Course Objectives
Chemistry 201
At the completion of this course, the successful student will be
adequately prepared to take the subsequent course: General Chemistry II
(Chemistry 203), and do the following:
1. Describe
the scientific method.
2.
Define and explain the terms: law, hypothesis, and theory.
3.
Use
significant figures as indicators of precision of measurements and
calculated values.
4.
Use exponential notation.
5.
Do mathematical calculations involving significant figures.
6.
Differentiate between mass and weight.
7.
Convert
from the English system to the metric system (& vise versa) common
units of
length, mass, volume, and temperature.
8.
Use the metric system in calculations.
9.
Differentiate between heat and temperature.
10.
Do simple calculations of heat changes using specific heat.
11.
Define
and use the terms standard state, standard enthalpy change, molar
enthalpy of formation.
12.
Solve problems using density as the relationship between mass and
volume.
13.
Use and define (describe or explain) basic
chemical concepts with respect to
properties of matter: physical
states of matter, physical and chemical
properties of matter, physical and chemical changes, the law of
conservation
of
mass, the law of conservation of energy, the law of definite composition,
classification of elements.
14.
Distinguish
between pure substances (elements and compounds) and mixtures
(homogeneous and heterogeneous).
15.
Know the names and chemical symbols of 48 elements.
16.
Distinguish between ionic and molecular compounds.
17.
Understand
chemical formulas of common substances in terms of the number
and
kind of atoms which have been bonded.
18.
Use basic chemical nomenclature for inorganic compounds.
19.
Write the formulas of binary ionic compounds,
common binary molecular
compounds, 12 common acids, 4 common bases, inorganic ternary
compounds using 15 common polyatomic ions.
20.
Use
oxidation numbers to distinguish oxidation states of metals in
compounds.
21.
Balance chemical equations given the formulas of the reactants and
products.
22.
Calculate
the oxidation number of each element, given the formulas of the
reactants and products.
23.
Balance redox equations using oxidation numbers.
24.
List the basic principles of Dalton’s atomic
theory and indicate how the
theory
has been further developed in this century.
25.
State
the basic properties of the subatomic particles: protons, neutrons, and
electrons.
26.
Describe the Rutherford atom.
27.
Define atomic number, mass number, and isotopes.
28.
Define the atomic mass unit and Advogardo’s number.
29.
Use the conversion factor from grams to amu in simple calculations.
30.
Be
able to calculate the average atomic weight from isotopic masses and
percent abundances.
31.
Be
able to apply the terms: metals,
nonmetals, alkali metals, alkaline earth
metals, metalloids, transition metals, noble gases, halogens, and
inner
transition metals to the arrangement of elements in the periodic table.
32.
Describe t the arrangement of the elements in the periodic table.
33.
Use the periodic table to predict formulas of compounds.
34.
Define the terms anion, cation, and polyatomic ion.
35.
Describe how ionic and covalent bonds are formed.
36.
Calculate the oxidation number of each element, given the formula.
37.
Calculate the percent composition of compounds, given the formulas.
38.
Calculate the empirical formula, given the percent composition.
39.
Calculate
the empirical formula of compound given the mass of the sample,
the
mass of CO2 and mass of H2O produced in a combustion
reaction.
40.
Distinguish between empirical and molecular formulas.
41.
Understand
the concepts of the chemical quantity, the mole, and relate it to
counting of atoms and molecules.
42.
Convert
mass in grams to moles, formula units, molecules (and/or atoms)
using
atomic weights, formula weights, and molecular weights.
43. Know the
basic rules which predict whether a salt is soluble in water.
44.
Be
able to write the balanced equations describing several examples of
combustion, acid-base, precipitation, and exchange reactions. Write the
equations in the molecular, total ionic and net ionic format.
45. Explain
the information given by chemical equations.
46. Perform
stoichiometric calculations from a given chemical equation.
47.
Use
calculations to show which the limiting reagent, how much excess
reagent is left, and what is the theoretical and percentage yield of
each
product.
48.
List
the properties of solutions and distinguish true solutions from
heterogeneous and colloidal mixtures.
49.
Define
solubility, percent concentration, molarity, mole fraction, and
molality.
50. Explain
factors affecting solubility and the rate of dissolving.
51.
Write
molecular, total ionic and net ionic equations which show that the
solution is the reaction medium.
52.
Use
percent concentration, molarity, and molality in stoichiometric
calculations.
53. List the
basic principles of the Kinetic Molecular Theory of gases.
