Course Outline: Principles of Chemistry (CHM 124)

Course Information

Credits: 3

Contact Hours:

  • Lecture: 3  
  • Lab: 3

Note on Laboratory: Both Lecture and Laboratory must be taken simultaneously; separate grades will not be given for either. Students must pass the laboratory section to receive a passing grade in the entire course.

Semesters Offered: Fall, Spring, & Summer

Prerequisites: MP 2 or MTH 015

Catalog Course Description:

A one semester survey of general chemistry. Emphasis is placed on quantitative applications of chemical concepts. Topics include: measurement, matter and energy, atomic structure, periodic table, chemical bonding, nomenclature, chemical stoichiometry, chemical equations, gases, liquids and solids, solutions, acids and bases, equilibrium and kinetics. This course will fulfill the requirement of certain science, health science, or pre-health programs that have an introductory chemistry course as a prerequisite. Note: the laboratory course CHM 124L is a part of your grade for this course. Attendance in the laboratory course is required. Approved eye-protection and a laboratory coat are required materials. A student must pass the laboratory course to receive a passing grade in the entire course.

Elective Course for: Ornamental Horticulture, Automotive Technology, Liberal Arts and Sciences.

General Education: This course satisfies 4 credits of the Natural Sciences competency area of the  General Education requirements at Farmingdale State College.

Course Texts: Foundations of College Chemistry 15th Edition by Hein and Arena.

Principles of Chemistry Laboratory Manual, CHEM 124, by Weiner; Special Edition by Cengage

Other Required Course Materials

Laboratory coat and goggles.  

Course Learning Objectives

I. Introduction 

Definitions, the scientific method, mathematical review, dimensional analysis.

At the end of this section, the student should be able to:

  1. Define chemistry, and recognize its importance in other fields of science.
  2. Understand how the scientific method is applied in chemistry, and how the steps of hypothesis, theory, and scientific law; lead time and again toward an understanding of natural phenomena.
  3. Understand the importance of problem solving to the science of chemistry.
  4. Express any number in scientific or exponential notation.  Multiply and divide exponential numbers.
  5. Solve elementary algebraic equations by rearrangement until the desired term is isolated.
  6. Utilize the method of dimensional analysis to set up and solve problems.  

II. Measurements and Calculations 

Metric units of length, mass, and volume. Dimensional analysis in measurement conversions. Mass and weight. Density. Temperature measurement. Significant figures. Rules for significant figures in addition and subtraction; multiplication and division. Accuracy and precision of measurements.

At the end of this section, the student should be able to:

  1. Know the metric units of length, mass, and volume.
  2. Know the numerical equivalent for the metric prefixes: deci, centi, milli, micro, and kilo.
  3. Convert measurements of length, mass, and volume from American to metric units, and vice versa; utilizing the method of dimensional analysis.
  4. Differentiate between mass and weight.  Indicate the instruments used to measure each.
  5. Make temperature conversions among the Fahrenheit, Celsius, and Kelvin scales.
  6. Calculate the density, mass, or volume of an object from the appropriate data.
  7. Express answers to calculations to the proper number of significant figures.
  8. Differentiate between the term’s accuracy and precision.

III. Matter and Energy 

Physical and chemical properties. Physical and chemical changes. Types of matter: atoms, elements, compounds, molecules, and mixtures. Homogeneous and heterogeneous mixtures. Pure substances and mixtures. States of matter: gases, liquids, and solids. Energy. Types of energy. Energy in chemical changes. Heat and its measurement. Conservation laws of mass and energy.

