PURDUE UNIVERSITY
CHM II STUDY GUIDE
(Chemistry 116/124)
This study guide describes the topics to be mastered prior to attempting the examination for
the second semester of chemistry at Purdue University. The material covered can be found in
current textbooks used for CHM 115/116, CHM 123/124 at Purdue University.
Suggested Textbooks
!
General Chemistry, W. R. Robinson, J. D. Odom and H. F. Holtzclaw, Jr., 10th Ed.;
Houghton Mifflin Company, ISBN 0-669-41861-7.
!
Chemistry: The Molecular Nature of Matter and Change, 2nd ed., Silberberg; McGraw-Hill
Publishers, ISBN 0-697-39597-9.
!
Chemistry: The Central Science, T.L. Brown, H.E. LeMay, Jr and B.E. Bursten, 7th or 8th
edition; Prentice-Hall Publishers.
Texts can be purchased at Local Book Stores:
!
University Book Store, 360 W. State Street, W. Lafayette, IN 47906.
!
University Book Store, 720 Northwestern, W. Lafayette, IN 47906.
!
Follett's Purdue West Book Store, 1400 W. State Street, W. Lafayette, IN 47906.
IMPORTANT!
1.
Read this material thoroughly if you contemplate trying for advanced
placement (and extra credit which counts toward graduation).
2.
Study all the material listed in the outline.
3.
Work many practice problems.
4.
When you feel prepared for it, take the sample examination.
5.
Come to the actual examination rested and confident.
2
Test Topics/Preparation for Exam:
The subject matter of General Chemistry II is varied. It usually deals with the following
topics: solutions, reduction-oxidation reactions, acids and bases, electrochemistry, chemical kinetics,
chemical equilibrium (e.g., gas-phase, acid-base, solubility and complex ion equilibria), chemical
thermodynamics and nuclear chemistry. Many of these topics are related. For example, many of
the problems associated with electrochemistry and thermodynamics depend on the concepts
associated with solution equilibria. Other relationships will become apparent as you study the
concepts presented in this outline.
You should study the topics listed in the attached outline prior to attempting the simple
examination included with this study guide. These topics are categorized to be consistent with the
content studied during the past several years in the various Chemistry II courses offered at Purdue.
In preparing for the examination, it is important to work many problems. The problems
should be used to measure your understanding of the concepts and principles involved. In the
section dealing with electrochemistry, for example, you must understand the relationship that exists
between the positions of various reduction-oxidation half reactions in an activity series and the
reactions occurring at the anode and cathode of an electrochemical cell. These relationships will
allow you to predict which electrode in the cell will be the cathode, what the reaction would be at
the cathode, and the direction of electron flow in the cell.
At the end of this study guide you will find a sample examination over this material. Allow
yourself 90 minutes to complete this sample exam. Naturally, it does not cover every topic, or
every aspect of a topic. No examination extending over a reasonable time period could do that.
However, if you have no difficulty with the sample examination, you should not have difficulty with
the actual one. Following the sample exam are the answers to the questions. A score of 65% would
be considered passing.
A Word of Advice
No one does well on an examination when he/she is excessively fatigued; therefore, you are
urged to provide yourself with an adequate rest period before taking the actual examination. If your
trip to campus necessitates travel into the late hours of the night or an extremely early departure
from your home, you should consider allowing yourself one night of rest in the Lafayette area before
you take the examination. Many students who are unsuccessful with the examination tell us that
failing to take the above precautions contributed strongly to their inability to complete the exam
successfully. Most of these students find that their first year in college was somewhat less rewarding
than it might have been because of the time spent retracing material studied in high school. Please
consult your advanced credit schedule for the actual time and place of the examination. It is usually
given both morning and afternoon.
Listed below are a set of major topics that was compiled from recent Chemistry II courses
offered at Purdue. Study the list carefully before attempting the sample examination. To use this
study guide effectively, you should preface each bulleted item with the words, “I should be able
to...”.
3
Solutions
!
describe a “solution”, a "solute" and a "solvent".
!
describe how intermolecular forces between solute-solute particles, solvent-solvent particles
and solute-solvent particles are responsible for determining whether or not a solute will
dissolve in a solvent form a solution.
!
describe a “hydrogen bond”.
!
describe why the solubility of molecules containing O-H, N-H or F-H groups in water
decreases as the size of the non-polar part of the molecule increases.
!
describe why the ions produced from the dissolution of a solid salt (such as NaCl) in water
do not recombine in the solution.
!
describe the "like dissolves like" rule and describe the limitations of this rule.
!
calculate the concentration of a solution in units of molarity, molality, percent by mass and
mole fraction given the concentration of the solution expressed in one of these units.
!
describe why the concentration of a solution expressed in units of molality, percent by mass
or mole fraction, but not molarity, can be used at a temperature different from the
temperature at which the solution was prepared.
!
describe the effects of temperature and pressure on solutions of gases in liquids.
!
list four colligative properties of solutions.
!
describe the factor(s) on which colligative properties depend.
