Chemistry 1A: Chapter 10 P a g e | 1 Chapter 10: Chemical Bonding II: Molecular Shapes; VSEPR, Valence Bond and Molecular Orbital Theories Homework: Read Chapter 10: Work out sample/practice exercises. Check for the MasteringChemistry.com assignment and complete before due date Molecular Shapes: Properties of molecular substances depend on its 3D structure Bonding neighbors, what is next to what (skeleton arrangement) Type of bonding; polar, nonpolar, ionic Shape and Polarity; overall do dipoles cancel or is there an overall dipole moment Limitations in Lewis Structures: • Lewis theory predicts the number of electron regions (lone pair or any bond; single double, triple), but does not determine actual bond angles. • Lewis theory predicts trends in properties, but does not give good numerical predictions of bond strength and bond length • Lewis theory cannot write one correct structure for molecules where resonance is important • Lewis theory often does not correctly predict magnetic behavior of molecules. Oxygen, O2, is paramagnetic, though the Lewis structure predicts it is diamagnetic Valence Shell Electron Pair Repulsion (VSEPR) Theory: Three-dimensional • Electron groups (all negatively charged) around the central atom are most stable when they are as far apart as possible –valence shell electron pair repulsion theory. Use all the information gained in the Lewis Dot Structure and convert it to a three dimensional model to predict electronic and molecular shapes, angles, and polarity of the molecule. VSEPR Guidelines: Start with information from a Lewis Dot Structure Electronic and Molecular 3D shapes Bonds angles: When electron groups attach to different size atoms the ideal bond angles are affected. Lone pairs (nonbonding) use more space. Polarity of whole substance (ionic, ion, nonpolar, polar molecule) Electronic and Molecular Geometry: Chemistry 1A: Chapter 10 P a g e | 2 Count the electron regions. Electron regions will give an electronic shape while the number of bonded versus nonbonded regions will give the molecular shape. # Electron 2 3 4 5 6 regions Electronic Linear trigonal planar tetrahedral trig. bipyramidal octahedral geometry 180˚ 120˚ 109.5˚ 90˚, 120˚, 180˚ 90˚, 180˚ five basic shapes molecular Linear Trig planar, Tetrahedral, trig.bipyramidal, octahedral, geometry bent Trig. pyramidal, see saw, square pyramidal, bent T-shaped, square planar linear Samples Imperfect Geometry: When electron groups attach to different size atoms the ideal bond angles are affected CH2O ideally should be trigonal planar with angles of 120° each. In reality the angle between the smaller H atoms is smaller. Lone pairs (nonbonding electrons) use more space. Ideally four regions should spread out to angles of 109.5°. Notice how the bond angles around the atoms are forced closer together as the unseen nonbonding electrons take more space. Website to try: ChemEdDL.org Click on molecules 360. This website shows the 3D structure of many chemicals and allows you to rotate in three dimensions, showing bonding, bond length, dipole arrows, dipole moment, etc. Writing 3D shapes on paper: May use lines and wedges. Chemistry 1A: Chapter 10 P a g e | 3 Multiple Central Atoms: Describe the shape around each central atom atom separately. Polarity of the Molecule: Polar: must have polar bonds (electronegativity difference between the neighbor atoms with a measureable bond dipole moment) and an unsymmetrical shape (lone pairs or varying atom neighbors) Polarity affects properties: boiling points, solubilities (like dissolves like) HCl and H2O are both polar CO2 is nonpolar Chemistry 1A: Chapter 10 P a g e | 4 Valence Bond (VB) Theory: Three-dimensional The Valence Bond theory is a quantum mechanical model that expands the previous two theories to describe the electronic nature of covalent bonds. • Valence bond theory applies principles of quantum mechanics to molecules • A chemical bond between atoms occurs when atomic orbitals and hybridized atomic orbitals interact with those in another atom to form a new molecular orbital with two electrons. If orbitals align along the axis between the nuclei, sigma bonds which directly overlap will form ( bonds). It is possible to rotate a sigma bond If orbitals align outside the axis, pi bonds form, which indirectly overlap above and below ( bonds). Unable to rotate without breaking bonds. This causes cis and trans structural isomers. VB Guidelines: Use all the information from a Lewis Dot Structure Hybridizing some orbitals allow for more bonds and more stability Visualize orbital picture using atomic (s, p, d, f) and hybridized (sp, sp2, sp3, sp3d, and sp3d2) orbitals Direct overlap orbitals, sigma () bonds Indirect overlap orbitals, pi () bonds All types of bonds have only