Learning Objectives for Chemistry 118
When you finish your study of
Chapter 1 you should be able to:
1. Convert numbers written in
ordinary decimal notation to scientific notation and numbers in scientific
notation to decimal notation.
2. Multiply, divide, add and
subtract numbers written in scientific notation.
3. Use your calculator to
carry out operations on numbers written in scientific notation. Refer to the
manual that came with your calculator. Different models do this in different
ways.
4. State the relationship of
the uncertainty in a measurement to the number of significant figures in the
measurement.
5. Determine the number of
significant figures in any measured number.
6. Carry out addition,
subtraction, multiplication and division operations on measured numbers and
retain the proper number of significant figures in the answer.
7. State the base units for
mass and length in the International System .
8. Explain how derived units
are obtained from base units. State the derived units of area, volume, and density.
9. Write the names, symbols,
and values for the SI prefixes found in table 1.2 .
12. Interconvert temperatures
in the Fahrenheit, Celsius and Kelvin scales.
13. Write unit factors from
statements of equality. For example the statement
12 in = 1 ft yields the unit
factors 12 in/1 ft and 1 ft/12 in, both of which are equal to 1.
14. Use the unit factor
method to carry out conversions from one set of units to another. It is very
important that you become fluent with this method as it will be used for
solving most numerical problems in the study of chemistry.
15. Be able to carry out
calculations involving ratios of units such as density and concentration.
When you finish your study of
chapter 2 you should be able to:
1. Discuss how the laws of
conservation of mass, constant composition, and multiple proportions along with
2. Describe in a general way
the experiments which led to the discovery of the electron and the proposition
of a nuclear model for the atom.
3. Define the term, isotope.
4. Given the symbol
representing a particular isotope of an atom or ion, determine the number of
electrons, protons and neutrons in that species.
5. Given the number of
electrons, protons and neutrons in a chemical species write the symbol
representing that particular isotope.
6. Describe what is meant by
radioactivity and how the emission of alpha and beta particles can lead to the
transmutation of elements.
7. Describe the general
construction of the modern periodic table, thedifference between periods and
groups (families). Distinguish among alkali metals, alkaline earth metals,
halogens, noble gases, transition elements, metals, nonmetals, and metalloids.
8. Describe how molecules can
be represented by molecular, structural,and condensed structural formulas.
9. Distinguish between
molecules and ions. Indicate how ions are formed.
10. Distinguish between
cations and anions. Given the charge on a cation or an anion, determine the
total number of electrons present in that ion.
11. Be able to state the
charge on simple cations and anions with a noble gas configuration.
12. State the names of simple
monatomic cations and anions. Note that such anions have names that end in ide
(e.g., oxide ion).
13. Memorize the names and
formulas (including charge) of the polyatomic ions in tables 2.2 and 2.3.
14. Be able to state the
names of binary molecular compounds or given the name of such a compound, state
its formula. Learn the common names and formulas near the bottom of page 43.
15. State the names and
formulas of simple binary and oxoacids.
When you finish your study of
chapter 3 you should be able to:
1. Point out that the atomic
mass scale used today is based on the assignment of a mass of exactly 12 amu
for the isotope carbon-12.
2. Describe how the relative
atomic masses (these are also called average atomic masses) of atoms found on
the periodic table are determined by considering both isotopic masses and the
relative abundance of the isotopes found in nature.
3. State the value of
Avogadro's number to 4 significant figures and point out that it represents the
number of atoms of carbon-12 in exactly 12 g of carbon-12, i.e., exactly 1 mole
of carbon-12 or in general, the number of atoms of an element in a smaple whose
mass in grams is numerically equal to the atomic mass of the element.
4. Use relative atomic
masses, such as those found on a periodic table to calculate molecular masses
and formula masses.
5. State that the mass in
grams of one mole of atoms of any element, or molecules of any compound, or
formula units of any salt is numerically equal to the relative atomic mass of
the element, the relative molecular mass of the compound, or the relative formula
mass of the salt.
6. Use the term molar mass to
describe the mass in grams of Avagadro's number (onemole) of atoms, molecules
or other entities.
7. Carry out mole-gram
calculations such as those found on page 57. Note that the unit factor method
previously employed for conversions is used in doing these kind ofproblems.