!
calculate the total vapor pressure above a solution containing one or more non-volatile
solutes given the vapor pressure of the pure solvent and solute concentration information.
!
calculate the partial pressures and total vapor pressure above a solution containing one or
more volatile solutes given the vapor pressures of the pure solvent and solutes and solute
concentration information.
!
calculate the mole fraction of solvent and one or more solutes in the vapor above a solution
given the partial pressures and total vapor pressure.
!
calculate the boiling point elevation, molal boiling point elevation constant or molality of
a solution using: )Tbp = kb•m.
!
calculate the freezing point depression, molal freezing point depression constant or molality
of a solution using: )Tfp = -kf•m.
!
calculate the concentration of a substance in a solution given the absorbance of the solution.
4
Chemical Thermodynamics
!
state "standard state" conditions and describe why they are needed.
!
state the First Law of Thermodynamics and describe it in terms of changes in internal energy
of the universe, system and surroundings.
!
describe the First Law of Thermodynamics in terms of the change in internal energy of the
system, and the amount of work done by or on the system and the heat gained or lost by the
system.
!
describe the "universe" as a sum of the "system" and the "surroundings".
!
state the sign conventions for changes in internal energy of the system when work is done
by or on the system and/or heat is gained or lost by the system.
!
describe the two kinds of "work" normally associated with chemical reactions.
!
calculate the amount of expansion work done given a change in volume at constant pressure.
!
describe a "state function" and list several examples of state functions.
!
state the definition of "spontaneous" in thermodynamic terms.
!
describe entropy in terms of randomness/disorder, probability and number of positions or
arrangements available a chemical system.
!
predict whether or not the entropy change in a chemical reaction/process is positive or
negative based on the physical states of matter present, and/or the number of moles of
reactants and products involved.
!
describe why the entropy of gases is greater than liquids and solids, and why the entropy of
liquids is greater than solids.
!
state the Second Law of Thermodynamics and describe it in terms of changes in entropy of
the universe, system and surroundings.
!
use the Second Law of Thermodynamics predict whether or not a chemical reaction/process
is spontaneous.
!
state the Third Law of Thermodynamics and describe its significance.
!
calculate the standard entropy change and the standard enthalpy change for a chemical
reaction or process given standard entropy values, S°, and standard heats of formation, )Hf°,
for the reactants and products.
!
identify and describe the two thermodynamic driving forces for chemical
reactions/processes.
!
describe the change in free energy for a chemical reaction/process in terms of changes in
5
enthalpy and entropy.
!
describe the relationship between the change in free energy and the maximum amount of
work that can be done by the system.
!
describe why a chemical reaction/process is spontaneous only if )G is negative.
!
describe why a chemical reaction/process is at equilibrium if )G is equal zero.
!
describe what a positive value for )G or )G° means.
!
describe the difference between )G and )G°.
!
predict whether or not a chemical reaction/process is spontaneous given the temperature and
the enthalpy and entropy changes.
!
!
calculate standard free energy changes using:
!
)G° = )H° – T)S°
!
the summation of two reactions
!
standard free energies of formation
use the equation, )G = )G° + RTlnQ calculate free energy changes under non-standard state
conditions.
!
use the equation, )G° = – RTlnK calculate equilibrium constants or standard free energy
changes.
!
use the equations, )G° = – nFE° and )G = – nFE, calculate cell potentials and free energy
changes.
Chemical Equilibrium
!
describe the difference between chemical reactions that go completion and chemical
reactions that come equilibrium.
!
write an equilibrium constant expression for a reversible reaction.
!
calculate the value of the equilibrium constant for a reversible chemical reaction given the
equilibrium concentrations of all reactants and products.
!
describe why the concentrations of the reactants and products do not change when a
chemical reaction has reached equilibrium.
!
describe chemical equilibrium in terms of the rates of the forward and reverse reactions.
!
calculate the value of the reaction quotient, Q, given the concentrations of reactants and
products at any moment in time.
6
!
describe how the value of the reaction quotient at any moment in time can be used to:
!
determine whether or not a chemical reaction is at equilibrium.
!
determine in which direction the chemical reaction must proceed reach
equilibrium.
!
describe LeChatelier's principle.
!
describe how the position of a chemical system at equilibrium will change when the
following are changed:
!
!
concentration(s)
!
partial pressure(s) of gases
!
total pressure (gas-phase reactions only)
!
temperature
calculate the equilibrium concentrations of all chemical species for gas-phase reactions given
values for either Kc or Kp.
!
calculate the value for the equilibrium constant for the overall reaction when two (or more)
reactions are summed, given the values of the equilibrium constants for the individual
reactions.
!
calculate the value for the equilibrium constant for the overall reaction when two (or more)
reactions are subtracted, given the values of the equilibrium constants for the individual
reactions.
Acids, Bases and Acid-Base Equilibria
!
describe the Arrhenius definition of acids and bases.
!
describe the Bronsted definition of acids and bases.
!
identify whether or not a chemical reaction is an acid-base reaction.
!
identify the chemical species that function as Bronsted acids and/or Bronsted bases given
an equation for an acid-base reaction.
