Ch.10 - Molecular Shapes & Valence Bond TheoryWorksheetSee all chapters
All Chapters
Ch.1 - Intro to General Chemistry
Ch.2 - Atoms & Elements
Ch.3 - Chemical Reactions
BONUS: Lab Techniques and Procedures
BONUS: Mathematical Operations and Functions
Ch.4 - Chemical Quantities & Aqueous Reactions
Ch.5 - Gases
Ch.6 - Thermochemistry
Ch.7 - Quantum Mechanics
Ch.8 - Periodic Properties of the Elements
Ch.9 - Bonding & Molecular Structure
Ch.10 - Molecular Shapes & Valence Bond Theory
Ch.11 - Liquids, Solids & Intermolecular Forces
Ch.12 - Solutions
Ch.13 - Chemical Kinetics
Ch.14 - Chemical Equilibrium
Ch.15 - Acid and Base Equilibrium
Ch.16 - Aqueous Equilibrium
Ch. 17 - Chemical Thermodynamics
Ch.18 - Electrochemistry
Ch.19 - Nuclear Chemistry
Ch.20 - Organic Chemistry
Ch.22 - Chemistry of the Nonmetals
Ch.23 - Transition Metals and Coordination Compounds

Solution: Use molecular orbital theory to predict whether or not each of the following molecules or ions should exist in a relatively stable form.H2 2 - 

Solution: Use molecular orbital theory to predict whether or not each of the following molecules or ions should exist in a relatively stable form.H2 2 - 

Problem

Use molecular orbital theory to predict whether or not each of the following molecules or ions should exist in a relatively stable form.

H2 2 - 

Solution

We’re being asked to predict whether or not H22- should exist in a relatively stable form


For this, we need to determine the bond order for each species. The bond order tells us the stability of a bond


The higher the bond order, the more electrons holding the atoms together, and therefore the greater the stability


We will do the following steps to solve the problem:

Step 1: Calculate the total number of valence electrons present.

Step 2: Draw the molecular orbital diagram.

Step 3: Calculate the bond order of the molecule/ion. 


Recall that the formula for bond order is:


Bond Order = 12[# of e- in bonding MO -# of e- in antibonding MO]


Solution BlurView Complete Written Solution