Prerequisite: CHEM 1110 with a grade of C or higher.

General Information: General College Chemistry is a transferable college level sequence which is required in many science programs including pre medical, pre dentistry, pre engineering, pre pharmacy, and pre veterinary medicine. As such, it is a comprehensive introduction to the entire field of chemistry with considerable stress placed on mathematical applications and problem solving. One of the most frequent difficulties in chemistry is an inability to perform simple arithmetic and algebraic operations. Therefore, a knowledge of basic algebra is a must if the student is to succeed in general chemistry.

Instructor:

General Education Goal: The general education goal of this course is to provide Scientific information and instruction in the thought processes involved in the scientific method of inquiry.

General Education Outcomes: As a result of completing this course successfully, students will have demonstrated an acceptable level of mastery of designated scientific facts, concepts, and principles and demonstrated an understanding of and ability to apply the scientific method of inquiry. Mastery of course contents will have indicated the acquisition of a foundation suitable for pursuing further course work in chemistry.

Other Goals: This course also seeks to provide opportunities to apply problem solving skills and to acquire critical skills for the assessment and evaluation of values. Additionally, this course will require effective communication skills in both receiving and giving information.

Outcome Statement: Upon completion of this course the student will have demonstrated the ability to:

1. Distinguish between gases, liquids, and solids on a molecular level.
2. Employ the kinetic molecular theory and concept of intermolecular attractions to explain the properties of each phase, such as surface tension, viscosity, vapor pressure, and boiling and melting points.
3. Explain the nature of the equilibria that may exist between phases.  Account for the enthalpy changes that accompany phase changes.
4. Explain the relation between pressure, vapor pressure, temperature, and boiling point.
5. Explain the meaning of the terms critical temperature and critical pressure and account for the variation in critical temperatures of different substances in terms of intermolecular forces.
6. Describe the various types of intermolecular attractive forces, indicate how each arises, and indicate the manner in which each varies with distance.
7. Predict, for any particular substance of known structure, which types of intermolecular forces may be operative and which particular type is of major importance.
8. Describe the nature of the hydrogen bond and distinguish those molecular systems in which hydrogen bonding is likely to be important.
9. Define molarity, molality, mole fraction, normality, and weight percentage and calculate concentrations in any of these concentration units.
10. Convert concentration in one concentration unit into any other (given the density of the solution where necessary).
11. Give definitions of the qualitative terms used to describe solutions: dilute, concentrated, saturated, unsaturated, and supersaturated.
12. Describe the solution process, including the molecular or ionic associations made and broken when a substance dissolves.
13. Describe the role of disorder (entropy) in the solution process.
14. Rationalize the solubility's of substances in various solvents in terms of their molecular structures and intermolecular forces.
15. Discuss the effects of pressure and temperature on solubility's of gases.
16. Predict which substances are electrolytes and which are non electrolytes.
17. Predict the course of reactions involving formation of a precipitate, gas, or non electrolyte, and write net ionic equations for these reactions.
18. Describe the effect of solute concentration on solvent vapor pressure.
19. Determine the molar mass of a solute from the magnitude of the effect of a known concentration of solute on one of the colligative properties of a solvent.
20. Explain the difference between the magnitude of changes in colligative properties caused by electrolytes compared to those caused by non electrolytes.
21. Calculate the average rate of a reaction, given the concentrations of a reactant or product at the beginning and end of a time interval.
22. Explain the meaning of the term rate constant and state the units associated with rate constants for first- and second order reactions.
23. Calculate rate, rate constants, or reactant concentration, given two of these together with the rate law.
24. Use the equation for first order reactions and for second order reactions to determine graphically whether the rate law for a reaction is first or second order.
25. Explain the concept of reaction half-life and describe the relationship between half-life and rate constant for first order reaction.
26. Describe the effect of a catalyst on the energy requirements for a reaction.
27. Explain the functions of a catalytic converter in an automobile exhaust system.
28. Write the equilibrium constant expression for a balanced chemical equation, whether heterogeneous or homogeneous.
29. Numerically evaluate K_c from a knowledge of the equilibrium concentrations of reactants or products or from the initial concentrations and the equilibrium concentration of at least one substance.
