Chapter 9 Notes AP CHEMISTRY Galster

Molecular Geometry VESPR Theory Abbreviation for valence shell electron pair repulsion. Valence electrons will arrange themselves to reduce repulsion i.e. they will be as far apart as possible 2 e-sets Bond angle: Shape name: AZx designation: 3 e- sets

4 e-sets Bond angle: Shape name: AZx designation: 5 e- sets 6 e-sets

Electron Set Geometry Name Molecular Geometry Example 3-D Shape AZ2 (2 e- domains) Linear AX2 CO2 AZ3 (3 e- domains) Trigonal planar AX3 BF3 AX2E Bent SO2

Electron Set Geometry Name Molecular Geometry Example 3-D Shape AZ4 (4 e- domains) Tetrahedral AX4 CH4 AX3E Trigonal pyramidal NH3 AX2E2 Bent H2O AXE3 Linear HCl

Electron Set Geometry Name Molecular Geometry Example 3-D Shape AZ5 (5 e- domains) Trigonal bipyramidal AX5 PCl5 AX4E See-Saw SF4 AX3E2 T-shaped ClF3 AX2E3 Linear XeF2

Electron Set Geometry Name Molecular Geometry Example 3-D Shape AZ6 (6 e- domains) Octahedral AX6 SF6 AX5E Square pyramidal BrF5 AX4E2 Square planar XeF4

Effect of non-bonding electrons on bond angles Order of Repulsion Nonbonding-nonbonding > nonbonding-bonding > bonding-bonding Trends Angle between nonbonding pairs increase Angles between bonding pairs decrease Examples 1. CH4 2. NH3 3 Water, H2O tetrahedral tetrahedral tetrahedral tetrahedral trig. Pyramidal bent

4. Sulfur tetrafluoride, SF4 5. COCl2

Polarity of Molecules Review of polar bonds Molecular polarity Polar bonds - ∆EN is less than or equal to 0.5 (but >1.7, ionic) Nonpolar bonds ∆EN is less than 0.5 Molecular polarity Nonpolar bonds Diatomic molecules or others with nonpolar bonds Ex: H2, Cl2, I2 No dipole moment So… molecules with all nonpolar bonds are nonpolar Cancellation of dipoles Depends on geometry “tug of war” Polar bonds – yes Dipole moment – no Ex: CCl4, CO2

Examples of Nonpolar Molecules CO2 SiCl5 AsCl5 ICl4-1

Examples of Polar Molecules Polar bonds: yes Dipole Moment? Yes H2O SF4 BrF5

Which of these are polar? SO2 I3-1 CHCl3 PCl3 XeCl4 SeBr4

Atomic Orbitals: Hybridization Atomic Orbitals – covalent bonds result when orbitals in valence level overlap **VESPR theory says that the number of unpaired electrons tells you the number of bonds the atom will form. Ex: H2 HCl Cl2

Hybrid Orbitals Mixing of orbitals Pure orbitals cannot always be used to explain bonding Basics Atomic orbitals are combined to create new hybrid orbitals The number of hybrid orbitals formed is equal to the number of atomic orbitals combined * We are not increasing the number of orbitals, just rearranging them. The names of the hybrid orbitals are based on the type and number of atomic orbitals combined. The hybrid orbitals are degenerate and have the same shape.

Process of hybridization A. methane, CH4 B. BF3 C. BeCl2

D. H2O E. NH3 F. XeF4

Hybridization and Electron Set Geometry sp = one s & one p = AZ2 linear sp2= one s & two p = AZ3 trigonal planar sp3= one s & three p = AZ4 tetrahedral sp3d = one s, three p, and one d = AZ5 trigonal bipyramidal sp3d2 = one s, three p, and two d = AZ6 octahedral

Multiple Bonds Multiple bonds do not affect the electron set geometry of a molecule Multiple bonds make attractions (i.e. bond) stronger and shorter Types of bonds Single bond = σ (sigma) bond Double bond = σ (sigma) bond – first bond π (pi) bond – 2nd or subsequent bonds Ex: C2H4

Triple bond Ex: C2H2 (Acetylene)

Sigma bonds are characterized by Head-to-head overlap Cylindrical symmetry of electron density about the internuclear axis Pi bonds are characterized by Side-to-side overlap Electron density above and below the internuclear axis

Delocalized Electrons: Resonance In reality, each of the four atoms in the nitrate ion has a p orbital. The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen. This means the electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.

Resonance The organic molecule benzene has six σ bonds and a p orbital on each carbon atom. In reality the π electrons in benzene are not localized but delocalized The even distribution of the π electrons in benzene makes the molecule unusually stable.

Molecular-Orbital (MO) Theory Through valence bond theory effectively conveys most observed properties of ions and molecules, there are some concepts better represented by molecular orbitals. In MO theory, we invoke the wave nature of electrons If waves interact constructively the resulting orbital is lower in energy: a bond

Molecular-Orbital (MO) Theory If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital.

MO Theory In H2 the two electrons go into the bonding molecular orbital The bond order is one half the difference between the number of bonding and antibonding electrons

MO Theory For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is ½ (2 – 0) = 1

MO Theory In the case of He2, the bond order would be ½ (2 - 2) = 0 Therefore, He2 does not exist.

MO Theory For atoms with both s and p, there are two types of interactions: The s and the p orbitals that face each other overlap in σ fashion. The other two sets of p orbitals overlap in π fashion.

MO Theory The resulting MO diagram looks like this : There are both σ and π bonding molecular orbitals and σ* and π* antibonding molecular orbitals

MO Theory The smaller p-block elements in the second period have a sizable interaction between the s and p orbitals. This flips the order of the σ and π molecular orbitals in these elements

Second-Row MO Diagrams