!
identify the conjugate acid and the conjugate base given an equation for an acid-base
reaction.
!
identify the conjugate base of a Bronsted acid given the molecular formula for the Bronsted
acid.
!
identify the conjugate acid of a Bronsted base given the molecular formula for the Bronsted
base.
write the equilibrium reaction for the dissociation of pure water produce H3O+ and OH– ions.
!
7
!
list the molar concentrations of H3O+ and OH– in pure water at 25°C.
!
describe why the equilibrium constant expression for the dissociation of pure water does not
contain a molar concentration term for H2O (i.e., [H2O]).
!
list the value of the water dissociation equilibrium constant, Kw, at 25°C.
!
describe pH.
!
describe a "p function".
!
calculate the pH, pOH, [H3O+]total and/or [OH – ]total for a solution given one of these values.
!
describe how the relative strengths of several given acids (or bases) can be evaluated using
the values of Ka (or Kb).
!
describe how the degree of dissociation of an acid (or base) is related the value of Ka (or Kb).
!
write the acid-dissociation equilibrium constant, Ka, for a weak acid dissolved in water.
!
write the base-dissociation equilibrium constant, Kb, for a weak base dissolved in water.
!
list five strong, monoprotic acids.
!
describe the sources of both the H3O+ ion and the OH – ion in an aqueous solution containing
a strong acid.
!
calculate the pH, pOH, [H3O+ ]tot, [OH – ]tot, [H3O+ ]water, and [OH – ]water in a solution
containing a strong acid given the initial concentration of the acid.
!
calculate the pH, pOH and equilibrium concentrations of all chemical species for the
following types of solutions given values for Ka and/or Kb:
!
!
Weak, monoprotic acid in water
!
Salt of a weak, monoprotic acid in water
!
Weak, monoprotic acid + salt of the weak, monoprotic acid in water
!
Weak, polyprotic acid in water
describe how the relative strengths of the conjugate bases of several given acids can be
evaluated using the values of Ka for the acids.
!
describe how the relative strengths of the conjugate acids of several given bases can be
evaluated using the values of Kb for the bases.
!
describe a "buffer" solution.
!
describe how an acidic buffer is prepared and how a basic buffer is prepared.
!
describe "buffer capacity".
!
determine the concentration and pKa (or pKb) for a weak acid (or base) given a pH titration
curve for a:
8
!
Weak acid titrated with a strong base
!
Weak base titrated with a strong acid
Solubility Equilibria
!
describe how the conductivity of a solution containing a soluble salt is different from a
solution containing a slightly soluble salt.
!
describe how the conductivity of a solution containing a slightly soluble salt changes as more
and more of the slightly soluble salt is added.
!
write a solubility product equilibrium constant expression for any given slightly soluble salt
dissolved in water.
!
identify which salts are soluble given the molecular formula of several different salts.
!
calculate the molar solubility of a slightly soluble salt given the value of Ksp.
!
calculate the Ksp of a slightly soluble salt given the molar solubility.
!
list one reason why the solubility of a slightly soluble salt, calculated using the value of Ksp,
may not agree with the experimentally measured solubility of the salt.
!
describe the relationship between the magnitude of Ksp and the solubility of a slightly
soluble salt.
!
calculate the solubility of a slightly soluble salt in a solution containing a common ion.
!
describe how the reaction quotient, Qsp, can be used determine the extent of saturation of a
solution containing a slightly soluble salt.
!
describe how solubility equilibria can be used separate two (or more) metal cations in a
solution.
!
calculate the concentration of a substance necessary initiate precipitation of a slightly soluble
salt.
!
calculate the concentration of a substance necessary precipitate a slightly soluble salt from
a solution containing several precipitable substances.
9
Complex Ion Equilibria
!
write a complex ion formation equilibrium constant expression for any given complex ion.
!
describe how the formation of a complex ion can sometimes be used dissolve a slightly
soluble salt.
!
describe how the addition of either acids or bases solutions containing slightly soluble salts
can sometimes be used dissolve a slightly soluble salt.
!
calculate the number of moles of a complexing agent necessary dissolve a given amount of
a slightly soluble salt.
Coordination Chemistry
!
describe why it is important know the structure of a coordination compound.
!
describe "paramagnetism" and "diamagnetism".
!
describe a Lewis acid and a Lewis base.
!
describe why transition metals act as Lewis acids.
!
describe a coordinate covalent bond.
!
describe the difference between a monodentate ligand and a bidentate ligand.
!
identify the coordination number of a coordination compound given either the molecular
formula or a two dimensional picture of the molecule.
!
list the three most common coordination numbers and the four most common shapes for
coordination compounds.
!
describe why water and halide ions are not bidentate ligands.
!
identify the complex ion (if present) from the molecular formula of the coordination
compound.
!
assign the oxidation number of the transition metal in a coordination compound given either
the molecular formula or a two dimensional picture of the molecule.
!
identify the donor atom(s) in a coordination compound given either the molecular formula
or a two dimensional picture of the molecule.
!
describe the difference(s) between structural isomers and stereoisomers.
