Chemistry 1156 Final Exam Study Guide
Exam will have 200 points.
You will have from 8:00 am – 10:30 am on Wednesday, May 9th
to complete the exam.
The exam will be a mix of multiple-choice and show your work problems,
following the format of the quizzes and mid-term exam.
Review the quizzes, mid-term exam, homework problems,
Chang web site, class notes, and textbook.
Don’t forget your calculator !!!
Chapter 13
1. Sketch the rate of reaction versus concentration of reactant for zero and first order
reactions.
2. Determine the time required for the concentration of a reactant to change a desired
amount given the initial concentration and the rate constant for a first order reaction.
3. Use the concept of half-life to determine concentration of reactants over time.
4. Be able to identify the order of a reaction by the units of the rate constant and be able
to solve kinetic problems given the following equations:
[A]o 
1
1
ln
 kt

 kt
[A]o [A]  kt
 [A] 
[A] [A]o
1
ln(2)
[A]o
 t 1/
 t 1/
 t 1/
2
2
2
[A]o k
k
2k
5. Describe the Collision Theory of Chemical Kinetics using the terms activation energy,
activated complex (transition state), potential energy profiles, endothermic and
exothermic reactions.
6. Use Arrhenius equation to determine the activation energy of a reaction.
7. Show that the sum of elementary steps is the overall reaction for a reaction mechanism
and that intermediates appear in the reaction mechanism but not in the overall reaction.
8. Derive the rate law from a reaction mechanism.
9. Know the difference between: activated complex, intermediate and catalyst.
10. Define molecularity of unimolecular, bimolecular, and termolecular reactions.
11. Relate the importance of the rate-determining step in determination of reaction
mechanisms.
12. Describe what a catalysis does, how it effects activation energy and the difference
between homogeneous and heterogeneous catalysis.
Chapter 14
1. Describe chemical equilibrium using the terms forward and reverse reactions and
dynamic process.
2. Write the equilibrium constant in terms of the equilibrium concentration of products and
reactants and their respective stoichiometric coefficients for both homogenous and
heterogeneous equilibria.
3. Know the standard units used in defining Kp and Kc and be able to convert between Kp
and Kc.
4. Determine equilibrium constant given equilibrium concentration data.
5. Show that if a reaction can be expressed as the sum of two or more reactions, the
equilibrium constant for the overall reaction is given by the product of the equilibrium
constants of the individual reactions.
6. Relate equilibrium constant to rate constants from chemical kinetics.
7. Describe the relationship between reaction quotient and equilibrium constant and
predict the direction a reaction will proceed to reach equilibrium.
8. Use the concepts of equilibrium to determine concentration of all species in a solution.
9. Use Le Châtelier’s Principle to describe how changing concentration, volume, pressure,
or temperature will shift the reaction so that a equilibrium will be maintained.
10. Describe the effect of a catalyst has on equilibrium concentrations.
Chapter 15
1.
2.
3.
4.
5.
6.
7.
8.
Compare and contrast Arrhenius, and Brønsted acids and bases.
Describe what is meant by conjugate acid-base pairs and give several examples.
Use Kw to determine [H+] and [OH-] of solutions.
Discuss the pH scale and calculate pH and pOH given either [H+] or [OH-].
Define strong and weak acids and bases and give several examples of each.
Relate properties of conjugate acid-base pairs.
Determine Ka from experimental data.
Calculate pH, [H+], weak acid concentration, and conjugate base concentration given
Ka and the initial concentration of the weak acid using the quadratic equation or using
an appropriate approximation, as needed.
9. Calculate percent ionization for a weak acid.
10. Calculate pH, [OH-], weak base concentration, and conjugate acid concentration given
Kb and the initial concentration of the weak base using the quadratic equation or using
an appropriate approximation, as needed.
11. Show the relationship between Ka,, Kb, and Kw.
12. Calculate concentrations of all species present at equilibrium for diprotic and polyprotic
acids.
13. Relate molecular structure and the strength of acids.
14. Predict the relative strengths of oxoacids.
15. Describe salt hydrolysis and explain how some salts produce neutral solutions, some
acidic solutions and others basic solutions.
16. Calculate the pH of salt solutions and determine the percent hydrolysis.
17. Describe acid-base properties of oxides and hydroxides.
18. Be able to define Lewis acids and bases and give several examples of Lewis acid-base
reactions.
