12/26/2024 9:07:18 PM	CHEM 1A Course Outline as of Fall 1999	Changed Course
CATALOG INFORMATION
Discipline and Nbr:  CHEM 1A	Title:  GENERAL CHEMISTRY
Full Title:  General Chemistry
Last Reviewed:5/13/2019

Units		Course Hours per Week			Nbr of Weeks	Course Hours Total
Maximum	5.00	Lecture Scheduled	4.00		17.5 max.	Lecture Scheduled	70.00
Minimum	5.00	Lab Scheduled	3.00		17.5 min.	Lab Scheduled	52.50
		Contact DHR	0			Contact DHR	0
		Contact Total	7.00			Contact Total	122.50
		Non-contact DHR	0			Non-contact DHR Total	0

Total Out of Class Hours:  140.00

Total Student Learning Hours: 262.50

Title 5 Category:  AA Degree Applicable
Grading:  Grade Only
Repeatability:  00 - Two Repeats if Grade was D, F, NC, or NP
Also Listed As: 
Formerly: 

Catalog Description:

An introduction to the fundamental facts and principles of chemistry including atomic structure, nomenclature, bonding, stoichiometry, oxidation-reduction reactions, properties of ideal and real gases, kinetic-molecular theory, properties of solutions, liquids and solids, phase equilibria, colligative properties, acid-base reactions, molecular geometry and chemical bonding theories.

Prerequisites/Corequisites:
Chem 51 (formerly Chem 110) or Chem 55 or placement on the Chemistry Diagnostic Test AND Math 155 or two years of high school algebra or equivalent.

Recommended Preparation:

Limits on Enrollment:

Schedule of Classes Information
Description:

First semester of a one year program of general chemistry.

(Grade Only)
Prerequisites:Chem 51 (formerly Chem 110) or Chem 55 or placement on the Chemistry Diagnostic Test AND Math 155 or two years of high school algebra or equivalent.
Recommended:
Limits on Enrollment:
Transfer Credit:CSU;UC.
Repeatability:00 - Two Repeats if Grade was D, F, NC, or NP

ARTICULATION, MAJOR, and CERTIFICATION INFORMATION

Associate Degree:		Effective:	Fall 1981	Inactive:
Area:	C	Natural Sciences
CSU GE:	Transfer Area			Effective:	Inactive:
	B1	Physical Science		Fall 1981
	B3	Laboratory Activity
IGETC:	Transfer Area			Effective:	Inactive:
	5A	Physical Sciences		Fall 1981
	5C	Fulfills Lab Requirement
CSU Transfer:	Transferable	Effective:	Fall 1981	Inactive:	Fall 2020
UC Transfer:	Transferable	Effective:	Fall 1981	Inactive:	Fall 2020
C-ID:
CID Descriptor: CHEM 110	General Chemistry for Science Majors I, with Lab	SRJC Equivalent Course(s): CHEM1A OR CHEM4A OR CHEM3A AND CHEM3AL
CID Descriptor: CHEM 120S	General Chemistry for Science Majors Sequence A	SRJC Equivalent Course(s): CHEM1A AND CHEM1B OR CHEM4A AND CHEM4B OR CHEM3A AND CHEM3AL AND CHEM3B

Certificate/Major Applicable: Not Certificate/Major Applicable

COURSE CONTENT

Outcomes and Objectives:
At the conclusion of this course, the student should be able to:

Upon completion of the course, the student should be able to:

1.  solve problems involving the concepts listed under Course Content;

2.  solve problems using SI units and dimensional analysis;

3.  write concise explanations describing various chemical phenomena

    studied;

4.  write and interpret balanced chemical equations;

5.  describe and identify various types of colloids;

6.  write balanced chemical equations for oxidation-reduction reactions;

7.  write balanced chemical equations for precipitation reactions;

8.  write balanced chemical equations for acid-base reactions;

9.  express numerical data and results to the proper number of

    significant figures;

10.  describe different models of atomic structure;

11.  use standard chemical notation and nomenclature;

12.  predict the outcomes of combination, decomposition, single-

    displacement, metathesis, and combustion reactions;

13.  identify strong electrolytes, weak electrolytes and nonelectrolytes;

14.  apply the Arrhenius, Bronsted-Lowry and Lewis models of acid-base

    theory;

15.  calculate the mass percentages of the elements from the formula of

    a compound;

16.  determine the empirical formula of a compound from elemental

    composition data;

17.  derive the molecular formula of a compound from the empirical

    formula;

18.  predict the amounts of reactants and products involved in a

    chemical reaction;

19.  solve limiting-reactant problems;

20.  calculate theoretical and actual percentage yields;

21.  perform calculations involving mass percentage, molarity, normality,

    molality and mole fraction units;