!
describe coordination isomers.
10
!
identify whether or not two (or more) coordination compounds are coordination isomers
given either the molecular formula or two dimensional pictures of the molecules.
!
describe linkage isomers.
!
identify whether or not two (or more) coordination compounds are linkage isomers given
either the molecular formula or two dimensional pictures of the molecules.
!
describe geometric isomers.
!
identify whether or not two (or more) coordination compounds are geometric isomers given
two dimensional pictures of the molecules.
!
describe optical isomers.
!
describe the difference(s) in physical properties and chemical properties of optical isomers.
!
determine whether or not a coordination compound can form geometric isomers given the
molecular formula for the compound.
!
write the electron configurations for all of the first-row transition metals and their ions.
!
draw a crystal field splitting diagram for a tetrahedral molecule or ion and identify the d
orbitals.
!
draw a crystal field splitting diagram for an octahedral molecule or ion and identify the d
orbitals.
!
describe the terms “high spin”, “low spin”, “strong field” and “weak field” with respect
crystal field theory.
!
determine the number of unpaired electrons in either a tetrahedral or octahedral complex ion
or coordination compound.
!
interconvert wavelength, frequency and energy for electromagnetic radiation.
Reduction/Oxidation (REDOX) Reactions
!
assign oxidation numbers atoms in elements, compounds and ions.
!
identify whether or not a chemical reaction is a redox reaction.
!
identify species that are "oxidized" and/or "reduced" and identify those species that are
"oxidizing agents" and/or "reducing agents" in a redox reaction.
!
balance reduction-oxidation reactions in acidic, basic or neutral solutions.
11
Electrochemistry
!
describe electrochemical processes in terms of redox reactions.
!
describe an overall redox reaction in terms of an oxidation half-reaction and a reduction halfreaction.
!
identify the two half-reactions that comprise an overall redox reaction using a Table of
Standard Reduction Potentials.
!
describe why it is necessary that the half-reactions in a voltaic cell be physically separated
in order for the cell do electrical work.
!
list and describe the components that are needed for a voltaic cell operate.
!
draw a diagram that illustrates how a voltaic cell with a positive cell potential would be
constructed given either two half-cell reactions or an overall cell reaction and a Table of
Standard Reduction Potentials.
!
identify the "anode" and "cathode" in a voltaic cell.
!
predict the direction of electron flow in a voltaic cell.
!
describe why it is necessary that a salt bridge (or porous disk) be added a voltaic cell in order
for the cell produce an electrical current.
!
describe "cell potential" as a measure of the driving force for the cell reaction.
!
state "standard state" conditions and describe why they are needed.
!
draw a diagram for the "standard hydrogen electrode" and describe its purpose.
!
predict whether or not a given redox reaction will occur using a Table of Standard Reduction
Potentials.
!
describe why multiplication of a half-reaction by a coefficient does not change the value of
the cell potential for the half-reaction.
!
calculate the standard cell potential for an overall voltaic cell reaction using a Table of
Standard Reduction Potentials.
!
use the Nernst equation calculate cell potentials under non-standard state conditions.
!
describe the difference between a cell potential and a standard cell potential.
!
describe "concentration cells" and explain how they are constructed.
!
describe galvanization, alloying and cathodic protection and explain how these methods
prevent corrosion.
!
identify the best oxidizing agent and/or reducing agent given a group of chemical species.
!
describe the difference between an "oxidation potential" and a "reduction potential".
12
!
draw a diagram of a galvanic cell that has been represented in line notation.
!
write galvanic cell reactions using line notation given the reactants and products of the
overall reaction.
!
calculate the value of the equilibrium constant or the standard cell potential given one of
these values for a galvanic cell reaction.
!
describe an electrolytic cell.
!
calculate the mass of product formed, time required or current necessary for an electrolytic
cell given one of these values.
Kinetics
!
describe the relationship between the value of the equilibrium constant and the rate of a
reversible chemical reaction.
!
describe a reaction rate in terms of a change in concentration divided by a change in time (at
constant volume).
!
describe the general form of a Differential Rate Law and describe how the rate of a chemical
reaction depends on the concentrations of species that appear in the rate law.
!
describe how rate laws are determined.
!
write a general form of the rate law for any chemical reaction.
!
describe the relationship between the order of a reactant and the stoichiometric coefficient
for the reactant in the overall balanced chemical equation.
!
describe how the order of each reactant appearing in the rate law is determined.
!
use the "Method of Instantaneous Rates" determine the rate law and the value for the rate
constant for a chemical reaction given experimental concentration versus time data.
!
describe how the rate of a chemical reaction changes as a function of time.
!
state the units for the rate of a chemical reaction.
!
determine the units for the rate constant given the rate law for the reaction.
!
write an expression that relates the rates of disappearance of reactants and the rates of
appearance of products for any chemical reaction using the overall balanced chemical
equation.
13
!
use the "Method of Initial Rates" determine the rate law and the value for the rate constant
for a chemical reaction given experimental concentration and initial rate data.
