change the name to -ite. ClO2- is chlorite. Two less oxygens? Put a hypo- on the front. ClO- is hypochlorite. (Hypo is from the Greek for too little or not enough.) So from above discussion we see that all acid radicals are anions but all anions are not acid radicals. Basic radical : The basic radical is the cation left after removal of OH or other alkaline group from the bases. Cations : Cations are atoms that have lost an electron to become positively charged. Sodium has one valence electron, one electron in its outer energy level, so tends to lose one electron, and to become an ion with a +1 charge. The same could be observed for lithium, potassium, rubidium, caesium and francium. Magnesium, along with the other elements in group 2 of the periodic table, has two valence electrons, so tends to become an ion with a +2 charge. Aluminium tends to become +3. What about the transition metals like lead, copper, tin and manganese? Electrons in the transition elements are packed in a way that, generally, places the additional electrons inside the outer energy level. Iron has six more electrons than calcium, but the additional electrons have less energy than the two on the outer. The transition elements tend to have either one or two loosely held (valence) electrons. The six electrons present in iron but absent in calcium are held much less loosely than those in the next level down, but more tightly than those in the outer level. Having removed two electrons from both iron and calcium, removing a third electron would be much, much, much harder from calcium than from iron, since calcium's next electron is both more tightly held (a lower energy level) and in a complete shell. Because of this ambiguity in the transition elements, it is sometimes hard to predict the charge for their corresponding cation. Copper, for example, frequently loses two electrons (Cu2+), but copper ions with a +1 charge (Cu+) are also well known. Tin most often looses two (Sn2+), but frequently loses four (Sn4+). Iron can lose two or three. Manganese? So many choices you don't want to know. For our purposes you will be able to tell which ion you are dealing with by the context it occurs in. You may be given the name of the ion. Sn2+ is named tin(II) [pronounced "tin-two"]. Pb4+ is named lead(IV). You may be given a formula including the ion. FeCl2 is called iron(II) chloride [iron two chloride], containing Fe2+. (Each of the chlorides have accepted one electron, similar to other Group VII elements.) FeCl3 is called iron(III) chloride, containing Fe3+. Thus all basic radicals are cations but all cations are not basic radicals.

Question:I need the info for the Alkaline Earth Metals. Atomic Radius Trend Electronegativity Trend Oxidation State Valence Electrons Reactivity What it bonds easily with Ionic/Covalent more likely? Thank U!!!!

Answers:They have 2 electrons in their outer shell and not surprisingly they tend to lose them so their oxidation number is 2+. Their electron configuration is that of a noble gas followed by a full s subshell ie with two electrons eg calcium is {Argon} 4s2, Barium is {Xenon} 6s2 Going down the group their atomic radii increase while their electronegativity decreases as the outer electrons become more distant from the nucleus and are held less strongly by it. But all of them are quite electropositive which means they have a tendency to give electrons and are therefore good reducing agents. They react with non metals which are much more very electronegative and as a result ionic compounds are formed with few exceptions.

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Valence Electrons - Part 2.avi :A Response to this Vlog: "Free Radicals" - www.youtube.com Part 1 didn't Record.

Huckel's rule: Aromatic vs. antiaromatic (12) :Organic chemistry: How to use Huckel's Rule to determine whether a molecule is aromatic, antiaromatic, or nonaromatic. These videos are offered on a "pay what you like" basis. You can pay for the use of the videos at my website: www.freelance-teacher.com I offer tutoring via Skype. For more information, go to my website. These videos are designed to help students who are finding the material difficult, so I go very slowly, with lots of repetition and examples. If you don't find this material difficult, you might be very bored by these videos and might prefer to learn straight from a textbook. Here is a playlist containing all the videos in this series: www.youtube.com (1) The rule for determining hybridization (2) The rule for determining hybridization, concluded (3) The exception to the rule for determining hybridization (4) What are the valence orbitals of hybridized atoms? (5) "Flat", "cyclic", "completely conjugated" (6) Counting pi electrons (7) Counting pi electrons--harder problems (8) Counting pi electrons--more problems (9) Counting pi electrons--more problems (10) Counting pi electrons--more problems (11) Counting pi electrons--even more problems (12) Counting pi electrons--rings with substituents (13) Counting pi electrons--polycyclics (14) Counting pi electrons--radicals; and a puzzle (15) Counting pi electrons--triple bonds (16) The lesson from triple bonds (17) A deeper look: why do the rules work? (18) A deeper look, continued (19) A deeper look: Why is ...