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Different Theories.pdf


Department
Chemistry
Course Code
CHEM 112
Professor
John Carran

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CHEM 151 Molecular Geometry: Fall 2009
Lewis Structures, VSEPR Theory, and Valence Bond Theory
#9 VSEPR/Molecular Geometry Rev W08AEM Winter 2009 Page 1 of 10
Fill-in, stamp the box on top of page 7. Name
Partner
Lecture Instructor Date
Turn in ONLY pages 7–10!!!!
LEARNING OBJECTIVES After completing this exercise, you should feel comfortable with:
Drawing Lewis structures for molecules that obey the octet rule.
Drawing Lewis structures for molecules that may violate the octet rule.
Visualizing the three-dimensional shape of a molecule based on a drawing.
Identifying electron pair geometries (VSEPR).
Applying VSEPR theory to determine the shapes of molecules.
Identifying molecular geometries.
Distinguishing between electronic and molecular geometries.
Applying valence bond theory to the hybridization of atomic orbitals.
TO EARN YOUR FINAL STAMP: The following items must be completed to earn the final
stamp. You may complete the entire assignment in lab or outside of the lab, this reflect the minimum
required to earn your final stamp.
Complete the entire worksheet. You may work on the worksheet outside of the lab, however
you MUST have it completed all tables to get a stamp. The lab instructor will check over
your worksheet when you get it stamped.
Introduction
Because atoms are too small to see with the eye, scientists use models to visualize the physical
arrangements of atoms in molecules and polyatomic ions. These three-dimensional models aid in
understanding the polarity, reactivity and interaction of molecules. In this laboratory exercise, you
will work with three related theories of molecular bonding and structure.
I. LEWIS STRUCTURES
Lewis Structures give information in two-dimensional representations that can be used to predict
the three dimensional shapes of molecules and ions.
A basic concept of the atomic theory is that the chemical and physical properties of a substance
are determined by the distribution of outer shell or valence electrons (highest "n" value) in its atoms
and by the arrangements of these atoms to each other. The extraordinary non-reactivity of the Noble
gases has been related to their common electronic configuration of eight valence electrons (an
"octet"). Many chemical reactions and molecular formulas can be related to the observation that
many elements would "like" to have the same electronic configuration as the Noble gases.

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#9 VSEPR/Molecular Geometry Rev F09NF Fall 2009 Page 2 of 10
The following rules and procedures are given as a guide for drawing Lewis Dot Structures.
1. Write the MOLECULAR FORMULA for the compound.
2. Determine the total number of VALENCE ELECTRONS available for bonding by:
a. counting the valence electrons from each element in the compound.
b. adding one electron for each negative charge or subtracting one electron for
each positive charge, for polyatomic ions
3. ARRANGE ATOMS. For small molecules and polyatomic ions, place the element with the
lowest electronegativity in the center and arrange the other atoms around this central atom
using the following rules:
a. Hydrogen is never the central atom.
b. For oxyacids, the hydrogen atoms are usually bonded to oxygen atoms that are
bonded to the less electronegative central atom.
4. CONNECT ATOMS. Attach the atoms together with a "—" to signify a two-electron bond
between the atoms.
5. SATISFY OCTET RULE. Place the remaining electrons, in pairs (lone pairs), around each
atom to satisfy the "octet" requirement (a "duet" for hydrogen). It is best to maximize
bonding; if all of the electrons have been used up without satisfying the octet rule then atoms
must share more electrons by forming another bond. The octet rule must apply for C, N, O
and F! For all others, it is a good guideline, but may be violated (see item #6)
6. EXCEPTIONS TO OCTET RULE: There are some compounds that contain elements with
exceptions to the octet rule. These may have more than eight electrons, typically when n 3
(PCl5), or less than eight electrons, typically Group IIA, (BeCl2), IIIA (BCl3), and, of course,
hydrogen. For most of the compounds, the atoms are bonded by single bonds consisting of
one electron from the central atom and one electron from the outer atom. If there are any
extra electrons, they are placed on the central atom as unshared, or lone, pairs.
II. VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR)
VSEPR theory states that regions of high electron density, such as bonding pairs or lone pairs of
electrons (a VSEPR or electron domain), will arrange themselves as far apart as possible around the
central atom. Each single, double, or triple bond or unshared pair is counted as an electron domain.
The VSEPR rules:
1. Write the LEWIS STRUCTURE.
2. Count the number of ELECTRON DOMAINS around the central atom (or any atom!).
(Single, double, triple bonds = 1 VSEPR domain)
(Non-bonding (lone) pairs of electrons = 1 VSEPR domain)
3. Pick the VSEPR/ELECTRONIC GEOMETRY. Note that this electronic geometry is based
only on the number of electron domains, regardless of what they are (triple bond, lone pair,
etc…). Consult Table 1 to assign these geometries and the corresponding bond angles.
(Arrange the VSEPR pairs to minimize repulsion)

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#9 VSEPR/Molecular Geometry Rev F09NF Fall 2009 Page 3 of 10
Table 1: VSEPR Electronic Geometries
Number of
Electron Domains
VSEPR/Electronic
Geometry
Ideal Bond Angle(s)
2
Linear
180°
3
Trigonal Planar
120°
4
Tetrahedral
109.5°
5
Trigonal Bipyramidal
180o/120°/90°
6
Octahedral
180o/90°
4. Determine the MOLECULAR GEOMETRY. The molecular geometry is determined by what
we can actually “see” – the atoms bonded to the central atoms, but not the lone pairs. The
molecular geometry is set by the electronic geometry. Consult Table 2 to assign these
geometries
(Account for any unfilled positions in the electron pair geometry, i.e. non-bonding pairs/lone pairs)
Table 2: VSEPR Molecular Geometries
Number of non-bonding or lone pairs on an atom
Total # of e-
domains
0
1
2
2
Linear
3
Trigonal Planar
Bent
Linear
4
Tetrahedral
Trigonal
Pyramidal
Bent
5
Trigonal
Bipyramidal
See-Saw
T-Shaped
6
Octahedral
Square Pyramidal
Square Planar
5. Adjust angles to Recognize STERIC (size) EFFECTS. This gives rise to slight deviations to
the ideal bond angles.
Multiple Bonds – double, triple bonds take up more space than single bonds, therefore
angles involving them will be somewhat larger.
Non-Bonding/Lone Pairs – Lone Pair electrons take up much more space than bonding
pairs, compressing the angles between other, bonding pairs.
The molecular geometries (the actual geometry of the atoms) might best be explained with the
diagrams in your text.
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