Medium VSEPR Geometry Practice Questions

Concept Explanation

VSEPR theory, or Valence Shell Electron Pair Repulsion theory, is a model used in chemistry to predict the three-dimensional geometry of individual molecules based on the electrostatic repulsion between electron pairs surrounding a central atom. The fundamental premise is that electron pairs, whether in chemical bonds or as lone pairs, naturally arrange themselves as far apart as possible to minimize repulsion, thereby determining the molecule's shape. To apply this theory, one must first determine the Lewis Structure of the molecule to count the number of bonding regions and lone pairs, collectively known as the steric number.

When analyzing VSEPR geometry, it is vital to distinguish between electron geometry (the arrangement of all electron groups) and molecular geometry (the arrangement of only the atoms). For instance, a molecule with four electron groups will always have a tetrahedral electron geometry, but if one of those groups is a lone pair, the molecular geometry is described as trigonal pyramidal. According to Wikipedia's overview of VSEPR, the strength of repulsion follows a specific hierarchy: Lone Pair-Lone Pair > Lone Pair-Bonding Pair > Bonding Pair-Bonding Pair. This hierarchy explains why lone pairs often compress the bond angles between adjacent atoms, deviating from the ideal angles of 109.5°, 120°, or 180°.

Steric Number Lone Pairs Molecular Geometry Ideal Bond Angle 3 1 Bent < 120° 4 2 Bent < 109.5° 5 1 Seesaw 90°, 120° 6 2 Square Planar 90°

Solved Examples

The following examples demonstrate how to move from a chemical formula to a specific VSEPR shape by evaluating steric numbers and lone pair counts.

1. Determine the molecular geometry of Sulfur Tetrafluoride (SF4).

   1. Count valence electrons: S (6) + 4F (4 × 7) = 34 electrons.

   2. Draw the Lewis structure: S is the central atom with 4 single bonds to F. This uses 32 electrons. The remaining 2 electrons form a lone pair on S.

   3. Calculate steric number: 4 bonding pairs + 1 lone pair = 5.

   4. Determine geometry: A steric number of 5 corresponds to a trigonal bipyramidal electron geometry. With one lone pair (occupying an equatorial position), the molecular geometry is Seesaw.

2. Determine the molecular geometry of the Triiodide ion (I3⁻).

   1. Count valence electrons: 3I (3 × 7) + 1 (charge) = 22 electrons.

   2. Draw the Lewis structure: A central Iodine connected to two other Iodine atoms. This uses 16 electrons for octets on outer atoms. The remaining 6 electrons are placed as 3 lone pairs on the central Iodine.

   3. Calculate steric number: 2 bonding pairs + 3 lone pairs = 5.

   4. Determine geometry: Trigonal bipyramidal electron geometry. Three lone pairs occupy the equatorial positions to minimize repulsion, leaving the atoms in a Linear molecular geometry.

3. Determine the molecular geometry of Xenon Tetrafluoride (XeF4).

   1. Count valence electrons: Xe (8) + 4F (4 × 7) = 36 electrons.

   2. Draw the Lewis structure: Xe is central with 4 bonds to F. After completing F octets, 4 electrons (2 lone pairs) remain for the Xe atom.

   3. Calculate steric number: 4 bonding pairs + 2 lone pairs = 6.

   4. Determine geometry: Octahedral electron geometry. Two lone pairs stay 180° apart to minimize repulsion, resulting in a Square Planar molecular geometry.

Practice Questions

Test your understanding of VSEPR geometry with these medium-level problems. You may need to reference hybridization concepts to fully understand the orbital arrangements.

