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CSUS Department of Chemistry Molecular Shapes Chem. 1A Chem. 1A Week 11 Discussion: Molecular Shapes Valence Shell Electron Pair Repulsion Theory (VSEPR) Section 8.6 in your text. One aspect of chemistry that often proves to be difficult for many students is the visualization of compounds, ions and molecules in three dimensional space. In lecture and homework you have seen how the octet rule is used to draw “Lewis dot structures” that are useful in predicting how groups of atoms combine to form molecules and polyatomic ions. You have also learned how to identify the molecular geometry, electron pair geometry and polarity of a molecule from these structures. As with most topics that prove to be difficult, all it takes is a bit of practice. In this exercise you will construct three‐dimensional models of various molecules and ions, then use these models to identify certain molecular properties. The following Lewis dot structures can be used to illustrate the usefulness of three‐dimensional models: H
C
: :
: Cl
H
H
: :
: :
C
H
H
Cl :
: Cl
H
C
:N
H
H
H
: Cl :
:
H
(1)
(2)
(3)
(4)
H
From VSEPR theory, structures 1−4 are all predicted to have tetrahedral electron pair geometries with 109° bond angles about the central atom. Structures 1−3 also have tetrahedral molecular geometries since there are no lone pairs on the central atom. Structure 4 is defined by a trigonal pyramidal molecular geometry with bond angles close to 109° (less than) due to the presence of the lone pair of electrons on the nitrogen atom. When you construct models of structures 1‐3 you will find that the square appearance of the two‐dimensional sketches shown here is not a very good representation of the actual 3‐dimensional structures. In fact, you will find that structures 2 and 3 are actually the same molecule, but viewed from different angles. You will also find that this molecule has a net dipole moment, despite its apparent symmetry as viewed from the angle of structure 2. Finally, you will note that structure 4 also has a net dipole moment. In more complex situations, it may be possible to draw more than one Lewis dot structure which follows the octet rule. Consider the four structures shown below: H
C
:
H
H
(5)
O
H
C
H
:Cl
C
Cl :
: Cl
C
: :
H
:
O
:
C
:O:
: :
H
:O :
H
: :
C
H
: :
H
:
H
Cl :
H
(6)
(7)
(8)
In determining which of two alternative dot structures is more stable, it is useful to consider the concept of “formal charge.” The formal charge on an atom in a molecule is defined as follows: Formal Charge = (# valence e− on free atom) − (# e− assigned to the atom in the molecule) The number of electrons assigned to the atom in the molecule is calculated by adding all non‐bonding electrons plus half the bonding (shared) electrons found around the atom. V = # of valence electrons on free atom ⎛B
⎞
FC = V − ⎜ + L ⎟ ⎝2
⎠
B = # of electrons participating in bonding L = # of lone pair electrons on the atom in the molecule None of the atoms in structures 5 or 6 has a formal charge, so these two structures are predicted to be stable. In this case, structures 5 and 6 are said to be “isomers,” both with the molecular formula C2H6O. (Structure 5 is Page 1 of 5 CSUS Department of Chemistry Molecular Shapes Chem. 1A ethyl alcohol and structure 6 is dimethyl ether.) None of the atoms in structure 7 has a formal charge, but in structure 8 the oxygen atom has a −1 formal charge and the double‐bonded chlorine atom has a +1 formal charge. In most cases, structures without formal charges are more stable, so we predict that COCl2 (phosgene) has the bonding predicted by structure 7. Molecular Polarity: Molecular polarity is a physical property of compounds which relates to other physical properties such as melting and boiling points, solubility, and intermolecular interactions between molecules. For the most part, there is a direct correlation between the polarity of a molecule and number and types of polar or non‐polar covalent bonds which are present. In a few cases, a molecule may have polar bonds, but in a symmetrical arrangement which then gives rise to a non‐polar molecule such as carbon dioxide. Polarity results from the uneven partial charge distribution between various atoms in a compound. Atoms, such as nitrogen, oxygen, and halogens, that are more electronegative have a tendency to have partial negative charges. Atoms that are less electronegative, such as carbon and hydrogen, have a tendency to be more neutral or have partial positive charges. Electrons in a polar covalent bond are unequally shared between the two bonded atoms, which results in partial positive and negative charges. The separation of the partial charges creates a dipole. The word dipole means two poles: the separated partial positive and negative charges. A polar molecule results when a molecule contains polar bonds in an unsymmetrical arrangement. Nonpolar molecules are of two types. Molecules whose atoms have equal or nearly equal electronegativities have zero or very small dipole moments. A second type of nonpolar molecule has polar bonds, but the molecular geometry is symmetrical allowing the bond dipoles to cancel one another through vector addition. Carbon dioxide is nonpolar because the dipoles that exist along the carbon/oxygen bonds are equal and in opposite directions. Water on the other hand is polar because the y−components of the dipole moments add together to produce a net dipole that points vertically up through the center of the molecule. One may also note that the symmetry of a molecule will indicate whether a molecule is polar. Note that in CO2, along or perpendicular to the bond axis, the molecule possesses a line (or plane) of reflection. In water, this is not the case. Along each O−H bond axis, there is no mirror image, thus indicating that the molecule is polar. Configurational Isomers: Certain Lewis structures (especially those containing double bonds) may be drawn such that the arrangement of the bonding atoms produces two or more structures that cannot be superimposed. These structures are know as "configurational isomers". Molecules that can be superimposed via a rotation are not isomers of one another. H
H
C
F
H
C
F
C
F
cis−1,2−difluoroethene F
F
C
H
C
H
trans−1,2−difluoroethene F
C
H
1,1−difluoroethene Each molecule has the same molar mass but different physical and chemical properties.
Page 2 of 5 CSUS Department of Chemistry Molecular Shapes Chem. 1A In Class Exercise: Each student should check out a molecular model kit from the lab instructor. For each molecular formula in the first part of the experiment, the following set of steps should be followed: (1) Draw a Lewis dot structure for the formula. Use formal charge to reject unstable possibilities. Note that carbon must always have four bonds in order to have a formal charge of zero. If the only stable dot structure has one or more formal charges, indicate these charges in your dot structure. (2) Choose appropriate parts from the kit to represent the necessary atoms, bonds, and non‐bonding pairs: H atom = small sphere Single bond to H atom = short straight stick C atom = black 4‐hole Single bond to other atoms = long straight sticks O atom = blue 4‐hole Double bond = two curved sticks Other atoms = red and green 4‐hole Triple bond = three curved sticks Non‐bonding pair = short straight stick Expanded Octet Atoms: 10 e‐ valence = black 5‐hole, 12 e‐ valence = gray with 6 square faces (3) Construct a three‐dimensional model corresponding to your dot structure. Use your model to answer the questions concerning geometry and polarity. Practice Molecules Formula Lewis Dot Structure (include all lone pairs) Electron Pair Geometry, Molecular Geometry, & Central Bond angles Polarity (yes or no, N/A for ions)
CO H2S CH3Cl Page 3 of 5 CSUS Department of Chemistry Molecular Shapes Formula Chem. 1A Lewis Dot Structure Molecular, Electron Pair Geometry & Central Bond angles Polarity (yes or no, N/A for ions)
BF3 NO2− PH3 I3− C2H2Cl2 FNO2 Page 4 of 5 CSUS Department of Chemistry Molecular Shapes Formula Chem. 1A Lewis Dot Structure (include all lone pairs) Molecular, Electron Pair Geometry & Central Bond angles Polarity (yes or no, N/A for ions)
SeCl6 XeH4 Additional Properties Among the practice molecules, which ones have configurational isomers? Draw structures of all isomers. Page 5 of 5