Chemistry 310 - Inorganic Chemistry - Spring 2015

Instructor: Tom Mallouk Office hours TR 1:30-2:30 PM 205 S. Frear

LabCell phone: 814-571-6115 Office phone: 814-863-9637 Course website: http://courses.chem.psu.edu/chem310 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 chemistryAcid-base and redox chemistrySolid state and materials chemistryNanomaterials and applications

 Chem 412: (Fall semesters only)

Symmetry and group theoryInorganic spectroscopyOrganometallic chemistryCatalysis and reaction mechanismsBio-inorganic chemistry

Big-picture:

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

Chapter 1 - Valence Bond Theory

Perspective

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)

NH3

Formal charge

NO3−

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

BF3

ONF

Resonance structures

In reality, there is often a combination of two or more resonance structures

Ozone

Ozone (continued)

What is the real structure?How would you study this experimentally?

Nitrate

Cyanate

ONF3

You perform a diffraction experiment to determine the crystal structure of ONF3 and find the following bond lengths:

N-F: 143.1 pmN-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 repulsions

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)

SO2

H2O

CO2

SF4

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

BrF4–

(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

XeF2

Bond strength / bond length correlation

What would you predict to be the correlation?

Bond strength / bond length correlation

Examples: carbon-carbon bonds

Anomalies

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)

X E(X-X) (kcal/mol)

F 38Cl 58Br 46I 36

Bond strength / bond length correlation

How is this manifested in the properties ofF2 vs. the other diatomic halogens?

https://www.youtube.com/watch?v=d2zP4jxTL4s

https://www.youtube.com/watch?v=DjbRH6bSIIw

https://www.youtube.com/watch?v=gFGwY_S4j_M

Related phenomena

C vs. Si

N vs. P

Polarity

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)

Exercise

Draw Lewis structures for the following molecules and compare/contrast

CO2

N2O

Allene (H2CCCH2)

Azide (N3–)

Which one of these is not like the others?

BF3

NF3

COF2

NO3–

CO32–

Isoelectronic principle (solids)

The isoelectronic principle also extends to solid-state materials

Al Si P S

Ga Ge As Se

In Sn Sb Te

Cd

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

catalysis?