Chemistry 1152 Final Exam Study Guide
Exam will have 200 points.
You will have from 8:00 am – 10:20 am on Monday, July 30th
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 15
1. 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.
2. Show the relationship between Ka,, Kb, and Kw.
3. Calculate concentrations of all species present at equilibrium for diprotic and polyprotic
acids.
4. Relate molecular structure and the strength of acids.
5. Predict the relative strengths of oxoacids.
6. Describe salt hydrolysis and explain how some salts produce neutral solutions, some
acidic solutions and others basic solutions.
7. Calculate the pH of salt solutions and determine the percent hydrolysis.
8. Describe acid-base properties of oxides and hydroxides.
9. 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.
Happy studying !!
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  
K
R
1


 T1 T2 
-14
Kw = 1.0 x 10
pH = pKa + log [base]
[acid]
G° = - RT ln K
G = G° + RT ln Q