30. Calculate the reaction quotient, Q_c, and comparison with the value of K_c to determine whether a reaction is at equilibrium.  If it is not at equilibrium, you should be able to predict in which direction it will shift to reach equilibrium.
31. Use the equilibrium constant to calculate equilibrium concentrations when either the equilibrium concentrations of all but one substance or the initial concentrations together with the equilibrium concentration of one substance is known.
32. Explain how the relative equilibrium concentrations of reactants and products are shifted by changes in temperature, pressure, or the concentrations of substances involved in the equilibrium reaction.
33. Explain how the change in equilibrium constant with change in temperature is related to the heat change in the reaction.
34. Describe the effect of a catalyst on a system as it approaches equilibrium.
35. Explain the process that occurs when an acid dissolves in water.
36. Describe the forms in which the proton exists in water.
37. Define an acid, base, conjugate acid, and conjugate base in terms of the Bronsted-Lowry theory of acids and bases.
38. Explain what is meant by the autoionization of water and write the ion product constant expression, K_w.
39. Explain what is meant by pH and calculate pH from a knowledge of [H⁺] or [OH^-]; also be able to perform the reverse operations.
40. Calculate [OH^-] from pOH and [H⁺] from pH, and be able to perform the reverse operations.
41. Memorize the common strong acids and bases.
42. Write the acid ionization constant expression for any weak acid in water, K_a.
43. Calculate [H⁺] for a weak acid solution in water, knowing acid concentration and K_a.
44. Write the base ionization constant expression for a weak base in water, K_b.
45. Calculate [H⁺] for any weak base solution in water, knowing base concentration and K_b.
46. Explain the relationship between an acid and its conjugate base or between a base and its conjugate acid and calculate K_b from a knowledge of K_a, or vice versa.
47. Predict whether a particular salt solution will be acidic, basic, or neutral.
48. Explain how acid strength relates in a general way to the nature of the H-X bond.
49. Predict the relative acid strengths of oxoacids.
50. Define an acid or base in terms of the Lewis acid base theory.
51. Predict qualitatively and calculate quantitatively the effect of an added common ion on the pH of an aqueous solution of a weak acid or base.
52. Calculate the concentrations of each species present in a solution formed by mixing an acid and a base.
53. Describe how a buffer solution of a particular pH is made and how it operates to control pH.
54. Calculate the change in pH of a simple buffer solution of known composition caused by adding a small amount of strong acid or base.
55. Set up the expression for the solubility product constant for a salt.
56. Calculate K_sp from solubility data and solubility from the value for K_sp.
57. Calculate the effect of an added common ion on the solubility of a slightly soluble salt.
58. Predict whether a precipitate will form when two solutions are mixed, given appropriate K_sp values.
59. Explain the effect of pH on a solubility equilibrium involving a basic or acidic ion.
60. Explain complex formation in relation to dissolving a slightly soluble solute.
61. Define the term spontaneity and apply it in identifying spontaneous processes.
62. Describe how entropy is related to randomness or disorder.
63. State the second law of thermodynamics.
64. Predict whether the entropy change in a given process is positive, negative, or near zero.
65. Describe how and why the entropy of a substance changes with increasing temperature or when a phase change occurs, starting with the substance as a pure solid at 0 K.
66. Calculate delta S° for any reaction from tabulated absolute entropy values, S°.
67. Define free energy in terms of enthalpy and entropy.
68. Explain how the sign of the free energy change, delta G, determines whether or not a process is spontaneous in the forward direction.
69. Calculate the standard free energy change at constant temperature and pressure, delta G°, for any process from tabulated values for the standard free energies of reactants and products.
70. List the usual conventions regarding standard states in setting the values for standard free energies.
71. Predict how delta G will change with temperature, given the signs for delta H and delta S.
72. Describe the relationship between delta G and the work that can be derived  from a spontaneous process.

The expected outcomes for the course will be assessed at frequent intervals by various pedagogical techniques including homework assignments, weekly laboratory reports, major unit tests and a comprehensive final examination. The laboratory activities afford the opportunity to assess manual, theoretical, and written communication skills. The student will also be encouraged to improve verbal communication skills. Unit tests will be balanced with computational skills and factual material, and the final examination will require that the student be able to assimilate and integrate information from various units. Multiple choices and other objective forms may be used to identify and define terms, ideas and concepts.

