Covalent compounds hydrogen atom

We need to learn the number of bonds that are commonly formed by the atoms that are important in organic chemistry. In the vast majority of cases, hydrogen forms one covalent bond so that it has two electrons in its outer shell. In covalent compounds the atoms of the second period (or row) of the periodic table need eight electrons in their outer shell. (This is called the octet rule.) We have already seen that to accomplish this carbon forms four bonds. In a similar fashion we can determine the number of bonds that some other second-period atoms prefer to form. The number of bonds that commonly occur in neutral molecules for the elements of most interest in organic chemistry are listed in Table 1.1. 

Ammonium salts constitute only one type of a large class of compounds known as onium salts/ The monopositive cation of these salts consists of a central atom of an element from Periodic Group V(N,P,As,Sb,Bi), VI(0,S,Se,Te), or VII(Cl,Br,I), which is bonded covalently to hydrogen atoms, organic radicals, or a combination of these. Onium salts, in which oxygen is the central atom, are known as oxonium salts. Although Werner made a number of important experimental studies of oxonium salts, and even discovered several new types of such compounds 78,90), his theoretical views are of primary interest here. 

The Lewis model for covalent bonding starts with the recognition that electrons are not transferred from one atom to another in a nonionic compound, but rather are shared between atoms to form covalent bonds. Hydrogen and chlorine combine, for example, to form the covalent compound hydrogen chloride. This result can be indicated with a Lewis diagram for the molecule of the product, in which the valence electrons from each atom are redistributed so that one electron from the hydrogen atom and one from the chlorine atom are now shared by the two atoms. The two dots that represent this electron pair are placed between the symbols for the two elements ... 

The reaction of beryllium metal with aqueous acids yields hydrogen and ionic compounds, such as BeCl2 4 H2O. In BeCl2 4 H2O, water molecules are covalently bonded to Be ions, producing the complex cations [Be(OH2)4] that, together with the anions d , form the crystal lattice. In covalent compounds, Be atoms appear to use hybrid orbitals—sp orbitals in BeCl2(g) and sp orbitals in BeCl2(s) (Fig. 21-12). 

Covalent. Formed by most of the non-metals and transition metals. This class includes such diverse compounds as methane, CH4 and iron carbonyl hydride, H2Fe(CO)4. In many compounds the hydrogen atoms act as bridges. Where there are more than one hydride sites there is often hydrogen exchange between the sites. Hydrogens may be inside metal clusters. 

The biochemical basis for the toxicity of mercury and mercury compounds results from its ability to form covalent bonds readily with sulfur. Prior to reaction with sulfur, however, the mercury must be metabolized to the divalent cation. When the sulfur is in the form of a sulfhydryl (— SH) group, divalent mercury replaces the hydrogen atom to form mercaptides, X—Hg— SR and Hg(SR)2, where X is an electronegative radical and R is protein (36). Sulfhydryl compounds are called mercaptans because of their ability to capture mercury. Even in low concentrations divalent mercury is capable of inactivating sulfhydryl enzymes and thus causes interference with cellular metaboHsm and function (31—34). Mercury also combines with other ligands of physiological importance such as phosphoryl, carboxyl, amide, and amine groups. It is unclear whether these latter interactions contribute to its toxicity (31,36). 

Covalent — refers to a chemical bond in which there is an equal/even sharing of bonding electron pairs between atoms. This is typical of the bonding between carbon atoms and between carbon and hydrogen atoms in organic compounds. 

Note that these compounds are covalently bonded compounds containing only hydrogen and carbon. The differences in their strucmral formulas are apparent the alkanes have only single bonds in their structural formulas, while the alkenes have one (and only one) double bond in their structural formulas. There are different numbers of hydrogen atoms in the two analogous series. This difference is due to the octet rule that carbon must satisfy. Since one pair of carbon atoms shares a double bond, this fact reduces the number of electrons the carbons need (collectively) by two, so there are two fewer hydrogen atoms in the alkene than in the corresponding alkane. 

Elements as well as compounds can exist as discrete molecules. In hydrogen gas, the basic building block is a molecule consisting of two hydrogen atoms joined by a covalent bond ... 

H3 PO4 Phosphoric acid is a covalent compound with a net charge of zero. Each hydrogen atom has an oxidation number of+1 (Guideline 3), and each oxygen has an oxidation number of-2 (Guideline 4). Now add the contributions from these atoms 3(+l) + 4(-2) = -5. For the oxidation numbers to sum to zero (Guideline 2), the phosphorus atom of phosphoric acid must have an oxidation number of +5. 

In (a), a double bond is needed to make the octet of sulfur. In (b), the extra pair of electrons makes the set of atoms an ion and eliminates the need for a double bond. In (c), that same ion is present, along with the two potassium ions to balance the charge. In (d). two hydrogen atoms arc covalently bonded to oxygen atoms to complete the compound. 

Some ionic compounds contain a combination of bonds. For instance, in polyatomic ions such as ammonium (NH4+), the hydrogen atoms are bonded to the nitrogen atom by polar covalent bonds. The ionic bond is thus between this covalently bonded moiety and another oppositely charged ion such as chloride (CT). 

This Lewis structure shows methane, the simplest organic compound. The carbon atom has four valence electrons, and it obtains four more electrons by forming four covalent bonds with the four hydrogen atoms. 

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