Chapter 1 - Review of chemical bonding

Chemistry 310 - Inorganic Chemistry - Spring 2015
Instructor: Tom Mallouk
Cell phone: 814-571-6115
Office hours TR 1:30-2:30 PM
205 S. Frear Lab
Office phone: 814-863-9637
Course website:
Text: Chem 310 Wikibook (Introduction to Inorganic Chemistry)
Other textbooks (recommended, not required):
Shriver & Atkins Inorganic Chemistry, 5th Ed (used in 412)
Miessler & Tarr, Inorganic Chemistry
Huheey, Inorganic Chemistry
Chem 310 is part of a two-semester sequence in
inorganic chemistry:
Chem 310: (Fall and Spring semesters)
Bonding models, coordination chemistry
Acid-base and redox chemistry
Solid state and materials chemistry
Nanomaterials and applications
Chem 412: (Fall semesters only)
Symmetry and group theory
Inorganic spectroscopy
Organometallic chemistry
Catalysis and reaction mechanisms
Bio-inorganic chemistry
Chapter 1 - Valence Bond Theory
When atoms combine to form molecules, their properties change. The structure of a
molecule and details of the bonding between its constituent atoms – which can be highly
interrelated – directly impact its properties and applications. It is therefore important to
understand molecular structure and bonding.
Learning goals:
• Be able to draw Lewis dot structures, assign formal charges, predict molecular
geometries (including bond angles), and calculate bond orders for molecules, including
hypervalent molecules and ions.
• Describe hypervalent molecules using no-bond resonance.
• Understand and articulate how predictions of molecular structure and bonding can be
experimentally verified.
• Use the isoelectronic principle to design new molecules and solids.
• Rationalize bond strength and chemical reactivity using bond polarity arguments.
• Interrelate bond length and bond strength.
Textbook Chapter 1
Conceptually parallels atomic structure – begin with simplest, usually
empirical guidelines, then increase the degree of sophistication and the
quantitative nature in order to more fully account for all observations
Fully explain what we are able to observe  predictions of new “stuff”
Historical development of understanding and describing bonding in
molecules – from more simple explanations to more sophisticated ones
Many of these concepts will be familiar – we will likely introduce them from
a different perspective, though, and for a more diverse range of systems.
We will also dive deeper into “reality” in these topics and systems than you
probably have previously.
Lewis structures
G.N. Lewis: The existence of many molecules can be rationalized by the
octet rule, where atoms share electrons so that they have eight valence
electrons – based on s2p6 configuration
Brief review of drawing Lewis structures
Formal charge
Formal charge (distinct from oxidation state or actual charge on the atom):
divide bonds equally between atoms, all formal charges add up to the total
charge on the molecule or ion
(Accounts for charge distribution within the molecule)
Formal charge
Formal charges must add up to the total charge on the molecule!
Formal charges gone bad…?
Ideal formal charges:
Measure charge distribution on ions or molecules – how?
Real charge distribution – closer to formal charge than oxidation state
Formal charge gives a sense of electron distribution in a molecule
Formal charge helps “check” whether a proposed structure is reasonable
Resonance structures
In reality, there is often a combination of two or more resonance structures
Ozone (continued)
What is the real structure?
How would you study this experimentally?
You perform a diffraction experiment to determine the crystal structure of
ONF3 and find the following bond lengths:
N-F: 143.1 pm
N-O: 115.8 pm
What does this suggest about the bonding in ONF3?
BH3 + CO
How might you probe this experimentally?
(How would you verify or refute the prediction from valence bond theory)
Hypervalent compounds: I3–
What would you predict the structure of I3– to be?
How would you verify or refute this prediction experimentally?
Hypervalent compounds: XeF2
What would you predict the structure of XeF2 to be?
How would you verify or refute this prediction experimentally?
Hypervalent compounds: SF6
What would you predict the structure of SF6 to be?
Hypervalent compounds: SF6 (and related)
How can we rationalize these structures?
Molecular shapes
The three-dimensional shapes that molecules adopt are central to their
function. Consider the role of cis-platin, for example:
Unfortunately it can be difficult to precisely predict molecular shapes, but a
very simple formalism – VSEPR theory – provides zero-order “guesses” that
are surprisingly useful
VSEPR theory (Review)
Valence Shell Electron Pair Repulsion theory uses electron-pair repulsions
to rationalize (and, at a zero-order, to predict) molecular shapes
1. Lone pairs and atoms/ligands attached to the central atom determine
the coordination number (“CN”)
2. Electron pairs and atoms orient themselves in 3D space to minimize
VSEPR theory
VSEPR theory
Using VSEPR theory
Electronic shape – determined by the 3D geometry of atoms and lone pairs
(both “containing electrons”) attached to the central atom
Molecular shape – same (subset of electronic shape), but ignore the lone
pairs (just the geometry defined by the actual attached atoms)
What is the shape of SF4?
VSEPR (coordination #5)
How do we
rationalize this?
Is there an
alternative option?
VSEPR lone pair / bonding pair hierarchy
Consider the various molecular geometry options for ClF3
(a) Draw BrF4– as an octet structure
(b) Draw BrF4– as it would be experimentally observed
(c) For (b), assign all bond orders, formal charges, and bond angles
(d) What is the electronic geometry of BrF4–?
(e) What is the molecular geometry of BrF4–?
If you finish early, do the same for XeF2
Bond strength / bond length correlation
What would you predict to be the correlation?
Bond strength / bond length correlation
Examples: carbon-carbon bonds
Some bond lengths and bond energies are anomalous, and these anomalies
are key contributors to the unique chemistries of such compounds
Example: the fluorine-fluorine bond length in F2 is 1.43 Å
Anomalies (F-F)
E(X-X) (kcal/mol)
Bond strength / bond length correlation
How is this manifested in the properties of
F2 vs. the other diatomic halogens?
Related phenomena
C vs. Si
N vs. P
One consequence of the way in which atoms are arranged around a central
atom is polarity – dipoles that depend on electronegativity differences
between atoms in asymmetric structures
Red = negative
blue = positive
(From Wikipedia, “Chemical Polarity”)
Electronegativity and bond strength
Polar bonds have higher bond energies than nonpolar bonds
Nonpolar, polar, ionic continuum
Electronegativity and bond strength
Bond polarity (from Δχ) helps us to understand reactivity
Example: Si-H vs. C-H
Si-H vs. C-H (continued)
Si-H is more hydride-like
Electrophilic substitution is more prevalent on Si-H than on C-H
Silanes react with strong acids to produce H2
Al-H is more reactive than Si-H (follows trend)
Draw Lewis structures for the following molecules and compare/contrast
N 2O
Allene (H2CCCH2)
Azide (N3–)
Which one of these is not like the others?
Isoelectronic principle (solids)
The isoelectronic principle also extends to solid-state materials
We can use this principle to design binary semiconductors that mimic Si
and Ge, but that have different band gaps (more on that later)
Isoelectronic principle (application)
Consider SiO2 … How could we use the isoelectronic principle to design new
solid-state materials that are related to SiO2, but that exhibit acidic sites for

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