Description:  A comprehensive study of chemical principles designed for students pursuing a career in chemistry or other scientific areas. Emphasis will be on atomic structure, bonding, formulas, equations, nomenclature, and stoichiometry. Also included are states of matter, hybridization, molecular geometry, and gas laws. The laboratory will consist of a number of quantitative experiments designed to teach basic laboratory techniques including the use of common laboratory instrumentation. Three lecture, and three laboratory hours per week.

Prerequisite:  One year of high school chemistry or CHEM 1030 with a grade of C or better and two years of high school algebra and an acceptable placement score, or DSPM 0850.

General Information:  General College Chemistry is a transferable college level sequence which is required in many science programs including pre-medicine, 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.

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

1. List the names, prefixes, and abbreviations of the basic metric units of measure.
2. Use the metric units to measure physical quantities.
3. Calculate physical quantities using metric units.
4. Inter convert metric and English system measurements using dimensional analysis.
5. Convert temperatures between the Fahrenheit, Celsius, and Kelvin scales.
6. Define kinetic energy.
7. Use the kinetic energy equation to calculate the energy of a moving particle.
8. Distinguish between kinetic energy and potential energy.
9. Give examples of different forms of energy and describe the relationship between heat and other forms of energy.
10. Define the important units in which energy is measured and be able to convert from one to another.
11. Explain the meaning of the law of conservation of energy. (First Law of Thermodynamics)
12. Distinguish between exothermic and endothermic processes.
13. Recognize the significant digits in a number.
14. Perform calculations of density.
15. Use the correct procedure for deciding the number of significant digits resulting from a calculation.
16. Round calculated numbers off properly.
17. Use the dimensions of measured quantities in a calculation to determine the unit for the quantity calculated.
18. Differentiate between the three states of matter.
19. Distinguish between elements, compounds, and mixtures.
20. Give the chemical symbols for most frequently encountered elements.
21. Distinguish between the physical and chemical properties of a substance.
22. Describe the composition of an atom in terms of protons, neutrons, and electrons.
23. Give the approximate size, relative mass, and charge of an atom, proton, neutron, and electron.
24. Write the chemical symbol for an element (for example, C), having been given its mass number and atomic number, and perform the reverse operation.
25. Write the symbol and charge for an atom or ion, having been given the number of protons, neutrons, and electrons, and perform the reverse operation.
26. Use the periodic table to predict the charges of monatomic ions.
27. Write the simplest formula for a compound, having been given the charges of the ions from which it is made.
28. Write the name of a simple inorganic compound, having been given its chemical formula, and perform the reverse operation.
29. Calculate the formula weight of a substance.
30. Calculate the per cent by weight of each element in a compound from the empirical formula.
31. Balance chemical equations.
32. Predict the products of a chemical reaction, having seen a suitable analogy.
33. Inter convert number of moles, mass in grams, and number of atoms, ions, or molecules.
34. Calculate the empirical formula of a compound, having been given appropriate analytical data such as elemental percentages or the quantity of CO₂ and H₂O produced by combustion.
35. Calculate the molecular formula, having been given the empirical formula and molecular weight.
36. Calculate the mass of a particular substance produced or used in a chemical reaction (mass-mass problems).
37. Determine the limiting reagent in a reaction.
38. Calculate theoretical yields and percent yields of reactions.
39. Define molarity.
40. Calculate the molarity of a solution using mass and volume data.
41. Solve problems involving inter conversions among molarity, solution volume, and number of moles of solute and serial dilutions.
42. State the current version of the Periodic law.
43. Give examples of elemental properties that are periodic.
44. Explain how electron configurations are related to the arrangement of elements on the periodic table.
45. Explain the terms electrolytes, non electrolytes, dissociation and ionization as they
46. refer to aqueous solutions.
47. Define the terms: "acids", "bases", and "salts".
48. Understand the central ideas of aqueous reactions, precipitation reactions, acid base reactions, oxidation- reduction reactions and displacement reactions.
49. Know the oxidation numbers of the most common elements.
50. Know how to name inorganic compounds.
51. Describe the wave nature of light and other electromagnetic radiation and cite the characteristic speed of these waves.
52. Explain what is meant by the term "spectrum."
53. Use the relationship ln = c, which relates the wavelength (l) and frequency (n) of radiant energy to its speed (c).
54. Describe how the wavelength and frequency differ in the various parts of the electromagnetic spectrum, such as the infrared, visible, and ultraviolet.
55. Explain the essential feature of Planck's quantum theory, namely, that the smallest increment, or quantum, of radiant energy of frequency n that can be emitted or absorbed is hn, where h is Planck's constant.
56. List the assumptions made by Bohr in his model of the hydrogen atom and explain how Bohr's model relates to Planck's quantum theory.
57. Explain the concept of an allowed energy state and how this concept is related to the quantum theory.
