1. Reviews for Chapter 8, 9 and 12 Exam 4 Must be familiar with concepts in Chapter 7 to better understand chapter 8 and 9

2. 8.1Chemical Bonds: A Preview Chemical bonds result from attractive and repulsive forces between electrons and nuclei of the 2 atoms forming molecules Ionic Bonds and Ionic Crystals (check eg. CaCl 2 ) Lewis Structures for Ionic Compounds: Electrons Transferred 8.2The Lewis Theory of Chemical Bonding: An Overview Lewis Structure is simply chemical symbol + # valence electrons around it A. Most atoms "want" eight electrons around them (octet rule) B. Exception: duet rule for hydrogen and helium C. A covalent bond is a shared par of electrons to get an octet around each atom also some have expanded octet Lewis Structures of Some Simple Molecules (H 2, F-F, O=O, N ≡N ). 1.Hydrogen,F atoms are rarely central atoms; they are almost always terminal atoms. 2.The central atom usually is the atom with the lowest Electronegativity (except for H that often has a relatively low Electronegativity). 3. In oxoacids, H usually is bonded to an oxygen atom.(HNO 3, H 2 SO 4 etc) 4.Molecules usually are clusters of atoms, rather than long chains. 5Understand bonding patterns of C, N, O and halogens. Chapter 8 Bonding and Molecular Structure

3. Strategies for Writing Lewis Structures A B. Lewis structures are easier to draw if one follows a systematic plan in constructing them. One such plan is as follows. 1.Sum the valence electrons in the structure. a.For each atom, the number of valence electrons equals that atom’s group number. b. Add electrons if the entire species is an anion, subtract them if the species is a cation. c. Take off electrons ( 2 electrons per bond) to form single bonds, 2.Arrange the atoms in the preferred skeletal structure (central atoms is lowest Electronegativity etc) and place a pair of electrons between each pair of bonded atoms. Subtract single bon elctrons 3.Place remaining electron pairs around terminal atoms so that each has an octet. 4.Place LEFT OVER electron pairs on the central atom. 5.Create multiple bonds as necessary to make sure each atom follows the octet rule.

4. 8.3 Learn to calculate formula charge sum of formal charges = charge of the species 8.4 Learn to draw plausible resonance structures 8.5Molecules that Don’t Follow the Octet Rule (SF 6, IF 5, XeF 4, BeCl 2, BH 3 etc.) Exceptions to octet rule 8.6 Molecular shapes see shape table, and visualize symmetric shapes 8.7Electronegativity (EN) and BOND POLARITY (Polar Covalent Bonds ) Electronegativity (tendency to pull shared electron towards it) 1. Increases as you move left to right along a Period HIGHEST is F 2. Decreases as you go down a group 3. Noble gases have very low EN Difference of Pure covalent vs polar covalent vs ionic bond based on EN difference Chemical bonds result from attractive and repulsive forces between electrons and nuclei of the 2 atoms forming molecules: Types of chemical bonds based on electronegativity difference between the bonding atoms  EN = 0 :Pure covalent bond;  EN = 0 to 0.39 Nonpolar Covalent  EN = 0.4-1.9 :Polar covalent bond;  EN ≥1.9 or Has a metal: Ionic

5. 8.8Molecualr polarity (net dipole moment µ) - A molecule may have polar bonds, BUT all cancelling out due to symmetry, results in NON POLAR molecule (e.g. CO 2, CH 4 ) - Combination of shapes, bond polarity additions and symmetry -Lone pair on central atom: Usually Polar (except XeF4, SeF 2 ) or linear shapes - Unsymmetrical molecule is Polar (e.g. CH 3 Cl, CH 2 Cl 2 )

6. 8.9 Bond order, Bond length and Bond energy/strength for Covalent Bonds Bond order = # of bonds for a pair of atoms e.g. Triple bond has Bond order =3 Bond length is inversely relationship to bond order Also depends on atomic sizes of atoms making up the bonds Bond strength also called as bond energy or bond dissociation energy is Highest for triple bond and least for single bond Bond energy is defined as the energy required to break covalent bonds holding up 2 atoms Calculate  H rxn =∑ Bond energies broken - ∑ Bond energies formed

7. 8, 9Bonding Theory and Molecular Structure 8.6 The Valence-Shell Electron-Pair Repulsion (VSEPR) Method -minimize electron –electron repulsion between bonded electrons pairs -Predicts shapes, bond angles and dipole moments -MEMORIZE Shape table, bond angles -REMEMBER we care only about electron groups on the central atoms and then -Predict Electron geometry and MOLECULAR GEOMETRY -9.1Atomic Orbital Overlap This region results from the overlap of the two atomic orbitals, one on each atom. The greater the degree of overlap, the stronger is the bond. Most often, the two atomic orbitals that overlap to form the bond are each half-filled 9.2Hybridization of Atomic Orbitals A.We hybridize not just to explain the bonding, but also to create the correct number of orbitals on the central atom, pointing in the correct direction, to create the molecule.

