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Use molecular orbital theory to predict if each molecule or ion exists in a relatively stable form. a. F22–
Ch.11 - Chemical Bonding II: Molecular Shapes, VSEPR & MO Theory
Tro6th EditionChemistry: A Molecular ApproachISBN: 9780137832217
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TOC.tableOfContents
- Ch.1 - Matter, Measurement & Problem Solving118
- Ch.2 - Atoms & Elements106
- Ch.3 - Molecules and Compounds131
- Ch.4 - Chemical Reactions and Chemical Quantities56
- Ch.5 - Introduction to Solutions and Aqueous Solutions82
- Ch.6 - Gases114
- Ch.7 - Thermochemistry96
- Ch.8 - The Quantum-Mechanical Model of the Atom60
- Ch.9 - Periodic Properties of the Elements90
- Ch.10 - Chemical Bonding I: The Lewis Model89
- Ch.11 - Chemical Bonding II: Molecular Shapes, VSEPR & MO Theory85
- Ch.12 - Liquids, Solids & Intermolecular Forces51
- Ch.13 - Solids & Modern Materials47
- Ch.14 - Solutions90
- Ch.15 - Chemical Kinetics115
- Ch.16 - Chemical Equilibrium68
- Ch.17 - Acids and Bases136
- Ch.18 - Aqueous Ionic Equilibrium153
- Ch.19 - Free Energy & Thermodynamics93
- Ch.20 - Electrochemistry106
- Ch.21 - Radioactivity & Nuclear Chemistry74
- Ch.22 - Organic Chemistry140
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Tro 6th EditionCh.11 - Chemical Bonding II: Molecular Shapes, VSEPR & MO TheoryProblema 81
Tro 6th EditionCh.11 - Chemical Bonding II: Molecular Shapes, VSEPR & MO TheoryProblema 81Capitolo 11, Problema 81
Use molecular orbital theory to predict if each molecule or ion exists in a relatively stable form. a. Li22+ b. Li2 c. Be22+ d. C22+
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Identify the number of electrons in each molecule or ion. For example, Li has 3 electrons, so Li2 would have 6 electrons, and Li2^2+ would have 4 electrons.
Write the molecular orbital (MO) configuration for each species. For diatomic molecules, the order is typically: \( \sigma_{1s}^2, \sigma_{1s}^*^2, \sigma_{2s}^2, \sigma_{2s}^*^2, \pi_{2p}^4, \sigma_{2p}^2, \pi_{2p}^*^4, \sigma_{2p}^*^2 \).
Fill the molecular orbitals with the electrons from each species, starting from the lowest energy level (\( \sigma_{1s} \)) and moving to higher levels.
Calculate the bond order using the formula: \( \text{Bond Order} = \frac{(\text{Number of electrons in bonding MOs} - \text{Number of electrons in antibonding MOs})}{2} \).
Determine the stability of each molecule or ion. A positive bond order indicates a stable molecule, while a bond order of zero or negative suggests instability.
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Molecular Orbital Theory
Molecular Orbital Theory (MOT) describes the behavior of electrons in molecules by combining atomic orbitals to form molecular orbitals. These orbitals can be bonding, antibonding, or non-bonding, and the distribution of electrons among them determines the stability of the molecule. A molecule is generally stable if it has more electrons in bonding orbitals than in antibonding orbitals.
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03:06
Molecular Orbital Theory
Bond Order
Bond order is a measure of the number of chemical bonds between a pair of atoms, calculated as the difference between the number of bonding and antibonding electrons divided by two. A higher bond order indicates a more stable molecule, as it suggests a greater number of bonding interactions. For example, a bond order of 1 corresponds to a single bond, while a bond order of 2 indicates a double bond.
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00:36Average Bond Order
Electron Configuration in Diatomic Molecules
The electron configuration of diatomic molecules involves filling molecular orbitals according to the Aufbau principle, Hund's rule, and the Pauli exclusion principle. Understanding how electrons fill these orbitals helps predict the stability of the molecule. For instance, in diatomic species like Li2 or C2, the arrangement of electrons in bonding and antibonding orbitals directly influences whether the molecule is stable or not.
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03:10Homonuclear Diatomic Molecules
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