58. Explain the concept of ionization energy.
59. Describe the uncertainty principle and explain the limitations it places on our ability to define simultaneously the location and momentum of a subatomic particle, particularly an electron.
60. Diagram energy levels and sub levels for an atom up through the fourth energy level.
61. Use s, p, d, and f notation to identify sub levels.
62. Explain the concepts of orbital, electron density, and probability as used in the quantum mechanical model of the atom.
63. Describe the three quantum numbers used to define an orbital in an atom and list the limitations placed on the values each may have.
64. Correctly write the four quantum numbers for a particular electron in an atom.
65. Describe the correlation between principal energy levels, sub levels and orbitals, and the quantum numbers used to represent these permitted electron positions in an atom.
66. Describe the shapes of s and p orbitals.
67. Recognize the difference in the energies of electrons which have different quantum numbers.
68. Explain the concepts of effective nuclear charge and the screening effect as they relate to the energies of electrons in atoms.
69. State the Pauli exclusion principle as it applies to ordinary particles and also as it applies to electrons in an atom.
70. State Hund's rule and demonstrate its application in orbital filling.
71. Describe the various blocks of elements in the periodic table in terms of the type of orbital being occupied by electrons in that block (s, p, d, and f blocks).
72. List the names and give the locations in the periodic table for the active metals (s-block), representative elements (p-block), transition metals (d-block), and inner transition metals (f-block).
73. Write the electron configuration for any element when its place in the periodic table is given.
74. Explain the effect of increasing nuclear charge on the atomic volume in many electron atoms.
75. Define "ionization energy" for an element.
76. Explain the general variations in first ionization energies among the elements.
77. Explain the observed changes in values of successive ionization energies for a given element.
78. Explain the variation in atomic radii for the elements in a period.
79. Explain the variation in atomic radii for the elements in a group or family.
80. Explain the concept of electron affinity and its relationship to ionization energy.
81. Write the orbital diagram representation for electron configurations of atoms using boxes with half arrows.
82. Determine the number of valence electrons for any atom and write its Lewis symbol.
83. Describe the origin of the energy terms that lead to stabilization of ionic lattices.
84. Predict on the basis of the periodic table the probable formulas of ionic substances formed between common metals and nonmetals.
85. Write the electron configurations of ions.
86. Describe the effects of gain or loss of electrons on atomic radii in producing ionic radii.
87. Explain the concept of an isoelectronic series and the origin of changes in ionic radius within such a series.
88. Describe the basis of the Lewis theory and predict the oxidation numbers of common nonmetallic elements from their position in the periodic table.
89. Write the Lewis structures for molecules and ions containing covalent bonds, using the periodic table.
90. Write resonance forms for molecules or poly atomic ions that are not adequately described by a single Lewis structure.
91. Explain the significance of electronegativity and in a general way relate the electronegativity of an element to its position in the periodic table.
92. Predict the relative polarities of bonds using either the periodic table or electronegativity values.
93. Relate bond energies to bond strengths.
94. Assign oxidation numbers to atoms in molecules and ions.
95. Give the meaning of the terms oxidation, reduction, and oxidation reduction reactions. Determine whether oxidation reduction has occurred in a reaction; if it has, be able to identify the substance that is oxidized and the one that is reduced.
96. Complete and balance redox equations using the method of half reactions.
97. Assign acceptable names to simple ionic compounds and ions.
98. Describe the general differences in physical properties between substances with ionic bonds and those with covalent bonds.
99. Describe how the water solubility of a metal oxide is related to cation size and charge.
100. Describe the reactions of oxides with water, metallic oxides with acids, and nonmetal oxides with bases.
101. Relate the number of electron pairs in the valence shell of an atom in a molecule to the geometrical arrangement around that atom.
102. Explain why unshared electron pairs exert a greater repulsive interaction on other pairs than do shared electron pairs.
103. Predict the geometrical structure of a molecule or ion from its Lewis structure.
104. Determine whether or not a molecule has a dipole moment.
105. Explain the concept of hybridization and its relationship to geometrical structure.
106. Assign a hybridization to the valence orbitals of an atom in a molecule, knowing the number and geometrical arrangement of the atoms to which it is bonded.
107. Formulate the bonding in a molecule in terms of p bonds and s bonds, from its Lewis structure.
108. Explain the concept of delocalization of electrons in a p bond system.
109. Explain the concept of orbital overlap.
110. Describe how covalent bonds are formed by overlap of atomic orbitals.
111. Using VSEPR theory and Molecular Orbital hybridization (sp, sp², sp³, sp³d, and sp³d²), explain the three dimensional shape of molecules.
112. Explain Arrhenius and Bronsted-Lowry theories of acid and bases.
113. Know the reactions of acids and bases and how they are prepared.
114. Explain Lewis theory of acids.
115. Balance oxidation reduction equations under acidic or basic conditions.
     
