Shapes

Writing Lewis Structures Step 1 Total number of valence electrons Add the total negative charge if you have an anion. Subtract the charge if you have a cation. Example: CO32- here it is 4 (for Carbon) + (3x6 oxygen's) + 2 charge = 24 total e-   Step 2 Number of electrons if each atom is to be happy: Atoms in our example will need 8 e (octet rule) or 2 e ( hydrogen)   Step 3 Draw the structure: The central atom is C ( usually the atom with least electro negativity will be in the center). The oxygen's surround it . Because there are four bonds and only three atoms, there will be one double bond. Step 4: Octet the periphery (2 for hydrogen) Left over for central atom = total (step 1) – Used (from step 3). = 24 – (3x8) = 0 Step 5  Double check your central atom has 8 electrons (except Boron), if not use multiple bonds for periphery to make it so .

Formal Charges: Writing Lewis Structures Then assign formal charges. For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. Subtract that from the number of valence electrons for that atom: the difference is its formal charge. © 2009, Prentice-Hall, Inc.

Writing Lewis Structures The best Lewis structure… …is the one with the fewest charges. …puts a negative charge on the most electronegative atom. © 2009, Prentice-Hall, Inc.

Resonance One Lewis structure cannot accurately depict a molecule like ozone. We use multiple structures, resonance structures, to describe the molecule.

Resonance Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.

Molecular Shapes The shape of a molecule plays an important role in its reactivity. By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule. © 2009, Prentice-Hall, Inc.

What Determines the Shape of a Molecule? Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule. © 2009, Prentice-Hall, Inc.

Electron Domains We can refer to the electron pairs as electron domains. In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain. The central atom in this molecule, A, has four electron domains. © 2009, Prentice-Hall, Inc.

Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.” © 2009, Prentice-Hall, Inc.

Electron-Domain Geometries These are the electron-domain geometries for two through six electron domains around a central atom. © 2009, Prentice-Hall, Inc.

Hybridization Refers to mixing of orbitals. Atomic orbitals of central atom undergo change to accommodate incoming atoms. Hybridization could be sp, sp2, sp3, sp3d and sp3d2. How do you tell the hybridization on the central atom?

9.1 – 9.2: V.S.E.P.R. Valence-shell electron-pair repulsion theory Because e- pairs repel, molecular shape adjusts so the valence e- pairs are as far apart as possible around the central atom. Electron domains: areas of valence e- density around the central atom; result in different molecular shapes Includes bonding e- pairs and nonbonding e- pairs A single, double, or triple bond counts as one domain Summary of LmABn (Tables 9.1 - 9.3): L = lone or non-bonding pairs A = central atom B = bonded atoms Bond angles notation used here: < xº means ~2-3º less than predicted << xº means ~4-6º less than predicted

Tables 9.1 - 9.3 # of e- domains & # and type of hybrid orbitals e- domain geometry Formula Molecular geometry Predicted bond angle(s) Example (Lewis structure with molecular shape) 2 Two sp hybrid orbitals Linear AB2 180º BeF2 CO2 |X X |B B A A

Three sp2 hybrid orbitals 3 Three sp2 hybrid orbitals Trigonal planar AB3 120º BF3 Cl-C-Cl << 120º Cl2CO LAB2 Bent < 120º NO21- |B B A |X X A |B : B A

Example: CH4 H | H—C—H Molecular shape = tetrahedral Bond angle = 109.5º 109.5º

Four sp3 hybrid orbitals 4 Four sp3 hybrid orbitals or Tetrahedral AB4 109.5º CH4 LAB3 Trigonal pyramidal < 109.5º Ex: NH3 = 107º NH3 L2AB2 Bent <<109.5º Ex: H2O = 104.5º H2O B A A X X : A B X X A : A B

: PCl5 :Cl: :Cl: \ / :Cl—P—Cl: | :Cl: \ / :Cl—P—Cl: | :Cl: : Molecular shape = trigonal bipyramidal Bond angles equatorial = 120º axial = 90º 90º 120º

Five sp3d hybrid orbitals 5 Five sp3d hybrid orbitals Trigonal bipyramidal AB5 Equatorial = 120º Axial = 90º PCl5 LAB4 Seesaw Equatorial < 120º Axial < 90º SF4 B A | A X | X X : B - A - B B

Five sp3d hybrid orbitals 5 Five sp3d hybrid orbitals Trigonal bipyramidal L2AB3 T-shaped Axial << 90º ClF3 L3AB2 Linear Axial = 180º XeF2 B : A | X A | : A B |

Six sp3d2 hybrid orbitals 6 Six sp3d2 hybrid orbitals or Octahedral AB6 90º SF6 LAB5 Square pyramidal < 90º BrF5 A X | A B | X X B B A B | .. B B A X | X X

Six sp3d2 hybrid orbitals 6 Six sp3d2 hybrid orbitals or L2AB4 Square planar 90º XeF4 L3AB3 T-shaped <90º KrCl31- A .. | .. A B | .. B B B A .. | .. A .. | B B .. B .. B