Lewis structures bonding pairs

Molecule or ion Lewis structure Bonding electron pairs Nonbonding electron pairs Shape... 

MO theory redefines bond order. In a Lewis structure, bond order is the number of electron pairs per linkage. The MO bond order is the number of electrons in bonding MOs minus the number in antibonding MOs, divided by two ... 

Similarly, the orbitals determined in both semiempirical and Hartree-Fock calculations may be transformed into natural bond orbitals (NBOs), which provide pictures of localized bonds and lone pairs that correspond closely with Lewis structure bonding models. " For example. Figure 4.53 shows the NBOs for ethene cch/ o ccr ai d ncc bonds. 

Lewis Theory An Overview—Lewis symbol represents the valence electrons of an atom by using dots placed around the chemical symbol. A Lewis structure is a combination of Lewis symbols used to represent chemical bonding. Normally, all the electrons in a Lewis structure are paired, and each atom in the structure acquires an octet—that is, there are eight electrons in the valence shell. In Lewis theory, chemical bonds are classified as ionic bonds, which are formed by electron transfer between atoms, or covalent bonds, which are formed by electrons shared between atoms. Most bonds, however, have partial ionic and partial covalent characteristics. 

Section 1 3 The most common kind of bonding involving carbon is covalent bond ing A covalent bond is the sharing of a pair of electrons between two atoms Lewis structures are written on the basis of the octet rule, which limits second row elements to no more than eight electrons m their valence shells In most of its compounds carbon has four bonds... 

Draw a Lewis structure for singlet methylene, CH2 (all the electrons in singlet methylene are spin-paired). Ho many electrons remain after all bonds have been formei Where are the extra electrons located, in the plane the molecule or perpendicular to the plane Examine t highest-occupied molecular orbital (HOMO) of methyle to tell. 

Compare and contrast the electrostatic potential map of a typical detergent with that of a typical soap (stearate). Which part of each molecule will be most water soluble (hydrophilic) Draw a Lewis structure that describes each molecule s water-soluble group (make sure you indicate all necessary formal charges and lone pairs). Which part(s) of each molecule will be most grease soluble (lipophilic) What kinds of atoms and bonds are found in these groups ... 

The development we have just gone through for NH3 is readily extended to the water molecule, H20. Here the Lewis structure shows that the central oxygen atom is surrounded by two single bonds and two unshared pairs ... 

Strategy Draw the Lewis structure for the molecule and determine the number of electron pairs (single bonds or unshared pairs) around the central atom. The possible hybridizations are sp (two pairs), sp2 (three pairs), sp3 (four pairs), sp3d (five pairs), and sp3d2 (six pairs). 

It is possible to write a simple Lewis structure for foe S042- ion, involving only single bonds, which follows foe octet rule. However, Linus Pauling and others have suggested an alternative structure, involving double bonds, in which foe sulfur atom is surrounded by six electron pairs. 

Some ligands have more than one atom with an unshared pair of electrons and hence can form more than one bond with a central metal atom. Ligands of this type are referred to as chelating agents the complexes formed are referred to as chelates (from the Greek chela, crab s claw). Two of the most common chelating agents are the oxalate anion (abbreviated ox) and the ethylenediamine molecule (abbreviated en), whose Lewis structures are... 

In Chapter 7, we used valence bond theory to explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of the overlap of atomic orbitals. By invoking hybridization, valence bond theory can account for the molecular geometries predicted by electron-pair repulsion. Where Lewis structures are inadequate, as in S02, the concept of resonance allows us to explain the observed properties. 

The Lewis structure of a molecule shows atoms by their chemical symbols, covalent bonds by lines, and lone pairs by pairs of dots. For example, the Lewis structure of HF is H - F . We shall see that Lewis structures are a great help in... 

Nonmetal atoms share electrons until each has completed its octet (or duplet) a Lewis structure shows the arrangement of electrons as lines (bonding pairs) and dots (lone pairs). 

