Warm Up – Name or Write the following compounds 1.Cr2(SO3)3 2.LiCN 3.Copper(II) chlorate 4.Iron(III) hydroxide 5.CS4

Writing Lewis Dot Structures In writing Lewis structures for covalent compounds, you must remember the following: 1. Determine the total number of valence electrons 2. The least electronegative element goes in the center H is NEVER a center atom Carbon is usually in the center

Writing Lewis Structures 3. Distribute 8 electrons around central atom 4. Place remaining atoms outside of central atom 5. Distribute any remaining electrons around outer atoms

Writing Lewis Structures 6. Use double or triple bonds to obey Octet Rule as needed 7. Check to make sure that every atom obeys the Octet Rule

Write Lewis Dot Structures for the following: 1. H2O 2. CO2 3. NH3 4. CBr4 5. SO2

Practice Draw the Lewis Structures for the Following: 1) GeF4 2) COS 3) SeO3 4) HOF

Practice Write Lewis structures for the following molecules: a. HCN d. PH3 b. CHCl3 e. SiO2 c. SeF2 f. CF4

Lewis Structures for Polyatomic Ions When writing dot structures for polyatomic ions, you must remember to add or subtract the amount of electrons represented by the charge When writing polyatomic ions, you must include the structure inside brackets, [ ], with the charge outside the bracket

Write Lewis Structures for the following polyatomic ions: 1. (SO4) 2- 2. (OH)- 3. (CN)- 4. (NH4)+

Warmup Practice Naming or writing Compounds Mn(ClO2)7 PbS NH4HSO4 Cd(NO3)2 Lead (IV) perchlorate Beryllium Acetate Copper (II) Bisulfate

Extended Octet Rule Examples: PCl5 or XeF4 “Expanded/Extended octet” refers to the Lewis structures where the central atom ends up with more than an octet Examples: PCl5 or XeF4

Limitation of Extended Octet Rule 1) The central atom with an expanded octet MUST have an atomic number larger than 10 (beyond neon). 2) Extra electrons should be first placed on the outside atoms. After the outside atoms have fulfilled the Octet Rule, and there are still extra electrons, start with placing them as lone pairs on the central atom.

Practice 1) BrF3 2) SF4 3) PCl5 4) XeF4 5) SbF5

VSEPR Theory VSEPR Theory- Valence Shell Electron Pair Repulsion Theory VSEPR Theory can be used to predict a molecule’s geometry Electrons repel each other Molecules adjust their shapes to account for this repulsion

VSEPR Theory Rules The adjusted shape allows for the minimum amount of repulsion between valence shell electrons Basically, a molecule will arrange itself in such a way that the valence electrons are as far apart from each other as possible

Using VSEPR Theory Unshared pairs of electrons play an important role in the molecule’s geometry Built off of Domains Single bonds Double Bonds Triple Bonds Lone Pairs of Electrons

Basic Geometric Shapes for Molecules Linear Trigonal Planar Tetrahedral Trigonal Pyramidal Trigonal bipyramidal Octahedral * The ones in yellow you must memorize

Linear

Trigonal Planar

Tetrahedral

Trigonal Pyramidal

Bent

Trigonal bypyramidal

Octahedral

Using the VSEPR Theory, Predict the Shape of the following Molecules 1. H2O 2. BF3 3. H2S 4. PH3

Practice VSEPR Theory: Draw the Lewis structures for the following compounds and predict the molecular geometry of the compound 1. SCl2 3. PI3 2. OCl2 4. NH2Cl

Draw the Lewis Structures and Predict the Molecular Geometry a. BeCl2 b. BCl3 c. (SeO4)2- d. SiCl3Br f. ONCl

Practice Draw the following Lewis structures for the following and predict the molecular geometry HCN SO2 XeF2 PCl5

Predict the following molecular geometry PBr3 N2H2 C2H4 OCl2 HC2F

Warm up Practice Drawing Lewis Structures and Predicting Molecular Geometry 1) SCl2 2) BeF2 3) HOF 4) COS

Origin of Polar Bonds Recall, electronegativity refers to the tendency of an atom to attract electrons Electronegativity really measures the strength that an atom uses to attract electrons The element with the higher electronegative value will attract electrons more strongly which means that it will pull the bonds closer

Origins of Polar Bonds The greater the difference between each atom’s electronegativity, the more likely it is that the bond is polar

Origins of Polar Bonds The closer the atoms are together, the closer the electronegative values are to each other Elements that have closer electronegative values form NONPOLAR BONDS The further the atoms are apart, the more likely they are to form POLAR BONDS

Polar Bonds Polar bonds result when shared electrons are not shared equally Polar bonds occur only with covalent compounds (means they have an ∆EN< 1.7 but > 0.5) This causes each atom to develop a temporary charge

NonPolar Bond Non-Polar bonds result when shared electrons are shared equally Non-Polar bonds occur only with covalent compounds with an ∆EN= 0 to .3

Using Electronegative Values to Predict Bond Types Predict if the following bonds will be: Polar Covalent, Nonpolar Covalent, or Ionic based on Electronegative Values a. H and O b. Cl and Br c. K and S d. F and F

Polar Molecules It is possible for a molecule to have polar bonds, but be nonpolar overall When polar bonds are in opposite directions, they “cancel” each other out Polar molecules are usually not symmetrical Example: CCl4 has polar bonds, but is a NONPOLAR MOLECULE

Predict if the following molecules are polar or nonpolar 1. CF4 2. CO2 3. H2O 4. (SO4)2- 1. NonPolar 2. Nonpolar 3. Polar 4. Nonpolar