Limitations of Lewis Dot Structure

3 Key excerpts on "Limitations of Lewis Dot Structure"

• 

  eBook - ePub

  Organic Chemistry

  Concepts and Applications

  • Allan D. Headley(Author)
  • 2019(Publication Date)
  • Wiley

    (Publisher)

  The subscript is written to the right of the atoms in the condensed formula. Given condensed formulas for compounds, the Lewis dot structures give fairly good descriptions of the arrangements of the atoms, covalently bonded and nonbonded electrons in molecules and ions, but molecular formulas or structural formulas are used frequently to represent organic molecules. The Lewis dot structure of a molecule or ion can be transformed easily into structural formulas. For structural formulas, lines represent the bonding electrons of the Lewis dot structures, and dots are still used for the nonbonded electrons. Examples of both representations are shown below. 1.4.1 Line‐Angle Representations of Molecules Throughout the course, there is yet another representation of molecular structures that will be used and is known as the line‐angle representation. In this representation, lines that intersect each other at an angle of 120° are used. Each intersection represents a carbon along with the appropriate number of hydrogens, but the hydrogens are not shown. For the line‐angle representation, the start of a line represents a carbon with three hydrogens; at an intersection of two lines, there are two hydrogens; at an intersection of three lines, there is one hydrogen; and at the intersection of four lines, there are no hydrogens

  Sign up to read

  Learn more about book
• 

  eBook - ePub

  Survival Guide to General Chemistry

  • Patrick E. McMahon, Rosemary McMahon, Bohdan Khomtchouk(Authors)
  • 2019(Publication Date)
  • CRC Press

    (Publisher)

  14 Alternate Methods for Visualizing and Constructing

  Lewis Structures of Covalent Molecules

  I	INTRODUCTION TO INTERPRETATION OF LEWIS STRUCTURES

  Lewis structures are pictorial representations, not exact theoretical analyses, of the role of all valence electrons (outer shell electrons for non-metals) in a covalent molecule. All valence electrons for all atoms in the molecule are accounted for as either covalent bonding electron pairs, shown as a line, or as electrons not used in bonding, termed lone electron pairs. A lone electron pair is usually shown as a pair of dots; occasionally, non-octet atoms will have an unpaired electron, shown as a single dot.

  The Lewis structure for an individual atom shows all valence electrons as dots around the symbol of the atom. The generally accepted method is to show the electrons as pairs or to show them symmetrically distributed around the symbol, especially when indicating bonding in molecules.

  Example: Acceptable for carbon with four valence electrons:

  Example: Acceptable for oxygen with six valence electrons:

  The standard method for constructing Lewis structures involves counting up all valence electrons for all atoms in the molecule, then distributing them into a bonding pattern according to a set of fixed rules. The methods described in this chapter adapts and extends the general bonding pattern information generated from this approach. The purpose is to generate techniques for understanding the electron behaviors responsible for bonding in typical covalent molecules. Corresponding techniques are then produced for determining valid atom connection patterns for larger covalent molecules without counting all valence electrons and in the absence of a provided bonding pattern.

  Sign up to read

  Learn more about book
• 

  eBook - ePub

  Understanding Molecules

  Lectures on Chemistry for Physicists and Engineers

  • Franco Battaglia, Thomas F. George(Authors)
  • 2018(Publication Date)
  • CRC Press

    (Publisher)

  This is Lewis scheme, for which we are going to present benefits and drawbacks. 6.2 LEWIS STRUCTURES In the years when quantum theory was still under development, but with a periodic table portrait already established, starting from the observed chemical inertia of the noble gases, 1 G. N. Lewis posed the conjecture that the formation of a stable molecule is determined by the fact that each of its atoms, sharing electrons with nearest-neighbor atoms, would reach a total number of electrons equal to that of the noble gas closest to it in the periodic table. More precisely, Lewis scheme defines for each atom the number of valence electrons as the difference between the total number of the electrons in the atom and the number of electrons in the noble gas which precedes the atom in the periodic table. Moreover, assuming that the valence electrons in an atom are organized in pairs, the scheme distinguishes in a molecule two types of electron pairs: bond pairs and lone pairs. The total number of electrons in bond and lone pairs pertaining to each atom in a molecule must then be 2, 8, or 18, because this is the number of valence electrons in a noble-gas atom (a circumstance which, in the light of the quantum theory of atoms as seen in the previous chapter, amounts to the number of electrons in a complete electron shell). 6.2.1 DIATOMIC MOLECULES For instance, given the hydrogen atom with only one electron, the H 2 molecule is represented with only one bond pair: H − H. When the electron pair is ascribed to each atom, it has two electrons altogether, as many as the closest noble gas, which is helium. A helium molecule is correctly predicted not to exist, and a beryllium molecule, according to Lewis rules, behaves as that of helium, i.e., does not exist; and lithium’s diatomic molecule, with a more liberal interpretation of Lewis rules, behaves as that of hydrogen

  Sign up to read

  Learn more about book