CHM 1025 - Introductory Chemistry

College of Natural Sciences

Credit(s): 3
Contact Hours: 47
Effective Term Fall 2025 (655)

Requisites

Permission of the Program or
((Prerequisite ENC 0025 with a minimum grade of C and
Prerequisite REA 0017 with a minimum grade of C) and
(Prerequisite MAT 1033 with a minimum grade of C or
Pre- or Co-requisite MAC 1105 with a minimum grade of C or
Pre- or Co-requisite Any higher-level math course with a MAC prefix (excluding liberal arts and statistic courses) with a minimum grade of C) and
Pre- or Co-requisite CHM 1025L with a minimum grade of C) or
(Prerequisite EAP 1695 with a minimum grade of C and
(Prerequisite MAT 1033 with a minimum grade of C or
Pre- or Co-requisite MAC 1105 with a minimum grade of C or
Pre- or Co-requisite Any higher-level math course with a MAC prefix (excluding liberal arts and statistic courses) with a minimum grade of C) and
Pre- or Co-requisite CHM 1025L with a minimum grade of C)

Course Description

This introductory course is a presentation of modern chemistry concepts, periodicity and atomic structure, states of matter, chemical formulas and nomenclature, chemical reactions, chemical calculations, and solutions. (Note: This course will prepare students for General Chemistry I but is not designed for credit toward a major in chemistry and may not be taken for credit subsequent to receiving a grade of "C" or better in CHM 2045, CHM 2045L or CHM 2046, CHM 2046L.)

