Ch 9 Molecular Geometry Bonding Theories Lewis structures

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. • We will first explore the molecular form ABn. This represents one central atom bonded to 1 -6 other atoms. • First 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 ABn 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).

Chapter 9

Electron Domain Geometry

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.

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

Molecular Geometry

Molecular Geometry

Molecular Geometry—Most Common Shapes • The most common shapes we deal with are as follows: -Tetrahedral, pyramidal, bent, linear, and trigonal planar. (It is to your advantage to know some common examples of each of these shapes!!) -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.

Shapes of Larger Molecules • In acetic acid, CH 3 COOH, there are three interior atoms: two C’s and one O. • We assign the molecular (and electron-domain) geometry about each interior atom separately: -The geometry around the first C is tetrahedral. -The geometry around the second C is trigonal planar. -The geometry around the O is bent (tetrahedral).

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 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. H–C≡C–H 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)

Molecular Orbitals • Some aspects of bonding are not explained by Lewis structures, VSEPR theory or hybridization. For example, why does O 2 interact with a magnetic field? • For these molecules, we use Molecular Orbital (MO) Theory. Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals. -Molecular Orbitals: 1) contain a maximum of 2 electrons. 2) have definite energies. 3) can be visualized with energy level diagrams. 4) are associated with an entire molecule.

Molecular Orbitals Let’s look at the formation of H 2… • 1 s (H) + 1 s (H) must result in two MO’s for H 2: - One MO has electron density between the nuclei. This is called a bonding MO. - The other MO has little or no electron density between nuclei. This is called an antibonding MO. • MO’s resulting from s-orbitals are MOs. • (bonding) MO is lower energy than * (antibonding) MO… Confused yet? Let’s look at a diagram!

Molecular Orbitals (caused by destructive interference) (caused by constructive interference)

Molecular Orbitals • Energy level diagrams or MO diagrams show the energies and electrons in an orbital. • The total number of electrons in all atoms are placed in the MO starting from lowest energy ( 1 s) and ending when you run out of electrons. • Note that electrons in MOs have opposite spins. (Sound familiar? !) H 2 has a total of two bonding electrons.

Molecular Orbitals • Let’s look at another example… He 2 • He 2 has two bonding electrons and two antibonding electrons. That is not a stable situation and therefore He 2 will not form. • In order to predict stability of a molecule from a MO diagram, we need to learn about the concept of “BOND ORDER”…

Bond Order = ½ (bonding electrons – antibonding electrons) - Bond order = 1 for single bond. - Bond order = 2 for double bond. - Bond order = 3 for triple bond. • Fractional bond orders are possible… This indicates that resonance structures are likely. • For H 2 the bond order = 1 which correctly predicts the single bond. • For He 2 the bond order is 0 which says that it does not form a bond.

Bond Order Here’s another example…O 2 How many total # of e-’s will we fill 8+8 = 16 in? _____ What is the bond order for O 2? _____________ ½ (10 -6) = 2…double bond Notice that there are 2 unpaired e- in the * 2 p orbitals. This leads to our last topic of discussion. ↑ ↑ ↑↓ ↑↓

Electron Configurations and Molecular Properties • Molecules have two types of magnetic behavior: - paramagnetism (unpaired electrons in molecule): strong attraction between magnetic field and molecule - diamagnetism (no unpaired electrons in molecule): weak repulsion between magnetic field and molecule. • Magnetic behavior is detected by determining the mass of a sample in the presence and absence of a magnetic field. (diamagnetic) (paramagnetic) small decrease in mass large increase in mass

Electron Configurations and Molecular Properties • O 2 has 2 unpaired electrons as shown in the MO diagram which explains why liquid oxygen is attracted to a magnetic field. • The Lewis structure for O 2 does correctly predict a double bond, but it does not show unpaired electrons, so it cannot explain this observed paramagnetic behavior. • That is why we have the MO theory. It predicts not only the double bond, but also the paramagnetism of such molecules! • Last but not least, MO diagrams can be used for heteronuclear diatomic molecules as well, that is, molecules of 2 different atoms bonding together. (For example, CO, carbon monoxide. ) • MO diagrams are filled out the same way, and bond order is calculated the same way as well.