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From Wikipedia


Phosphine (IUPAC name: phosphane) is the compound with the chemical formula PH3. It is a colorless, flammable, toxic gas. Pure phosphine is odourless, but technical grade samples have a highly unpleasant odor like garlic or rotting fish, due to the presence of substituted phosphine and diphosphine (P2H4). Phosphines are also a group of organophosphorus compounds with the formula R3P (R = organic derivative). Organophosphines are important in catalysts where they complex to various metal ions; complexes derived from a chiral phosphine can catalyze reactions to give chiral products.


Perhaps because of its strong association with elemental phosphorus, phosphine was once regarded as a gaseous form of the element but Lavoisier (1789) recognised it as a combination of phosphorus with hydrogen by describing it as “hydruyet of phosphorus, or phosphuret of hydrogen�.

Thénard (1845) used a cold trap to separate diphosphine from phosphine that had been generated from calcium phosphide, thereby demonstrating that P2H4 is responsible for spontaneous flammability associated with PH3, and also for the characteristic orange/brown colour that can form on surfaces, which is a polymerisation product. He considered diphosphine’s formula to be PH2, and thus an intermediate between elemental phosphorus, the higher polymers, and phosphine. Calcium phosphide (nominally Ca3P2) produces more P2H4 than other phosphides because of the preponderance of P-P bonds in the starting material.

Structure and properties

PH3 is a trigonal pyramidal molecule with C3vmolecular symmetry. The length of the P-H bond 1.42 Å, the H-P-H bond angles are 107°. The dipole moment is 0.58 D, which increases with substitution of methyl groups in the series: CH3PH2, 1.10 D; (CH3)2PH, 1.23 D; (CH3)3P, 1.19 D. In contrast, the dipole moments of amines decrease with substitution, starting with ammonia, which has a dipole moment of 1.47 D. The low dipole moment and almost orthogonal bond angles lead to the conclusion that in PH3 the P-H bonds are almost entirely pσ(P) – sσ(H) and the lone pair contributes only a little to the molecular orbitals. The high positive chemical shift of the P atom in31P NMR spectrum accords with the conclusion that the lone pair electrons occupy the 3s orbital and so are close to the P atom (Fluck, 1973). This electronic structure leads to a lack of nucleophilicity and an inability to form hydrogen bonds.

The aqueous solubility of PH3 is slight; 0.22 mL of gas dissolve in 1 mL of water. Phosphine dissolves more readily in non-polar solvents than in water because of the non-polar P-H bonds. It acts as neither an acid nor a base in water. Proton exchange proceeds via a phosphonium (PH4+) ion in acidic solutions and via PH2− at high pH, with equilibrium constants Kb = 4 × 10−28 and Kz = 41.6 × 10−29.

Preparation and occurrence

Phosphine may be prepared in a variety of ways. Industrially it can be made by the reaction of white phosphorus with sodium hydroxide, producing sodium hypophosphite and sodium phosphite as a by-product. Alternatively the acid-catalyzed disproportioning of white phosphorus may be used, which yields phosphoric acid and phosphine. Both routes have industrial significance; the acid route is preferred method if further reaction of the phosphine to substituted phosphines is needed. The acid route requires purification and pressurizing. It can also be made (as described above) by the hydrolysis of a metal phosphide such as aluminium phosphide or calcium phosphide. Pure samples of phosphine, free from P2H4, may be prepared using the action of potassium hydroxide on phosphonium iodide (PH4I).

Phosphine is probably a constituent of the atmosphere at very low and highly variable concentrations and hence may contribute to the global phosphorus biochemical cycle. The origin(s) of atmospheric phosphine is not certain. Possible sources include bacterial reduction of phosphate in decaying organic matter and the corrosion of phosphorus-containing metals.


Related to a PH3 is the class of organophosphorus compounds commonly called "phosphines." These alkyl and aryl derivatives of phosphine are analogous to organic amines. Common examples include triphenylphosphine ((C6H5)3P) and BINAP, both used as molecule or atom is the energy change when an electron is added to the neutral atom to form a negative ion. This property can only be measured in an atom in gaseous state.

X + e−→ X−

The electron affinity, Eea, is defined as positive when the resulting ion has a lower energy, i.e. it is an exothermic process that releases energy:

Eea = Einitial âˆ’ Efinal

Alternately, electron affinity is often described as the amount of energy required to detach an electron from a singly chargednegative ion, i.e. the energy change for the process

X−→ X + e−

A molecule or atom that has a positive electron affinity is often called an electron acceptor and may undergo charge-transfer reactions.

