1. REVIEW III - CHAPTERS 8,10, 11, 12

2. 8.1 Stoichiometry—What is it?- Quantities in a chemical reactions 8.2 Mole: Counting Atoms by the Gram A. Pair =2; dozen = 12; 1 mole = 6.02 x 10 23 B. 1 mole of Au atoms = 6.02 x 10 23 Au atoms C. The periodic table gives the mass of 1 mole of that atom in grams 1 mole of S atoms = 32.06 g S ; 1 mole of C atoms = 12.01 g C D. Unit conversions between grams and moles of atoms Extend idea to Molecules by the Gram A. Mass of 1 mole H 2 O = mass of 2 moles of H + 1 mole O B. Calculations performed in a similar fashion Chemical Formulas as Conversion Factors A. Molecular formula gives ratios of atoms in molecule B. Converting moles of atoms to moles of molecules C. Converting grams of atoms to grams of molecules Conversions and Reactions Moles  moles of reactants and products Moles  Grams of reactants and products Grams  Grams of reactants and products STOICHIOMETRY and MOLE Chapter-8

3. 8.4 OMIT 8.5 Learn the % composition Mass Percent Composition of Compounds A. Percentage of molecule's total mass that is due to atom X B. Convert between mass of element and mass of compound C. Mass Percent Composition from a Chemical Formula 8.6 Calculating Empirical Formulas for Compounds A. Empirical formula is the smallest whole number ratio of atoms B. Mostly different from molecular formula C. Many different compounds may have the same empirical formula Calculating Molecular Formulas for Compounds A. Empirical formula molar mass B. Compare molar mass to empirical formula mass C. Ratio R = (Molar mass/emp. Mass) = (molecular formula/Emp. Formula) D. Ratio R is the nearest whole number

4. 10.1 Why Does Matter Exist in Different Phases? Interactions between Molecules Intermolecular forces Thermal energy = energy of motion A.Properties of solids Highest Intermolecular forces Definite shape and Volume B. Properties of liquids Weaker Intermolecular forces compared to solids, stronger than gases Definite Volume NO shape C. Properties of Gases Weakest Intermolecular forces No Definite Volume or shape, Lowest Density Evaporation and Condensation A. Evaporation and vaporization – liquid to gas B. Condensation – gas to liquid C. Boiling, Boiling Point D. Energy change going from liquid to gas is to overcome stronger intermolecular forces in liquids Intermolecular Forces Chapter 10

5. 10.2 Types of Intermolecular Forces: Dispersion, Dipole-Dipole, and Hydrogen Bonding A. Dispersion force, also known as London forces 1. Weakest intermolecular force due to Instantaneous dipole 2. Involved in every molecule/atom intermolecular interaction 3. Increases with molar mass 4. Only operative force for non-polar molecules B. Dipole-dipole force- IN POLAR MOLECULES ONLY 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 D. Ionic Bonding Both inter and Intra molecular force, stronger than rest of the intermolecular forces mentioned above 10.3 Types of Crystalline Solids: Molecular, Ionic, and Atomic A. Molecular solids- Lower melting points 1. Composite units are molecules 2. Examples, dry ice( CO 2 (s) ), ice ( H 2 O(s) ) B. Ionic solids-High melting points 1. Composite units are formula units 2. Examples, NaCl, CaF 2 C. Atomic solids- 1. Covalent atomic solids-High melting points 2. Nonbonding atomic solids –very low melting points 3. Metallic atomic solids-Variable melting STRONGER intermolecular forces are Higher is the Boiling and Melting points

6. 11.1 Describing the Ideal Gas Kinetic Molecular Theory: A Model for Gases A. A gas is a collection of particles in constant straight line motion B. Particles neither attract nor repel one another C. Between species is mostly empty space D. Average kinetic energy proportional to absolute (Kelvin) temperature 11.2 Pressure: The Result of Constant Molecular Collisions on the wall A. Pressure = force/area B. Units of pressure: Atmosphere (atm), Millimeter of mercury (mm Hg) Torr; Pounds per square inch (psi); Pascal (Pa) Boyle's Law: Pressure and Volume A. P  1/V or P is inversly proportional to V B. Temperature and number of moles must be constant Gases (Omit Section 11.3) (No difficult math problems) P 1 V 1 = P 2 V 2 Chapter 11

7. Charles's Law: Volume and Temperature A. V  T B. Pressure and number of moles of gas must be constant Avogadro's Law: Volume and Moles A. V  n B. Pressure and temperature must be constant Rule 1: P is proportional to 1/V P α 1/V Rule 2: P is proportional to T Rule 3: P is proportional to n Overall P α nT/V PV = nRT IDEAL GAS LAW, where P is pressure, V is Volume, T is Temperature, n is # of moles, R is gas constant

8. 13.1 What is Solution?: Homogeneous Mixtures A. Any homogenous mixture is a solution B. Solvent: major component C. Solute: minor component, can be more than one 13.2 Energetics in Solution?: Dissociation in solution into Ions, Ion-Dipole forces 13.3OMIIT 13.4 Solutions of Solids Dissolved in Water A. Solubility and Saturation 1. Saturated: solvent holding as much solute as it can 2. Unsaturated: solvent can hold more solute 3. Supersaturated: solvent holding more than the maximum amount of solute B. Solubility is temperature dependent (Solid in liquid, solubility increases with temperature) C. How to make saturated solutions D.Miscibility Solutions of Gases in Water: A. Solubility increases with pressure ; and DECRESES with temperature B. Dilute vs. concentrated 12.5 Getting Unlikes to Dissolve-Soaps and Detergents –Soaps Soaps have polar and non-polar ends which can make non polar grease get washed away by water SOLUTIONS Chapter 13

9. 12.7 Specifying Solution Concentration: Mass Percent A.Mass percent = B.B. Conversion factor between amount of solution and amount of a particular solute 12.6 Specifying Solution Concentration: Molarity A. Molarity (M) = B. Volume of solution, not just solvent C. Units always mol/L 12.8 Solution Stoichiometry -TITRATIONS A. Balanced chemical equations give molar ratios only At end point moles of H+ = moles of OH- B. Convert volume to moles using molarity, then use balanced chemical equation Or At end point moles of H+ = moles of OH- 1:1 Stoichiometry Ma Va = Mb Vb (eg. HCl vs NaOH) DiproticAcid : 1hydroxybase 2* Ma Va = Mb Vb (H 2 SO 4 vs NaOH) DiproticAcid/dibasic : 2hydroxybase 2* Ma Va = 2*Mb Vb (H2SO 4 vs Ba(OH 2 )) DiproticAcid : tribasic: 2* Ma Va = 3*Mb Vb (H 2 SO 4 vs Al(OH) 3 ) Remember that Mass of solution = mass solvent + mass solute

10. CH19 grade calculation for my sections only Exams (4)100 (60%) Finals100 (15%) Home works100 (3%) Lab (12 out of 14 labs) Averaged to 100% 100 (25%) Attendance/Class participation/Curving 100 (3%) Total 106 Grades curved at the end!!!!!!!!!!!!!!!!!!!!!! Below 60%F 60-64%D 65-69% D+ 70-74% C 75-79%C+ 80-84%B 85-89%B+ Above 90%A