Learning Outcomes and Objectives

  1. The student will apply principles involved in measurement and problem solving by:
    1. defining: mass, weight, significant figures, heat, temperature, density, specific gravity.
    2. explaining the difference between mass and weight.
    3. listing the basic metric units of mass, length, and volume.
    4. listing the equivalents of the metric prefixes in exponential notation.
    5. stating the number of significant figures in any number.
    6. expressing the results of arithmetic operations to the proper number of significant figures.
    7. expressing any number in exponential notation.
    8. setting up factors to convert from one unit to another with dimensional analysis.
    9. converting measurements within the metric system.
    10. listing the conversion factors between non-metric and metric units for mass, length, and volume.
    11. converting between the English and metric systems.
    12. making temperature conversions between Fahrenheit, Celsius, and Kelvin.
    13. explaining the differences between heat and temperature.
    14. calculating the density, volume, or mass of a substance from a given set of data.
    15. calculating specific gravity from density and vice versa.
  2. The student will explain the nature and variety of forms of matter and energy found in the universe by:
    1. defining: matter, homogeneous, heterogeneous, phase, substance, mixture, kinetic energy, potential energy, reactant, and product.
    2. listing and distinguishing the three physical states of matter.
    3. classifying properties as physical or chemical.
    4. classifying changes as physical or chemical.
    5. stating the laws of conservation of matter and energy.
    6. calculating percent composition of compounds from masses of elements involved in a chemical reaction.
    7. defining: element, atom, compound, molecule, ion, chemical formula, chemical equation, mixture, metal, nonmetal, and metalloid.
    8. classifying common materials as compounds, elements or mixtures.
    9. writing the symbols for all the most common elements.
    10. naming the most common elements when given their symbols.
    11. stating the law of definite composition.
    12. interpreting chemical formulas in terms of number of atoms of each element present.
    13. writing formulas for compounds when given the number of atoms of each element in the compound.
    14. listing the characteristics of metals and nonmetals.
    15. naming binary compounds when given the formulas.
    16. listing the elements that occur as diatomic molecules.
  3. The student will explain the structure of atoms and will apply the periodic law to predict chemical and physical properties of the elements by:
    1. defining: nucleus, orbital, atomic number, electron shell, noble gas, isotopes, atomic mass unit, atomic weight, gram-atomic-weight, Avogadro's number, and mole.
    2. listing the major points of Dalton's Atomic Theory.
    3. listing the electrical charge and relative mass for each of the three primary subatomic particles.
    4. describing the atom as conceived by Rutherford following his alpha particle scattering experiment.
    5. describing the atom as conceived by Niels Bohr.
    6. calculating the maximum number of electrons that can exist in any given energy level.
    7. drawing an s orbital and a p orbital, and recognizing d orbitals.
    8. stating the sublevel electron structure (1s22s22p6, etc.) for any of the first 56 elements, or identifying the element when given the sublevel electron configuration.
    9. drawing basic atomic orbital diagrams while applying the concepts of electron spin, the Pauli exclusion principle, and Hund’s rule.
    10. diagramming the atomic structure showing the composition of the nucleus and the number of electrons in each principle energy level for any element.
    11. stating the electron dot structure for any element falling in an A group in the periodic table.
    12. naming the three isotopes of hydrogen and giving the number of protons, neutrons, and electrons in each.
    13. listing the number of protons, neutrons, and electrons for any element when given the atomic number and atomic weight.
    14. calculating the number of atoms, moles, or grams from appropriate data.
    15. defining: periods of elements, groups or families of elements, and transition elements.
    16. stating the periodic law.
    17. indicating the location on a periodic table of the metals, the nonmetals, the metalloids, and the noble gasses.
    18. indicating on the periodic table areas in which s, p, d, and f sublevels of electrons are being filled.
    19. describing the change in the atomic radius in moving across a period and in moving down a family on the periodic table.
    20. describing the change in outer-energy level electron structures in moving across a period and in moving down a group on the periodic table.
    21. predicting the formulas of simple binary compounds for Group A elements using the periodic table.
    22. describing the electronic configuration of transition elements.
  4. The student will explain the nature of compounds, their formation, composition and nomenclature by:
    1. defining ionization energy, valence electrons, electro negativity, chemical bond, electrovalent bond, ionic bond, covalent bond, non-polar covalent bond, polar covalent bond, polyatomic ion, oxidation number, oxidation, reduction.
    2. describing the variation of the ionization energies of the elements with respect to position in the periodic table and with respect to removal of successive electrons.
    3. describing the formation of ions by electron transfer between two elements and the nature of the ionic bond formed.
    4. predicting the formulas of the monatomic ions formed from group A elements.
    5. showing pictorially in the form of a chemical equation with electron dot structures the formation of an ionic compound from atoms.
    6. describing the relative sizes of atoms compared to their ions.
    7. drawing electron dot structures for common covalent compounds.
    8. describing the change in electro negativity in moving across a period and in moving down a family on the periodic table.
    9. predicting whether a covalent bond will be polar.
    10. predicting whether molecules will be dipoles.
    11. classifying the bonding in a compound as primarily ionic or primarily covalent.
    12. drawing the dot structures for simple polyatomic ions.
    13. stating the names or formulas of the common ions.
    14. writing formulas of compounds which are simple combinations of common ions.
    15. assigning oxidation numbers to each element in a compound or ion.
    16. stating the name or formula for inorganic binary compounds in which the metal has only one common oxidation state.
    17. stating the name or formula for inorganic binary compounds containing metals of variable oxidation state, using either the stock system or classical nomenclature.
    18. stating the name or formula for inorganic binary compounds containing two nonmetals.
    19. stating the name or formula for binary acids.
    20. stating the name or formula for ternary inorganic acids.
    21. stating the name or formula for ternary salts.
    22. stating the name or formula for salts containing more than one positive ion.
    23. stating the name or formula for inorganic bases.
    24. stating how each of the following is used in naming inorganic compounds: -ide, -ous, -ic, hypo-, per-, -ite, -ate, and Roman numerals.
    25. stating the formula for familiar substances as identified by the instructor.
    26. defining: formula weight, molecular weight, gram-formula weight, gram-molecular-weight, empirical formula, and molecular formula.
    27. determining the formula weight or molecular weight of a compound when given the formula.
    28. calculating moles, gram-formula weights, gram-molecular weights, molecules, or grams from appropriate data.
    29. calculating the percentage composition by weight of a compound when given the formula.
    30. explaining the relationship between an empirical formula and a molecular formula.
    31. calculating the empirical formula of a compound from its percentage composition.
    32. calculating the molecular formula of a compound from its percentage composition and molecular weight.
  5. The student will analyze chemical equations and use them in stoichiometric calculations by:
    1. defining: chemical equation, word equation, reactant, product, balanced equation, combination reaction, decomposition reaction, single replacement reaction, double replacement reaction, combustion reaction, exothermic reaction, and endothermic reaction.
    2. identifying and using common symbols in writing chemical equations.
    3. balancing chemical equations.
    4. interpreting a balanced equation in terms of molecules, atoms, grams, or moles of each substance used or produced.
    5. classifying reactions as combination, decomposition, single replacement, or double replacement.
    6. completing and balancing simple combination, decomposition, single replacement, and double replacement reactions.
    7. interpreting a chemical equation in terms of exothermic or endothermic heat effect.
    8. defining: stoichiometry, mole ratio, limiting reagent, excess reagent, theoretical yield, and actual yield.
    9. giving mole ratios involving any two specified substances when given chemical equations.
    10. calculating the number of moles of a substance involved in a chemical reaction from the mass of another substance used or produced in the reaction.
    11. calculating the mass of a substance involved in a chemical reaction from a given mass of another substance used or produced in the reaction.
    12. calculating the mass of a substance involved in a chemical reaction from a given mass of another substance used or produced in the reaction.
    13. deducing the limiting and excess reagents when given masses of each and a balanced chemical equation.
    14. applying theoretical yield or actual yield data in stoichiometric calculations.
  6. The student will apply the principles of gas behavior in ideal systems by:
    1. defining: pressure, ideal gas, atmospheric pressure, barometer, one atmosphere, standard conditions, and molar volume.
    2. listing the principle assumptions of kinetic molecular theory.
    3. describing how a gas exerts pressure.
    4. describing how a barometer works.
    5. expressing one atmosphere in terms of mm of Hg, inches of Hg, torr, and lbs/in2.
    6. stating Boyle's law.
    7. stating Charles law.
    8. using the combined gas laws to find the volume of a gas when both the temperature and pressure change.
    9. using the ideal gas law to find pressure, volume, number of moles, and temperature of a gas that is not at standard temperature and pressure (STP).
    10. using the molar volume of a gas in conjunction with the combined gas laws to solve for gram-molecular-mass, mass, or volume of a gas.
    11. calculating the density of an ideal gas at Standard Temperature Pressure (STP).
    12. calculating the specific gravity of a gas at STP.
    13. stating Dalton's law of partial pressures in determining the pressures of component gases in a mixture of gases.
  7. The student will describe the properties of aqueous solution systems and the theories describing the behavior of acids and bases in aqueous systems by:
    1. explaining the water molecule with respect to electron dot structure, and polarity.
    2. completing and balancing equations for neutralization.
    3. identifying hydrates as such, writing balanced equations for their decomposition reactions to water and the anhydride.
    4. defining: solution, solute, solvent, solubility, miscible, immiscible, concentration of a solution, dilute solution, concentrated solution, saturated solution, unsaturated solution, supersaturated solution, mass-percent, molarity.
    5. qualitatively predicting the effect of temperature change on the solubility of solids and gases in liquids.
    6. calculating the mass-percent concentration of a solution.
    7. calculating the mass or volume of solute, or mass or volume of solution when given the mass-percent or volume percent concentration.
    8. calculating the molarity of a solution.
    9. calculating the moles or the mass of solute, or volume of solution when given the molarity and other appropriate data.
    10. calculating the resulting molarity when a solution of known molarity is diluted with water or mixed with another solution.
    11. relating mass, moles, solution volume, or gas volume of substances in a chemical reaction when given the chemical equation.
    12. defining: salt, hydronium ion, amphoteric, electrolyte, nonelectrolyte, dissociation, ionization, strong electrolyte, weak electrolyte, pH, neutralization, titration, spectator ions.
    13. stating the Arrhenius definitions of acids and bases.
    14. classifying common compounds as electrolytes or nonelectrolytes.
    15. classifying common acids, based, and salts as strong or weak electrolytes.
    16. solving for the concentration of an unknown solution when given titration data.

