HTC Unit 4 Molecular Geometry and IMF

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Information about HTC Unit 4 Molecular Geometry and IMF

Published on February 20, 2008

Author: Peppar


Lewis Structures & Molecular Geometry:  Lewis Structures & Molecular Geometry Honors Theoretical Chemistry Chapter 6 Link to Shapes of Molecules Website Electron Distribution:  Electron Distribution 8 electrons = full outer shell = stable (except H and He which only hold a duet of 2) Shared pair – a pair of electrons shared by 2 atoms and thus bonding them together Lone pair – nonbonding electrons Rules for Writing Lewis Structures:  Rules for Writing Lewis Structures 1. Count the valence electrons. Put them in your bucket. 2. ID the central atom. Look for an atom that can form lots of bonds -- NOT hydrogen. 3. Bond each of the other atoms to the central atom with single bonds. 4. Complete valence shells of outside atoms. 5. Use remaining electrons to complete valence shell of central atom. 6. If you run out of electrons, form double or triple bonds. 7. If you have extra electrons, put them on the central atom. Draw Arsenic Triiodide:  Draw Arsenic Triiodide Arsenic Triiodide has the formula AsI3 How many total valence electrons? Let’s draw the Lewis structure How many shared pairs of electrons does As have? How many unshared pairs does As have? Draw these Lewis Structures:  Draw these Lewis Structures Cl2 O2 N2 CH3I C2H4 CH2O VSEPR valence shell electron pair repulsion:  VSEPR valence shell electron pair repulsion Repulsions between valence electrons causes them to be as far apart as possible in three dimensional space. These spatial relationships can be used to determine the shape of a given molecule. 5 Shapes To Know (and love?):  5 Shapes To Know (and love?) See Table 5, p.200 in Modern Chemistry Linear Trigonal Planar Bent Tetrahedral Trigonal Pyramidal Linear:  Linear BeCl2 CO2 Trigonal Planar:  Trigonal Planar BF3 Bent:  Bent SF2 Bent:  Bent H2O Tetrahedral:  Tetrahedral CH4 Trigonal Pyramidal:  Trigonal Pyramidal NH3 Quick Review: Rules for Drawing Lewis Structures:  Quick Review: Rules for Drawing Lewis Structures Count all the valence electrons Place the most electronegative atom in the center and add the other atoms around it Draw bonds between the atoms Add pairs of electrons around each atom until the atom has 8 electrons; do not exceed the total number of valence electrons Identify the geometry of these molecules::  Identify the geometry of these molecules: PCl3 ClO3- CO2 SF2 CCl2F2 Polar Molecules (Dipoles):  Polar Molecules (Dipoles) In a bond, the atom with the higher EN will attract the electrons more strongly The will cause that end of the molecule to have a partial negative charge (-), the other end will have a partial positive charge (+) Lewis structures which have a central atom with lone pairs of electrons indicate polar molecules. Intramolecular forces:  Intramolecular forces Bonds within a molecule that hold the atoms together Covalent Bonds Ionic Bonds Intermolecular Forces:  Intermolecular Forces Forces between molecules that hold them near each other Dipole-Dipole Forces – between polar molecules Hydrogen Bonding – between molecules that contain H bonded to O, N or F London Dispersion Forces – between large nonpolar molecules MC:  MC Which molecule contains a double bond? a. COCl2 b. C2H6 c. CF4 d. SF2 MC:  MC Which molecule is polar? a. CCl4 b. CO2 c. SO3 d. None of these MC:  MC What is the molecular geometry of CH3+? a. tetrahedral b. pyramidal c. bent d. trigonal planar Lewis Structures:  Lewis Structures The Lewis structure for HCN contains one double bond and one single bond. Draw it. IMF:  IMF Naphthalene, C10H8, is a nonpolar molecule with a boiling point of 208 oC. Acetic acid, CH3COOH, is a polar molecule with a boiling point of 118 oC. Which has stronger intermolecular forces? Polarity:  Polarity Predict which molecule has the more polar bonds: PCl3 or AsCl3 SnO or SrO SF2 or GeF4 SiCl4 or SCl2 Br2 or HBr Review:  Review Why is K2S ionic but H2S molecular? Write the electron configuration for the Zn2+ ion. Bonds:  Bonds Certain forces hold the water molecules in ice together. However, at 0 °C ice melts. When water reaches 100 °C, boiling occurs and the water molecules finally break free of each other. Even in the vapor state forces persist between the oxygen and hydrogen atoms. At several thousand degrees Celsius, the hydrogen atoms break free and water molecules no longer exist. Finally, at tens of thousands of degrees Celsius, a new and highly charged state of matter emerges.  Retell the above story using specific names for the forces and bonds that are involved at all levels and include the name for the last state of matter. Bonds:  Bonds When ice melts to form water, do chemical bonds break? ____________ When NaCl dissolves in water, do chemical bonds break? ____________  Polarity:  Polarity Identify the polar molecules H – C ≡ N: S = C = S   Slide29:  Why is H2O a dipole but CO2 is not, yet they both have 2 polar covalent bonds? Which intermolecular forces are at work between the noble gas atoms? Slide30:  Would these most likely be formed by polar or nonpolar molecules? a)    solid at room temperature b)    gas at room temperature c)    liquid with high boiling point d)    liquid with low boiling point MC:  MC Molecular compounds are easy to melt because a. covalent bonds are generally weak. b. forces between molecules are generally weak c. they are soluble in water d. none of the above MC:  MC Low electronegativity is characteristic of a. metals b. Group V c. metalloids d. nonmetals Slide33:  When an atom loses one or more electrons a(n) ______ is produced. In which compound are the bonds most polar? a. SbBr3 b. SbCl3 c. SbF3 d. SbI3 MC:  MC How does the strength of an intramolecular bond compare to the strength of intermolecular attractions? a. Intramolecular bonds are weaker. b. Intramolecular bonds are stronger. c. They are about the same. d. No generalization is possible. MC:  MC In order for arsenic (As) to form a stable ion it must a. lose 5 electrons. b. lose 3 electrons. c. gain 5 electrons. d. gain 3 electrons.

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