Main Arts and Sciences Science Department General Chemistry Lecture I CHEM-151 Fall 2014 Section 12350 4 Credits 09/18/2014 to 12/14/2014 Modified 08/05/2014 Description The first of two semester courses designed to provide an in-depth introduction to general chemistry for students who plan careers in the health professions, physical sciences, biological sciences, or engineering. Topics discussed include measurement, aqueous reactions, stoichiometry, thermochemistry, atomic structure, bonding, and acids and bases. (F,Sp,Su) Requisites Prerequisite: Reading Level 5 and Writing Level 6 and (Math Level 6 or MATH 112 concurrently) Recommended: (Minimum 2.0 in CHEM 125 or High School Chemistry) and CHEM 161 concurrently Contact Hours Lecture 64 Lab 0 Other 0 Total Hrs 64 Student Learning Outcomes Upon successful completion of this course, students should be able to: A. BASIC CONCEPTS 1. Differentiate among the three states of matter: gas, liquid, and solid. 2. Distinguish between a physical or chemical property or change. 3. Distinguish among pure substances, compounds, elements, mixtures, heterogeneous, homogeneous and solutions. 4. Using the density formula, solve for density, mass or volume given the other two. 5. Interconvert degrees Fahrenheit, Celsius and Kelvin. 6. Knowing the meaning of the prefixes used in the metric system, convert a value in one metric unit to another. 7. Knowing that 2.54 cm = 1 in., 1.06 quarts = 1 liter and 1 lb = 454 grams, convert between English and Metric units. 8. Distinguish among hypothesis, theory, and law. 9. Define and distinguish between precision and accuracy. 10. Determine the correct number of significant figures in a measured or calculated quantity. 11. Express numbers and do calculations in exponential notation. 12. Round off numbers to the correct number of significant figures. 13. Differentiate between exact numbers and measurements. 14. Use dimensional analysis to solve unit conversion problems. B. ATOMS, MOLECULES AND IONS 1. Apply the Law of conservation of mass. 2. State the basic postulates of Dalton's Atomic Theory. 3. Describe the properties of electrons as seen in cathode rays. 4. Describe the studies made by Thomson, Millikan and Rutherford which helped develop our understanding of the subatomic nature of matter. 5. Describe the composition of the atom in terms of protons, neutrons, and electrons; describe the properties of these subatomic particles. 1 of 13
6. Relate the chemical symbol of an element or ion to its mass number, atomic number and charge as well as to the number of protons, neutrons and electrons. 7. Given a periodic table, identify elements in the following sets: metals, nonmetals, transition metals, halogens, noble gases, alkali metals, and alkaline earth metals. 8. Given the chemical formula of any atom, ion or compound listed in the ""Formula and Nomenclature"" handout, give the name and vice versa. 9. Identify the elements that exist as diatomic molecules. 10. Given a molecular formula, indicate how many atoms of a given element are in the molecule. C. STOICHIOMETRY 1. Write a balanced chemical equation from a word description of it. 2. Identify a chemical reaction as combination, combustion, or decomposition. 3. Given the atomic masses of all naturally occurring isotopes of an element, calculate the average atomic mass of the element in amu. 4. Given the name or formula of an element or compound, determine its molar mass. 5. Given a chemical formula or name, convert among moles, grams, amu, molecules and atoms. 6. Determine the empirical formula of a compound from a molecular formula or percent composition. 7. Given the chemical formula of a substance, calculate its percent composition. 8. Given the empirical formula and molar mass of a compound determine its molecular formula. 9. Calculate the mass of a particular substance produced or used in a chemical reaction. 10. Given the amounts of two or more reactants and the chemical equation, determine the limiting reagent and the theoretical yield of the product. 