!
list one advantage and one disadvantage of using the "Method of Instantaneous Rates".
!
list one advantage and one disadvantage of using the "Method of Initial Rates".
!
determine the "overall reaction order" for a chemical reaction using the differential rate law.
!
use the "Integrated Rate Laws" determine the rate law and the value for the rate constant for
a chemical reaction given experimental concentration versus time data.
!
describe the "half-life" of a chemical reaction and calculate its value for chemical reactions
that are zero, first or second order.
!
describe the relationship between the rate of a chemical reaction and the frequency with
which reactant molecules collide.
!
describe why reactant molecules must have a certain minimum amount of kinetic energy
when they collide in order for a chemical reaction occur.
!
describe how the collision frequency, kinetic energy and orientation of colliding reactant
molecules affect the rate of a chemical reaction.
!
describe "activation energy".
!
describe an "activated complex".
!
describe how the activation energy for a chemical reaction can be experimentally
determined.
!
calculate the frequency factor, activation energy, rate constant or temperature for a chemical
reaction using the various forms of the Arrhenius Equation:
!
!
k = Ae –Ea/RT
!
ln k = –(Ea/RT) + ln A
!
ln (k1/k2) = Ea/R (1/T2 – 1/T1)
calculate the value of the activation energy for a chemical reaction given values for the rate
constant at several different temperatures.
!
use the Collision Model of Chemical Kinetics describe how changes in concentration or
temperature affect rates of chemical reactions.
!
describe how a catalyst increases the rate of a chemical reaction.
14
Nuclear Chemistry
!
determine the number of protons and neutrons in a nucleus given the atomic number (or
symbol) and the mass number.
!
describe and use the terms: isotope, nuclide and nucleon.
!
describe and use the three representations which are used symbolize nuclides.
!
describe radioactive decay as a first-order kinetic process.
!
use the first-order integrated rate law, and the first-order half-life expression, calculate
concentrations, half-lives, and rate constants for radioactive decay processes.
!
describe the six most common types of radioactive decay: (1) alpha decay, (2) beta decay,
(3) positron decay, (4) electron capture, (5) gamma decay and (6) spontaneous fission.
!
describe the differences between alpha particles, beta particles, positrons, gamma rays and
neutrons.
!
write balanced nuclear equations.
!
predict whether beta decay or positron decay is most likely given the neutron proton ratio
for an unstable nuclide.
!
describe how the neutron proton ratios of the stable nuclides change as their mass number
increases.
!
describe "binding energy".
!
calculate the binding energy per nucleon in units of MeV given the atomic mass of a nuclide.
!
describe why fission and fusion are favorable processes in terms of the binding energies of
the reactant nuclide(s) and product nuclide(s).
!
describe the source of the energy produced in fission and fusion reactions.
!
describe the function of the fuel, moderator, coolant and control rods in a fission power
plant.
!
describe the basic design of a Geiger counter and what types of radiation it detects.
!
describe how the potential for biological damage is affected by: (1) the energy of the
radiation, (2) the penetrating ability of the particles, (3) the ionizing ability of the particles,
(4) the chemical properties of the nuclide and (5) whether the radiation exposure is external
or internal the organism.
15
CHM II TEST-OUT PRACTICE EXAM
_____ 1.
Calculate the mole fraction of CCl4 (MM = 154 g/mol) in a solution prepared by dissolving
32 g of CCl4 in 75 g of C6H6 (MM = 78 g/mol).
(a)
_____ 2.
(c)
0.30
(d)
0.82
solubility increases with increasing pressure and increasing temperature.
solubility increases with increasing pressure and decreasing temperature.
solubility increases with decreasing pressure and increasing temperature.
lower than
(b)
higher than
(c)
the same as
Calculate the change in freezing point (in °C) of a solution prepared by dissolving 1.11 kg
of calcium chloride (MM = 111 g/mol) in 10.0 kg of water. For water, kf = 1.86 °C/m.
(a)
_____ 7.
0.22
When a solute is added a solvent the freezing point of the solution will be (5) the freezing
point of the pure solvent, the vapor pressure of the solution will be (6) that of the pure
solvent, and the boiling point of the solution will be (7) that of the pure solvent.
(a)
_____ 6.
(b)
The correct relationship for the solubility of a gas in a liquid is:
(a)
(b)
(c)
_____ 3-5.
0.18
1.56 °C
(b)
3.02 °C
(c)
5.58 °C
(d)
15.6 °C
Assume that the reaction quotient, Qc, for the following reaction at 25 °C is 1.0 x 10–8:
2 NO2(g)
» 2 NO(g) + O2(g)
Kc = 7.4 x 10 –16 @ 25 °C
From this we can conclude:
(a)
(b)
(c)
(d)
(e)
the reaction is at equilibrium.
without any reaction taking place, equilibrium could be reached by adding enough NO
or O2 the system.
the reaction must proceed from left right reach equilibrium.
the reaction must proceed from right left reach equilibrium.
the reaction can never reach equilibrium.