8. Given the formula of a
compound, calculate its percentage composition.
9. Describe what is meant by
the simplest or empirical formula of a substance and, given its molecular
formula, write its empirical formula.
10. Given experimental data
for the composition of a compound, determine its empirical formula.
11. Given the molar mass and
empirical formula of a substance, determine its true molecular formula.
12. Given the products and reactants
in a chemical system, write a balanced equation representing the reaction.
13. Carry out the type of
stoichiometric calculations described on pages 64 and 65.
14. Define the term, limiting
reactant, and given the amounts (masses or volumes) of reactants in a chemical
reaction, determine which of those reactants is the limiting reactant.
15. Calculate the theoretical
yield of a reaction based on complete consumption of the limiting reagent.
16. Calculate the percentage
yield of a reaction based on its theoretical and actual yields.
When you finish your study of
chapter 4 you should be able to:
1. Define what is meant by
the molarity of a solution.
2. Be able to use the
molarity of a solution to calculate the number of moles of solute in a given volume
of solution and the volume of solution containing a given number of moles
solute.
3. Write equations describing
what happens when an ionic strong electrlyte dissolves in water. Given the
molar concentration of such a solution, calculate the concentration of all ions
present in the solution.
4. State in a general way
what sorts of ionic compounds are most likely to be water soluble.
5. Use a precipitation
diagram or solubility table to write net ionic equations for precipitation
reactions. Describe what is meant by spectator ions.
6. Carry out limiting
reactant problems involving net ionic equations.
7. Describe what is meant by
the terms acid and base.
8. Describe the difference
between strong and weak acids. Memorize the names and formulas of the strong
acids in table 4.1. Describe the differnce between strong and weak bases.
Memorize the names and formulas of the strong bases in table 4.1
9. Write net ionic equations
for acid-base neutralization reactions.
10. Carry out calculations
based on net ionic acid-base reactions (titrations) as shown on page 86.
11. Describe the terms
oxidation and reduction in terms of electron loss and gain.
12. Describe what is meant by
the terms oxidizing agent and reducing agent.
13. Calculate the average
oxidation number of an element in a chemical substance.
14. Describe the processes of
oxidation and reduction in terms of change in oxidation numbers.
15. Given a balanced redox
equation carry out stoichiometric calculations for that system.
When you finish your study of
chapter 7 you should be able to:
1. State the number of
valence electrons for an atom or monatomic ion of any substance.
2. Describe how a Lewis
structure gives information about shared and unshared (lone pair) electrons in
a substance with covalent bonds. Note that covalent bonds normally occur
between atoms of nonmetallic atoms.
3. Draw Lewis structures for
species which obey the so-called octet rule.
4. State what is meant by the
concept of resonance. Draw resonance strutures for such molecules as sulfur
dioxide or ions as the carbonate ion.
5. Calculate the formal
charge of any atom in a Lewis structure. Note that formal charge is the charge
an atom would have if valence electrons in bonds were distributed evenly.
Further note that oxidation number is the charge an atom would have if all
bonds in a molecule were ionic.
6. Draw Lewis structures for
molecules which do not obey the octet rule.
7. Distinguish between the
electron-pair geometry and molecular geometry of chemical species. For example,
the electron-pair geometry of water is tetrahedral while its molecular geometry
is bent.
8. State the electron-pair
geometry for any species with two, three or four "electron pairs"
around a central atom. Note that the electrons in a double or triple bond are
treated as if they are a single pair of electrons.
9. State the molecular
geometry for any species with two, three or four electron pairs around a
central atom.
10. Describe what is meant by
the term, electronegativity and describe in a general way periodic variance in
electronegativity. This information is found on page 154 of your text.
11. Describe the difference
between polar and nonpolar bonds. For a polar bond, state which end of the bond
is more negative and which end is more positive.
12. Describe what is meant by
molecular polarity. Use the shape of a molecule to determine whether it is
polar or nonpolar. Note that in a nonpolar molecule, the center of positive
charge is in the same location as the center of negative charge.
When you finish your study of
section 9.3 you should be able to:
1. State the properties of
molecular substances.