11. Given the theoretical yield and the actual yield, calculate the percent yield. D. AQUEOUS REACTIONS 1. Write equations to show ion formation from electrolytes when they ionize or dissociate. 2. Differentiate between the terms acid and base, strong acid and weak acid. 3. Define and identify substances as strong, weak and non electrolytes. 4. Identify the common strong acids and bases. 5. Given the chemical formula of any atom, ion, or compound listed in the ""Formula and Nomenclature"" handout, give the name and vice versa. 6. Calculate molarity, solution volume, or number of moles of solute, given any two of these quantities. 7. Calculate the concentration of ions present for the solutions of strong electrolytes. 8. Solve problems involving dilution of solutions. 9. Use solubility rules to predict the water solubility of salts. 10. Knowing the solubility rules, predict whether a precipitation reaction will occur. 11. Write balanced molecular, complete ionic and/or net ionic equations. 12. Define and identify end point, Stoichiometric point, standard solution and indicator. 13. Calculate the mass of a substance produced or used in precipitation or neutralization reaction. 14. Identify a chemical reaction as dissociation, combination, decomposition, combustion, neutralization, single replacements, or double replacement. E. GASES 1. Describe how a barometer can be used to measure gas pressure. 2. Convert among various units of pressure, including torr, atmosphere, mm of Hg and Pascals. 3. Solve problems relating to Boyle's Law, Charles' Law, the combined gas law, and Avogadro's Law. 4. Solve problems using to ideal gas equation. 5. Given temperature and pressure, convert between molar mass and gas density. 6. Solve problems involving Dalton's Law of Partial Pressures. 7. Using a balanced chemical equation, convert from a given mass, moles or volume of one gas to mass, moles or volume of another. 8. Calculate the mass of a reactant given the total volume, pressure and temperature of a gas collected over water. 9. Describe the relationship between temperature, average velocity, and kinetic energy. 10. List and explain the principle points of the kinetic molecular theory of gases and describe the major factors responsible for real gases deviating from ideal gas behavior. 11. State Graham's Law and solve problems using Graham's Law. 12. Explain the origin of the correction terms to P and V which appear in Van der Waals equation. F. THERMOCHEMISTRY 2 of 13
1. Define the terms energy, work, kinetic energy, potential energy, power, system, surroundings, state function, enthalpy, internal energy. 2. State and apply the first law of thermodynamics. 3. Relate changes in enthalpy to changes in internal energy. 4. Distinguish between exothermic and endothermic processes. 5. Given the change in enthalpy for a reaction, calculate the heat transferred for a given amount of reactant. 6. Calculate DH for a process given DH values for processes that can be combined (Hess's Law) to yield the process of interest. 7. Calculate the change in enthalpy for a reaction given the standard free energy of formation of reactants and products. 8. Write equations that represent the standard enthalpy of formation of a substance. 9. Define and apply the definition of standard heat of formation. 10. Using the relationship: heat = (specific heat) * (mass) * (Dtemperature), determine the values of one of the variables given the values of the others. 11. Calculate the heat capacity of a calorimeter given the temperature change and quantity of material involved; also, calculate the heat evolved or absorbed in a process from a knowledge of the heat capacity of the system and its temperature change. G. ELECTRONIC STRUCTURE & PERIODIC PROPERTIES 1. Describe the wave nature of electromagnetic radiation. 2. Convert between wavelength and frequency of electromagnetic radiation. 3. Describe the essential features of Planck's quantum theory. 4. Describe the photo electric effect and Einstein's theoretical explanation of it. 5. Explain the origin of line spectra. 