VSEPR Theory Study Notes
VSEPR Theory: Student Learning Outcomes
Learn how to draw Lewis structures for atoms that violate the octet rule.
Use Lewis structures and VSEPR to predict the shapes of molecules.
Utilize the shape of a molecule to predict its polarity.
Experimental Goals
The purpose of this lab activity is to:
Predict the VSEPR shape of the molecules.
Determine if the molecules are polar.
Materials
Each group of students will need a molecular modeling kit.
Introduction
Key Information for 3D Models:
Electrons share the same charge; like charges repel each other.
Electron pairs in bonds will orient as far apart as possible from one another in a molecule.
Valence electrons occupy regions of space for even distribution around the central atom.
Chemistry Theories on Bonding:
Various theories lead to predictions of molecular shapes, focusing on:
Equal distribution of electron pairs around a central atom (VSEPR theory).
Modifications of 2D Lewis structures into 3D models.
Start by discussing exceptions to the rules learned in Section 2.2.
VSEPR Theory Overview
2.2 Exceptions: Resonance Structures
Definition:
When one Lewis structure isn’t sufficient, resonance structures demonstrate multiple valid configurations for a molecule.
Example: Ozone (O₃)
Contains 18 valence electrons (3 d7 6).
Central oxygen connected to two terminal oxygens.
Utilizes bonds that leave one lone pair on the central oxygen.
Satisfies octet rule by making double bonds:
Structure 2: Double bond between left O and central O.
Structure 3: Double bond between right O and central O.
Both structures (2 and 3) represent the molecule's configuration, generating resonance structures.
Actual molecule averages these structures, leading to two identical oxygen-oxygen bonds that are 1.5 in nature (not strictly single or double).
Indicates delocalization of charges, stabilizing the molecule.
Comparison with water (H₂O):
Lone pairs on O are localized, unlike the delocalized pairs in resonance structures, which contribute to overall stability in organic chemistry.
Multi-Center Molecules
Rule:
If a double bond can be formed in multiple locations, draw each corresponding resonance structure separated by arrows.
Example: Carbonate ion (CO₃²⁻).
Examples of Molecular Structures:
Ethane (C₂H₆), Ethylene (C₂H₄), Ethanol (CH₃CH₂OH).
Violations of the Octet Rule
Category: Species with fewer or more than eight electrons around central atom:
Electron Deficient Species: Beryllium (Be), Boron (B), Nitrogen (N), Aluminum (Al).
Fewer than eight electrons; reactive (often act as Lewis acids).
Free Radicals: Contain an odd number of valence electrons; very reactive.
Expanded Valence Shells: Found in nonmetals from period 3 or higher (e.g., P, S, Cl), while period 2 elements cannot exceed eight electrons:
Example:
Boron trifluoride (BF₃): 24 valence electrons; octet rule not fully satisfied for B.
Nitric oxide (NO).
Molecular Geometry and Predicting Shapes Using VSEPR Model
Lewis Structures to Shape:
Draw the Lewis structure and count electron groups around the central atom:
Definitions of Electron Groups:
A single bond, double bond, triple bond (count as one group).
Lone pairs.
Unpaired electrons.
Steps to Determine Shape:
Draw Lewis structure and count electron groups surrounding the central atom.
Predict arrangement of groups to maximize distance.
Larger molecules' shapes are composites of the shapes of constituent atoms.
Molecular Shapes Based on Number of Electron Groups
Two Electron Groups
Configuration:
2 bonds, 0 lone pairs
Shape: Linear
Bond Angles: 180°
Three Electron Groups
Configurations:
3 bonds, 0 lone pairs
Shape: Trigonal planar
Bond Angles: 120°
2 bonds, 1 lone pair
Shape: Bent
Bond Angles: < 120°
Note: Lone pairs take up more space, reducing bond angles.
Four Electron Groups
Configurations:
4 bonds, 0 lone pairs
Shape: Tetrahedral
Bond Angles: 109.5°
3 bonds, 1 lone pair
Shape: Trigonal pyramidal
Bond Angles: < 109.5°
2 bonds, 2 lone pairs
Shape: Bent or angular
Bond Angles: < 109.5°
Five Electron Groups
Configurations:
5 bonds, 0 lone pairs
Shape: Trigonal bipyramidal
Bond Angles: 120° (equatorial), 90° (axial)
4 bonds, 1 lone pair
Shape: Seesaw or sawhorse
Bond Angles: <120° (equatorial), <90° (axial)
3 bonds, 2 lone pairs
Shape: T-shaped
Bond Angles: <90°
2 bonds, 3 lone pairs
Shape: Linear
Bond Angles: 180°
Six Electron Groups
Configurations:
6 bonds, 0 lone pairs
Shape: Octahedral
Bond Angles: 90°
5 bonds, 1 lone pair
Shape: Square pyramidal
Bond Angles: <90°
4 bonds, 2 lone pairs
Shape: Square planar
Bond Angles: 90°
Polarity of Molecules
Determining Polarity
Bond Polarity:
Difference in electronegativity allows categorizing bonds as polar (0.5 < difference < 1.69) or nonpolar (difference < 0.5).
Unequal sharing of electrons leads to partial charges (b1).
Example: In HCl, Cl has a higher electronegativity and thus a slight negative charge (δ-), while H acquires a slight positive charge (δ+).
Molecular Shape and Overall Polarity
In diatomic molecules:
The bond polarity determines molecular polarity (polar bond = polar molecule, nonpolar bond = nonpolar molecule).
In polyatomic molecules:
Shape and bond polarity contribute to overall polarity.
If polar bonds cancel due to symmetrical arrangement (like CO₂), the overall molecule is nonpolar.
Example of Water (H₂O): Polar due to bent shape, creating a dipole.
Examples of Polarity Based on Structure
Nonpolar molecules include those with symmetrical arrangements.
Examples:
Nonpolar: Carbon tetrachloride (CCl₄)
Polar: Chloroform (CHCl₃), Methyl chloride (CH₃Cl)
Conclusion: Putting It All Together
Understand and visualize shapes and polarities for several chemical structures (E.g., Phosgene COCl₂, Ethanol CH₃CH₂OH).
Assignment/Lab Activity
Use the molecular modeling kit to build specified molecules. Create a table defining:
Shape
Number of Bonding Pairs
Number of Lone Electron Pairs
Molecular Geometry
Examples with polar character defined.