Grades:  The grades in all chemistry courses are based on the following scale:  

One requirement of the course is that every student take the final examination.  Failure to take the final examination will result in a grade of F for the course. In the case where a final examination is missed and the instructor has been notified in advance, it may be possible (at the discretion of the instructor) for the student to receive a grade of I. However, a grade of I must be converted to another letter grade by completing work prior to the end of the seventh week of the succeeding semester, otherwise the I will be automatically converted to a grade of F.  Students will not be allowed to register for chemistry courses on an Audit basis.

Attendance:  Attendance at all lecture and laboratory meetings is expected. Persistent unexcused absences exceeding 30% of the lecture meetings may result upon approval of the instructor and with approval of the Vice President of Academic Affairs in the Administrative Withdrawal of the student from that class. See the College Catalog for the last day to withdraw from the course or the College without penalty, and for a further explanation of the Administrative Withdrawal Policy.  If you are receiving Title IV financial assistance (Pell Grant, Student Loan or SEOG Grant), you must regularly attend class (a minimum of the first full week) or be subject to repay PART or ALL of the Federal Financial Aid you received for the semester.

Make-up Exams:   There are no makeup exams.  However, you will have the opportunity to drop your lowest exam score.  Any other missed exams will result in a grade of zero (0) for that exam.

Make-up Labs:  There are no makeup labs.  However, you will have the opportunity to drop your lowest lab score.  You can use this to drop your lowest grade or to replace a lab you were not able to attend.  Any other missed labs will result in a grade of zero (0) for that lab.

Deficiencies: A student who has an average of D or F on work completed and evaluated up to mid semester will receive a deficiency slip by mail indicating the need for improvement if a course grade of C or better is to be achieved. If a student receives a deficiency slip, he/she should explore with the instructor the wisdom of dropping or continuing in the course.

Cheating: Cheating on a test or exam will incur a grade of zero (0) on that test or exam.

Recommended Problems: These problems and exercises may be handed in for grading upon the direction of your instructor. It is not unusual for questions or problems similar to those assigned to appear on tests or examinations.

Other Regulations: A student is bound by all rules and regulations appearing in the Student Handbook.

Chemistry 1120 Topical Lecture Outline:

• Liquids and Solids
• Kinetic Molecular Model
• Intermolecular Forces
• Properties of the Liquid and Solid States
• Phase Diagrams
• Solutions
• The Dissolution Process
• Factors Affecting Dissolution and Solubility
• Colligative Properties
• Colloids
• Chemical Thermodynamics
• The First Law of Thermodynamics
• Enthalpy
• Standard States and Molar Enthalpies
• Hess's Law
• Bond Energies
• Relationship between DH and DE
• Spontaneity
• The Second Law of Thermodynamics
• Entropy and Free Energy
• Chemical Kinetics
• Rates of Reactions
• The Effects of Temperature and Concentration on Rate
• Collision Theory
• Reaction Mechanisms
• Catalysts
• Chemical Equilibrium
• Equilibrium constant
• Reaction Quotient
• Heterogeneous Equilibria
• Factors affecting Equilibria
• Relationship between K_p and K_c
• Relationship between DG° and the Equilibrium Constant
• Ionic Equilibria I:  Acids and Bases
• Autoionization of Water
• pH and pOH Scales
• Ionization Constants for Weak Acids and Weak Bases
• Acid Base Indicators
• Common Ion Effect and Buffer Solutions
• Polyprotic Acids
• Ionic Equilibria II:  Hydrolysis and Titrations
• Salts of Strong Soluble Bases and Strong Acids
• Salts of Strong Soluble Bases and Weak Acids
• Salts of Weak Bases and Strong Bases
• Salts of Weak Bases and Weak Bases
• Titration of Strong Acids Strong bases, Weak Acids Strong Bases, Weak Acids Weak Bases
• Ionic Equilibria III:  The solubility Product Principle
• Solubility Product Constants and Their Uses
• Dissolving Precipitates
• Electrochemistry
• Galvanic cells
• Electrochemical cells
• Batteries
  

Equal Opportunity Statement: Volunteer State Community College is an equal opportunity Affirmative Action Educational Institution.  No person shall be excluded from participation in, be denied the benefit of, or be subjected to discrimination under any program or activity of the College because of race, color, national origin, age, or handicap.  The College also complies with the Age Discrimination in Employment Act of 1967, as amended and with the Vietnam Era Veteran's Readjustment Act of 1974.  The commitment to equal opportunity applies to all aspects of the recruitment, employment and education of individuals at all levels throughout the College.

(Updated 1/10/08, C.R. Snelling)