6. Describe the Bohr model of the hydrogen atom. 7. Calculate the energy differences between any two electronic states of a hydrogen atom and the wavelength of frequency of the photon involved in the transition. 8. Calculate the wavelength of a particle from its mass and velocity. 9. Describe the Heisenberg uncertainty principle. 10. Explain the significance of the square of the wave function of an electron including the concepts of orbital, electron density and probability. 11. Explain the physical significance of each of the quantum numbers. 12. Describe the shapes of s and p atomic orbitals. 13. Use the aufbau principle, the Pauli exclusion principle and Hund's rule to predict how electrons will fill orbitals in an atom. 14. Give the electronic configuration of selected atoms. 15. Predict and explain the relative sizes of atoms based on their positions in the periodic table. 16. Predict the relationship between the size of a neutral atom and its ions. 17. Define ionization energy and electron affinity. 18. Given a periodic table, predict relative ionization energies of a series of elements. 19. Describe the periodic trends in metallic and nonmetallic character. H. BASIC CONCEPTS OF CHEMICAL BONDING 1. Define the terms: ionic bond, covalent bond, lattice energy, anion, cation, and octet rule. 2. Relate bond order to bond strength and bond length. 3. Write Lewis Dot Structures of ionic and covalent compounds and polyatomic ions. 4. Write the electronic configurations of ions of s- and p- block elements. 5. Recognize examples of molecules on ions that are exceptions to the octet rule. 6. Define resonance, recognize when resonance can exist and draw all possible resonance forms of a particular compound. 7. Calculate formal charges of atoms in a molecule or ion. 8. Given a table of bond energies, estimate the enthalpy change for a specified reaction. 9. Define electronegativity and relate the electronegativity of an atom to its position in the periodic table. 10. Predict relative bond polarities. 11. Relate the number of electron pairs in the valence shell of an atom in a molecule to their geometrical arrangement around the atom. 12. Using Valence Sheel Electron Pair Repulsion (VSEPR) Theory predict the geometry of a molecule or polyatomic ion. 13. Explain how nonbonding electron pairs compress the angles between bonding pairs. 14. Predict whether or not a molecule is polar. I. BONDING THEORY 1. Explain the essential features of the valance bond theory. 2. Using dot structures, determine the number of sigma and/or pi bonds in a molecule or polyatomic ion given the chemical 3 of 13
formula. 3. Explain the concept of hybrid orbitals and its relationship to geometrical structures. 4. Based on Lewis dot structures, describe the hybridization, bonding and geometry about the central atom of an ion or molecule. 5. Explain the concept of delocalization in pi bonds. J. LIQUIDS AND SOLIDS 1. Explain the origin of the various types of intermolecular forces and indicate the kinds of intermolecular forces expected for a substance given its molecular structure. 2. Distinguish between intermolecular and intramolecular forces. 3. Explain the terms viscosity, surface tension, capillary action, meniscus, adhesion and cohesion. 4. List the different phase changes and identify them as endothermic or exothermic. 5. Solve problems involving enthalpy of fusion and enthalpy of vaporization. 6. Describe how the vapor pressure of a substance is effected by temperature and intermolecular forces of attraction. 7. Define boiling point and normal boiling point and know the effect of pressure, intermolecular forces, and molar mass on each. 8. Understand the factors that influence the rate of evaporation and the relationship between boiling and evaporation. 9. Given the critical temperature and pressure of a substance, describe the conditions necessary to convert it from a gas to a liquid. 