Electron affinities of the elements

Although Eea varies greatly across the periodic table, some patterns emerge. Generally, nonmetals have more positive Eea than metals. Atoms whose anions are more stable than neutral atoms have a greater Eea. Chlorine most strongly attracts extra electrons; mercury most weakly attracts an extra electron. The electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative values.

Eea generally increases across a period (row) in the periodic table. This is caused by the filling of the valence shell of the atom; a group 7A atom releases more energy than a group 1A atom on gaining an electron because it obtains a filled valence shell and therefore is more stable.

A trend of decreasing Eea going down the groups in the periodic table would be expected. The additional electron will be entering an orbital farther away from the nucleus, and thus would experience a lesser effective nuclear charge. However, a clear counterexample to this trend can be found in group 2A, and this trend only applies to group 1A atoms. Electron affinity follows the trend of electronegativity. Fluorine (F) has a higher electron affinity than oxygen and so on.

The following data are quoted in kJ/mol. Elements marked with an asterisk are expected to have electron affinities close to zero on quantum mechanical grounds. Elements marked with a dotted box are synthetically made elements—elements not found naturally in the environment.

Molecular electron affinities

The electron affinity of molecules is a complicated function of their electronic structure. For instance the electron affinity for benzene is negative, as is that of naphthalene, while those of anthracene, phenanthrene and pyrene are positive. In silicoexperiments show that the electron affinity ofhexacyanobenzene surpasses that of fullerene.

Electron affinity of Surfaces

The electron affinity measured from a material's surface is a function of the bulk material as well as the surface condition. Often negative electron affinity is desired to obtain efficient cathodes that can supply electrons to the vacuum with little energy loss. The observed electron yield as a function of various parameters such as bias voltage or illumination conditions can be used to describe these structures with band diagrams in which the electron affinity is one parameter. For one illustration of the apparent effect of surface termination on electron emission, see Figure 3 in Marchywka Effect.

Structural formula

The structural formula of a chemical compound is a graphical representation of the molecular structure, showing how the atoms are arranged. The chemical bonding within the molecule is also shown, either explicitly or implicitly. There are several common representations used in publications. These are described below. Also several other formats are used, as in chemical databases, such as SMILES, InChI and CML.

Unlike chemical formulas or chemical names, structural formulas provide a representation of the molecular structure. Chemists nearly always describe a chemical reaction or synthesis using structural formulas rather than chemical names, because the structural formulas allow the chemist to visualize the molecules and the changes that occur.

Many chemical compounds exist in different isomeric forms, which have different structures but the same overall chemical formula. A structural formula indicates the arrangements of atoms in a way that a chemical formula cannot.

Lewis structures

Lewis structures (or "Lewis dot structures") are flat graphical formulas that show atom connectivity and lone pair or unpaired electrons, but not three-dimensional structure. This notation is mostly used for small molecules. Each line represents the two electrons of a single bond. Two or three parallel lines between pairs of atoms represent double or triple bonds, respectively. Alternatively, pairs of dots may used to represent bonding pairs. In addition, all non-bonded electrons (paired or unpaired) and any formal charges on atoms are indicated.

Condensed formulas

In early organic-chemistry publications, where use of graphics was severely limited, a typographic system arose to describe organic structures in a line of text. Although this system tends to be problematic in application to cyclic compounds, it remains a convenient way to represent simple structures:

CH3CH2OH (ethanol)

Parentheses are used to indicate multiple identical groups, indicating attachment to the nearest non-hydrogen atom on the left when appearing within a formula, or to the atom on the right when appearing at the start of a formula:

(CH3)2CHOH or CH(CH3)2OH (2-propanol)

In all cases, all atoms are shown, including hydrogen atoms.

Skeletal formulas

Skeletal formulas are the standard notation for more complex organic molecules. First used by the organic chemist Friedrich August Kekulé von Stradonitz the carbon atoms in this type of diagram are implied to be located at the vertices (corners) and termini of line segments rather than being indicated with the atomic symbol C. Hydrogen atoms attached to carbon atoms are not indicated: each carbon atom is understood to be associated with enough hydrogen atoms to give the carbon atom four bonds. The presence of a positive or negative charge at a carbon atom takes the place of one of the implied hydrogen atoms. Hydrogen atoms attached to atoms other than carbon must be written explicitly.

Indication of stereochemistry

Several methods exist to picture the three-dimensional arrangement of atoms in a molecule (stereochemistry).

Stereochemistry in skeletal formulas

Chirality in skeletal formulas is indicated by the Natta projection method. Solid or dashed wedged bonds represent bonds pointing above-the-plane or below-the-plane of the paper, respectively.