2. Describe how boiling (and
melting) points are related to the strength of intermolecular forces.
3. Describe what is meant by
dispersion forces and indicate their relationship to molecular structure.
4. Describe how dipole forces
operate between polar molecules.
5. Describe the conditions
required for hydrogen bonding to occur.
6. Given the formula of a
molecular substance, state which type of intermolecular forces should be the
most important in determining the physical behavior of that substance.
When you finish your study of
chapter 12 you should be able to:
1. Describe what is meant by
the partial pressure of a gas.
2. Point out that for a
system at equilibrium, the forward and reverse reactions are taking place at
the same rate.
3. Write the equilibrium
constant expression for a general system involving only gases.
4. Describe how the value of
the equilibrium constant changes as the coefficients of the balnced equation
change.
5. Describe how the
equilibrium constant for a reaction that can be expressed as the sum of two or
more reactions is the product of the equilibrium constants for the individual
reactions.
6. Write the equilibrium constant
expression for a heterogeneous reaction system.
7. Given initial and
equilibrium partial pressure, calculate the value of the equilibrium onstant.
This may involve the use of a table such as is found on page 330. This type of
table is extremely useful. Be sure that you understand its construction.
8. Describe the difference
between the equilibrium constant, K, and the reaction quotient, Q. Describe how
the relative values of the two can tell you whether or not a system is in a
state of equilibrium, and if it is not, what changes the system must undergo in
order to establish a state of equilibrium.
9. Given the value of the
equilibrium constant and initial partial pressures, calculate equilibrium
partial pressures.
10. State Le Chatelier's
principle and indicate the effect of adding or removing products or reactants,
compressing or expanding the system, and changing the temperature on the
position of equilibrium.
When you finish your study of
chapter 13 you should be able to:
1. State the Bronsted-Lowry
definitions of acids and bases, and given an acid-base reaction equation, pick
out the acid, base, conjugate acid, and conjugate base.
2. Describe what is meant by
an amphoteric species.
3. Write the equation for the
self ionization of water and the ion product constant expression for water.
Memorize its value at 25 degrees.
4. Define what is meant by pH
and pOH. Describe the pH and pOH of acidic, basic, and neutral solutions.
5. Given either hydronium
(hydrogen) ion concentration, hydroxide ion concentration, pH or pOH of an
aqueous solution, calculate the other three values.
6. Given the concentration of
a strong acid or base, calculate its hydrogen and hydroxide ion concentrations
and its pH or pOH.
7. Recognize that weak acids
are either molecules containing ionizable hydrogen atoms or certain cations.
Acidic cations can be viewed as hydrogen ion donors if they are bonded to water
molecules in an aqueous solution. Recognize that the conjugate bases of weak
acids are weak bases.
8. Write the form of the acid
equilibrium constant, Ka, for a weak acid.
9. Note that as acids get
weaker, the value of their Ka decreases. Also note that as they get weaker, the
value of Kb for their conjugate base increases; that is, their conjugate bases
get stronger. This type of information is summarized in tables such as table 13.2.
10. Calculate the hydrogen
ion concentration and pH, pOH, or hydroxide concentration of an aqueous weak
monoprotic acid solution given the concentration of the acid.
11. Recognize that weak bases
are either molecules capable of accepting hydrogen ions or certain anions (the
conjugate bases of weak acids).
12. Write the form of the
base equilibrium constant, Kb, for a weak base.
13. Calculate the hydroxide
ion concentration and pH, pOH, or hydrogen ion concentration of an aqueous weak
base solution given the concentration of the base.
14. State whether a
particular cation should be a weak acid or a spectator. State whether a
particular anion should be a weak base or a spectator. (Note that few anions
may be amphoteric)
15. State whether an aqueous
solution of a given salt should be acidic, basic, or neutral.
When you finish your study of
chapter 14 you should be able to:
1. Define the properties of a
buffer.
2. Calculate the hydrogen ion
concentration in a buffer solution.
3. Choose a buffer system
that would be appropriate to maintain a particular pH.
4. Describe how the buffer
capacity of a particular buffer is related to its composition.
5. Describe how an acid-base
indicator functions. Note that its color changes when the pH of a system in
which it is present is approximately equal to the pKa of the indicator.