CHM II TEST-OUT PRACTICE EXAM
_____ 8.
16
Consider the following reaction,
2 SO3(g)
» 2 SO2(g) + O2(g)
Kc = 1.4 x 10 –11 @ 500K
Calculate the equilibrium concentration (in M) of SO2(g) if 0.10 mole of SO3(g) is initially
placed in a 1.0 L flask and the reaction is allowed reach equilibrium.
_____ 9.
(a)
3.7 x 10–7 M
(c)
6.6 x 10–5 M
(b)
3.3 x 10–5 M
(d)
0.10 M
Calculate the pH of a solution prepared by dissolving 2 x 10–3 moles of HCl in enough water
produce 1.0 L of solution.
(a)
_____ 10.
_____ 11.
_____ 13.
(b)
2.3
(c)
2.7
(d)
3.3
Calculate the H3O+ ion concentration (in M) in a solution with a pH = 7.80.
(a)
1.6 x 10–8 M
(c)
8.0 x 10–7 M
(b)
2.0 x 10–7 M
(d)
0.89 M
Ammonium chloride is used as an electrolyte in dry cells. Which of the following statements
about a 0.10 M solution of NH4Cl is correct?
(a)
(b)
(c)
(d)
_____ 12.
–2.7
The solution is basic.
The solution is neutral.
The solution is acidic.
The values for Ka and Kb for the species in solution must be known before a prediction
can be made.
Calculate the [H3O+] concentration (in M) in a 0.1 M aqueous solution of NH3.
[Kb = 1.8 x 10–5 ]
(a)
7.5 x 10–12 M
(c)
1.8 x 10–6 M
(b)
3.0 x 10–10 M
(d)
1.3 x 10–3 M
A pH buffer is best described as a solution containing:
(a)
(b)
(c)
(d)
a weak acid.
a strong acid.
a mixture of a weak acid and a strong acid.
a mixture of a weak acid and the salt of a weak acid.
CHM II TEST-OUT PRACTICE EXAM
_____ 14.
Calculate the pH of a solution prepared by dissolving 0.50 moles of acetic acid (HOAc,
Ka = 1.8 x 10–5 ) and 0.020 moles of sodium acetate (NaOAc) in enough water produce 1.0 L
of solution.
(a)
_____ 15.
_____ 17.
_____ 18.
_____ 19.
2.52
(b)
3.35
(c)
6.14
(d)
7.71
Calculate the pH of a 0.10 M aqueous solution of sodium acetate (MM = 82 g/mol). For acetic
acid, Ka = 1.8 x 10–5.
(a)
_____ 16.
17
5.13
(b)
8.87
(c)
9.37
(d)
10.25
Which of the following compounds has the GREATEST molar solubility?
(a)
MnS (Ksp = 5.6 x 10–16)
(c)
Sn(OH)2 (Ksp = 5.0 x 10–26)
(b)
NiS (Ksp = 3.0 x 10–21)
(d)
Zn(OH)2 (Ksp = 4.5 x 10–17)
Calculate the volume of water (in L) required dissolve 1.00 g of NiCO3 (MM = 119 g/mol).
For NiCO3, Ksp = 1.36 x 10–7.
(a)
3.69 x 10–4 L
(d)
2710 L
(b)
1.00 L
(e)
6.18 x 104 L
(c)
22.8 L
Calculate the concentration (in M) of Al3+ ion that must be present in a solution that is
2.51 x 10–9 M in OH – in order initiate precipitation of Al(OH)3. For Al(OH)3, Ksp = 1.9 x
10–33.
(a)
2.5 x 10–25 M
(d)
1.2 x 10–7 M
(b)
7.6 x 10–25 M
(e)
1.8 M
(c)
4.4 x 10–9 M
The solubility product constant for calcium fluoride, CaF2, in water is equal 4.0 x 10–11.
Calculate the molar solubility of CaF2 in water.
(a)
2.0 x 10–11 M
(c)
2.2 x 10–4 M
(b)
6.3 x 10–6 M
(d)
3.4 x 10–4 M
CHM II TEST-OUT PRACTICE EXAM
18
Questions 20-21 refer the unbalanced equation shown below.
CrO42 – + HSnO2 –
_____ 20.
2
(b)
3
(c)
4
(d)
5
(d)
6
What is the coefficient of H2O in the final balanced equation?
(a)
_____ 22.
HSnO3 – + CrO2 – (basic solution)
How many hydroxide ions are involved in the balanced half-reaction involving HSnO2– ?
(a)
_____ 21.
ÿ
1
(b)
3
(c)
4
I2 (MM = 254 g/mol) can be produced by passing an electric current through a solution of KI.
Calculate the number of minutes a current of 10.0 A would have flow in order produce 6.0 g
of I2.