22.  solve solution stoichiometry problems;

23.  derive predictions of total pressure partial pressures, volume,

    temperature, moles or mass utilizing the ideal gas laws;

24.  relate properties of gases to the kinetic-molecular theory;

25.  predict deviations from ideal behavior in real gases;

26.  calculate molecular weights of gases from Graham's Law and the

    Dumas method;

27.  apply the First Law of Thermodynamics;

28.  describe colligative properties of solutions;

29.  write and interpret thermochemical equations;

30.  use Hess's Law to calculate enthalpies of reaction from standard

    heat of formation;

31.  derive enthalpies of reaction from calorimetric data;

32.  calculate wavelength, frequency, speed and energy of electromagnetic

    radiation;

33.  describe quantum effects of atoms and photons;

34.  describe the energy level diagram and spectral series for atomic

    hydrogen;

35.  calculate wavelength and momentum using the deBroglie relationship;

36.  use the Heisenberg principle to predict uncertainty in position or

    momentum;

37.  explain the basis of operation of chromatographic separation, mass

    spectrometry, scanning tunneling microscopy, x-ray diffractometry,

    nuclear magnetic resonance, and infrared spectroscopy;

38.  describe the significance of the four quantum numbers n, l, m and s;

39.  use the Aufbau Principle to derive the ground-state electronic

    configurations of the elements;

40.  apply Hund's Rule and predict the number of unpaired electrons in

    an atom;

41.  distinguish between diamagnetic and paramagnetic behavior;

42.  describe the relationship between electronic configuration and atomic

    radius, ionization energy, electron affinity and electronegativity;

43.  define ionic bonding, and apply the Born-Haber cycle to predict the

    stability of ionic crystalline solids;

44.  define covalent bonding utilizing Lewis dot structures;

45.  predict the existence of polar bonds and dipole moments in molecules;

46.  describe delocalized bonding and resonance structures;

47.  explain the factors that affect solubility;

48.  calculate formal charges, bond orders, oxidation numbers and

    coordination numbers;

49.  calculate enthalpies of reaction using bond dissociation energies;

50.  apply the valence-shell electron-pair repulsion model to predict

    molecular geometries;

51.  describe covalent bonding using the Valence Bond Theory;

52.  use Molecular Orbital Theory to describe bonding in homonuclear

    diatomic molecules;

53.  describe the nature of solids, liquids, gases and phase changes;

54.  construct and interpret phase diagrams;

55.  describe intermolecular forces;

56.  relate atomic radii to unit cell dimensions;

57.  describe the metallic bonding, covalent network solids and

    semiconductors;

58.  recognize the shapes of the unit cells for the seven fundamental

    crystal systems;

In the laboratory upon completion of the course, the student should be

able to:

1.  observe all of the fundamental safety procedures;

2.  properly dispose of waste chemicals;

3.  manipulate standard laboratory apparatus;

4.  perform gravimetric and titrimetric analyses

5.  collect and analyze scientific data using graphical and statistical

    methods;

6.  summarize lab results in both formal and informal report formats;

7.  use a Macintosh personal computer (or equiv) to perform word

    processing, spreadsheet computations, graphing and statistical

    calculations for lab reports.

Topics and Scope

LECTURE MATERIAL

1.  Introduction to Scientific Method & Measurements

    a.  Significant figures

    b.  SI units

    c.  Factor-label method of problem solving

2.  Atoms, Molecules & Ions

    a.  atomic theory & structure

    b.  isotopes

    c.  atomic weights

    d.  periodic table

    e.  chemical formulas

    f.  nomenclature of binary compounds

    g.  balancing chemical equations

3.  Types of Chemical Reactions

    a.  molecular & net-ionic equations

    b.  oxidation-reduction reactions

    c.  acid-base neutralization reactions

    d.  precipation reactions

    e.  gas formation reactions

    f.  combustion reactions

4.  Calculations with Chemical Fromulas and Equations

    a.  molecular and formula weights

    b.  the mole concept

    c.  mass percentages

    d.  empirical and molecular formulas

    e.  stoichiometric calculations

    f.  limiting reactants & theoretical and percentage yields

    g.  molarity & dilution

    h.  solution stoichiometry

    i.  equivalents & normality

5.  The Gaseous State of Matter

    a.  Measurement of gas pressure

    b.  Fundamental gas laws (Boyle, Charles, Gay-Lussac, Avogadro)

    c.  Ideal gas law

    d.  Stoichiometry involving gases

    e.  kinetic-molecular theory & gaseous diffusion and effusion

    f.  real gases

6.  Thermochemistry

    a.  energy units and measuring heats of chemical reactions

    b.  enthalpy and enthalpy change

    c.  thermochemical equations and stoichiometry

    d.  Hess's Law & standard enthalpies of formation

7.  Quantum Theory of the Atom

    a.  wave nature of light & particles

    b.  quantum effects (photolectric effect, black body radiation &

        atomic spectra)