10. Distinguish between crystalline and amorphous solids and give an example of each. 11. Define unit cell and describe the number of atoms and their location in a simple cubic, body-centered cubic and facecentered cubic unit cell. 12. Use a phase diagram to determine the physical state of a substance. K. SOLUTIONS 1. Explain the solution process (solvation). Include a distinction between electrolyte and nonelectrolyte solution. 2. Solve problems involving mass percent, parts per million (ppm), parts per billion (ppb), mole fraction, molarity and molality. 3. Describe the energy changes that occur during the solvation and relate these to particle interactions. 4. Describe the role of disordering the solution process. 5. Define the terms solubility, miscible, and immiscible. 6. Use the phrase ""like dissolves like"" to predict whether two substances will be soluble in each other. 7. Describe the effects of pressure and temperature on solubilities. 8. Given the solubility of a gas in a liquid at one pressure, determine its solubility at a different pressure. 9. Describe the colligative effects of solute particles on the vapor pressure, boiling point, freezing point and osmotic pressure of a solution. 10. Use Raoult's Law to calculate the vapor pressure of a solution. 11. Calculate the boiling point and freezing point of a solution from colligative properties data. 12. Calculate molar masses and molarities from colligative properties data. 13. Define osmotic pressure and calculate osmotic pressure, molarity, or the Van't Hoff factor given information about the solution. L. CHEMICAL EQUILIBRIUM 1. Given the balanced equation of a reversible reaction, write the equilibrium constant expression. 2. Relate the magnitude of the equilibrium constant to the position of the equilibrium and the relative equilibrium concentrations of reactants and products. 3. Given the equilibrium constant for one reaction, determine the value of the equilibrium constant for a related reaction. 4. Calculate Kc or Kp from each of the following equilibrium concentrations or partial pressure of all species 5. Initial concentrations and one equilibrium concentration. 6. Given the concentrations of all recating species, calculate the reaction quotient, Q, and predict the direction of reaction. 7. Use the equilibrium constant to calculate equilibrium concentrations. 8. Based on LeChatelier's principle explain how relative equilibrium quantities of reactants and products are shifted by changes to the equilibrium conditions. M. ACID-BASE EQUILIBRIA 1. List some general properties of acids and bases. 2. Given the chemical formula of any acid or base in the "Formula & Nomenclature" handout, give the name and vice versa. 3. Given an acid/base reaction, identify the Bronsted, Arrhenius and Lewis acids or bases present. 4. Given an acid or base, write the conjugate base or acid. 5. Describe the autoionization of water and write the Kw expression. 4 of 13
6. Calculate [H+], ph and/or [OH-], poh given the value of any one of the variables or the concentration of a strong acid or base. 7. Write equations demonstrating amphoterism. 8. Calculate all equilibrium concentrations and ph or poh in a solution of a weak acid or base given the Ka or Kb or vice versa. 9. Calculate percent ionization of an acid or base from initial concentrations and Ka or Kb values. 10. Explain the relationship between the strength of an acid and the strength of its conjugate base. 11. Interconvert between the Ka of an acid and the Kb of its conjugate base. 12. Write the Ka expression for the stepwise ionization of a polyprotic acid and understand the significance of the difference in magnitude of Ka1 and Ka2. 13. Write simple acid-base reactions for: a. Ionization of a strong acid b. Dissociation of a weak acid/base c. The addition of a strong acid/base to a weak base/acid d. The reactions of salts with water e. Lewis acid-base reactions 14. Recognize the common strong acids: HCl(aq), HBr(aq), HI(aq), HNO, H SO, HCl10 and HClO. 15. Write the Ka/Kb expression for a weak acid/base. 16. Calculate the Ka/Kb for weak acid/base given the ph of a solution of known concentration. 17. Using a balanced chemical equation for both the cation and anion, predict whether a salt solution will be acidic, basic, or neutral. 18. Use knowledge of chemical structure to predict relative acid-base strength. Materials Tools, Equipment or Apparel (Required of the Student) Scientific calculator Evaluation Criteria 3 2 4 3 4 Type Weight Topic Notes Assignments 8-9% Up to 30 bonus points for homework problems Exams or Tests 57-65% Final Exam 23-26% Quizzes 0-11% 5 of 13