Unspecified stereochemistry

Wavy single bonds represent unknown or unspecified stereochemistry or a mixture of isomers. For example the diagram below shows the fructose molecule with a wavy bond to the HOCH2- group at the left. In this case the two possible ring structures are in chemical equilibrium with each other and also with the open-chain structure. The ring continually opens and closes, sometimes closing with one stereochemistry and sometimes with the other.

Perspective drawings

Newman projection and sawhorse projection

The Newman projection and the sawhorse projection are used to depict specific conformers or to distinguish vicinal stereochemistry. In both cases, two specific carbon atoms and their connecting bond are the center of attention. The only difference is a slightly different perspective: the Newman projection looking straight down the bond of interest, the sawhorse projection looking at the same bond but from a somewhat oblique vantage point. In the Newman projection, a circle is used to represent a plane perpendicular to the bond, distinguishing the substituents on the front carbon from the substituents on the back carbon. In the sawhorse projection, the front carbon is usually on the left and is always slightly lower:

Cyclohexane conformations

Certain conformations of cyclohexane and other small-ring compounds can be shown using a standard convention. For example, the standard chair conformation of cyclohexane involves a perspective view from slightly above the average plane of the carbon atoms and indicates clearly which groups are axial and which are

From Yahoo Answers


Answers:Group 3(13) has 3 electrons in the valence level

Question:2. Which of the following substances contains a nonpolar covalent bond? NaCl, MgF2, N2, H2O, or NH3? 3. How many grams of glucose (C6H12O6) are in 3.55 moles of glucose? 4. In the reaction of silver nitrate with sodium chloride, how many grams of silver chloride will be produced from 100. g of silver nitrate when it is mixed with an excess of sodium chloride? AgNO3 + NaCl AgCl + NaNO3 5. Any reaction that absorbs 150 kcal of energy can be classified as oxidation, exothermic, activated, reduction, or endothermic? 6. Which one of the following compounds contains an ion with a 3+ charge? CuCl, KCl, MgCl2, Na2, or FeCl3.

Answers:Ok, but I am not doing the math for you. 1) 3 2) N2 3) 3.55 x molecular mass of glucose 4) 100 / molecular mass of Silver Nitrate x molecular mass of silver chloride 5) Endothermic 6) FeCl3

Question:Does the lewis dot structure for Aluminum Oxide (Al2O3) look like this: O=Al-O-Al=O? If not, then how does it look? Also, could you help me make a lewis dot structure for Carbon Dioxide, Calcium Oxide, Diphosphorous Pentoxide, and Diiodine Pentoxide. I don't get this at all really. I don't know how to tell whether or not it's ionic or convalent. Thanks for your help.

Answers:The electron dot structure of the Al is just Al3+. This is because its three valence electrons are gone. The structure of the O is O(2-) with ** on three sides and oo on the fourth side. The *'s are O electrons and the o's are electrons gained from Al. Put down 2 such Al's for every 3 O's. Calcium oxide: The same as for aluminum oxide, except Ca is 2+ and there is one Ca2+ for every O=. Carbon dioxide: Draw O C O and leave a space between each. Draw ** on the top and bottom of each O and between each O and C. Draw oo between C and each O. *'s are O valence electrons; o's are C valence electrons. Phosphorus pentoxide: Draw O...............O ....P...O...P O...............O Put *o between the center O and each P. Put oo between each P and each of the four other O's. Put ** on the top and bottom of the center O and the remaining three sides of the other four O's. Notice that P provides both electrons to the pairs that are shared between it and the outer four O's. These are coordinate covalent bonds. Iodine pentoxide: The same as for P2O5. Just remember that you have to have an extra oo on one side of each iodine, because they have 7 valence electrons.

Question:ok i cant find this crap, please help me and i will give you "best answer"

Answers:6 come from the S, 6 from each O and 2 from the 2-. Now you can work it out.

From Youtube

Lewis dot structure :How to make a Lewis dot structure steps- ~ predict the location of the atom's (H is always terminal and the element furthest left on the periodic table is usually central). ~find the amount of valence electrons. ~Divide the valence electrons by 2 to get the number of electron pairs. ~place a single bond between the central atom and each of the terminal atoms ~ Subtract the # of pairs used in step 4 from the total # of (e-) to get the remaining number of (e-)or (electrons). ~Place lone pairs around each terminal atom bounded to the central atom to satisfy the octet rule(8) then place any remaining pairs around the central atom ~If the central atom doesn't have 8 electrons then u must convert some of the lone pairs from terminal toms to shared pairs.

Lewis Dot Structures of Ionic and Covalent Compounds :The following lesson looks at drawing Electron Dot or Lewis Dot diagrams of various ionic and covalent compounds, including polyatomic compounds.