1. Predict the molecular geometry and bond angles for the chlorite ion (ClO2⁻).

2. What is the molecular geometry of Phosphorus Pentachloride (PCl5)?

3. Identify the molecular geometry of Bromine Trifluoride (BrF3).

4. Compare the bond angles in H2O and NH3. Which is smaller and why?

5. Determine the geometry of Carbonyl Sulfide (OCS). Note: Carbon is the central atom.

6. What is the molecular geometry of the Carbonate ion (CO3²⁻)?

7. Predict the shape of Selenium Hexafluoride (SeF6).

8. Determine the geometry of the Hydronium ion (H3O⁺).

9. Analyze the molecule ClF5. What is its molecular geometry and electron geometry?

10. For the molecule SO2, determine the electron geometry and the approximate bond angle.

Answers & Explanations

1. ClO2⁻: Bent. The central Chlorine has 2 bonding pairs and 2 lone pairs (steric number 4). The electron geometry is tetrahedral, but the presence of two lone pairs makes the molecular shape bent with an angle < 109.5°.

2. PCl5: Trigonal Bipyramidal. Phosphorus has 5 valence electrons and bonds with 5 Chlorine atoms. With no lone pairs, the molecular geometry matches the electron geometry, featuring 90° and 120° angles.

3. BrF3: T-shaped. Bromine has 7 valence electrons. Three are used for bonding, leaving 2 lone pairs. Steric number = 5. To minimize repulsion, lone pairs occupy equatorial positions, resulting in a T-shape.

4. H2O is smaller. Both have tetrahedral electron geometry. However, H2O has 2 lone pairs while NH3 has only 1. Since lone pairs exert more repulsion than bonding pairs, the two lone pairs in water compress the H-O-H angle (~104.5°) more than the single lone pair in ammonia (~107°).

5. OCS: Linear. Carbon is double-bonded to Oxygen and double-bonded to Sulfur. With 2 bonding regions and 0 lone pairs on the central carbon, the angle is 180°.

6. CO3²⁻: Trigonal Planar. The central Carbon has three resonance structures where it is bonded to three Oxygen atoms. With a steric number of 3 and no lone pairs on Carbon, the shape is trigonal planar with 120° angles.

7. SeF6: Octahedral. Selenium has 6 valence electrons, all of which are used to bond with 6 Fluorine atoms. Steric number 6 with 0 lone pairs results in a perfect octahedral shape.

8. H3O⁺: Trigonal Pyramidal. Oxygen has 6 valence electrons, but the positive charge removes one (total 5). Three are used for bonds, leaving 1 lone pair. Steric number 4 with 1 lone pair gives a trigonal pyramidal shape.

9. ClF5: Square Pyramidal. Chlorine has 5 bonding pairs and 1 lone pair (steric number 6). The electron geometry is octahedral, and the molecular geometry is square pyramidal.

10. SO2: Trigonal Planar (Electron) / Bent (Molecular). Sulfur has 2 double bonds and 1 lone pair. The steric number is 3. The lone pair pushes the bonds together, resulting in a bent shape with an angle of approximately 117°.

Quick Quiz

Frequently Asked Questions

What is the difference between electron geometry and molecular geometry?

Electron geometry considers the spatial arrangement of all electron groups (bonds and lone pairs) around a central atom. Molecular geometry specifically describes the positions of the nuclei (atoms) only, though it is still influenced by the presence of lone pairs.

Why do lone pairs reduce bond angles?

Lone pairs are held closer to the nucleus than bonding pairs and occupy more space. This increased volume causes them to exert greater repulsive forces on neighboring electron pairs, pushing the bonding pairs closer together and reducing the bond angles.

Can VSEPR theory predict the polarity of a molecule?

Yes, by determining the molecular geometry, you can see if the individual bond dipoles cancel out or combine. You can practice this further with polarity determination practice questions to see how shape influences molecular dipoles.

What is a steric number in VSEPR?

The steric number is the sum of the number of atoms bonded to a central atom and the number of lone pairs on that central atom. It determines the base electron geometry from which the molecular shape is derived.

Does VSEPR apply to transition metal complexes?

VSEPR is most effective for main-group elements. While it provides a basic framework, transition metal complexes are better described by Crystal Field Theory or Ligand Field Theory due to the involvement of d-orbitals.

How do multiple bonds affect VSEPR shapes?

In VSEPR theory, double and triple bonds are treated as a single "electron group" or bonding region. However, the high electron density in multiple bonds can exert slightly more repulsion than a single bond, occasionally causing minor deviations in bond angles.

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