Each atom in a polyatomic molecule completes its octet (or duplet for hydrogen) by sharing pairs of electrons with its immediate neighbors. Each shared pair counts as one covalent bond and is represented by a line between the two atoms. A Lewis structure does not portray the shape of a polyatomic molecule it simply displays which atoms are bonded together and which atoms have lone pairs. 

A single shared pair of electrons is called a single bond. Two electron pairs shared between two atoms constitute a double bond, and three shared electron pairs constitute a triple bond. A double bond, such as C 0, is written C=0 in a Lewis structure. Similarly, a triple bond, such as C C, is written G C. Double and triple bonds are collectively called multiple bonds. The bond order is the number of bonds that link a specific pair of atoms. The bond order in H, is 1 in the group C=0, it is 2 and, for O C in a molecule such as ethyne, C2H2, the bond order is 3. 

The Lewis structure of a polyatomic species is obtained by using all the valence electrons to complete the octets (or duplets) of the atoms present by forming single or multiple bonds and leaving some electrons as lone pairs. 

To assign a formal charge, we establish the ownership of the valence electrons of an atom in a molecule and compare that ownership with the free atom. An atom owns one electron of each bonding pair attached to it and owns its lone pairs completely. The most plausible Lewis structure will be the one in which the formal charges of the atoms are closest to zero. 

The Lewis structure of the product, a white molecular solid, is shown in (32). In this reaction, the lone pair on the nitrogen atom of ammonia, H3N , can be regarded as completing boron s octet in BF3 by forming a coordinate covalent bond. 

Draw the most important Lewis structure for each of the following ring molecules (which have been drawn without showing the locations of the double bonds). Show all lone pairs and nonzero formal charges. If there are equivalent resonance... 

The Lewis structures encountered in Chapter 2 are two-dimensional representations of the links between atoms—their connectivity—and except in the simplest cases do not depict the arrangement of atoms in space. The valence-shell electron-pair repulsion model (VSEPR model) extends Lewis s theory of bonding to account for molecular shapes by adding rules that account for bond angles. The model starts from the idea that because electrons repel one another, the shapes of simple molecules correspond to arrangements in which pairs of bonding electrons lie as far apart as possible. Specifically ... 

A molecule with only two atoms attached to the central atom is BeCl2. The Lewis structure is CI — Be — CE, and there are no lone pairs on the central atom. To be as far apart as possible, the two bonding pairs lie on opposite sides of the Be atom, and so the electron arrangement is linear. Because a Cl atom is attached by each bonding pair, the VSEPR model predicts a linear shape for the BeCL molecule, with a bond angle of 180° (4). That shape is confirmed by experiment. 

Now consider the alkynes, hydrocarbons with carbon-carbon triple bonds. The Lewis structure of the linear molecule ethyne (acetylene) is H—O C- H. To describe the bonding in a linear molecule, we need a hybridization scheme that produces two equivalent orbitals at 180° from each other this is sp hybridization. Each C atom has one electron in each of its two sp hybrid orbitals and one electron in each of its two perpendicular unhybridized 2p-orbitals (43). The electrons in the sp hybrid orbitals on the two carbon atoms pair and form a carbon—carbon tr-bond. The electrons in the remaining sp hybrid orbitals pair with hydrogen Ls-elec-trons to form two carbon—hydrogen o-bonds. The electrons in the two perpendicular sets of 2/z-orbitals pair with a side-by-side overlap, forming two ir-honds at 90° to each other. As in the N2 molecule, the electron density in the o-bonds forms a cylinder about the C—C bond axis. The resulting bonding pattern is shown in Fig. 3.23. 

According to Lewis s approach and valence-bond theory, we should describe the bonding in 02 as having all the electrons paired. However, oxygen is a paramagnetic gas (Fig. 3.24 and Box 3.2), and paramagnetism is a property of unpaired electrons. The paramagnetism of 02 therefore contradicts both the Lewis structure and the valence-bond description of the molecule. 

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