Download CHM – 124 Principles of Chemistry

FARMINGDALE STATE COLLEGE
DEPARTMENT OF CHEMISTRY
COURSE OUTLINE:
Prepared by: Dr. J. Ursino, Jr.
June 2014
COURSE TITLE:
Principles of Chemistry
COURSE CODE:
CHM 124
CREDITS:
4
CONTACT HOURS:
Lecture: 3
CATALOG 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.
Laboratory: 3
This course will fulfill the requirement of certain
science, health science, or pre-health programs
that have an introductory chemistry course as a
prerequisite.
PREREQUISITE:
High School Sequential (Integrated) Math 1
IMPORTANT NOTE:
BOTH THEORY AND LABORATORY PARTS OF
THIS
COURSE
MUST
BE
TAKEN
CONCURRENTLY IN ORDER TO RECEIVE
CREDIT.
GENERAL EDUCATION:
This course satisfies 4 credits of the Natural
Sciences competency area of the General
Education requirements at Farmingdale State
College.
ELECTIVE FOR:
Ornamental Horticulture, Automotive Technology,
Liberal Arts and Sciences.
2
REQUIRED TEXT:
Lecture Text: Foundations of College Chemistry
14th Edition by Hein and Arena.
Laboratory Text:
Principles of Chemistry
Laboratory Manual CHM 124, by Weiner.
OPTIONAL TEXT:
None.
REQUIRED SUPPLIES:
Calculator, lab coat and safety glasses.
3
FARMINGDALE STATE COLLEGE
DEPARTMENT OF CHEMISTRY
CHM-124
Principles of Chemistry
Lecture Schedule
Student achievement objectives are listed after each section below:
I.
Introduction
Definitions, the scientific method, mathematical review, dimensional
analysis.
Section I - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
II.
Define chemistry, and recognize its importance in other
fields of science.
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.
Understand the importance of problem solving to the
science of chemistry.
Express any number in scientific or exponential notation.
Multiply and divide exponential numbers.
Solve elementary algebraic equations by rearrangement
until the desired term is isolated.
Utilize the method of dimensional analysis to set up and
solve problems.
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.
4
Section II. - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
7.
8.
III.
Know the metric units of length, mass, and volume.
Know the numerical equivalent for the metric prefixes:
deci, centi, milli, micro, and kilo.
Convert measurements of length, mass, and volume from
American to metric units, and vice versa; utilizing the
method of dimensional analysis.
Differentiate between mass and weight. Indicate the
instruments used to measure each.
Make temperature conversions among the Fahrenheit,
Celsius, and Kelvin scales.
Calculate the density, mass, or volume of an object from
the appropriate data.
Express answers to calculations to the proper number of
significant figures.
Differentiate between the term’s accuracy and precision.
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.
Section III. - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
7.
8.
9.
List the physical properties used to characterize a
substance.
Distinguish between the physical and chemical properties
of matter.
Classify changes undergone by matter as either physical
or chemical.
Classify common materials as elements, compounds, or
mixtures.
Distinguish between pure substances and mixtures.
Distinguish between homogeneous and heterogeneous
mixtures.
Identify the three physical states of matter.
Distinguish between kinetic and potential energy.
Differentiate between endothermic and exothermic
reactions.
5
10.
11.
12.
13.
State the law of Conservation of Mass.
State the Law of Conservation of Energy.
Differentiate clearly between heat and temperature.
Make calculations using the equation:
energy
t)
IV.
=
(mass)
X
(specific heat)
X
(

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.
Section IV. - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
Give the names, symbols, and relative masses of the
three principal subatomic particles.
Describe the atom as conceived by Ernest Rutherford
after the alpha-scattering experiment.
Determine the atomic number, mass number, or number
of neutrons of an isotope when given the values of any
two of these three items.
Calculate the average atomic mass of an element, given
the isotopic masses and the abundance of its isotopes.
Determine the number of protons, neutrons, and
electrons from the atomic number and atomic mass of an
atom.
Describe the atom as conceived by Niels Bohr.
Explain what is meant by an electron orbital.
Explain how the Pauli exclusion principle, Hund’s rule,
and the “Aufbau” order are used to write electron
configurations for elements.
State the periodic law and explain how the periodic table
of the elements is based on it.
Determine the number of valence electrons in any atom in
the Group A elements.
Distinguish between representative and transition
elements.
Identify groups of elements by their special family names.
6
13.
14.
15.
16.
17.
Explain the relationship between group number and the
number of outer shell electrons for the representative
elements.
Indicate the locations of the metals, nonmetals,
metalloids, and noble gases in the periodic table.
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.
Describe how atomic radii varies (a) from left to right in a
period, and (b) from top to bottom in a group.
List the characteristics of metals, nonmetals, and
metalloids.
UNIT EXAM I
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.
Section V - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
VI.
Write the symbols when given the names, or write the
names when given the symbols of common elements.
List the elements that occur as diatomic molecules.
Understand how symbols, including subscripts and
parentheses, are used to write chemical formulas.
Explain how compounds follow the Law of Definite
Composition.
Differentiate between the structural units that make up
molecular and ionic compounds.