At the end of this section, the student should be able to:

  1. List the physical properties used to characterize a substance.
  2. Distinguish between the physical and chemical properties of matter.
  3. Classify changes undergone by matter as either physical or chemical.
  4. Classify common materials as elements, compounds, or mixtures.
  5. Distinguish between pure substances and mixtures.
  6. Distinguish between homogeneous and heterogeneous mixtures.
  7. Identify the three physical states of matter.
  8. Distinguish between kinetic and potential energy.
  9. Differentiate between endothermic and exothermic reactions.
  10. State the law of Conservation of Mass.
  11. State the Law of Conservation of Energy.
  12. Differentiate clearly between heat and temperature.
  13. Make calculations using the equation: energy = (mass) X (specific heat) X (Δt)

 IV. Atomic Structure and the Periodic Table 

Subatomic particles. The nuclear atom. Atomic number. Mass number, atomic weight, and isotopes. The Bohr atom. Modern concept of atomic structure. Atomic orbitals, Pauli exclusion principle, Hund’s rule, and the “Aufbau” order.  Electron configuration. The periodic law and the periodic table. Trends in the periodic table.

At the end of this section, the student should be able to:

  1. Give the names, symbols, and relative masses of the three principal subatomic particles. 
  2. Describe the atom as conceived by Ernest Rutherford after the alpha-scattering experiment.
  3. Determine the atomic number, mass number, or number of neutrons of an isotope when given the values of any two of  these three items.
  4. Calculate the average atomic mass of an element, given the isotopic masses and the abundance of its isotopes.
  5. Determine the number of protons, neutrons, and electrons from the atomic number and atomic mass of an atom.
  6. Describe the atom as conceived by Niels Bohr.
  7. Explain what is meant by an electron orbital.
  8. Explain how the Pauli exclusion principle, Hund’s rule, and the “Aufbau” order are used to write electron configurations for elements.
  9. State the periodic law and explain how the periodic table of the elements is based on it.
  10. Determine the number of valence electrons in any atom in the Group A elements.
  11. Distinguish between representative and transition elements.
  12. Identify groups of elements by their special family names.
  13. Explain the relationship between group number and the number of outer shell electrons for the representative elements.
  14. Indicate the locations of the metals, nonmetals, metalloids, and noble gases in the periodic table.
  15. Describe how the ionization energies of the elements vary with respect to (a) their position in the periodic table and (b) the removal of successive electrons.
  16. Describe how atomic radii varies (a) from left to right in a period, and (b) from top to bottom in a group.
  17. List the characteristics of metals, nonmetals, and metalloids.

V. Chemical Formulas for Elements and Compounds 

Symbols of the elements. Names of the elements. The structure of elements. Kinds of elements. Compounds. Types of compounds. Law of Definite Composition. Chemical formulas for: acids, bases, salts, and hydrates.

At the end of this section, the student should be able to:

  1. Write the symbols when given the names, or write the names when given the symbols of common elements.
  2. List the elements that occur as diatomic molecules.
  3. Understand how symbols, including subscripts and parentheses, are used to write chemical formulas.
  4. Explain how compounds follow the Law of Definite Composition.
  5. Differentiate between the structural units that make up molecular and ionic compounds.
  6. Recognize the chemical formulas for acids, bases, salts and  hydrates.

VI. Quantitative Composition of Compounds 

The mole. Molar mass of elements and compounds. Percent composition of compounds. Empirical formula versus molecular formula. Calculation of empirical formulas. Calculation of the molecular formula from the empirical formula.

At the end of this section, the student should be able to:

  1. Explain the meaning of the mole.
  2. Discuss the relationship between a mole and Avogadro’s number.
  3. Convert grams, atoms, molecules, and molar masses to moles, and vice versa.
  4. Determine the molar mass of a compound from the formula.
  5. Calculate the percent composition of a compound from its formula.
  6. Calculate the percent composition of a compound from experimental data.
  7. Explain the relationship between an empirical formula and a molecular formula.
  8. Determine the empirical formula for a compound from its percent composition.
  9. Calculate the molecular formula of a compound from its percent composition and molar mass.

VII. Chemical Reactions and Equations 

The chemical equation. Writing and balancing chemical equations. Symbols used in chemical equations. Information provided by chemical equations. Types of chemical equations.

At the end of this section, the student should be able to:

  1. Know the format used in setting up chemical equations.
  2. Recognize the various symbols commonly used in writing chemical equations.
  3. Be able to balance simple chemical equations.
  4. Interpret a balanced equation in terms of the relative numbers or amounts of molecules, atoms, grams, or moles of each substance represented.
  5. Classify equations as combination, decomposition, single-displacement, or double-displacement reactions.

