Download Unit 9 – Behavior of Gases

2011-2012 CHEMISTRY FINAL EXAM REVIEW
Put your answers on another sheet of paper to be turned in!!
Unit 7: Periodic Table
1. Which subatomic particle plays the largest role in determining the physical and chemical properties
of an element? Explain.
2. Describe Bohr’s model of the atom.
3. Write the full electron configuration for the following:
a. potassium atom
b. sulfur atom
c. Chlorine ion
d. Calcium ion
4. Create an orbital electron filling diagram for the following atoms.
a. oxygen atom
b. silicon atom
5. a. Calculate the wavelength, in meters, of green light that has a frequency of 5.0x1014 s-1? What is
the energy of this light? (h = 6.626 x 10-34 Js)
6. What happens when an electron drops from a higher to a lower energy level?
7. Where do you find the following on the periodic table: groups, periods, Alkali metals, Alkaline Earth
metals, Halogens, Noble Gases, transition metals, metalloids?
8. Do the short-hand electron configuration for:
Arsenic
b. Barium
9. Determine the # of valence electrons for the following elements. Then, tell what the charge on their
ion would be when they form an ionic compound.
a. oxygen
b. neon
c. strontium d. potassium
e. silicon
f. aluminum
10. Define each term below. What is the general periodic trend (going across left to right; down)?
a. atomic radius (size) b. ionization energy
c. electronegativity
Unit 8: Chemical Bonding
11. Explain the difference between an ionic and covalent bond. Give an example of each.
12. Explain the difference between a polar covalent and nonpolar covalent bond. Give an example of
each.
13. Explain the difference (including properties) between an ionic compound and molecular compound.
Give an example of each.
14. Fill out the table below for the following compounds or element.
Formula
Electron
Structural
# lone pairs on
Molecular (3-D)
Polarity
Dot
Formula
central atom
Shape
(polar/
Structure
(Sketch and name)
nonpolar)
CF4
PH3
H2O
N2
AlCl3
15. What is the octet rule? What are the exceptions to the octet rule?
16. Categorize each type of bond as nonpolar covalent, polar covalent or ionic.
a. N2
b. CaCl2
c. NaF
17. Which bonding pair has a more polar bond? a. H—F vs. H—I
b. S—O vs. S—Cl
18. Which has a higher boiling point, CH4 or C2H6? Explain your choice.
Unit 9 – Behavior of Gases
19. What happens to the average kinetic energy of particles in a sample as the temperature increases?
20. Convert the following pressures given that: 1atm = 101.3kPa = 760mmHg.
21. 257mmHg = ? atm
b. 2.7atm = ? kPa
c. 93.6kPa = ? mmHg
22. Calculate the volume of a gas at a pressure of 100. kPa if its volume at 102 kPa is 1.50 x 103 mL.
23. The gas in a closed container has a pressure of 300. kPa at 30oC (303K). What will the pressure be if
the temperature is lowered to –127oC?
24. A sealed cylinder of gas contains nitrogen gas at 1.00 x 103 kPa pressure and a temperature of 20.0oC.
The cylinder is left in the sun and the temperature of the gas increases to 70.0oC. What is the new
pressure in the cylinder?
25. If 4.50 g of methane gas (CH4) is introduced to an empty 2.02 L container at 35.3oC, what is the
pressure in the container?
26. What is the volume of NO gas at STP that contains 2.2 x 1023 molecules?
27. A breathing mixture used by deep-sea divers contains helium, oxygen and carbon dioxide gases.
What is the partial pressure of oxygen if PHe = 84 kPa ,PCO2 = 0.1 kPa; the total pressure is 101.3 kPa.
28. Explain Avogadro’s Hypothesis.
29. Which of the following molecules would have the greatest average speed at STP? CO2, N2, or CH4
Unit 10 – Aqueous Solutions
30. Write how the following ionic compounds will dissociate in water. For each compound, what is the
total number of ions represented in the formula? a. CaCl2 b. CaCO3
c. Mg3(PO4)2
31. Write the net ionic equation for the reaction between calcium chloride and potassium carbonate.
32. What is the molarity of a 250mL salt solution made with 5.0 g NaCl?
33. How many grams of NaOH are required to make 0.50 L of a 2.6M solution?
34. Which of the following substances dissolve to a significant extent in water?
a.
CH4
b. KCl
c. sucrose (C12H22O11)
d. NH3 (ammonia)
35. Which of the substances listed above in #34 would be considered electrolytes?
36. How does adding a solute to a solution affect the solution’s:
a. freezing point?
b. boiling point?
c. vapor pressure?
37. How would the freezing point be different if CaCl2 were used instead of NaCl
38. How do you make a supersaturated solution?
Unit 11/12 –Acids and Bases and Equilibrium
39. List the following characteristics for an acid and a base.
a.
pH
c.
feel (touch)
e. reaction with metal
b.
litmus paper color
d.
taste
f. g. color in phenolphalein
40. Name each or write the formula.
a.
HCl
f. H3PO4
b.
KOH
g. H2SO4
c.
hydrobromic acid
h. iron(II) hydroxide
d.
nitric acid
i. lithium hydroxide
a.
HF
j. nitrous acid
41. Find the pH:
a. [H+]= 0.0010M
b. [OH-]= 0.0010M
c. pOH = 12
42. Find the pOH:
a.
[OH-]= 0.05M
b. [H+]= 1.5x10-3M
c. pH= 3
+
-6
43. If a solution has a [H ] of 4.5 x 10 M, what is the [OH ] in this solution? (Kw = 1.0 x 10-14)
44. Write the products for the folloiwng neutralization reaction.
HCl + NaOH 
45. How many milliliters of 0.45M hydrochloric acid must be added to 25.0mL of 1.00M potassium
hydroxide to make a neutral solution?
46. Explain how you used titration to determine the molarity of an unknown acid.
47. Explain in terms of collision theory how each of the following would affect the rate of a reaction:
a. Temperature
b. surface area
c. concentration
d. catalyst
48. Consider the following equation: 2 CO (g) + O2 (g) ↔ 2CO2 (g) ΔH = -580 kJ
a.
Write the equilibrium expression
b.
If the product is favored, what would be true of the value of the equilibrium constant?
c.
In which direction would the reaction shift if more CO gas was added?
d. Pressure increased?
e. Temperature increased?
49. The Ksp calcium sulfate is 9.1 x 10-6. What is the molar concentration of CaSO4 in a saturated
solution?
50. How can you determine the strength of an acid from the Ka?