Recognize the chemical formulas for acids, bases, salts
and
hydrates.
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.
Section VI - At the end of this section, the student should be able to:
7
1.
2.
3.
4.
5.
6.
7.
8.
9.
VII.
Explain the meaning of the mole.
Discuss the relationship between a mole and Avogadro’s
number.
Convert grams, atoms, molecules, and molar masses to
moles, and vice versa.
Determine the molar mass of a compound from the
formula.
Calculate the percent composition of a compound from
its formula.
Calculate the percent composition of a compound from
experimental data.
Explain the relationship between an empirical formula
and a molecular formula.
Determine the empirical formula for a compound from its
percent composition.
Calculate the molecular formula of a compound from its
percent composition and molar mass.
Chemical Reactions and Equations
The chemical equation. Writing and balancing chemical equations.
Symbols used in chemical equations.
Information provided bychemicalequations. Types of chemical
equations.
Section VII - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
VIII.
Know the format used in setting up chemical equations.
Recognize the various symbols commonly used in writing
chemical equations.
Be able to balance simple chemical equations.
Interpret a balanced equation in terms of the relative
numbers or amounts of molecules, atoms, grams, or
moles of each substance represented.
Classify equations as combination, decomposition,
single-displacement, or double-displacement reactions.
Calculations from Chemical Equations
Introduction to stoichiometry: the mole-ratio method.
Mole-mole molecalculations.
Mole-mass calculations. Mass-mass calculations. Limiting reactant and
percent yield calculations.
Section VIII -
At the end of this section, the student should be able to:
8
1.
2.
3.
4.
5.
6.
7.
IX.
Write mole ratios for any two substances involved in a
chemical reaction.
Outline the mole or mole-ratio method for making
stoichiometric calculations.
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).
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 massmole calculations).
Calculate the mass of a desired substance involved in a
chemical reaction from a given mass of a starting
substance
(mass-mass calculation).
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).
Calculate the percent yield of a substance from a
chemical reaction using its theoretical yield and actual
yield.
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.
Section IX - At the end of this section, the student should be able to:
1.
2.
3.
Write Lewis structures for the representative elements
from their position in the periodic table.
Describe (a) the formation of ions by electron transfer
and (b) the nature of the chemical bond formed by
electron transfer.
Show by means of Lewis structures the formation of an
ionic compound from atoms.
9
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
Describe the covalent bond and predict whether a given
covalent bond will be polar or nonpolar.
Draw Lewis structures for the diatomic elements.
Identify single, double, and triple covalent bonds.
Describe the formation of a coordinate covalent bond.
Provide examples of compounds that are exceptions to
the octet rule.
Describe the changes in electronegativity in (a) moving
across a period and (b) moving down a group in the
periodic table.
Describe the effect of electronegativity on the type of
chemical bonds in a compound.
Draw Lewis structures for molecules of covalent
compounds.
Describe the difference between polar and nonpolar
bonds.
Distinguish clearly between ionic and molecular
substances.
Predict whether the bonding in a compound will be
primarily ionic or covalent.
Describe the VSEPR model for molecular shape.
Use the VSEPR model to determine molecular structure
from the Lewis structure of a given compound.
UNIT EXAM II
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
containingpolyatomic ions. Naming nonoxygen acids and ternary oxyacids.
Naming inorganic bases.
Section X - At the end of this section, the student should be able to:
1.
2.
3.
Write the formulas of ionic compounds formed by
combining their ions in the correct ratios.
Write the names or formulas for inorganic binary
compounds in which the metal has only one type of
cation.
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.
10
4.
5.
6.
7.
8.
XI.
Write the names or formulas for inorganic binary
compounds that contain two nonmetals.
Write the names or formulas for inorganic ternary
compounds with metals that have either one type of
cation or multiple types of cations.
Write the names or formulas for nonoxygen acids.
Write the names or formulas for ternary oxyacids.
Write the names or formulas for inorganic bases.
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.
Section XI - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
XII.
Explain the nature of the gaseous state.
Describe the characteristics of an ideal gas.
State two reasons why real gases may deviate from the
behavior predicted for an ideal gas.
Sketch and explain the operation of a mercury barometer.
List two factors that determine gas pressure in a vessel
of fixed volume.
State Boyle’s, Charles’, and Gay-Lussac’s laws. Use all
of them in problems.
State the combined gas law. Indicate when it is used.
State the ideal gas law. Solve problems involving its use.
State the molar volume concept.
Calculate the molar mass of a gas from its density at STP
and the molar volume concept.
Make mole-volume, mass-volume, and volume-volume
stoichiometric calculations from balanced chemical
equations.
State Dalton’s law. Use it to solve problems involving
mixtures of gases.
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.
11
Physical properties of solids: melting point, molar heat of fusion. Types of
crystalline solids.
Section XII - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
7.
8.
XIII.
Explain the nature of the liquid and solid states. Explain
how they are different from gases.
Explain how the three types of intermolecular forces arise
between molecules.
Explain the process of evaporation.
Relate vapor pressure and rate of evaporation.
Describe how unbalanced intermolecular forces bring
about the surface tension of liquids.