VIII. Calculations from Chemical Equations 

Introduction to stoichiometry: the mole-ratio method. Mole-mole calculations. Mole-mass calculations. Mass-mass calculations. Limiting reactant and percent yield calculations.

  1. Write mole ratios for any two substances involved in a chemical reaction.
  2. Outline the mole or mole-ratio method for making stoichiometric calculations.
  3. Calculate the number of moles of a desired substance obtainable from a given number of moles of a starting substance in a chemical reaction (mole-mole calculations).
  4. Calculate the mass of a desired substance obtainable from a given number of moles of a starting substance in a chemical reaction, and vice versa (mole-mass and mass-mole calculations).  
  5. Calculate the mass of a desired substance involved in a chemical reaction from a given mass of a starting substance (mass-mass calculation).
  6. Deduce the limiting reactant or reagent when given the amounts of starting substances, and then calculate the moles or mass of the desired substance obtainable from a given chemical reaction (limiting reactant calculation).
  7. Calculate the percent yield of a substance from a chemical reaction using its theoretical yield and actual yield.

IX. Chemical Bonding

The octet theory. Lewis symbols of atoms. The ionic bond. Elements that  form ionic bonds. 

Predicting formulas of ionic compounds. The covalent bond. Elements that form covalent bonds. 

Multiple electron pair bonds. Exceptions to the octet rule. Coordinate covalent bonds. Electronegativity. Polar covalent bonds. Polar molecules and net molecular polarity. Shapes of covalent molecules.

At the end of this section, the student should be able to:

  1. Write Lewis structures for the representative elements from their position in the periodic table.
  2. Describe (a) the formation of ions by electron transfer and (b) the nature of the chemical bond formed by electron transfer.
  3. Show by means of Lewis structures the formation of an ionic compound from atoms.
  4. Describe the covalent bond and predict whether a given covalent bond will be polar or nonpolar.
  5. Draw Lewis structures for the diatomic elements.
  6. Identify single, double, and triple covalent bonds.
  7. Describe the formation of a coordinate covalent bond.
  8. Provide examples of compounds that are exceptions to the octet rule.
  9. Describe the changes in electronegativity in (a) moving across a period and (b) moving down a group in the periodic table.
  10. Describe the effect of electronegativity on the type of chemical bonds in a compound.
  11. Draw Lewis structures for molecules of covalent compounds.
  12. Describe the difference between polar and nonpolar bonds.
  13. Distinguish clearly between ionic and molecular substances.
  14. Predict whether the bonding in a compound will be primarily ionic or covalent.
  15. Describe the VSEPR model for molecular shape.
  16. Use the VSEPR model to determine molecular structure from the Lewis structure of a given compound.

X. Inorganic Nomenclature 

Formulas and names of monatomic and polyatomic ions. Writing formulas for ionic compounds. 

Naming binary ionic and binary covalent compounds.  Writing formulas from names of compounds. Naming compounds containing polyatomic ions. Naming nonoxygen acids and ternary oxyacids.  Naming inorganic bases.

At the end of this section, the student should be able to:

  1. Write the formulas of ionic compounds formed by combining their ions in the correct ratios.
  2. Write the names or formulas for inorganic binary compounds in which the metal has only one type of cation.
  3. Write the names or formulas for inorganic binary compounds that contain metals with multiple types of cations, using the Stock System and the –ous, -ic system.
  4. Write the names or formulas for inorganic binary compounds that contain two nonmetals.
  5. Write the names or formulas for inorganic ternary compounds with metals that have either one type of cation or multiple types of cations.
  6. Write the names or formulas for nonoxygen acids.
  7. Write the names or formulas for ternary oxyacids.
  8. Write the names or formulas for inorganic bases.

XI. The Gaseous State

The nature of the gaseous state. The ideal gas model. Gas measurements. Boyle’s law. Absolute temperature. Charles’ law.  Gay-Lussac’s law. The combined gas law.  The ideal gas law.  The molar volume concept and standard temperature and pressure. Gas stoichiometry. Dalton’s law of partial pressures.