Type Weight Topic Notes Additional Information Examinations: There will be five 100 point examinations given during the semester, as well as a two-part final exam. Part A of the final will cover the same material as the first three hourly exams, and part B will cover material from the last two hourly exams. Each part of the final will contribute 100 points to a student's course total. In addition, if the student has taken all five hourly exams, his/her score on the appropriate part of the final can replace the lowest score from one hourly exam. All exams are timed exams with 60 minutes allotted. Quizzes: Quizzes may also be given but their cumulative value cannot exceed 100 points. Point Distribution Quizzes 0-100 points Homework 70 points Hourly Exams (5) 500 points Two Part Final 200 points Total 770-870 points Breakdown Department/Program Specified 90-100% = 4.0 --- Excellent 84-89.9% = 3.5 --- --------- 78-83.9% = 3.0 --- Good 72-77.9% = 2.5 --- --------- 66-71.9% = 2.0 --- Satisfactory 60-65.9% = 1.5 --- --------- 55-59.9% = 1.0 --- Poor 0-54.9% = 0.0 --- -------- Course Policies Class Attendance/Participation Late Tests and Assignments Other Extra Credit Extra Credit is not available for this course. 6 of 13
Institutional Policies Transfer Potential For transfer information, please consult the LCC website at http://www.lcc.edu/transfer. The MACRAO Transfer Agreement simplifies the transfer of students from one Michigan institution to another. The most current MACRAO Transfer Agreement information can be found at http://www.lcc.edu/transfer/macrao_agreement.aspx. The MACRAO Transfer Agreement will be replaced by the Michigan Transfer Agreement (MTA) which will take effect for students entering Fall 2014 or later. Students who started prior to Fall 2014 will be able to complete the MACRAO Transfer Agreement through Summer 2019, or they may complete the MTA requirements. For additional transfer information contact the LCC Academic Advising Center, (517) 483-1904. Disability Statement Students with disabilities who believe that they may need accommodations in this class are encouraged to contact the Center for Student Access, Gannon Building, Star Zone - Campus Resources (http://lcc.edu/odss) or by calling (517) 483-1924 [TTY (517) 483-1207] as soon as possible to better ensure that such accommodations are implemented in a timely fashion. Student Code of Conduct and General Rules and Guidelines LCC supports a positive educational environment that will benefit student success. In order to ensure this vision, the College has established the LCC Student Code of Conduct and the Student General Rules and Guidelines to ensure the protection of student rights and the health and safety of the College community, as well as to support the efficient operation of College programs. In addition, the College has established guidelines for the redress of grievances by individuals accused in such proceedings. A copy of the most current Code can be found on the College s website at http://www.lcc.edu/catalog/policies_procedures/studentrulesguidelines.aspx#code. It is the responsibility of the student to be familiar with, and abide by, the Student Code of Conduct, as well as the General Rules and Guidelines. Furthermore, the instructor may establish reasonable guidelines within the classroom environment. Violations of the Student Code may be reported to the Office of Student Compliance. Enrollment Verification Class attendance and participation are essential to student success. Class rosters will be updated by the end of the second week of the semester (50% refund period) to accurately reflect student enrollment in this course. Students who