Describe the process of boiling and the relationships
among boiling point, vapor pressure, and the
surrounding atmospheric pressure.
Distinguish between crystalline and amorphous solids.
Distinguish among the following types of crystalline
solids: ionic, macro-molecular, molecular, and metallic.
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.
Section XIII - At the end of this section, the student should be able to:
1.
2.
3.
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.
Calculate the heat flow when given (a) the mass of a pure
substance, (b) its specific heat and (c) its temperature
change.
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.
12
4.
5.
6.
Calculate the heat flow when given the quantity of a pure
substance changing between the liquid and vapor states,
and the heat of vaporization.
Calculate the heat flow when given the quantity of a pure
substance changing between the solid and liquid states,
and the heat of fusion.
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)Hvap
7.
8.
and/or Hfus of the substance and (c) the average
specific heat of the substance in the solid, liquid and/or
vapor state.
Write the thermochemical equation in two forms when
given a chemical equation and the heat (enthalpy) of
reaction.
Calculate the amount of heat evolved or absorbed for a
given amount of reactant or product, when given a
thermochemical equation.
UNIT EXAM III
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.
Section XIV -
At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
Describe the types of solutions.
List and define the terms associated with solution
terminology.
Describe and illustrate the process by which an ionic
substance dissolves in water.
Indicate the effects of temperature and pressure on the
solubility of solids and gases in liquids.
Identify and explain the factors affecting the rate at which
a solid dissolves in a liquid.
Describe a simple test to determine whether a solution is
saturated, unsaturated, or supersaturated at a given
temperature.
13
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
XV.
Calculate a solute’s percent by weight in a given solution.
Calculate the amount of solute in a given quantity of a
solution when given the percent by weight of a solution.
Calculate the molarity of a solution from the volume and
the mass, or moles, of solute.
Calculate the mass of a substance necessary to prepare a
solution of specified volume and molarity.
Determine the resulting molarity in a typical dilution
problem.
Calculate the molarity of solution from the mass of the
solvent and the mass, or moles, of solute.
Use the concepts of equivalent mass and normality in
calculations.
Explain the effect of a solute on the vapor pressure of a
solvent.
Explain the effect of a solute on the boiling point and
freezing point of a solution.
Calculate the boiling and freezing points of a solution
from concentration data.
Calculate the molality and molar mass of a solute from
boiling/freezing point data.
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.
Section XV - At the end of this section, the student should be able to:
1.
2.
3.
4.
5.
6.
Describe a reversible reaction.
Explain why the rate of the forward reaction decreases
and the rate of the reverse reaction increases as a
chemical reaction approaches equilibrium.
Describe the qualitative effect of Le Chatelier’s principle.
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.
Write the equilibrium constant expression for a chemical
reaction from a balanced chemical equation.
Explain the meaning of the numerical constant, Keq, when
given the concentration of the reactants and products in
equilibrium.
14
XVI.
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.
Section XVI able to:
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
UNIT EXAM IV
At the end of this section, the student should be
State the general characteristics of acids and bases.
Define an acid and base in terms of the Arrhenius,
Bronsted-Lowry, and Lewis theories.
Identify acid-base conjugate pairs in a reaction.
When given the reactants, complete and balance
equations for the reactions of acids with bases.
Classify common compounds as electrolytes or
nonelectrolytes.
Distinguish between strong and weak electrolytes.
Explain the process of dissociation and ionization.
Indicate how they differ.
Write the equations for the dissociation and/or ionization
of acids, bases, and salts in water.
Describe and write equations for the ionization of water.
Explain how pH expresses hydrogen-ion concentration or
hydronium ion concentration.
+
Given pH as an integer, calculate the H molarity, and vice
versa.
Explain the process of acid-base neutralization.
Compare the relative strength of acids by using their
ionization constants.
Use the ion product constant for water, Kw, to calculate
+
[H ], [OH ], pH, and pOH when given any one of these
quantities.
15
FARMINGDALE STATE COLLEGE
DEPARTMENT OF CHEMISTRY
CHM – 124 Principles of Chemistry
Required Text:
Principles of Chemistry Laboratory Manual
by Weiner
Required Safety Equipment:
Safety glasses and lab coat. Other equipment may
be required by the instructor.
LABORATORY SCHEDULE
Lab
Period
Expt.
No.
Page
No.
1
2
3
4
5
6
7
8
__
2
1
3
4
5
6
9
__
11
1
25
35
49
59
93
9
10
11
7
__
__
69
Handout
Handout
12
13
14
__
8
__
Handout
79
__
Experiment Title
Check in, Safety Discussion, use of balance
Densities of Liquids and Solids.
Separations of Cations.
Calorimetry.
Hydrates, Part 1 in duplicate.
Percent 02 in Potassium Chlorate.
Empirical Formula, Option 3.
Qualitative Analysis, Test both Known and
Unknown simultaneously.
Molar Weight of a Volatile Liquid
Neutralization of Acids and Bases (HandOut).
Spectrophotometric Determination of
Concentration. (Hand-Out).
Molar Weight by Freezing Point Depression.
Chemical Equilibrium.
Check Out.