Download 2nd Semester Final Exam Review

CHEMISTRY – 2ND SEMESTER FINAL EXAM REVIEW QUESTIONS
Chapter 11- Gases
1. Determine the number of grams in .055 mol of oxygen gas.
2. When working with gas problems which conversion must always be performed?
3. Determine the partial pressure of a gas collected over 25°C water if the total pressure of the system was 765 mmHg.
You’ll have to use “Water Vapor Pressure” table on page 859 of your Chemistry book (or at end of this WS).
4. Find the volume of 2.50 mol of a gas at STP conditions.
5. Calculate the number of moles that would occupy 450.0 ml at a temperature of 35°C and a pressure of 785 mmHg.
Chapter 12 & 13 – Solutions
1. Define: solute, solvent, solution
2. Define precipitate
3. If a 35.00% solution of NaCl contained 90.0 grams of NaCl, how many grams of water was it dissolved in?
4. How many grams of KBr are needed to make 750.0 ml of a .500 M solution?
5. What is the effect on the number of dissolved particles on: vapor pressure, freezing point, and boiling point?
Colligative properties
6. Which of the following will have the higher boiling point, .100M NaNO 3 or .100 M MgBr2?
7. Write the dissociation equations for the following compounds in water: calcium chloride, sodium chloride, glucose
(C6H12O6), hydrobromic acid (HBr, but you should know this).
Once you’ve written the equations, which of these would have the lowest freezing point? (assuming that all are
1.0 M solutions).
8. Which 3 of the above solutions will conduct electricity? Why?
9. Will a precipitate form if sodium chloride is mixed with barium hydroxide? If so, what’s the ppt?
Use the ppt chart in back of lab manual or on p. 860 fo your textbook.
10. Will a precipitate form if silver nitrate is mixed with calcium iodide?
11. Define: saturated, unsaturated, supersaturated as they pertain to solutions.
12. In the term “like dissolves like”, the word “like” refers to what?
13. What is the main determinant of a substance’s boiling point or melting point?
14. How many moles of ions are produced when 1 mole of Fe(OH)3 is dissolved in water?
Chapter 14 & 15 - Acids and Bases
15. Name all of the properties of acids.
16. Name all of the properties of bases.
17. If the pH of a solution is 8, what is the pOH?
18. What’s the pH of a .000010 M HBr solution?
19. Given a .00500M solution of NaOH, calculate the pH.
20. A 35.0 ml sample of an unknown concentration of base was titrated with a .0100 M solution of HCl. The initial
buret reading of the acid was .2 ml and the final reading was 47.3 ml. Calculate the concentration of the unknown
base.
21. A 35 ml sample of .0317 M NaOH was used to titrate a .200 g sample of unknown acid. Calculate the molar mass
of the acid.
22. Name four pieces of equipment used in a titration.
23. Write the net ionic equation for reaction between any strong acid with a strong base.
24. If the pH of a solution drops from 8 to 5, then the [H+] has ____________ by a factor of _______.
25. Identify the 8 strong bases.
Chapter 16 & 17 - Energy and Kinetics
26. If the H in a chemical reaction is 25.7 kcal/mol, is the reaction endothermic or exothermic?
27. The melting of 1 mole of H2O takes 1.44 kcal of energy. Calculate the energy involved if only 3.15 grams of ice
were melted. q= H x n
28. In the problem above did the entropy increase, decrease, or not change? Explain.
29. If 45.0 g of water is heated and the temp. rose from 20.6 oC to 30.0 oC. How much heat was absorbed by the
water? Leave your answer in calories.
30. Calculate the G (Gibbs’ Free Energy) for a rxn where H is 135 kJ/mol and the S
is 5.75 kJ/mol K at 35 oC.
31. List all of the ways to speed up a chemical reaction.
32. How does a catalyst act to speed up a reaction?
33. Draw an energy diagram showing a reaction that has an E a of 50 kcal and H of -20 kcal.
34. Using the heat of formation table on p. 862 of your textbook or in the back of your lab manual, determine the
for the following reaction: CO(g) + ½ O2(g)  CO2(g) Hess’ Law
35. Given the following diagram, write the phase or phase change in each of the numbered areas:
#5
Temperature
#4
#2
#3
#1
Chapter 18 - Equilibrium
36. Write the equilibrium constant expression for the following reaction:
H2CO3(aq)  H2O(l) + CO2(g) given that H is -45 kJ
37. How could you increase the yield of the products in the above reaction? There are 5 ways.
38. The molar solubility of NiS is 3.162x10-11 M. Calculate the Ksp.
39. Calculate the molar solubility for Ca(NO3)2 if the Ksp = 2.56x10-7.
40. Calculate the Ka of a .0200 M solution of C2H5COOH if the equilibrium concentration of the H+
was .0100 M.
41. Which of the following is the strongest acid?
HCHO2 (Ka=1.8x10-4), HOC6H5 (Ka=1.6x10-10), or HClO2 (Ka=1.2x10-2)
42. How does the value of an ionic compound’s Ksp reflect the solubility of that substance?
Chapter 21 - Organic Chemistry
43. Give the chemical formulas for: pentane, pentene, pentyne, and cyclopentane.
44. Define isomer. In the problem above, which two are isomers?
45. Classify the following organic compounds:
CH3CH2-OH
CH2=CHCH3
C2H4 (there are two!)
C4H10
C6H10
You will also be supplied with the following information on your final exam:
1) These formulas from Chapter 10/11.
Ptot = P1 + P2 + …
PV=nRT
P1V1 = P2V2
T1
T2
K = 273 + °C
2) A Periodic Table
3) Chart of Solubilities in Water (for precipitate problems).
It’s in back of Lab Manual. There’s also a version on table A-12 p. 860 of your text book.
4) Vapor pressure of water table. Table A-8 p. 859 of text book.
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R = .0821 L•atm
mol •K