Lewis Structure Handout

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Lewis Structure Handout for Chemistry Students An electron dot diagram, also known as a Lewis Structure, is a representation of valence electrons in a single atom, and can be further utilized to depict the bonds that form based on the available electrons for covalent bonding between multiple atoms. -Each atom has a characteristic dot diagram based on its position in the periodic table and this reflects its potential for fulfillment of the octet rule. As a general rule, the noble gases have a filled octet (column VIII with eight valence electrons) and each preceding column has successively one fewer. • For example, Carbon is in column IV and has four valence electrons. It can therefore be represented as: • Each of these "lone" electrons can form a bond by sharing the orbital with another unbonded electron. Hydrogen, being in column I, has a single valence electron, and will therefore satisfy carbon's octet thusly: • When a compound has bonded electrons (as with CH4) it is most accurate to diagram it with lines representing each of the single bonds. • Using carbon again for a slightly more complicated example, it can also form double bonds. This is the case when carbon bonds to two oxygen atoms, resulting in CO2. The oxygen atoms (with six valence electrons in column VI) are each represented as: • Because carbon is the least electronegative of the atoms involved, it is placed centrally, and its valence electrons will migrate toward the more electronegative oxygens to form two double bonds and a linear structure. This can be written as: where oxygen's remaining lone pairs are still shown. Practice Problems: Draw each of the following with all valence electrons (bonded or unbounded) represented Single Atom 1) Mg 2) Br 3) N Compounds: 1) BCl3 2) N2H2 3) HClO Chemistry Last Edited 7/2016 .

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Valence Bond Theory
UMass Boston, Chem 103, Spring 2006 CHEM 103 Molecular Geometry and Valence Bond Theory Lecture Notes May 2, 2006 Prof. Sevian Announcements z The final exam is scheduled for Monday, May 15, 8:00- 11:00am It will NOT be in our regularly scheduled lecture hall (S- 1-006). The final exam location has been changed to Snowden Auditorium (W-1-088). © 2006 H. Sevian 1 UMass Boston, Chem 103, Spring 2006 More announcements Information you need for registering for the second semester of general chemistry z If you will take it in the summer: z Look for chem 104 in the summer schedule (includes lecture and lab) z If you will take it in the fall: z Look for chem 116 (lecture) and chem 118 (lab). These courses are co-requisites. z If you plan to re-take chem 103, in the summer it will be listed as chem 103 (lecture + lab). In the fall it will be listed as chem 115 (lecture) + chem 117 (lab), which are co-requisites. z Note: you are only eligible for a lab exemption if you previously passed the course. Agenda z Results of Exam 3 z Molecular geometries observed z How Lewis structure theory predicts them z Valence shell electron pair repulsion (VSEPR) theory z Valence bond theory z Bonds are formed by overlap of atomic orbitals z Before atoms bond, their atomic orbitals can hybridize to prepare for bonding z Molecular geometry arises from hybridization of atomic orbitals z σ and π bonding orbitals © 2006 H. Sevian 2 UMass Boston, Chem 103, Spring 2006 Molecular Geometries Observed Tetrahedral See-saw Square planar Square pyramid Lewis Structure Theory
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Lewis Structures
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Draw Three Resonance Structures for the Chlorate Ion, Clo3
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Molecular Orbital Theory
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Localized Electron Model for Bonding What Does Hybridization Mean?
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Understanding Hypervalency
Created by Gerard Rowe ([email protected]) and posted on VIPEr (www.ionicviper.org) on Nov 5, 2013. Copyright Gerard Rowe 2013. This work is licensed under the Creative Commons Attribution-NonCommerical-ShareAlike 3.0 Unported License. To view a copy of this license visit http://creativecommons.org/about/license/. Understanding Hypervalency In the world of quantum chemistry, there are two prevailing theories that explain chemical bonding: valence bond (VB) theory and molecular orbital (MO) theory. By this point in time, you have had exposure to both. The concepts of equivalent hybrid orbitals and resonance structures are part of VB theory, and VB orbitals generally only sit between two atoms. MO theory, on the other hand, uses orbitals that are delocalized; that is, they can span multiple atoms, or even the entire molecule. The interesting thing about these two theories is that they are both equally valid, and make the exact same predictions. Which one you use depends on the situation. The utility of VB hybrid orbitals lies in their simplicity in explaining structural features such as the four equivalent C-H bonds of methane (as does MO theory in a more complex manner). However, when one looks at the electronic spectrum of methane, two major peaks are observed. Hybrid orbital theory and molecular orbital theory predict very different orbital energy levels for methane (Figure 1). Using MO theory, it is very easy to see why there would be two peaks; electrons jump from the HOMO to each of the two higher energy levels (Figure 1b). With VB theory, it takes a lot of extra math to explain how there could be two peaks.
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Chem 115 Quiz #4
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Chapter 12 Arenes and Aromaticity
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