2.  Basic Terminology  Molecule  Group of atoms held together by covalent bonds  Bond Length  Distance between 2 bonded atoms  Bond Energy  Energy required/needed to break a chemical bond  Bond Angles  Determined based on bonding and nonbonding regions of the central atom in a molecule

3.  Molecular Shape  3-D structure of a molecule  Predicted from the Lewis dot diagram of a chemical compound  Exact shape experimentally determined

4.  Valence-Shell Electron-Pair Repulsion(“VSEPR”) Method  Valence electrons repel each other in chemical bond  Valence electron pairs are arranged so the distance between them is MAXIMIZED !  Valence electron pair number (bonded vs. lone)  Influence atom arrangement around central atom  Indirect influence on shape  Focus on central atom  Shared electrons  Unshared electrons

5.   Group of valence electrons “hanging out” around central atom, orientation around central atom  Composition: unpaired, lone pair, one/two/three bound pairs  Orientations based on electron groups  Linear—2  Trigonal planar—3  Tetrahedral—4  Trigonal bipyramidal—5  Octahedral--6 1) Electron Groups

6.  What about unshared electrons on the central atom?  Lone/unshared electrons  Occupy space around central atom  Not involved in chemical bonding between atoms  These pairs are NOT taken into account for a chemical molecule’s shape !  Require MORE space than shared electrons, indirect contribution to shape

7.   Orientation of BOUND atoms around a central atom 2) Molecular Geometry

8.  1)Draw Lewis structure 2)Identify number of electron groups—bound or lone pair? 1)Determine electron group geometry 2)Figure out molecular geometry  How many bonded atoms are around the central atom  Is this geometry the same as electron group geometry? How do we draw molecular shapes?

9.  Table 10.1 (pp. 390-391) Handout

10.   Molecular geometry = Electron group geometry !!! 1)Linear (AX 2 )  Bond angle = 180°  2 electron groups  Examples: MgCl 2, CO 2 No Lone Pair Electrons

11.  2) Trigonal Planar (AX 3 )  Bond angle = 120°  3 electron groups  Ex. BF 3 3) Tetrahedral (AX 4 )  Bond angle = 109. 5°  4 electron groups  Ex. CH 4, CCl 4 No Lone Pair Electrons

12.  4) Trigonal bipyramidal (AX 5 )  5 electron groups  Only seen with expanded valence shell, period 3+ 5) Octahedral (AX 6 )  6 electron groups  Only seen with expanded valence shell, period 3+ **Refer to Table 10.1 pp. 390-391 No Lone Pair Electrons

13.  1)Trigonal planar with 1 lone pair (AX 2 E)  3 electron groups, one lone pair  How many bound atoms?  Bond angle= 120°  Known as “angular”  Ex. SO 2 Lone Pair Electrons Present

14.  2) Tetrahedral with 1 lone pair (AX 3 E)  4 electron groups (3 pairs, 1 lone pair)  How many bound atoms?  Known as “trigonal pyramidal”  Bond angle = 109.5°, 107°(actual)  Ex. NH 3 Lone Pair Electrons Present

15.  2) Tetrahedral with 2 lone pair (AX 2 E 2 )  4 electron groups (2 pairs, 2 lone pair)  How many bound atoms?  Known as “angular/bent”  Bond angle ≈ 104/105°  Ex. H 2 O Lone Pair Electrons Present

16.   More than one central atoms is present  Figure out molecular geometry for each central atom  Combine BOTH geometries together to get the overall shape  Ex. HNO 3 Complex Molecules

17.   1) Repulsion between electron groups increases as the groups’ distance decreases.  2) Lone pair electrons have more repulsion than bonded electrons.  Charge can spread  Need more space in shape VSEPR Method “Take Home” Points