1. Chemistry Unit 4 Chapter 8

2.  Molecule  A neutral group of atoms joined together by covalent bonds  Molecular Compound  Tend to have lower melting and boiling points than ionic compounds  Covalent Bond  Sharing of electrons between two or more nonmetals  Between molecules, acids, and polyatomic ions

3.  Molecular Formulas  Represents how many atoms of each element are in the compound Cannot be determined by the criss-cross method – why? Can be determined by creating a structural (Lewis dot) formula  Octet Rule Atoms still tend to attain the electron configuration for a noble gas by sharing electrons

4.  Structural Formula  Shows the arrangement of covalently bonded atoms using (dashes) as bonds (shared electrons) and (dots) as unshared electrons  Example – Hydrogen H + H  H H (Lewis Dot Structure)  Creates an electron configuration of Helium (noble gas) for both hydrogens by sharing the electrons H H (Structural Formula)  Shared electrons represented by a dash

5.  Drawing Structural Formulas  Steps 1) Count the total number of valence electrons for the compound ex. CH 4 (methane) C = 4 e - × 1 atom = 4 H = 1 e - × 4 atoms = 4 Total = 8 electrons

6.  Drawing Structural Formulas  Steps 2) Arrange the atoms Hydrogen is NEVER in the middle Carbon is ALWAYS in the middle Elements with only one atom are typically in the middle ex. CH 4 (methane) H H C H H

7.  Drawing Structural Formulas  Steps 3) Fill in electrons, starting around the central atom(s) until all your valence electrons are used. ex. CH 4 (methane) H H C H H

8.  Drawing Structural Formulas  Steps 4) Draw lines (dashes) for any shared electron pairs and leave all unshared electrons where they are. ex. CH 4 (methane) H H C H H

9.  Drawing Structural Formulas  Now Try These H2OH2O F2F2 NH 3

10.  Single Bond  Two atoms held together by sharing 2 electrons (1 pair)  Double Bond  Two atoms held together by sharing 4 electrons (2 pairs)  Triple Bond  Two atoms held together by sharing 6 electrons (3 pairs) Double and triple bonds NEVER form with Hydrogen or the Halogens!!

11.  Examples  Double and triple bonds are needed sometimes to satisfy the octet rule for all the atoms in the compound.

12.  Molecular Shape  Compounds are THREE-DIMENTIONAL  VSEPR  Valence-Shell Electron-Pair Repulsion Theory  The negative charge on electrons repel each other creating molecular shapes where the electrons are as far away from each other as possible.  Tetrahedral Angle – 109.5 ⁰ The shape where all the electron pairs are as far apart as possible

13.  Common Shapes  Linear – two atoms or two double bonds  Trigonal Planar Must have one double bond  Bent  Pyramidal  Tetrahedral