Ch. 9 Molecular Geometry & Bonding Theories Lewis structures tell us which atoms are bonded together, but we will now explore the geometric shapes of these molecules. –Overall shape is determined by bond angles. –Bond angles are determined by the VSEPR theory. Electrons repel & will try to get as far away from each other as possible Nonbonded electron pairs take up more space than bonded electrons. You must determine the # of electron domains on the central atom. –An electron domain is a region of electrons that are either bonded or non-bonded (lone pairs). A double or triple bond only counts as one domain.

Electron Domain Geometry The arrangement of electron domains about the central atom of an AB n molecule is its electron-domain geometry. There are five different electron-domain geometries: linear --(2 electron domains) trigonal planar --(3 domains) tetrahedral --(4 domains) trigonal bipyramidal --(5 domains) octahedral --(6 domains).

VSEPR – Valence Shell Electron Pair Repulsion X + E Overall Structure Forms 2 LinearAX 2 3 Trigonal PlanarAX 3, AX 2 E 4 TetrahedralAX 4, AX 3 E, AX 2 E 2 5 Trigonal bipyramidalAX 5, AX 4 E, AX 3 E 2, AX 2 E 3 6 OctahedralAX 6, AX 5 E, AX 4 E 2 A = central atom X = atoms bonded to A E = nonbonding electron pairs on A

Electron Domain Geometry

For example… :O=C=O: There are 2 electron domains on carbon…Its shape must therefore be linear. H–O –H There are 4 electron domains on oxygen….Its shape is based on the tetrahedral. Next, we will look at the molecular geometry!..

Molecular Geometry The molecular geometry is the arrangement of the atoms in space. To determine the shape of a molecule we will distinguish between bonding pairs and lone pairs. -Count the # of bonding domains vs. nonbonding domains. H-O-H Oxygen has 2 bonding and 2 nonbonding domains With this information, we can determine the molecular geometry…bent (as we know already!)..

Molecular Geometry

According to VSEPR theory, if there are three electron domains in the valence shell of an atom, they will be arranged in a(n) _____ geometry –A. octahedral –B. linear –C. tetrahedral –D. trigonal planar –E. trigonal bipyramidal

The electron-domain geometry of the central atom in OF 2 is _________. –linear –trigonal planar –tetrahedral –trigonal bipyramidal

Molecular Geometry—Most Common Shapes The most common shapes we deal with are as follows: -Tetrahedral, pyramidal, bent, linear, and trigonal planar. ( -The “ideal” bond angle between the central atom and the other atoms should be noted… Linear= 180º Tetrahedral = 109.5º Trigonal Planar =120º -Due to the lone pairs of electrons on pyramidal and bent shapes, the ideal bond angles will be less than 109.5º

Molecular Geometry— e - repulsion In general, multiple bonds repel more as do lone pairs.

Molecular Shape and Molecular Polarity When there is a difference in electronegativity between two atoms, then the bond between them is polar. It is possible for a molecule to contain polar bonds, but not be polar. - For example, the bond dipoles in CO 2 cancel each other because CO 2 is linear.

Molecular Shape and Molecular Polarity In water, the molecule is not linear and the bond dipoles do not cancel each other. Therefore, water is a polar molecule.

Molecular Shape and Molecular Polarity The overall polarity of a molecule depends on its molecular geometry.

Why do bonds form? Bonds form when orbitals on atoms overlap. There are two electrons of opposite spin in the overlapping orbitals.

Why do bonds form? The overlapping of the orbitals will lower the overall energy of the 2 atoms, therefore it is more stable.

Hybrid Orbitals A hybrid orbital is simply a mixing of different orbitals together to form a new “hybridized orbital”. We need the concept of hybrid orbitals to explain molecular shapes. (Let’s try to keep it simple…) When you mix n atomic orbitals we must get n hybrid orbitals. Example: If you mix one “s” orbital and three “p” orbitals you will get four “sp 3” hybrid orbitals that all have exactly the same energies.

Hybrid Orbitals The # of electron domains on the atom will indicate the hybridization needed. Example: H 2 C=CH 2 ( Carbon has 3 e - domains so its hybridization must be sp 2 which has 3 hybrid orbital domains as well.)

Sigma and Pi Bonds Overlapping orbitals come in 2 varieties…  -Bonds: electron density lies on the axis between the nuclei. - All single bonds are  -bonds.

Sigma and Pi Bonds  -Bonds: electron density lies above and below the plane of the nuclei. -A double bond consists of one  -bond and one  -bond. -A triple bond has one  -bond and two  - bonds. Often, the p-orbitals involved in  -bonding come from unhybridized orbitals. H 2 C=CH 2 A total of 5  -bonds are formed from the overlapping sp 2 hybrid orbitals of carbon, and the  -bond is from the unhybridized overlapping p-orbitals on each carbon.

Sigma and Pi Bonds H 2 C=O H–C≡C–HH–C≡C–H C and O both have sp 2 hybridization and each has an unhybridized p-orbital available to make the  -bond portion of the double bond. In this case, C has sp hybridization. One  -bond and two  -bonds form the triple bond between the carbon atoms.

Delocalized Pi Bonds Simply put, if there are resonance structures, the  -bond is delocalized or “smeared” between the 2 resonance structures. (By the way,  -bonds are never delocalized!) Example: Benzene (C 6 H 6 )