have not attended by the end of week two will be administratively dropped and responsible for any required tuition and fee charges. This policy only applies to 16-week (full semester), first 12-week, and first 8-week sections. Additional Items Course Practices An Academic Honors Option is a form of instructional agreement whereby a student can earn an Honors designation on his or her transcript by completing one or more approved enrichment assignments in a non-honors course. This course allows qualified students to complete Academic Honors Options. The Academic Honors Option assignment(s) designed by faculty for this course are explained below. If you elect to complete this work and do so successfully, your achievement will be noted on your official LCC transcript. Please be aware that Academic Honors Options DO NOT earn Honors course credit. Students seeking honors section will complete all the course work assigned to the regular sections. In addition to that, they will solve a set of problems in Mastering Chemistry which will challenge them to think about chemistry and apply their critical thinking skills. The Mastering Chemistry problems include more advanced conceptual and/or mathematical skills often utilizing topics covered in different chapters or exam units. Honors option students will be spending an average of 1 to 2 hours per week more than the students enrolled in the regular section. Detailed Outline of Course Content and Sequencing 7 of 13
CHEMICAL FOUNDATIONS A. Scientific Method 1. Experiment and Observation 2. Hypothesis and Theory 3. Law B. Measurement 1. Units 2. Accuracy, Precision 3. Significant Figures/Exact Numbers 4. Scientific (Exponential) Notation 5. Problem Solving with Dimensional Analysis 6. Temperature Conversions 7. Density C. Classification of Matter 1. States of Matter 2. Pure Substances a. Elements b. Compounds 3. Mixtures a. Homogeneous b. Heterogeneous 4. Physical Change & Chemical Change ATOMS, MOLECULES, AND IONS A. Fundamental Chemical Laws 1. Law of Conservation of Mass 2. Law of Definite Proportions 3. Law of Multiple Proportions B. Dalton s Atomic Theory C. Early Experiments to Characterize the Atom 1. Thomson & the Cathode Ray Experiment 2. Millikan & the Oil Drop Experiment 3. Rutherford & the Gold Foil Experiment D. Atomic Structure 1. Properties of Protons, Electrons and Neutrons 2. Concept of the Nuclear atom 3. Nuclear Symbols 4. Mass Number and Atomic Number E. Molecules and Ions 1. Chemical Formula 2. Structural Formula 3. Covalent and Ionic Bonding F. Periodic Table 1. Metals & Nonmetals 2. Group (Family) & Period 3. Family Names G. Nomenclature (See the nomenclature handout, sections I-III and VII) 1. Binary Compounds a. Molecular b. Ionic 2. Ternary Salts 3. Diatomic Molecules STOICHIOMETRY 8 of 13
A. Atomic Masses & Mass Spectrometry B. The Mole 1. Avogadro s Number 2. Atomic Mass Units 3. Atom, Mole, & Mass problems C. Molar Mass D. Chemical Equations 1. Symbols Used in Equation Writing 2. Balancing Chemical Equations E. Per Cent Composition of Compounds F. Empirical Formulas G. Empirical Formula to Molecular Formula H. Stoichiometric Calculations 1. Mass-Mass Problems 2. Limiting Reagents Problems 3. Theoretical and Percent Yield TYPES of CHEMICAL REACTIONS A. Electrolytes 1. Water as a Solvent 2. Strong, Weak, and Non Electrolytes 3. Strong and Weak Acids and Bases B. Solution Composition 1. Molarity 2. Dilution Problems C. Precipitation Reactions 1. Solubility Rules 2. Describing Reactions a. Molecular Equations b. Complete Ionic Equations c. Net Ionic Equations & Spectator Ions 3. Solution Stoichiometry D. Nomenclature 1. Review sections I-III & VII 2. Acids (binary and ternary) 3. Acid Salts 4. Hydrates E. Acid-Base Reactions 1. Neutralization 2. Titration a. Indicator b. Stoichiometric Point & End Point GASES A. Measurement of Gas Pressure 1. Barometer 2. Pressure Units B. Gas Laws 1. Boyle's Law 2. Charles' Law 3. Avogadro's Law 4. Combined Gas Law 5. Ideal Gas Law a. Molar Mass 9 of 13
