Chem1 Chemical Bonding
The twenty lessons in the Chemical Bonding set cover most
of the aspects of this subject that are usually taught in first-year college
and university Chemistry courses. The intention, as in all of the Chem1
lessons, is to get students actively engaged in thinking about the basic
concepts by asking questions and forcing them to make decisions. An important
theme that runs through all of these lessons is that our working theories
of chemical bonding are based on models which have varying degrees of applicability
to different kinds of systems. Although no special prerequisites are assumed,
some previous exposure to the fundamentals of quantum mechanics, particularly
the Uncertainty Principle, would be useful.
The lessons are divided into four groups, each of which can be used independently.
The lessons in each group are selected by a menu. The menu in the first
group allows access to the other groups.
1 - Introduction to Chemical Bonding.
- Introduction. Observable properties of chemical bonds. What
is a chemical bond? Kinetic and potential energy; potential energy curves
and stability in relation to thermal energy; bond energy, bond length,
and stretching frequencies; infrared absorption.
- Basic theory of the shared-electron bond. The H2+ molecule;
binding and antibinding effects of the electron depending on its location;
review of the Uncertainty Principle; bonding as a consequence of electron
tunnelling between atoms (see S. Nordholm, J. Chem. Education 65(7) 581-584
- Bond polarity and dipole moments. Consequences of asymmetric
electron distri-bution in HF. Definition of the dipole moment; simple calculations.
Bond polarity and dipole moments in CO2, H2O and CH4. Comparison of dipole
moments in CH3Cl, CH2Cl2, and CHCl3.
- Polar covalence. The Pauling electronegativity scale; electronegativity
trends in the Periodic Table. Contrast between electronegativity and electron
affinity. Using a table of electronegativities to predict bond polarities.
Periodic trends in the halogen halides and in CH4, NH3, H2O. Highly-polar
and "ionic" bonds; ionic solids.
- The Octet Rule and Lewis electron-dot structures. Constructing
a theory of valence for C, N, O, and F. The Octet Rule and noble gas structures.
The Octet Rule as a natural consequence of energy gaps between p- and d-block
elements. Applications: predicting formulas of molecules and monatomic
ions. Bonding and non-bonding electron pairs. Examples: Cl2, H2S, CH3Cl.
- Multiple bonds and resonance structures. Electron-dot structures
of CO2, CO, C2H6, C2H4, and C2H2. Comparison of C-C bond lengths in the
latter three compounds. Equivalent structures and bond orders in SO2, and
SO3; comparison with SO32-. Bond order in the NO3- ion.
- Formal charge. Formal charges in NH4+, H2O. Assessing the relative
importance of different bonding arrangements in NO and SCN-.
2 - Molecular shapes
- Introduction to VSEPR theory. Orbital repulsion, axial and equatorial
positions in octahedral structures.
- Tetrahedral coordination. Distinction between coordination and
molecular geometry when non-bonding electrons are present.
- Higher coordinations. Octahedral and trigonal bipyramidal coordination
- Drill exercise. Ten examples randomly selected from 13, covering
all major shapes.
3 - Hybrid Atomic Orbital model of chemical bonding
Although hybridization no longer gets the play it once did in textbooks,
it is still a concept that students are expected to be familiar with in
subsequent courses. These lessons might also serve as a review for students
beginning courses in organic chemistry.
- Introduction. How to think about hybrid orbitals. Inadequacy
of s- and p- atomic orbitals in explaining bonding in CH4; sp3 hybrids
as a more useful set of equivalent orbitals. Analogy with standing waves;
in-phase and out-of-phase combinations of s- and p-orbitals; shapes of
the resulting hybrids.
- Single bonds. General principles of constructing and predicting
shapes of hybrid orbitals. CH4, BF3, BeH2.
- Multiple bonds. Ethane, ethylene and acetylene; role of atomic-p
orbitals; s-and p- orbitals. "Bent-bond" model as an alternative
approach to multiple bonding. Hybrid orbital description of the nitrate
- Hybrids involving d-orbitals. Nonbonding electrons in
hybrid orbitals; examples of NH3, H2O, NO3-, [PtCl4]2-, PCl3 and PCl5,
SF6, Zn(NH3)6++ and coordinate covalent bonds; configuration of Fe(III);
inner- and outer- transition metal complexes. The magnetic balance.
4 - Molecular Orbitals
This group of six lessons covers the basics of the molecular orbital
model at the level commonly required in first-year courses. The major objective
is to get the student to be able to predict the bond orders in diatomic
molecules of the first- and second-row elements. An introductory lesson
develops the basic idea of the splitting of atomic orbitals into bonding
and antibonding pairs. Two final lessons show how this basic idea can be
applied to transition metal complexes and to metals and semiconductors,
without attempting to treat either of these large subjects thoroughly.
- Introduction. Attractive and repulsive forces in the simplest
molecule, H2+. Identifying in- and out-of-phase combinations of hydrogen
1s wave functions, and bonding/antibonding orbitals. Simple m.o.'s from
atomic-s orbitals. Bond orders and relative bond energies of H2 and H2+.
Working out orbital occupancies and bond orders in He2, He2+, Li2, Be2,
LiH. Review based on Li and Be.
- Survey of second-row homonuclear diatomics. B2, C2, N2 and its
cation and anion, O2, and F2.
- Transition metal ions. Shapes of d orbitals; identifying those
most strongly affected by an octahedral field. Consequences of ligand field
- Metals and semiconductors. Qualitative introduction to band
theory of metals and semiconductors. Molecular orbitals in Li2, Li3, and
Li4; splitting into filled and unfilled levels; extension to Lin. Mechanism
of electrical and thermal conduction by excitation of electrons to contiguous
unoccupied levels. Overlap of 2s and 2p bands in Group 2 metals, development
of band gap within periods; semiconductors. Temperature coefficient of
conductivity in semiconductors; insulators.
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