    c.  Bohr's theory of the hydrogen atom

    d.  quantum mechanics & atomic orbitals

8.  Electron Configuration and Periodicity

    a.  electron spin & Pauli exclusion principle

    b.  Hund's rule & Aufbau principle

    c.  predicting electronic configurations

9.  Ionic and Covalent Bonding

    a.  electronic configurations of ions

    b.  ionic bond formation

    c.  ionic radii

    d.  covalent bond formation

    e.  electronegativity & polar covalent bonds

    f.  Lewis electron-dot formulas and formal charge

    g.  resonance structures & delocalized bonding

    h.  bond length, bond order & bond energy

10.  Molecular Geometry and Chemical Bonding Theories

    a.  delocalized bonding and MO theory

    b.  valence-shell electron-pair repulsion model

    c.  molecular geometry & dipole moments

    d.  valence bond theory

    e.  multiple bond descriptions

    f.  molecular orbital theory

    g.  homonuclear and heteronuclear diatomic molecules

11.  Solid and Liquid States of Matter

    a.  x-ray diffraction & crystal structure

    a.  phase transitions and equilibria

    b.  phase diagrams

    c.  intermolecular forces & properties of liquids

    d.  classification of solids (molecular, network, metallic, ionic)

    e.  crystal lattices & unit cells

    f.  calculations involving unit cell dimensions

12.  Solutions

    a.  colligative properties

    b.  types of solutions

    c.  effects of temperature & pressure on solubility

    d.  mass percentage, mole fraction and molality

    e.  colloids

LABORATORY MATERIAL

1.  Laboratory safety, techniques and maintaining data notebooks

2.  Limiting reactant

3.  Empirical formula of a compound

4.  Nomenclature

5.  Qualitative analysis

6.  Graphing experimental data and statistical analysis

7.  Atomic emission spectroscopy

8.  Writing laboratory reports

9.  Intro to MS Word, MS Excel & CricketGraph for the Macintosh computer

10.  Gravimetric analysis

11.  Titrimetric analysis (acid-base and oxidation-reduction)

12.  Calorimetry

13.  Gas laws

14.  Molecular geometry

Assignments:

1.  Specific reading and study assignments from the lecture textbook

   (averaging 25-30 pages per week)

2.  Completion of recommended end-of-chapter problems (averaging 15-20

   per week).

3.  Writing an average of one laboratory report per week, previewing

   the upcoming laboratory experiment, and completing the required

   pre-laboratory assignment.

Methods of Evaluation/Basis of Grade.

Writing: Assessment tools that demonstrate writing skill and/or require students to select, organize and explain ideas in writing.		Writing 10 - 30%
Written homework, Lab reports, Essay exams
Problem solving: Assessment tools, other than exams, that demonstrate competence in computational or non-computational problem solving skills.		Problem Solving 40 - 70%
Homework problems, Lab reports, Exams
Skill Demonstrations: All skill-based and physical demonstrations used for assessment purposes including skill performance exams.		Skill Demonstrations 5 - 20%
Class performances, LAB SKILL TECH/ACCUR LAB RSLTS
Exams: All forms of formal testing, other than skill performance exams.		Exams 15 - 25%
Multiple choice, Completion, PROB SOLVING & SHORT ESSAY
Other: Includes any assessment tools that do not logically fit into the above categories.		Other Category 0 - 5%
ATTENDANCE, ASSIGNMENTS SUBMITTED ON TIME, IMPROVEMENT DEMONSTRATED ON FINAL EXAM.

Representative Textbooks and Materials:

LECTURE MANUALS

GENERAL CHEMISTRY by Darrell Ebbing, Houghton Mifflin, 1996.

CHEMISTRY by Steven Zumdahl, D.C. Heath, 1997.

CHEMISTRY:  SCIENCE OF CHANGE by Oxtoby, Nachtrieb & Freeman, Saunders,

1994.

CHEMISTRY by Chang, McGraw-Hill, 1998.

LABORATORY MANUALS

CHEMISTRY IN THE LABORATORY by Jo Beran, Wiley, 1993.

EXPERIMENTS IN GENERAL CHEMISTRY by R. Wentworth, Houghton Mifflin, 1993.

EXPERIMENTS IN THE LABORATORY by Roberts, Hollenberg, and Postma,

Freeman, 1997.

EXPERIMENTAL CHEMISTRY by James F. Hall, D.C. Heath, 1993.

SPECIAL STUDENT MATERIALS

 Safety goggles

 Laboratory apron

 Scientific calculator

 Laboratory data notebook

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