54. Describe
the measurement of pressure using a barometer.
55. Use 4
kinds of pressure units in calculations and convert from one to another.
56.
Calculate
pressure, volumes, and temperatures of gases using Boyle’s law,
Charles’ law, the Combined gas law, and Dalton’s law of partial
pressures.
57.
Calculate Kelvin temperatures from Centigrade and vise versa.
58. Define
standard conditions of temperature and pressure.
59. Use the
/Ideal gas law to calculate density and molecular weight of a gas.
60. Use the
gas laws in chemical stoichiometric calculations.
61. Define
and distinguish between diffusion and effusion.
62.
Define
and explain the terms electromagnetic radiation, wavelength, frequency, wave
amplitude, spectrum, and nodes.
63.
Describe
the Bohr hydrogen atom; describe the hydrogen atom in terms of simple quantum
mechanics.
64.
Perform
calculations using the equation lu = c.
65.
Explain
the source of the atomic line spectra.
66.
Know
and understand the properties of light.
67.
Be
able to find and use electronic configurations of the first 50 elements; show
the diagrams of their electronic structure, and indicate the spin of each
electron.
68.
Sketch
the shape of the s, p and d orbitals.
69.
Be
able to identify the 4 quantum numbers for any electron in an atom.
70.
From
the electronic configurations predict which atoms or ions are paramagnetic and
which are diamagnetic.
71.
State
the Pauli exclusion principle, Hund’s rule, and the Aufbau principle.
72.
Define
ionization energy and be able to rank using the periodic table.
73.
Use
ionization energy trends to predict the stability of electronic configurations
and the tendency for outer shell electrons to undergo changes in order to form
compounds.
74.
Define
electronegativity: show how it varies with respect to the periodic table.
75.
Use
electronegativity to estimate the polarity of bonds.
76.
Show
the trends of atomic and ionic sizes on the periodic table.
77.
State
the octet rule, including exclusions.
78.
Write
Lewis electron dot structures for simple covalent compounds and polyatomic
ions.
79.
Use
double and triple bonds to show structures of molecules and ions; use resonance
to describe equivalent bonds.
80.
Use
the Valence Shell Electron Pair Repulsion theory to describe electron pairs
geometry, molecular geometry, hybridization, and bond angles.
81.
Predict
the polarity of bonds and molecules.
82.
Define
bond order and bond dissociation energy; use bond energies to estimate reaction enthalpies.
83.
Calculate
the formal change of an atom in a
molecule or ion, and use it to predict which are the most reasonable resonance
structures.
84.
Explain
the difference between oxidation number and formal change.
85.
Explain
simple valence bond theory.
86.
Use
the concepts of orbital overlap, sigma and pi bonds, hybrid orbitals to explain
the strength and orientation of covalent bonds.
87.
Use
molarity in calculations concerning the dilution of solutions.
88.
Explain
at least two examples of colligative properties.
89.
Calculate
the freezing point depression and the boiling point elevation due to the
addition of a nonvolatile molecular solute to a pure solvent.
90.
List
at least four properties each for acids and bases.
91.
Explain
the behavior of acids and bases in terms of the Arrhenius and Bronsted/Lowry
theories.
92.
Write
equations for acids and bases showing conjugated acid/base pairs.
93.
List
at least five common strong acids and five common strong bases.
94.
Given
the conjugated acid, write the formula of the conjugated base, and vise versa.
95.
Write
complete equations for at least two examples of each of the following
reactions: acid + base, acid + metal, acid + metal oxide, acid + carbonate.
96.
Given
the formula of a salt, write the formulas of the acid and the base which would
react to form the salt.
97.
Distinguish
between electrolytes and non-electrolytes, strong and weak electrolytes. List at least three examples of each.
98.
Define
pH. Given a pH value, state whether the
solution is acidic, basic, or neutral.
99.
Given
a pH value calculate the H+ concentration, and vise
versa.
100. Given a pOH value calculate the OH-
concentration, and vise versa.
101. Be able to convert from H+ concentration
to pH then to pOH then to OH-
concentration.
102. Perform simple tasks in the laboratory. Perform 12 laboratory experiments.
Perform the
necessary calculations, prepare any required graphs and answer
the
questions for each experiment.
103. Achieve a grade of at least 50% for the final
comprehensive examination.
104. Record all data in ink directly onto the data sheet.
105. Prepare a lab report including a summary.
106. On quizzes and exams answer short essay questions.
_____________________________
Catherine
Schwab Ph.D.