6-5: Molecular Geometry

1. 6-5: Molecular Geometry Objective: Explain VSEPR theory Predict the shapes of molecules or polyatomic ions using VSEPR theory Explain how the shapes of molecules are accounted for by hybridization theory Describe the dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces Explain what determines molecular polarity

2. Properties of molecules not only depend on the bonding of atoms but also on molecular geometry • Molecular geometry: the three-dimensional arrangement of a molecule’s atoms in space

3. VSEPR Theory • “Valence-Shell Electron-Pair Repulsion” • Refers to the repulsion between pairs of valence electrons of the atoms in a molecule • VSEPR theory: states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible • Electron pairs, shared or unshared, stay as far apart as possible to minimize repulsion • VSEPR theory is used to predict the shapes of molecules based on the fact that electron pairs strongly repel each other and tend to be oriented as far apart as possible

4. VSEPR and Unshared Electron Pairs: • VSEPR theory suggests that the lone pair occupies space around an atom just as bonding pairs do • Actual shape of a molecule is determined by the positions of the atoms only • Example: NH3  • In VSEPR theory, double and triple bonds are treated in the same way as single bonds

5. Common Molecular Shapes • Linear • Atoms bonded to central atom: 2 • Lone pairs of electrons: 0 • Bond angles: 180° • AB2

6. Trigonal Planar • Atoms bonded to central atom: 3 • Lone pairs of electrons: 0 • Bond angles: 120° • AX3

7. Bent • Atoms bonded to central atom: 2 • Lone pairs of electrons: 1 • Bond angles: <120° • AX2E

8. Tetrahedral • Atoms bonded to central atom: 4 • Lone pairs of electrons: 0 • Bond angles: 109.5° • AX4

9. Trigonal-Pyramidal • Atoms bonded to central atom: 3 • Lone pairs of electrons: 1 • Bond angles: 107° • AX3E

10. Bent • Atoms bonded to central atom: 2 • Lone pairs of electrons: 2 • Bond angles: 104.5° • AX2E2

11. Trigonal-Bipyramidal • Atoms bonded to central atom: 5 • Lone pairs of electrons: 0 • Bond angles: 120°, 90° • AX5

12. Octahedral • Atoms bonded to central atom: 6 • Lone pairs of electrons: 0 • Bond angles: 90° • AX6

13. Sample Problem • Use VSEPR theory to predict the molecular geometry of AlCl3. • 1. Draw the Lewis structure for the molecule. • 2. Look at the number of bonds and lone electron pairs made in the molecule. • How many bonds are attached to the central atom? • Three • How many lone electron pairs are there? • None • 3. Look at the type of molecule it is and refer to the chart to see what type of molecular shape it has. • AlCl3 = AX3 = Trigonal-planar

14. You Try! • Use VSEPR theory to predict the molecular geometry of the following molecules: • A) HI • Linear • B) CBr4 • Tetrahedral • C) AlBr3 • Trigonal-planar • D) CH2Cl2 • Tetrahedral • E) CO2 [double bond] • Linear

15. Hybridization • To explain how the orbitals of an atom can become arranged when the atoms form covalent bond, we use hybridization theory • Hybridization: the mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies • Hybrid orbitals: orbitals of equal energy produced by the combination of two or more orbitals on the same atom • The number of orbitals involved in hybridization determines the geometry of a molecule • There must be space around the atoms for every hybrid orbital Each orbital will be as far apart as possible from every other orbital

17. Factors that affect the geometry of a molecule: • The number of bonds formed by each atom in the molecule • The number of lone pairs of electrons on the atoms • The sizes of the various types of atoms • The hybridization of some of the atoms’ orbitals

18. Intermolecular Forces • Intermolecular forces: the forces of attraction between molecules • Boiling point is a good measure of the force of attraction between particles of a liquid  The higher the boiling point, the stronger the forces between particles

19. Molecular Polarity and Dipole-Dipole Forces • The strongest intermolecular forces exist between polar molecules • Polar molecules act as tiny dipoles because of their uneven charge distribution • Dipole: created by equal but opposite charges that are separated by a short distance • Arrow points toward the more electronegative atom + - H Cl

20. Polar molecules contain permanent dipoles • Dipole-dipole forces: the forces of attraction between polar molecules • These forces are short-range forces acting only between nearby molecules • A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons

21. Hydrogen Bonding: • Hydrogen bonding: the intermolecular forces in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule • For a hydrogen bond to form, hydrogen must be bonded to either a fluorine, oxygen, or nitrogen atom • The large electronegativity differences between hydrogen atoms and these atoms make the bonds connecting them highly polar

22. London Dispersion Forces • In any atom or molecule—polar or nonpolar– the electrons are in continuous motions • At any instant the electron distribution may be slightly uneven  creates a positive pole in one part of the atom and a negative pole in another • London dispersion forces: the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles • London forces act between all atoms and molecules • They are only intermolecular forces acting among noble-gas atoms and nonpolar molecules