Molecular Modeling with Spartan: Polyatomic Molecules, VSEPR, Localized and Delocalized Bonding

Introduction
Spartan Student Edition was used to construct and examine molecular diagrams of H2O, NH3, CH4, and SF4. The properties examined include electron density, dipole moments, electrostatic potential maps, and equilibrium geometries. The geometries calculated with spartan were compared to the VSPER calculations. Spartan was used to examine the electron density, electrostatic potentials, and electrostatic charges of NO2 , NO2+, and NO2-, and compare them to the predictions made by using Lewis structures.
Procedure
The H2O molecule was created and the electrostatic charges were displayed for the atoms. All equilibrium geometry calculations were done at the B3LYP level in a 6-31G* basis set using Spartan Student Edition. The 0.002 electron/Å3 isodensity surface was calculated and displayed for H2O. Then a 0.08 electron/Å3 isodensity surface was calculated, displayed, and used to observe the electron density. An electrostatic potential map was created, made transparent, and legend displayed. To visualize the volume taken up by the lone pairs, a potential isodensity surface was calculated at the value of -83.68 kJ/mol. The angles and distances between the atoms were calculated. This procedure was repeated on the NH3 and CH4 molecules.
An electrostatic potential map was created and displayed for SF4. The bond angles and lengths were found and created. The molecules energy (au) was found and recorded. Another SF4 molecule was created which had three fluorine atoms in equatorial positions and one in an axial position, and the same procedure was done with it.
An electrostatic potential map was created for the NO2 molecule and the electrostatic charges displayed. An 0.08 electron/Å3 isodensity map was displayed. The bond lengths and angles were measured and recorded. The same procedure was followed on NO2+ and NO2- as was NO2.
Detailed procedures can be found in reference 1.

Results
The electron density of water was not distributed evenly throughout the molecule, as can be seen in Figure 1. The highest electron density was near the oxygen and the bonds where the red and green areas were.

Figure 1. Electrostatic Potential Diagram of Water Molecule.





The electron density of ammonia was not distributed evenly throughout the molecule, as can be seen in Figure 2.

Figure 2. Electrostatic Potential Diagram of Ammonia.
The electron density of Methane was not distributed evenly throughout the molecule but was symmetrical, as can be seen in Figure 3.

Figure 3. Electrostatic Potential Diagram of Methane.
The bond angles, bond lengths, and dipole moments were calculated using spartan, and the bond lengths predicted by the VSEPR theory. This data can be found in Table 1.
Table 1. The bond angles, bond lengths, and dipole moments calculated using spartan, and the bond lengths predicted by the VSEPR theory.
Molecule Angles (degrees) VSEPR Angles (degrees) Bond Length (Å) Dipole Moment (D)
H2O 103.70° 109.5° 0.969 2.09
NH3 105.76° 109.5° 1.019 1.91
CH4 109.47° 109.5° 1.093 0.00
SF4 equilateral: 102.09°
axial: 87.29° 120°
90° 1.595
1.672 0.89
SF4 (w/ 3 F atoms in equilateral positions) equilateral: 119.14°
axial: 84.66° 120°
90° 1.679
1.067 0.85
NO2 133.82° 120° 1.203 -
NO2+ 179.99° 180° 1.129 -
NO2- 129.98° 120° 1.305 -
1.1 The bond angels in the molecules increase from water, to ammonia, to methane. This is because the bond angles decrease when more lone pairs are present. The lone pairs push the other bonds because they take up more space than bonded electrons.
1.2 The bond angles of the three molecules, found using the VSPER theory are all 109.5° because they are tetrahedrals. The bond angles found using Spartan are different because it into account the effects of lone pairs and double bonds.
1.3 The bond lengths are shorter in water than ammonia, and longer in methane than ammonia. This is because the electronegativity of the central atom, the greater the electronegativity the closer it pulls the bonding atom.
1.4 Electrostatic charges exist in water, ammonia and methane because each atom has a partial charge. Electronegativity causes the electrons in the molecule to be attracted to certain atoms, giving the atoms different partial charges and making the atom polar.
1.5 The dipole moment in water is larger than the one in ammonia, and methane doesn’t have a dipole moment. Each atom’s electronegativity along with the molecules lone pairs and the dimensions of the molecule effect the molecules electrostatic charge. The electrostatic charge can be represented by a dipole moment, which shows the size and direction of the force.
1.6 A molecular dipole moment of zero does not mean that there is not a separation of charge. There are often positively charged atoms and negatively charged atoms, but the charges are distributed evenly throughout the molecule.

The bond angles in SF4 found using Spartan were significantly different than than the angles predicted using the VSEPR theory. The bond angles in SF4 with three fluorine atoms in equilateral positions are close to the angles predicted using the VSEPR theory.

The electron density in the NO2/NO2+/NO2- molecules appeared different in each molecule and were consistent with the predictions made using the lewis structures. All of these predictions were close the calculations made using spartan. In the three molecules, NO2+ has the shortest bond length followed by NO2 and NO2- with the longest. NO2+ has the largest angle, followed by NO2, and NO2- has the smallest.

Discussion
The bond lengths and angels in the molecules increase from water, to ammonia, to methane. The bond lengths act in that manner because the bond angles get smaller when there are more lone pairs in the molecule, which repel the other atoms. The lone pairs push the other bonds because they are not localized and take up more space than a bond. The VSEPR theory defines all three of these molecules tetrahedrals and doesn’t take into account the lone pairs, which is why it’s angle predictions are not precise in many cases. The bond lengths are different because the electronegativity of the central atoms are not the same. The greater the electronegativity, the closer it pulls the bonding electron. Electronegativity also the causes these molecules to have electrostatic charges. In some molecules the charges were identical but went in opposite directions and canceled out. The difference in the equilateral and axial bond lengths were due to the atoms positions around the central atom, and the repulsion from the lone pairs. The differences in the bond angles and lengths from spartan and VSEPR demonstrate the importance of considering factors such as repulsion from lone pairs and atom sizes.

References
1. General Chemistry Experiments: A Manual for Chemistry 204, 205, and 206, Southern Oregon University: Ashland OR, Fall 2010

2. Brown, LeMay, Burnsten, Murphy. Chemistry: The Central Science. 11th Edition. Pearson Prentice Hall. Upper Saddle River, NJ. 2009.