At the end of this section, the student should be able to:

  1. Explain the nature of the gaseous state.
  2. Describe the characteristics of an ideal gas.
  3. State two reasons why real gases may deviate from the behavior predicted for an ideal gas.
  4. Sketch and explain the operation of a mercury barometer.
  5. List two factors that determine gas pressure in a vessel of fixed volume.
  6. State Boyle’s, Charles’, and Gay-Lussac’s laws.  Use all of them in problems.
  7. State the combined gas law.  Indicate when it is used.
  8. State the ideal gas law.  Solve problems involving its use.
  9. State the molar volume concept.
  10. Calculate the molar mass of a gas from its density at STP and the molar volume concept.
  11. Make mole-volume, mass-volume, and volume-volume stoichiometric calculations from balanced chemical equations.
  12. State Dalton’s law.  Use it to solve problems involving mixtures of gases

XII. Liquids and Solids

The nature of the liquid state. Types of intermolecular forces. Physical properties of liquids: viscosity, evaporation and vapor pressure, boiling point, molar heat of vaporization, surface tension.  The nature of the solid state.  Physical properties of solids: melting point, molar heat of fusion. Types of crystalline solids. 

At the end of this section, the student should be able to:

  1. Explain the nature of the liquid and solid states.  Explain how they are different from gases.
  2. Explain how the three types of intermolecular forces arise between molecules.
  3. Explain the process of evaporation.
  4. Relate vapor pressure and rate of evaporation.
  5. Describe how unbalanced intermolecular forces bring about the surface tension of liquids.
  6. Describe the process of boiling and the relationships among boiling point, vapor pressure, and the surrounding atmospheric pressure.
  7. Distinguish between crystalline and amorphous solids.
  8. Distinguish among the following types of crystalline solids: ionic, macromolecular, molecular, and metallic.

XIII. Energy in Physical and Chemical Changes

Energy and changes of state. Heating curves. Energy and change of temperature: specific heat. Energy and change of state: heat of fusion and heat of vaporization.  Energy and change in temperature plus change in state. Thermochemical equations. Enthalpy and change in enthalpy. Thermochemical stoichiometry. 

At the end of this section, the student should be able to:

  1. Sketch, interpret and/or identify regions in a graph of temperature versus energy for a pure substance over a temperature range from below the melting point to above the boiling point.
  2. Calculate the heat flow when given (a) the mass of a pure substance, (b) its specific heat and (c) its temperature change.
  3. Calculate the specific heat of the substance when given the amount of heat flow to or from a known mass of a substance, and its temperature change.
  4. Calculate the heat flow when given the quantity of a pure substance changing between the liquid and vapor states, and the heat of vaporization.
  5. Calculate the heat flow when given the quantity of a pure substance changing between the solid and liquid states, and the heat of fusion.
  6. Calculate the total heat flow in going from one state and temperature to another state and temperature, when given (a) the quantity of a pure substance, (b)ΔHvap and/or ΔHfus  of the substance and (c) the average specific heat of the substance in the solid, liquid and/or vapor state.
  7. Write the thermochemical equation in two forms when given a chemical equation and the heat (enthalpy) of reaction.
  8. Calculate the amount of heat evolved or absorbed for a given amount of reactant or product, when given a thermochemical equation.

XIV. Solutions 

The characteristics of a solution. The components of a solution. Types of solutions. Solution terminology. The solution process. Factors that increase the rate at which solids dissolve. Factors that determine solubility. Solution concentration: percent by weight, molarity, molality, and normality. Equivalent weights of acids and bases.  Acid-base titration.  Colligative properties of solutions.

