|Modules||Bonding and Molecules|
|Concepts||Lewis structures: octet stability, partial charge, bonding and nonbonding electrons, electronegativity: polar bonds and polar molecules, ionic character of covalent bonds, Coulomb’s law|
|Keywords||cation, anion, Madelung constant, enthalpy, valence electron, Gilbert Lewis, ionization, isoelectronic, metal, nonmetal, ionic bond, electron transfer, electron sharing, covalent bond, percent ionic character, homonuclear bond, heteronuclear bond, triple bond, dative bond, s and p orbitals, Lewis structures, Linus Pauling, hybrid orbital, crystallization energy, bond energy, charge displacement, dipole moment, polar covalency, electronegativity, polar bond, polar molecule|
|Chemical Substances||sodium (Na), chloride (Cl), nitrogen (N), oxygen (O), lithium (Li), beryllium (Be), magnesium (Mg), aluminum (Al), silicon (Si), hydrogen (H), helium (He), sulfuryl chloride (SO2Cl2), methane (CH4), magnesium chloride (MgCl2), hydrogen fluoride (HF), hydrogen chloride (HCl), sodium chloride (NaCl), Freon-12|
|Applications||capacitors, refrigerant, compressor design|
Before starting this session, you should be familiar with:
- Hybridized bonding in molecules, VSEPR, properties of covalent bond, electron domain theory (Session 8)
Prof. Sadoway discusses hybridized and molecular orbitals along with paramagnetism (Session 10).
After completing this session, you should be able to:
- Sketch the Lewis structure for a given compound.
- Explain how octet stability is satisfied by electron sharing and electron transfer.
- Understand how electron states can be mixed to form hybrid orbitals.
- Define electronegativity and dipole moment.
- Calculate the percent ionic character of a heteronuclear bond.
- Explain how polar bonds may be present in polar and nonpolar molecules.
Archived Lecture Notes #2 (PDF), Section 3
|[Saylor] 7.3, “Energetics of Ion Formation.”||Ionization energies; electron affinities; electronegativity|
|[Saylor] 8.4, “Introduction to Lewis Dot Structures.”||Creating a Lewis dot symbol; the octet rule|
|[Saylor] 8.5, “Lewis Structures and Covalent Bonding.”||Using electron structures to describe covalent bonding; using Lewis electron structures to explain stoichiometry; using formal charges to distinguish between Lewis structures; resonance structures|
|[Saylor] 8.6, “Exceptions to the Octet Rule.”||Odd number of electrons; more than an octet of electrons; fewer than an octet of electrons|
|[Saylor] 8.8, “Properties of Covalent Bonds.”||Bond order; the relationship between bond order and bond energy; the relationship between molecular structure and bond energy|
|[Saylor] 8.9, “Polar Covalent Bonds.”||Bond polarity; dipole moments|
Prof. Sadoway discusses the following concepts:
- Problems with ionic bonding for diatomic molecules
- G. N. Lewis – shell filling by electron sharing
- Lewis dot notation
- Cooperative use of valence electrons to achieve octet stability = covalent bonds
- Ionic bond = electron transfer
- Covalent bond = electron sharing (directional)
- s-orbitals “merge” with p-orbitals – sp3 hybridized
- Results in 4 unpaired electrons, ready to bond
- Energy of heteronuclear bonds
- Percent ionic character
- Polar bonding
|[Saylor] 7.3, “Energetics of Ion Formation.”||1, 10||none|
|[Saylor] 8.2, “Ionic Bonding.”||none||3|
|[Saylor] 8.6, “Exceptions to the Octet Rule.”||none||1, 7|
|[Saylor] 8.9, “Polar Covalent Bonds.”||none||2, 6, 7, 8|
|[Saylor] 9.1, “Predicting the Geometry of Molecules and Polyatomic Ions.”||none||7|
For Further Study
Molina, Mario J. and Rowland, F. S. “Stratospheric sink for chlorofluoromethanes: chlorine atomc-atalysed destruction of ozone.” Nature 249 (June 28, 1974): 810-812.
Other OCW and OER Content
|5.111 Principles of Chemical Science||MIT OpenCourseWare||Undergraduate (first-year)|