one bond. Double bonds have 1 and 1 and triple bonds have 1 and 2 bonds Valence Bond (Bubble) Pictures draw the orbitals in balloon type pictures Delocalized bonding occurs in substances with resonance Chemistry 1A: Chapter 10 P a g e | 5 Chemical bonds between atoms occur when atomic orbitals interact with those in another atom to form a new molecular orbital with two electrons. Sigma bonds (direct overlap) are stronger than pi bonds (indirect overlap). Chemistry 1A: Chapter 10 P a g e | 6 Double bond: CH2O C 3 sp2 hybridized orbitals and 1 p unhybridized orbital H 1 s orbital on each O 1 s unhybridized orbital and 3 p unhybridized orbitals Triple bond: C2H2 C 2 sp hybridized orbitals and 2 p unhybridized orbitals H 1 s orbital on each Limitations in Valence Bond Theory: Valence Bond theory predicts bond strengths, bond lengths, and bond rigidity better than Lewis theory. • Other properties, such as the magnetic behavior of O2, of molecules are not predicted well. • VB theory views electrons as localized in overlapping atomic orbitals and it doesn’t account for delocalization Chemistry 1A: Chapter 10 P a g e | 7 Molecular Orbital (MO) Theory: The Molecular Orbital Theory is separate from the first three. This theory explains the paramagnetic behavior found in O2 gas molecules. • In MO theory, Schrödinger’s wave equation is applied to the molecule to calculate a set of molecular orbitals • Electrons and orbitals belong to the whole molecule – Delocalization • A Bonding Molecular Orbital forms when wave functions combine constructively, resulting in a molecular orbital with lower energy than the original atomic orbitals. Most of the electron density is between the nuclei. Lower energy-stabilizing • The Antibonding* Molecular Orbital forms when wave functions combine destructively, resulting in a molecular orbital with more energy than the original atomic orbitals. Most of the electron density is outside the nuclei creating nodes between nuclei. Higher energy-unstable Sigma () 1s molecular orbitals (2s looks the same, but a bit bigger) Sigma () px molecular orbitals Chemistry 1A: Chapter 10 P a g e | 8 Pi () py or pz molecular orbitals MO Guidelines: Electrons belong to the molecule, not the individual atoms For this class, limit most of the discussion and examples to diatomic species -1 such as: H2, O2, CN , HF. Occasionally this gives a more accurate electronic structure than VB Combination of two atomic orbitals makes a molecular orbital Bonding orbitals are sigma or pi orbitals. Sigma orbitals directly overlap and pi orbitals indirectly overlap Antibonding* sigma or pi orbitals create a node between the atoms with no overlap Two atomic s orbitals combine to form a lower energy bonding and a higher energy * antibonding* orbital six atomic p orbitals combine to form lower energy bonding orbitals, and 2 degenerate orbitals and higher energy antibonding* orbitals,and 2 degenerate orbitals Predicts paramagnetic or diamagnetic behavior Predicts bond order Compares bond lengths and bond strengths For diatomic molecules with fewer than 15 total electrons like N2, energy increases as follows: s, 1s*, 2s, 2s*, 2p,2p, 2p, 2p*, 2p*, 2p* Chemistry 1A: Chapter 10 P a g e | 9 For diatomic molecules with 15 or more total electrons like O2, energy increases as follows: s, 1s*, 2s, 2s*, 2p,2p,2p, 2p*, 2p*, 2p* Magnetic behavior of O2 Diatomic oxygen is attracted to a magnetic field, indicating paramagnetic behavior, so it has unpaired electron(s) Chemistry 1A: Chapter 10 P a g e | 10 Heteronuclear Diatomic Elements and Ions: • The more electronegative an atom is, the lower in energy are its orbitals • Lower energy atomic orbitals contribute more to the bonding MOs • Higher energy atomic orbitals contribute more to the antibonding MOs • Nonbonding MOs remain localized on the atom donating its atomic orbitals Chemistry 1A: Chapter 10 P a g e | 11 Polyatomic Molecular Orbitals: • Atomic orbitals of all the atoms in a molecule, even those with 3 or more atoms, combine to make a set of molecular orbitals, delocalized over the entire molecule • Predictions made using molecular orbital theory, (especially resonance molecules and predicting magnetic properties), match the real molecule properties better than either Lewis or Valence bond theories. Ozone, O3: MO theory predicts equivalent bond lengths due to the delocalized electrons. Chemistry 1A: Chapter 10 P a g e | 12 Molecular Shapes, Handedness and Drugs: The shapes of molecules can dramatically change its characteristics. Mirror images have different biological properties due to the specific shapes of receptor sites in the body. For a molecule to exhibit handedness it needs four different groups attached to a carbon. Identify the electronic and molecular geometries,