# AP Chapter 9 Molecular Geometry and Bonding Theories HW: 3 4 7 15 21 25 35 41 55 76 79 87.

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AP Chapter 9 Molecular Geometry and Bonding Theories HW: 3 4 7 15 21 25 35 41 55 76 79 87

9.1 – Molecular Geometry (Shapes) Molecular Geometry = The 3-dimensional arrangement of atoms in a molecule Bond Angle = Angles made by the lines joining the nuclei of atoms in the molecule Determined by the VSEPR Theory = Valence Shell Electron Pair Repulsion – Attached atoms and lone pairs repel each other The shape of a molecule plays an important role in its reactivity. By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

9.2 – The VSEPR Model Electron Areas / Domains Memorize information from the geometry charts Determine the number of “electron areas / domains” around a central atom

“Electron Area / Domain” Geometries: (Table 9.1 – pg 345) The electron-domain geometry is often not the shape of the molecule, however. The molecular geometry is defined by the positions of only the atoms in the molecules, not the nonbonding pairs. You must take the “Electron Area” Geometry and then determine the true Molecular Geometry

Molecular Geometry Within each electron domain, then, there might be more than one molecular geometry. Example – FOUR “electron areas”: …Details to come soon…

No Central; 2 Areas; 3 Areas No central atom = LINEAR – Example = HCl, N 2

4 Areas

Five Areas (Trigonal Bipyramidal)Electron Domain There are two distinct positions in this geometry: – Axial – Equatorial

Five Areas - Trigonal Bipyramidal Electron Domain Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

5 Areas

Six Areas All positions are equivalent in the octahedral domain

Molecules with More Than 1 Central Atom In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole.

Bond Angles – Lone Pairs Nonbonding pairs (lone pairs) are physically larger than bonding pairs. Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

Bond Angles - Multiple Bonds Double and triple bonds place greater electron density on one side of the central atom than do single bonds. (Multiple bonds create more push) Therefore, they also affect bond angles.

9.3 – Molecular Polarity But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

Polarity By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

Polarity Draw the overall dipole moments if applicable…

Polar (Dipole) Molecules Has partial charge and can be sliced (has a resultant dipole moment – Example: Water

Nonpolar Molecules True Nonpolars = Have all nonpolar covalent bonds Nonpolars (with polar sites) – CALLED NONPOLAR!! – Have partial charge, but cannot be sliced (have no resultant dipole moment) – Example: carbon dioxide

9.4 – Orbital Overlap – Valence Bond Theory Bond = An overlap of orbitals There is a distance of overlap of orbitals which is most stable

9.5 – Valence Bond Theory (Hybrid Orbitals) Lewis Structures alone do not explain all properties Example: H 2 vs Cl 2 – They both have single, nonpolar covalent bonds, but they are very different Example: If C has 2s 2 2p 2 valence electrons, why is it a tetrahedral shape?

Two Theories There are 2 main theories to explain bonding and properties of molecules Neither is perfect alone! – 1. Valence Bond Theory (the one AP Chemistry focuses on) – 2. Molecular Orbital Theory

9.5 - Valence Bond Theory (Hybrid Orbitals) e-s occupy atomic orbitals of individual atoms. When atoms come together to form molecules, HYBRIDIZATION occurs

2. Molecular Orbital Theory Atomic orbitals combine to form completely NEW MOLECULAR orbitals

But it’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize. 9.5 - Hybridization (Valence Bond Theory)

Hybrid Orbitals Consider beryllium: – In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.

Hybrid Orbitals But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds. The 2s orbital and the 2p orbital combine to form TWO equal energy HYBRID ORBITALS They are called “sp hybrids” because they are formed from an s and a p orbital

Hybrid Orbitals Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals. – These sp hybrid orbitals have two lobes like a p orbital. – One of the lobes is larger and more rounded as is the s orbital.

Hybrid Orbitals These two degenerate orbitals would align themselves 180  from each other. This is consistent with the observed geometry of beryllium compounds: linear.

Hybrid Orbitals Using a similar model for boron leads to… Called sp 2 hybrids because they are made from one s and two p orbitals

Hybrid Orbitals …three degenerate sp 2 orbitals.

Hybrid Orbitals With carbon we get…

Hybrid Orbitals …four degenerate sp 3 orbitals.

Hybrid Orbitals For geometries involving expanded octets on the central atom, we must use d orbitals in our hybrids. Elements must be in Period 3+ for this to occur. Example: Phosphorus:

Hybrid Orbitals This leads to five degenerate sp 3 d orbitals for 5 areas around the central atom And for 6 areas around the central atom - from s 2 p 4 to form six degenerate sp 3 d 2 orbitals.

Hybrid Orbitals -The moving of the electron to the higher energy orbital does require energy, but it is more than recovered during bond formation. -Once you know the electron-domain geometry, you know the hybridization state of the atom. 2345623456 Areas

9.6 - Multiple Bonds Formaldenhyde (H 2 CO) is trigonal planar, therefore carbon has sp 2 orbitals. The H, H and first bond to Oxygen bond with the hybrids. The second bond to Oxygen occurs with the left over p orbital. Trigonal planar

Multiple Bonds Acetylene (C 2 H 2 ), has a linear arrangement, therefore has sp hybrid orbitals. This leaves two p orbitals with one electron each to form the triple bond between the carbon atoms. Linear

Sigma (  ) Bonds Sigma bonds are characterized by – Head-to-head overlap. – Cylindrical symmetry of electron density about the internuclear axis. – Occurs in ALL single bonds – Occurs in ONE bond of a double or triple bond

Pi (  ) Bonds Pi bonds are characterized by – Side-to-side overlap. – Electron density above and below the internuclear axis. – Occurs in one bond of a double bond or two bonds of a triple bond

Sigma and Pi Bonds Summary = -single bond = one sigma -double bond = one sigma & one pi -triple = one sigma & two pi

Bond Order Single = 1 Double = 2 Triple = 3

Delocalized Electrons: Resonance When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.

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.

Delocalized Electrons: Resonance 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  bonds and a p orbital on each carbon atom.

Resonance In reality the  electrons in benzene are not localized, but delocalized. The even distribution of the  electrons in benzene makes the molecule unusually stable.

9.7 – Molecular Orbital Theory (MO) New molecular orbitals form as atoms bond Better for explaining some properties (for example: magnetic properties)

Molecular Orbital (MO) Theory In MO theory, we use the wave nature of electrons. If waves interact constructively, the resulting orbital is lower in energy: a bonding molecular orbital. If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital. No longer asked on AP Exam….

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