116. Know how to perform a titration of an acid and a base and how to perform the necessary calculation to ascertain molarity and normality.
117. Understand the properties of a primary standard.
118. List the properties of gases and compare gases, liquids, and solids.
119. Use the Kinetic Molecular Model to describe pressure and how it is measured.
120. Describe the relationships among pressure, volume, temperature, and amount of gas (Boyle's Law, Charles's Law, Avogadro's Law, and the combined Gas Law).
121. Use Boyle's, Charles's, Avogadro's, and the Combined Gas Laws to calculate changes in pressure, volume, temperature, and amount of gas.
122. Use the ideal gas equation to perform pressure, volume, temperature, and mole calculations.
123. Describe how mixtures of gases behave and predict their properties (Dalton's Law of Partial Pressures).
124. Describe how the Kinetic Molecular Model is consistent with the observed gas laws.
125. Describe the molecular features that are responsible for the non ideal behavior of real gases and explain when this non ideal behavior is important.
     

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:  

The student's grade in this chemistry course will be determined according to the following weighting scheme:  

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.  

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.  Your instructor will inform you how a missed exam will be handled.

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 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 1110 Topical Lecture Outline:

• Historical Perspectives
• Basic Concepts: Definitions, Measurement, Metric System, Dimensional Analysis
• Description of Matter: States, Elements, Compounds and Mixtures, Atoms, Molecules and Ions
• The Atomic Theory, Atomic Structure
• Periodic Table
• Nomenclature
• Stoichiometry: Conservation of Mass, Chemical Equations, Atomic and Molecular Weights, The Mole Concept, Formulas for Compounds, Calculations from Equations, Solution Stoichiometry
• Electron Theory of Atomic Structure: Nature of Radiant Energy , The Quantum Theory, Bohr's Model, Wave Mechanical Model, Shapes of Orbitals, Orbital Filling Rules, Electron Shells in Atoms
• Periodic Functions: Ionization Energy, Atomic Volume, Elements and the Periodic Table
• Historical Development : Modern Periodic Table, Metals & Non-Metals, Representative Metals, Representative Non-Metals
• Basic Concepts of Bonding: Octet Rule; Symbols, Ionic Bonding, Covalent Bonding, Lewis Structure, Bond Polarity, Electronegativity, Oxidation Numbers, Geometry of Molecules, VSEPR, Valance Bond Theory, Multiple Bonding, Resonance.
• Reactions in Aqueous Solutions, Acids, Bases, Salts, Molarity, Normality, and Titrations.
  
• Gases and the Kinetic-Molecular Theory
  

(Updated 8/20/08, C.R. Snelling)