b. Density 6. Gas Stoichiometry 7. Dalton s Law of Partial Pressures a. Mole Fraction b. Collecting a gas over water C. Kinetic Molecular Theory of Gases 1. Postulates of Kinetic Theory 2. Root Mean Square Velocity a. The Distribution of Molecular Velocities b. Effusion, Diffusion, and Graham's Law 3. Real Gases a. Deviations from Ideality 4. Van der Waal s equation THERMOCHEMISTRY A. Energy and the First Law 1. The Nature of Energy a. Kinetic/Potential 2. System/Surroundings 3. Exothermic/Endothermic 4. The First Law a. Internal Energy b. Heat and Work c. State Functions B. Enthalpy and Calorimetry 1. Enthalpy of Reaction 2. Heat Capacity 3. Specific Heat C. Hess s Law D. Standard Enthalpy of Formation ATOMIC STRUCTURE AND PERIODICITY A. Electromagnetic Radiation B. The Nature of Matter 1. Planck's Theory 2. Dual Nature of Light 3. The de Broglie Equation 4. The Atomic Spectrum of Hydrogen C. Bohr Theory D. Quantum Mechanics 1. Wave Function 2. The Heisenberg Uncertainty Principle 3. Quantum Numbers 4. Orbital Shapes and Energies a. The Pauli Exclusion Principle b. The Aufbau Principle and the Periodic Table c. Hund s Rule d. Electronic Configurations e. Valence and Core Electrons E. Periodic Trends in Atomic Properties 1. Ionization Energy 2. Electron Affinity 3. Atomic Radius 4. Metals and Nonmetals 10 of 13
BONDING GENERAL CONCEPTS A. Types of Bonds: Ionic, Covalent, and Polar Covalent 1. Bond Energy 2. Bond Length 3. Electronegativity 4. Bond Polarity and Dipole Moment B. Ionic Bonding 1. Electron Configurations of Ions 2. Ionic Sizes/Isoelectric Series 3. Lattice Energy Definition C. Covalent Bonding 1. Bond Energy 2. Bond Orders 3. Bond Energy and Enthalpy Calculations D. Drawing Lewis Dot Structures 1. Octet Rule 2. Exceptions to the Octet Rule 3. Resonance 4. Formal Charge E. Molecular Structure: VSPER 1. Molecular Polarity CHEMICAL BONDING: ORBITALS A. Hybridization(Valence Bond) Theory 1. Orbital Overlap 2. Hybridization 3. Sigma (s) Bonds 4. Pi (p) Bonds LIQUIDS and SOLIDS A. Intermolecular Forces 1. Dipole-Dipole 2. Hydrogen Bonding 3. London Dispersion Forces B. Properties of Liquids 1. Surface Tension 2. Capillary Action 3. Formation of a Meniscus 4. Viscosity C. Solids 1. Crystalline and Amorphous Solids 2. Unit Cells D. Changes of State 1. Vapor Pressure/Evaporation & Condensation 2. Energy Changes 3. Critical Temperature and Pressure 4. Boiling point /Normal Boiling Point 5. Phase Diagrams PROPERTIES of SOLUTIONS A. The Solution Process 1. Expressing Concentration a. Molarity 11 of 13
b. Mole Fraction c. Mass Percent d. Molality 2. Energy Changes/Particle Interactions 3. Enthalpy of Solution 4. Factors Effecting Solubility 5. ppm and ppb B. Colligative Properties 1. Vapor Pressure Lowering 2. Boiling Point Elevation 3. Freezing Point Depression 4. Determination of Molar Mass 5. Osmotic Pressure 6. Van't Hoff Factors CHEMICAL EQUILIBRIUM A. The Concept of Equilibrium B. The Equilibrium Constant Expression, K 1. The Law of Mass Action 2. Equilibrium Expressions Involving Pressure 3. Applications of the Equilibrium Constant a. Extent of Reaction b. Reaction Quotient c. Calculation Involving Equilibrium Constants C. Le Chatelier's Principle 1. Change in Reactant or Product Concentrations 2. Effect of Volume and Pressure Changes 3. Effect of Temperature Changes 4. Catalysts ACIDS & BASES A. The Nature of Acids & Bases 1. General Properties 2. Arrhenius Model of Acids and Bases 3. Bronsted-Lowery Acids and Bases a. Conjugate Acid-Base Pairs b. Acid Dissociation Constants 4. Conjugate Acid-Base Strengths 5. Amphoteric 6. Ion Product of Water - K B. ph Scale C. Strong Acids/Calculating ph D. Weak Acids 1. K a 2. Calculation of ph 3. Percent Dissociation 4. Calculation of Ka from Percent Dissociation E. Bases 1. Strong Bases 2. Weak Bases - Calculation of ph F. Relationship between K and K G. Polyprotic Acids H. Acid-Base Properties of Salt Solutions 1. Hydrolysis a W b 12 of 13
2. Predicting Acidity/Basicity 3. Calculating ph I. Acid-Base Strength Vs Molecular Structure 1. Binary Acids 2. Oxyacids J. Lewis Acids and Bases Additional Outcomes (Optional) 13 of 13