At the end of this section, the student should be able to:

  1. Describe the types of solutions.
  2. List and define the terms associated with solution terminology.
  3. Describe and illustrate the process by which an ionic substance dissolves in water.
  4. Indicate the effects of temperature and pressure on the solubility of solids and gases in liquids.
  5. Identify and explain the factors affecting the rate at which a solid dissolves in a liquid.
  6. Describe a simple test to determine whether a solution is saturated, unsaturated, or supersaturated at a given temperature.
  7. Calculate a solute’s percent by weight in a given solution.
  8. Calculate the amount of solute in a given quantity of a solution when given the percent by weight of a solution.
  9. Calculate the molarity of a solution from the volume and the mass, or moles, of solute.
  10. Calculate the mass of a substance necessary to prepare a solution of specified volume and molarity.
  11. Determine the resulting molarity in a typical dilution problem.
  12. Calculate the molarity of solution from the mass of the solvent and the mass, or moles, of solute.
  13. Use the concepts of equivalent mass and normality in calculations.
  14. Explain the effect of a solute on the vapor pressure of a solvent.
  15. Explain the effect of a solute on the boiling point and freezing point of a solution.
  16. Calculate the boiling and freezing points of a solution from concentration data.
  17. Calculate the molality and molar mass of a solute from boiling/freezing point data.

XV. Chemical Equilibrium

Reversible reactions. Rates of reaction. Chemical equilibrium. Principle of Le Chatelier. Effect of concentration, pressure, temperature, and catalysts, on reaction rate and equilibrium. Equilibrium constants.

At the end of this section, the student should be able to:

  1. Describe a reversible reaction.
  2. Explain why the rate of the forward reaction decreases and the rate of the reverse reaction increases as a chemical reaction approaches equilibrium.
  3. Describe the qualitative effect of Le Chatelier’s principle.
  4. Predict how the rate of a chemical reaction is affected by (a) changes in the concentration of reactants, (b) changes in the pressure of gaseous reactants, (c) changes in temperature, and (d) the presence of a catalyst.
  5. Write the equilibrium constant expression for a chemical reaction from a balanced chemical equation. 
  6. Explain the meaning of the numerical constant, Keq, when given the concentration of the reactants and products in equilibrium.

XV. Acids, Bases, and Salts 

Acids and bases.  Reactions of acids.  Reaction of bases.  Neutralization. Salts.  Electrolytes and nonelectrolytes.  Dissociation and ionization of electrolytes.  Strong and weak electrolytes.  Ionization and dissociation constants.  Ionization of water.  Introduction to pH. Ion product constant for water.

At the end of this section, the student should be able to:

  1. State the general characteristics of acids and bases.
  2. Define an acid and base in terms of the Arrhenius, Bronsted-Lowry, and Lewis theories.
  3. Identify acid-base conjugate pairs in a reaction.
  4. When given the reactants, complete and balance equations for the reactions of acids with bases.
  5. Classify common compounds as electrolytes or nonelectrolytes.
  6. Distinguish between strong and weak electrolytes.
  7. Explain the process of dissociation and ionization.  Indicate how they differ.
  8. Write the equations for the dissociation and/or ionization of acids, bases, and salts in water.
  9. Describe and write equations for the ionization of water.
  10. Explain how pH expresses hydrogen-ion concentration or hydronium ion concentration.
  11. Given pH as an integer, calculate the H+molarity, and vice versa.
  12. Explain the process of acid-base neutralization.
  13. Compare the relative strength of acids by using their ionization constants.
  14. Use the ion product constant for water, Kw, to calculate [H+], [OH-], pH, and pOH when given any one of these quantities.

Laboratory Schedule

  1. Safety Discussion & Use of Balance

  2.  Densities of Liquids and Solids

  3. Separations of Cations

  4. Calorimetry

  5. Hydrates

  6. Percent O2 in Potassium Chlorate

  7. Empirical Formula

  8. Qualitative Analysis (Known and Unknown)

  9. Molar Weight of a Volatile Liquid

  10. Neutralization of Acids and Bases

  11. Spectrophotometric Determination of Concentration

  12. Molar Weight by Freezing Point Depression

  13. Chemical Equilibrium

NOTE: This course outline supersedes any course syllabus provided by a professor.

Farmingdale State College

Monday-Friday 8:30am-4:30pm