8. 9.2Hybrid Orbitals and Multiple Covalent Bonds B.General distinctions between sigma and Pi Sigma bonds a.Result from the end-to-end overlap of two orbitals. b.Form between all varieties of orbitals. c.Are the first bond that forms in a multiple bond, the only bond that forms in a single bond. d.Produce the most effective overlap, the strongest bond, at a given internuclear distance. Pi bonds (ALSo results in geometric isomerism) a.Result from the side-to-side overlap of p orbitals. b.Only form between two p orbitals c.Are the second and third bonds of multiple bonds d. Produce weak overlap and restrict rotation around the bond. Learn to recognize sigma and pi in a molecule

9. 12.1 Properties of Solids and Liquids A. Properties of liquids 1. High densities compared to gases 2. Indefinite shape, fluid 3. Definite volume B. Properties of solids 1. High densities compared to gases 2. Definite shape 3. Definite volume 4. May be crystalline or amorphous 12.2 Interaction between ions Energy Changes in Ionic Compound, Coloumb’s law Ionic bond energy is directly proportional to Magnitude of Charges Ionic bond energy is inversly proportional to distance separting the ions Recollect Ion sizes 12.3-12.5 Interactions between Molecules A. Intermolecular forces, ion dipole forces, solvation energy etc B. Thermal energy = energy of motion Chapter 12

10. 12.4 Types of Intermolecular Forces: Dispersion, Dipole-Dipole, and Hydrogen Bonding A. Dispersion force, also known as London forces 1. Weakest intermolecular force 2. Instantaneous dipole 3. Involved in every intermolecular interaction 4. Increases with size and molar mass of compounds/atoms B. Dipole-dipole force- ONLY IN POLAR MOLECULE 1. Stronger force than dispersion forces 2. Strength a function of dipole moment and structure C. Hydrogen bonding 1. Strongest intermolecular forces 2. Only involving F, O, or N bonded to a hydrogen atom within a molecule Intermolecular forces affect state, boiling and melting points

11. EXAM 4- 100 POINTS Part 1 Multiple Choice –Show calculations for partial/full credit 10 questions Part 2 1) VSEPR, Lewis structure predict MOLECULAR geometry, Dipole moment, hybridization around central atom, resonance 2) Sigma and pi bonds eg. N≡N how many sigma and how many pi bonds? 3) Enthalpy of reaction from bond energies STOP HERE for EXAM 4 review SLIDES ARE BEING CONTINUED for finals review

12. Intermolecular Forces in Action: Surface Tension and Viscosity A. Surface tension – why some things more dense than water can still float B. Viscosity – resistance of liquid to flow 12.5 Evaporation and Condensation A. Evaporation and vaporization – liquid to gas B. Condensation – gas to liquid C. Dynamic equilibrium D. Boiling E. Energy change = heat of vaporization 12.6 Melting, Freezing, and Sublimation A. Melting – solid to liquid B. Freezing – liquid to solid C. Energy change = heat of fusion D. Sublimation – solid to gas Chapter 13 Phase diagrams

13. Finals Chapter 1 Conversion e.g. cm-> meter->kilometer, liters-> ml etc. Density calculations, density lab, pay attention to sig fig!! Chapter 2 No. protons, no. neutrons, no. electrons, charge problem Isotopes, writing in the 14 C convention, atomic mass Calculation Correct formulas vs names for chemicals Names of compounds molecular and ionic compounds Finals on Tuesday Dec 20th 9:30-11:30am Room T-109 40 multiple choice questions Good Luck and Happy Holidays!! REVIEW CHAPTERS 1 through 9, 11-13

14. Chapter 3 how many moles, number of atoms/molecules ? Empirical formula- easy one; Molecular formulas- Ratio Balancing equation- write the coefficients?-Sum Calculations in chemical reaction how much reactant will produce how much product, moles grams etc % yield, mass % Chapter 4 Molarity, Dilution, Acid base titration net ionic equation(single-double lab) precipitation reaction stoichiometry Chapter 11 Gases Boyle’s law relation of P and V Charles law relation of V and T Avogadro's law relation of V and N combined gas law PVT relationship PV =nRT, gas laws Gas reactions-stoichiometry-simple one like e.g. in book Molar mass, density, partial pressure-Exam 2 e.g.

15. Chapter Exothermic? Endothermic?- Exam 2! Specific heat –Simple Moles calculation with  H reaction-see quiz 6 Hess’s law Chapter 6 Quantum Bohr’s H-atom -Lab Allowed quantum number orbital, how many electrons can sub-shells have? S-2, p-6 etc. Chapter 7 Atomic structure electronic configuration of atoms and ions, unpaired electrons valence electrons, periodic properties V. important (IE, size) Chapter 8, 9 Bonding, Bond energy calculations Correct Lewis structure Molecular geometry –important-2 questions! Type of bonding? Ionic, Polar covalent, non polar covalent etc. Polar- non polar??, Hybridization-EASY, resonance, VSEPR

16. Chapter 12,13 Vapor pressure- High vapor pressure means low boiling!! Critical tempertaure Triple point, CO 2 and H 2 O phase diagram Polar dissolves polar and non polar dissolves non polar heats of vaporization high implies high boiling point and low vapor pressure etc H-bonding INTER molecular bond between H and O or F or N Electrolytes and metallic bonding; London dispersion forces

17. 6.023x10 23 ATOMS = 1 mole OF ATOMS weighs At. Mass in g 6.023x10 23 MOLECULES = 1 mole OF MOLECULES weighs Molar mass in g Suppose the molecule has formula A 2 B then 1 mole of A 2 B consists of 6.023x10 23 MOLECULES of A 2 B and 2 X 6.023x10 23 ATOMS of A weighs 2 X At. Mass of A in g 1 X 6.023x10 23 ATOMS of B weighs 1 X At. Mass of B in g

18. Important Formulas to memorize Make a list!!