2 I–
(a)
(b)
ÿ
3.8 min
4.6 min
I2 + 2 e –
(c)
(d)
7.6 min
76 min
You may need the following list of Standard Reduction Potentials for questions 23-26.
half-reaction
Mg2+ + 2e – ÿ Mg
Mn2+ + 2e – ÿ Mn
Zn2+ + 2e – ÿ Zn
Cr3+ + 3e – ÿ Cr
Ni2+ + 2e –
ÿ Ni
ÿ H2
2H+ + 2e –
Cu2+ + 2e – ÿ Cu
Ag+ + e –
ÿ Ag
+
14H + Cr2O72– + 6e – ÿ 2Cr3+ + 7H2O
8H+ + MnO4– + 5e – ÿ Mn2+ + 4H2O
_____ 23.
E°, V
!2.38
!1.03
!0.76
!0.74
!0.23
0.00
+0.34
+0.80
+1.33
+1.49
Which one of the following species is the best reducing agent?
(a)
Mn2+
(b)
MnO4 –
(c)
Mg
(d)
Mg2+
CHM II TEST-OUT PRACTICE EXAM
_____ 24.
19
Calculate the value for the equilibrium constant for the following reaction under standard
conditions.
Zn * Zn2+ 2 Cu2+ * Cu
_____ 25.
(a)
1.9 x 10–37
(c)
4.4 x 1018
(b)
1.6 x 1016
(d)
1.9 x 1037
The potential for the cell shown below is 0.31 V. What is the Mn2+ concentration?
Mn * Mn2+(?) 2 Zn2+(1.5 M) * Zn
(a)
_____ 26.
0.066 M
(c)
0.32 M
(d)
34 M
Mn Zn
(b)
Zn Mn
Addition of a strong acid a solution in which Ag+, AgCl, Ag(NH3)+, ammonia and Cl – are at
equilibrium will cause:
(a)
(b)
(c)
(d)
_____ 28.
(b)
For the cell described in question 25, electrons will flow from _____.
(a)
_____ 27.
0.029 M
more AgCl dissolve.
some AgCl precipitate from solution.
more Ag(NH3)+ form.
the concentrations of Ag+, Ag(NH3)+ and Cl – increase.
Silver ion, Ag+, reacts with thiosulfate ion, S2O3 –, in two steps form Ag(S2O3)– in the first step
and Ag(S2O3)23 – in the second step. If the stepwise formation constants are K f 1 = 6.6 x 108
and K f 2 = 4.4 x 104, what is the overall formation constant, K f , of Ag(S2O3)23 – ?
_____ 29.
(a)
6.7 x 10 –5
(c)
6.6 x 108
(b)
1.5 x 104
(d)
2.9 x 1013
What is the oxidation state of the metal ion in the coordination compound, [Ni(NH3)6]Cl2?
(a)
(b)
(c)
–1
0
+1
(d)
(e)
+2
+3
CHM II TEST-OUT PRACTICE EXAM
_____ 30.
_____ 31.
Which of the following compounds would be paramagnetic?
(a)
Sc(NH3)63+ (high spin)
(b)
Zn(OH)42– (tetrahedral)
(c)
Co(NH3)63+ (high spin)
(d)
Fe(CN)64 – (low spin)
How many total isomers (structural isomers and stereoisomers) exist for the complex ion
[Co(NH3)5Cl]2+ ?
(a)
_____ 32.
_____ 33.
1
(b)
2
(c)
3
(d)
4
Which one of the following is the correct electron configuration for the Fe3+ ion?
(a)
[Ar] 4s1 3d5
(d)
[Ar] 3d6
(b)
[Ar] 4s2 3d3
(e)
[Ar] 3d5
(c)
[Ar] 4s1 3d4
Which one of the following best explains why water is a monodentate ligand?
(a)
(b)
(c)
(d)
(e)
_____ 34.
20
The oxygen atom in a water molecule only has one lone pair of electrons that it can use
form a coordinate covalent bond a metal atom.
Each hydrogen atom in a water molecule has only one electron with which form a
coordinate covalent bond a metal atom.
The oxygen atom in a water molecule has two lone pairs of electrons, but both pairs
are used form a single coordinate covalent bond a metal atom.
The oxygen atom in a water molecule has two lone pairs of electrons, but the second
lone pair is not close enough a second coordination site form a coordinate covalent
bond.
None of these.
Which one of the following thermodynamic properties is not a state function of a system?
(a)
(b)
(c)
(d)
a transfer of heat
a change in temperature
a change in internal energy
a change in free energy
CHM II TEST-OUT PRACTICE EXAM
_____ 35.
Which of the following statements is TRUE?
(a)
(b)
(c)
(d)
(e)
_____ 36.
An exothermic process will always be spontaneous.
A process in which the entropy of the system increases will always be spontaneous.
An endothermic process can never be spontaneous.
A process in which the entropy of the surroundings increases will always be
spontaneous.
An exothermic process that is accompanied by an increase in the entropy of the
system will always be spontaneous.
Acetic acid, CH3COOH, has an enthalpy of vaporization equal 52.25 kJ and an entropy of
vaporization equal 122 J K–1 at its boiling point. Calculate the boiling point (in K) of acetic
acid.
