8: Chemical Bonding I - The Lewis Model
- Page ID
- 465547
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The Lewis Theory used observations from chemists and physicists to form a theory about chemical bonding. This work was essentially a compilation of the knowledge at the time. It revolved around the importance of valence electrons in chemical bonding. These are the electrons that are in the outermost shell. For example Na may have 11 electrons, but only one is a valence electron, the one in 3s1. Meanwhile P has 15 electrons, but has five valence electrons, 3s2 and 3p2. The bonding of an element is based on how they fill their octets i.e. achieve a noble gas electron configurations. Lewis went on to explain how certain elements such as Boron did not necessarily follow these same rules. By drawing schematics of molecules called Lewis Dot Structures, important characteristics of molecules can be described.
Chapter Sections
- 8.1: Lewis Symbols
- Lewis dot symbols can be used to predict the number of bonds formed by most elements in their compounds. Lewis electron dot symbols, which consist of the chemical symbol for an element surrounded by dots that represent its valence electrons, grouped into pairs often placed above, below, and to the left and right of the symbol. The structures reflect the fact that the elements in period 2 and beyond tend to gain, lose, or share electrons to reach a total of 8 valence electrons in their compounds.
- 8.2: Covalent Bonding and Lewis Dot Structures
- Lewis dot symbols provide a simple rationalization of why elements form compounds with the observed stoichiometries. A plot of the overall energy of a covalent bond as a function of internuclear distance is identical to a plot of an ionic pair because both result from attractive and repulsive forces between charged entities. Lewis structures are an attempt to rationalize why certain stoichiometries are commonly observed for the elements of particular families.
- 8.3: Resonance Structures and Formal Charge
- Some molecules have two or more chemically equivalent Lewis electron structures, called resonance structures. Resonance is a mental exercise and method within the Valence Bond Theory of bonding that describes the delocalization of electrons within molecules. These structures are written with a double-headed arrow between them, indicating that none of the Lewis structures accurately describes the bonding but that the actual structure is an average of the individual resonance structures.
- 8.4: Polar Covalent Bonds
- Bond polarity and ionic character increase with an increasing difference in electronegativity. The electronegativity (χ) of an element is the relative ability of an atom to attract electrons to itself in a chemical compound and increases diagonally from the lower left of the periodic table to the upper right. The Pauling electronegativity scale is based on measurements of the strengths of covalent bonds between different atoms, whereas the Mulliken electronegativity of an element is the average
- 8.5: Bond Energies, Strengths, and Lengths
- Bond order is the number of electron pairs that hold two atoms together. Single bonds have a bond order of one, and multiple bonds with bond orders of two (a double bond) and three (a triple bond) are quite common. The bond with the highest bond order is both the shortest and the strongest. In bonds with the same bond order between different atoms, trends are observed that, with few exceptions, result in the strongest single bonds being formed between the smallest atoms.
References:
- Lewis, Gilbert Newton. Valence and the Structure of Atoms and Molecules,. New York: Chemical Catalog, 1923. Print.
- Pauling, Linus. The Nature of the Chemical Bond: and the Structure of Molecules and Crystals : Introduction to Modern Structural Chemistry. Ithaca, NY: Cornell UP, 1960. Print.
- Petrucci, Ralph H. General Chemistry: Principles and Modern Applications. Toronto, Ont.: Pearson Canada, 2011. Print.
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Contributors
- Nilpa Shah (UCD)