Criteria Performance Standard

Upon successful completion of the course the student will, with a minimum of 70% accuracy, demonstrate mastery of each of the above stated objectives through classroom measures developed by individual course instructors.

History of Changes

Revised 7/25/83 Revised 8/84 DBT 2/86 Effective Session 19861 SCN Change 11/5/86 Effective Session 19871 DBT 5/15/90 Effective Session 19901 3 YR C&I Review 1993-94 DBT 5/17/94 Effective Session 19941 C&I 3/17/98; DBT 4/20/98 Effective Session 19981 C&I 12/1/98; DBT 12/14/98 Effective Session 19991 C&I 11/9/99; DBT 12/14/99 Effective Session 19992 Online effective session 2000 1. Effective Sess 20011 C&I 11/11/03, BOT 12/16/03, eff20032. C&I 11/8/05, BOT 12/20/05, Effective 20052(0360). Effective 20102(0435). Amended prerequisites effective 20112(0450). C&I Approval: 10/12/2010, BOT Approval: 11/16/2010, Effective Term: Spring 2012 (450). C&I Approval: , BOT Approval: , Effective Term: Fall 2018 (550). C&I Approval: 02/21/2020, BOT Approval: 03/17/2020, Effective Term: Fall 2020 (580).
C&I Approval: , BOT Approval: , Effective Term: Fall 2025 (655)

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