(a)
(b)
(c)
_____ 37.
21
0.43 K
2.3 K
100 K
(d)
(e)
428 K
563 K
Use the given data at 298 K calculate )G° for the reaction:
2 Cl2(g) + SO2(g) ÷ SOCl2(g) + Cl2O(g)
Substance
)H°f (kJ/mol)
S° (J/K@mol)
(a)
(b)
(c)
(d)
_____ 38.
Cl2(g)
0
223.0
SO2(g) SOCl2(g)
Cl2O(g)
!296.8
!212.5
80.3
248.1
309.77 266.1
129.3 kJ
133.6 kJ
196.0 kJ
199.8 kJ
Which of the following results in a decrease in entropy?
(a)
O2(g), 300 K ÷ O2(g), 400 K
(b)
H2O(s), 0°C ÷ H2O(Ä), 0°C
(c)
N2(g), 25°C ÷ N2(aq), 25°C
(d)
NH3(Ä), !34.5°C ÷ NH3(g), !34.5°C
CHM II TEST-OUT PRACTICE EXAM
22
The following information applies for questions 39-40. Consider the following reaction:
aA + bB ÿ
C
It was experimentally shown that when the concentration of A is tripled and the concentration
of B is held constant, the reaction rate increases by a factor of nine. When the concentrations
of both A and B are doubled, the reaction rate increases by a factor of eight.
_____ 39.
What is the order of the reaction with respect A?
(a)
_____ 40.
0
(b)
1
(c)
2
(d)
3
2
(d)
3
What is the order of the reaction with respect B?
(a)
0
(b)
1
(c)
Use the following information answer questions 41 and 42.
A + B
experiment
1
2
3
_____ 41.
C
A, M
B, M
initial rate, M min–1
0.30
0.90
0.30
0.30
0.30
0.60
5.250
15.75
21.00
Calculate the value for the rate constant for the reaction shown above.
(a)
_____ 42.
»
58.3
(b)
194
(c)
648
(d)
2160
Calculate the initial rate (in M min–1) for the above reaction if the initial concentration of both
A and B is 0.90 M.
(a)
(b)
15.75 M min–1
31.50 M min–1
(c)
(d)
94.50 M min–1
141.75 M min–1
CHM II TEST-OUT PRACTICE EXAM
_____ 43.
The decomposition of compound AB form A2 and B2 is a first-order reaction with a rate
constant of 0.037 s–1. At 25 °C, it was found that 7.8 s were required for 25% decomposition
of AB. What is the half-life (in s) for this reaction?
(a)
_____ 44.
23
0.053 s
(b)
0.074 s
(c)
15.6 s
(d)
18.7 s
The following reaction is second order in A and second order overall with a rate constant of
0.079 L mol–1 s–1.
A ÿ B + C
How long (in s) will it take an initial concentration of A of 0.75 M decrease 0.65 M?
(a)
_____ 45.
(c)
2.1 s
(d)
2.6 s
8.8 s
(b)
13 s
(c)
26 s
(d)
77 s
In-114
Sn-115
Cd-115
Sn-116
Calculate the binding energy per nucleon (in MeV) of boron-11 if the atomic mass of boron-11
is 11.00931 amu.
(a)
(b)
(c)
(d)
_____ 48.
1.8 s
What is the final, stable nuclide, X, if indium-115 decays by the emission of a beta particle?
(a)
(b)
(c)
(d)
_____ 47.
(b)
What is the half-life (in s) of the above reaction if the initial concentration of A is 0.75 M?
(a)
_____ 46.
1.3 s
3.54 MeV
5.67 MeV
6.94 MeV
7.82 MeV
A pure sample of tritium, 3H, was prepared and stored for a number of years. Tritium
undergoes $ decay with a half-life of 12.32 years. How long has the container been sealed if
analysis of the contents shows there are 5.25 mol of 3H amd 6.35 mol of 3He?
(a)
(b)
(c)
(d)
2.34 y
3.38 y
9.77 y
14.1 y
CHM II TEST-OUT PRACTICE EXAM
_____ 49.
The coordination compound, Pt(NH3)2Cl2, has a square-planar structure and is used
medicinally as an anti-cancer drug. The square-planar shape of this molecule is determined
primarily by the number and arrangement of which of the following?
(a)
(b)
(c)
(d)
(e)
_____ 50.
24
positrons
protons
neutrons
electrons
quarks
Which of the following characteristics of particles produced by radioactive decay are
important for assessing the potential for biological damage living systems?
(a)
(b)
(c)
(d)
(e)
mass
charge
penetrating ability
kinetic energy
All of the above.
25
ANSWERS CHEM II SAMPLE TEST
1.
a
26.
a
2.
b
27.
b
3.
a
28.
d
4.
a
29.
d
5.
b
30.
c
6.
c
31.
a
7.
d
32.
e
8.
c
33.
d
9.
c
34.
a
10.
a
35.
e
11.
c
36.
d
12.
a
37.
d
13.
d
38.
c
14.
b
39.
c
15.
b
40.
b
16.
d
41.
b
17.
c
42.
d
18.
d
43.
d
19.
c
44.
d
20.
a
45.
d
21.
a
46.
b
22.
c
47.
c
23.
c
48.
d
24.
d
49.
d
25.
b
50.
e