Chapter 16
1. Describe the common ion effect as a special case of Le Châtelier’s principle.
2. Use the Henderson-Hasselbalch equation to determine the pH of a solution containing
a weak acid (weak base) and its conjugate base (conjugate acid).
3. Describe what a buffer solution is and its importance in chemical and biological
systems.
4. Calculate the pH of a buffer solution using the Henderson-Hasselbalch equation.
5. Calculate the pH of a buffer solution after the addition of H+ or OHˉ.
6. Describe how to prepare a buffer of a desired pH.
7. Predict the pH profile of a strong acid-strong base titration and calculate the pH at any
stage of the titration.
8. Predict the pH profile of a strong acid-weak base (or strong base-weak acid) titration
and calculate the pH at any stage.
9. Describe common acid-base indicators and suggest the correct method of selection for
a specific titration.
10. Use the concepts of equilibrium to relate ion product, Q, with Ksp to predict if a solution
is unsaturated, saturated or supersaturated.
11. Distinguish between solubility product, molar solubility and solubility.
12. Calculate the solubility of an insoluble ion when a common ion is present.
13. Describe how changing pH can effect solubility.
14. Use the concept of equilibrium, formation constant (Kf) and complex ion formation to
predict ion concentration in solutions.
Chapter 18
1. State the three laws of thermodynamics.
a. 1st Law – Energy can be converted from one form to another but energy cannot
be created or destroyed.
b. 2nd Law –The entropy of the universe increases in a spontaneous process and
remains unchanged in an equilibrium process.
c. 3rd Law - The entropy of a perfect crystalline substance is zero at the absolute
zero of temperature.
2. Provide several examples of spontaneous processes.
3. Give examples of quantities that are state functions and are not state functions.
4. Define the standard conditions under which thermodynamic properties are measured
(25°C, 1 Molar and 1 atm).
5. Define entropy using the terms disorder or randomness.
6. Justify why entropy of one mole of steam is greater then the entropy of one mole of
water.
7. Predict the sign on the change in entropy (S) for common processes and chemical
reactions.
8. Describe what is meant by Suniverse , Ssystem, and Ssurroundings.
9. Relate Suniverse for spontaneous processes (Suniv = Ssys + Ssurr > 0) and for
processes at equilibrium(Suniv = Ssys + Ssurr = 0).
10. Use thermodynamic tables to determine Sreaction – Hess’ Law.
11. Predict if a reaction is spontaneous, spontaneous in the reverse direction or at
equilibrium from the sign on G.
12. Use thermodynamic tables to determine G of a reaction, given G°f, H°f, and S°’s
– Hess’ Law.
13. Rationalize the direction of a spontaneous reaction given H, S and T.
14. Calculate S for phase changes given Hfusion or Hvaporization and the melting and
boiling point temperatures.
15. Be able to calculate quantities using the relationship between G and G°, the gas
constant, temperature and the reaction quotient (G = G0 + RT lnQ).
16. Relate G° and K, equilibrium constant (G0 =  RT lnK).
17. Calculate K, equilibrium constant, using data from thermodynamic tables.
Chapter 19
1. Describe the concept of redox reactions using such terms as reduction, oxidation,
reducing agents, oxidizing agents and oxidation numbers.
2. Be able to calculate oxidation numbers.
3. Balance redox equations in acidic, neutral and basic solutions.
4. Describe an electrochemical cell using such terms as oxidation, reduction, galvanic
cell, electrolytic cell, anode, cathode, half-cell reactions, salt bridge, cell voltage and
emf.
5. Use standard cell diagrams to describe an electrochemical cell.
6. Use standard reduction potentials to predict the emf of a cell.
7. Define what SHE (standard hydrogen electrode) means and relate its significance to
the standard reduction potential table.
8. Predict the outcome of reactions based on standard reduction potentials and E° overall.
Happy studying !!
Optional Review Sessions: Monday, May 7th and Tuesday, May 8th
1:30 pm – 3:00 pm, John Barry 204
Good Luck !!
Potentially Useful Information:
1 atm = 760 torr = 760 mm Hg
K = °C + 273.15
No = 6.022 x 1023
.
R = 0.08206 L atm/(mole.K) = 8.314 J/(mole.K) = 0.008314 kJ/(mol.K)
 P  Hvap  1 1 
  
ln  2  
R  T1 T2 
 P1 
Kp = (RT)ΔnKc
1 year = 31,557,600 s
1 year = 365.25 days
1 day = 24 hours 1 hour = 60 min
[A]o 
ln
 kt
 [A] 
1
1

 kt
[A] [A]o
[A]o [A]  kt
ln(2)
 t 1/
2
k
1
 t 1/
2
[A]o k
[A]o
 t 1/
2
2k
k  E  1
1 

ln 2   a  
k1  R T1 T 2 
x=
G
1 min = 60 s
-b ± √ b2 – 4ac
2a
= H - TS
 K  H   1 1 
  
ln  2  
R  T1 T2 
 K1 
-14
Kw = 1.0 x 10
pH = pKa + log [base]
[acid]
G° = - RT ln K
G = G° + RT ln Q