6. Describe the general
features of acid-base titration curves.
When you finish your study of
the kinetic theory of gases (section 5.6) and chapter 11 you should be able to:
1. State the general
assumptions of the kinetic theory.
2. Describe in a
nonmathematical way what happens to the speed and kinetic energies of molecules
as their temperature changes.
3. Describe what is meant by
the average, instantaneous, and initial rates of a chemical reaction and
indicate how one might determine them given experimental data.
4. Relate the rate of
formation for a particular product to the rates of formation of other products
and the rates of consumption of reactants in a chemical reaction.
5. Describe how experimental
data can lead to the writing of a rate law expression.
6. Describe how integrated
rate law expressions give information about the
relationship of reactant
concentrations and time.
7. Describe the relationship
of the rate constant and the half-life for a first order reaction.
8. Describe how a reaction
mechanism can be viewed as a series of elementary steps which add up to the
overall chemical process under study.
9. Determine which species
are intermediates in a reaction mechanism.
10. Determine which species
are catalysts in a reaction mechanism.
11. Describe what is meant by
the rate determining step of a mechanism and show how a particular mechanism
may or may not be consistent with a particular experimentally determined rate
law expression.
12. Describe what is meant by
a chain reaction and differentiate among initiation, propagation and
termination steps in the mechanism.
13. Describe the dependence
of the rate constant on temperature and activation energy.
14. Describe what is meant by
activation energy.
15. Describe what factors are
included in the Arrhenius coefficient, A.
16. Describe what is meant by
a transition state.
17. Determine the activation
energy and the energy change for a reaction using a transition state diagram
such as figure 11.8.
18. Draw a reaction pathway
for a catalyzed reaction on a transition state diagram.
19. Distinguish between
homogeneous and heterogeneous catalysts and give an example of each.
When you finish your study of
chapters 8 and 17 you should be able to:
1. Describe what are meant by
heat capacity, molar heat capacity andspecific heat capacity, and given
information about the temperature change of a material or system calculate
each.
2. Recognize that heat lost
by a system is equal to heat gained by its surroundings and apply this concept
to calorimetric calculations.
3. Describe what is meant by
enthalpy and note that it is a state property.
4. Distinguish between
endothermic and exothermic reactions in terms of the enthalpy changes of both
system and surroundings.
5. Given the ethalpy change
for a chemical process, calculate the enthalpy change for the reaction of any
amount of reactant or the production of any amount of product that is part of
that process. (See examples 8.4-8.7)
6. Combine two or more
chemical equations to give a desired equation and combine the enthalpy changes
for each of those reactions to obtain the desired enthalpy change--that is,
apply
Hess's Law. (See example 8.7)
7. Describe what is meant by
standard state and formation.
8. Use standard enthalpies of
formation to calculate the enthalpy change for a desired chemical reaction.
9. Describe what is meant by
bond enthalpies and use them to determine the enthalpy change for a desired
chemical reaction.
10. State the first law of
thermodynamics in terms of heat and work.
11. Determine the energy
change of a system in terms of heat lost or gained by the system and work done
on or by the system.
12. State the relationship of
energy change and enthalpy change for a system.
13. Note that exothermic
reactions are generally, but not always spontaneous, and that endothermic
reactions may be spontaneous.
14. Describe what is meant by
entropy and what its relationship to disorder is.
15. State the third law of
thermodynamics and show how standard molar entropies can be used to determine
the entropy change for a chemical reaction.
16. State the second law of
thermodynamics and indicate the relationship of entropy change to spontaneity
of chemical reactions.
17. Describe what is meant by
Gibbs free energy and its relationship to the spontaneity of chemical
reactions.
18. Given the entropy and
enthalpy changes for a system calculate the free energy change and indicate
whether the system should undergo a spontaneous change.
19. Use standard molar free
energy values to calculate the free energy change for a chemical reaction.
20. Determine for both
positive and negative values of entropy change and enthalpy change whether a
chemical reaction should be spontaneous or nonspontaneous, indicate whether
temperature changes should affect that spontaneity, and, if so, in what way.
(see section 17.5)
22. Note the relationship of
standard free energy change and the value of the equilibrium constant.