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Molecular geometry, also known as the molecular structure, is the three-dimensional structure or arrangement of atoms in a molecule. Understanding the molecular structure of a compound can help determine the polarity, reactivity, phase of matter, color, magnetism, as well as the biological activity.

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Introduction

To determine the shapes of molecules, we must become acquainted with the Lewis electron dot structure. Although the Lewis theory does not determine the shapes of molecules, it is the first step in predicting shapes of molecules. The Lewis structure helps us identify the bond pairs and the lone pairs. Then, with the Lewis structure, we apply the valence-shell electron-pair repulsion (VSPER) theory to determine the molecular geometry and the electron-group geometry.

To identify and have a complete description of the three-dimensional shape of a molecule, we need to know also learn about state the bond angle as well. Lewis Electron Dot Structures play crucial role in determining the geometry of molecules because it helps us identify the valence electrons. To learn how to draw a Lewis electron dot structure click the link above.


Valence-Shell Electron-Pair Repulsion Theory

Now that we have a background in the Lewis electron dot structure we can use it to locate the the valence electrons of the center atom. The valence-shell electron-pair repulsion (VSEPR) theory states that electron pairs repel each other whether or not they are in bond pairs or in lone pairs. Thus, electron pairs will spread themselves as far from each other as possible to minimize repulsion. VSEPR focuses not only on electron pairs, but it also focus on electron groups as a whole. An electron group can be an electron pair, a lone pair, a single unpaired electron, a double bond or a triple bond on the center atom. Using the VSEPR theory, the electron bond pairs and lone pairs on the center atom will help us predict the shape of a molecule.

The shape of a molecule is determined by the location of the nuclei and its electrons. The electrons and the nuclei settle into positions that minimize repulsion and maximize attraction. Thus, the molecule"s shape reflects its equilibrium state in which it has the lowest possible energy in the system. Although VSEPR theory predicts the distribution of the electrons, we have to take in consideration of the actual determinant of the molecular shape. We separate this into two categories, the electron-group geometry and the molecular geometry.

Electron-group geometry is determined by the number of electron groups.

Number of electron groupsName of electron group geometry
2linear
3trigonal-planar
4tetrahedral
5trigonal-bipyramidal
6octahedral

Molecular geometry, on the other hand, depends on not only on the number of electron groups, but also on the number of lone pairs. When the electron groups are all bond pairs, they are named exactly like the electron-group geometry. See the chart below for more information on how they are named depending on the number of lone pairs the molecule has.


VSEPR Notation

As stated above, molecular geometry and electron-group geometry are the same when there are no lone pairs. The VSEPR notation for these molecules are AXn. "A" represents the central atom and n represents the number of bonds with the central atom. When lone pairs are present, the letter Ex is added. The x represents the number of lone pairs present in the molecule. For example, a molecule with two bond pairs and two lone pairs would have this notation: AX2E2.

Geometry of Molecules ChartNumber of Electron GroupsElectron-Group GeometryNumber of Lone PairsVSEPR NotationMolecular GeometryIdeal Bond AnglesExamples
2 linear1AX2
*
180° BeH2
3 trigonal-planar0AX3
*
120° CO32-
1AX2E
*
120°O3
4 tetrahedral0AX4

*

Tetrahedral

109.5° S042-
1AX3E
*
109.5° H3O+
2AX2E2
*
109.5° H2O
5 trigonal-bipyramidal0 AX5
*
90°, 120° PF5
1 AX4Eb
*
90°, 120° TeCl4
2 AX3E2
*
90°ClF3
3 AX2E3
*

octahedral

90° PF6-
1 AX5E
*
90°SbCl52-
2 AX4E2
*


Dipole Moments

A molecule is polar when the electrons are not distributed equally and the molecule has two poles. The more electronegative end of the molecule is the negative end and the less electronegative end is the positive end. A common example is HCl. Using the capital sigma + or - as a symbol to show the the positive end and the negative end we can draw the net dipole. So sigma + would be on the hydrogen atom and sigma - would be on the Chlorine atom. Using the cross bow arrow shown below we can show that it has a net dipole. The net dipole is the measurable, which is called the dipole moment. Dipole moment is equal to the product of the partial charge and the distance. The equation for dipole moment is as follows.

\< \mu = \delta \times d\>

with

µ = dipole moment (debye) δ = partial charge (C)d = distance (m)

The units for dipole is expressed in debye which is also known as Coulombs x meter (C x m)

Example of a Dipole

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On the cross-base arrow, the cross represents the positive charge and the arrow represents the negative charge.

Here"s another way to determine dipole moments. We need to comprehend electronegativity which is abbreviated EN. What is EN? Well, EN is how much an element really wants an electron. Think about basketball and how two players pass the ball to each other. Each player represent an element and the ball represents the electron. Let"s say one player is a ball hog. The player that is the ball hog is more electronegative because he or she wants the ball more.

Here is a link that has all the EN listed: www.green-planet-solar-energy...electroneg.gif

What if we are not given EN? Luckily, there is a trend in the periodic table for EN. From bottom to the top, EN will increase. From left to right, EN will increase. The most electronegative element is Flourine with 4.0.

Now, we are ready to apply EN to determine whether or not molecules are polar. We look back at the picture of H2O above. The EN is given. What do we do with all the EN? We compare the EN between each bond. Oxygen has a greater EN than Hydrogen. Therefore, we can draw a cross bow arrow towards Oxygen. We have two arrows because Oxygen is bonded to two Hydrogens. Since both arrows point toward Oxygen, we can say that there is a net EN. We added the arrows that point to Oxygen and we end up with a new, bigger arrow. This is examplified in the picture above. If arrows are drawn away from each other like , then we are more likely to have no net EN because the molecule is symmetrical. Refer back to the Lewis dot diagram of CO2. The shape is linear and the EN arrows point towards Oxygen. The arrows are opposite of each other and have the same EN difference. Therefore, we have no net charge and the molecule is non-polar.

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Total # of electrons: 1+(3x6)+7=26electron group geometry: tetrahedralmolecular: trigonal pyramidalideal angle: 109.5°polar, has a dipole moment

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