Download PREP Chemistry 2008 Final Exam Review Problems

TWHS Chemistry 2012 SPRING Final Exam Review Problems
CHEMICAL REACTIONS
1.
Write and balance the following chemical equations, showing the states, and identify the type; S,
D, C, SR, DR, redox. Can be more than one.
a. Solid calcium reacts with solid sulfur to produce solid calcium sulfide.
b. Solid aluminum metal reacts with aqueous zinc chloride to produce solid zinc metal and
aqueous aluminum chloride.
c. Aqueous aluminum sulfate reacts with aqueous barium hydroxide to produce aluminum
hydroxide and barium sulfate. Use solubility rules to determine states of products.
d. Propane (C3H8) reacts with oxygen gas to form carbon dioxide gas and water vapor.
e. Solid silver oxide decomposes to produce solid silver and oxygen gas
2.
Predict the products of the following reactions, balance the final reactions, classify the reactions
using as many types as needed. If no reaction occurs write NR. Be able to read Activity Series and
determine solubility for the products of double replacement reactions.
a. F2(g) + NaI (aq)
b. Liquid pentane burns (C5H12).
c. Aluminun Sulfate + Sodium Phosphate
3.
If 45.98 g of sodium combines with an excess of chlorine gas to form 116.89 g sodium chloride,
what mass of chlorine gas is used in the reaction?
STOICHIOMETRY
4.
When an antacid tablet dissolves in water, the fizz is due to a reaction between sodium hydrogen
carbonate (NaHCO3) and citric acid (H3C6H5O7).
NaHCO3 + H3C6H5O7 → CO2 + Na3C6H5O7 + H2O
How many moles of carbon dioxide can be produced if one tablet containing 0.0119 mol NaHCO 3
is dissolved?
5.
How many grams of copper metal will be needed to react with 100.0 g of silver nitrate.
Cu + 2 AgNO3 → Cu(NO3)2 + Ag
6.
To prevent corrosion and make paints adhere better, some aluminum products are treated with
chromium (III) phosphate before finishing. Chromium (III) phosphate (CrPO4) is commercially
produced by treating chromium metal with phosphoric acid (H3PO4).
a. Balance the following equation for the reaction.
Cr(s) +
H3PO4(aq) 
H2(g) +
CrPO4(s)
b. How many moles of chromium metal are needed to produce 855 g of chromium (III)
phosphate?
c. The reaction of 206 g chromium will release how many moles of hydrogen gas?
7.
Sand (silicon dioxide) and coke (carbon) are combined to form silicon carbide (SiC), a compound
used in high-strength ceramic materials.
a. Balance the following equation for the reaction.
SiO2(s) +
C(s) 
SiC(s) +
CO(g)
b. What mass of silicon carbide will be produced from the reaction of 352 g silicon dioxide?
c. If 1.00 g of carbon is reacted, what mass of carbon monoxide is released?
8.
Two compounds of nitrogen, nitrogen tetroxide (N2O4) and hydrazine (N2H4), are used as rocket
fuels. When the two compounds are mixed, they ignite spontaneously and produce nitrogen gas
and water.
a. Balance the following equation for the reaction.
N2O4 +
N2H4

N2
+
H2O
b. If 8.00 g Nitrogen tetroxide and 4.00 g hydrazine are mixed, determine the following
quantities.
1.
mass of product (N2)
2.
limiting reactant
9.
One step in the industrial refining of nickel is the decomposition of nickel carbonyl (Ni(CO)4) into
nickel and carbon monoxide. In a laboratory reaction, 25.0 g nickel carbonyl yielded 5.34 g nickel.
a. Balance the following equation for the reaction.
Ni(CO)4(g) 
Ni(s)
+
CO(g)
b. Determine the theoretical yield of nickel.
c. Determine the percent yield.
SOLUTIONS, ACIDS & BASES
10. List the factors that affect solubility and indicate how.
11. What is the molarity of a solution that contains 20.45 g of sodium chloride (NaCl) dissolved in
700.0 mL of solution?
12. Calculate the molarity of 0.205 L of a solution that contains 156.5 g of sucrose (C12H22O11)?
13. A 0.600 L sample of a 2.50M solution of potassium iodide (KI) contains what mass of KI?
14. What mass of ammonium chloride (NH4Cl) would you use to prepare 85.0 mL of a 1.20 M
solution NH4Cl?
15. How would you correctly prepare 125 mL of a 0.30 M solution of copper(II) sulfate (CuSO4) from
a 2.00 M solution of CuSO4?
16. A 22.0 mL sample of 12M H2SO4 is diluted to a volume of 1200.0 mL. What is the molarity of the
diluted solution?
17. How many liters of a 0.340 M HCl solution are required if you need 0.25 moles of HCl?
18. Write a balanced chemical equation for each of the following reactions that involve acids and
bases. Identify the types of reactions for each (list all that apply: combustion, SR, DR, synthesis,
decomposition, redox, neutralization, precipitation reaction.)
a. Aluminum and hydrochloric acid
b. potassium hydroxide and sulfuric acid
19. A solution has a [OH-] of 3.6 x 10-7 M.
a. What is its [H+]?
b. What is its pH?
c. What is its pOH?
d. Is it acidic or basic?
20. A solution has a pH of 5.79.
a. What is its [H+]?
b. What is its pH?
c. What is its pOH?
d. Is it acidic or basic?
ELECTRONS AND PERIODIC TRENDS
21. Orange light has a frequency of 4.8 x1014s-1. What is the energy of one quantum of orange light?
22. A radio station emits radiation at a wavelength of 2.90 m. What is the station's frequency in Hz?
23. List the sequence in which the following orbitals fill up: 1s, 2s, 3s, 4s, 5s, 6s, 2p, 3p, 4p, 5p, 3d,
4d.
24. Which element has the ground-state [Kr] 5s24d105p4?
25. Identify the following elements:
a. 1s22s22p63s23p64s23d4
b. 1s22s22p63s23p64s23d104p65s24d105p4
26. What is the ground-state electron configuration of each of the following atoms?
a. Selenium
b. krypton
c. chlorine
27. Write the electron configuration for an element that contains 14 electrons.
28. Rank the following atoms in order of decreasing atomic radii.
a. Al, Na, P, S
b. Al, Ga, In
c. As, Ge, Ga
d. Br, Ca, Cl, K
29. In the following pairs, circle the smallest particle.
a. I, Ib.
K, K+
c. Al, Al3+
30. Rank the following atoms in order of decreasing electronegativity.
a. Na, Li, K
b. K, Sc, Ca
c. As, Sn, S
31. Rank the following atoms in order of increasing ionization energy.
a. Na, Li, K
b. K, Sc, Ca
c. As, Sn, S
COVALENT BONDS AND METALLIC BONDS
32. Draw the Lewis structures for the following substances, name the substances and identify the
shape and polarity of the molecules:
a. BH3
b. NH3
c. H2O
d. CO2
e. CCl4
33.
34.
35.
36.
Identify if BOND is Polar, Non-Polar, or Ionic using Electronegativity Chart
a. H-F
b. H-H
c. C-O
d. C-Cl
Draw the Lewis Dot structure of the following elements:
a. Cl
b. Fe
c. N
d. Al
Write the prefixes used for the following numbers:
a. 1
b. 3
c. 5
d. 2
e. 10
f. 8
g. 4
h. 6
i. 7
j. 9
Name the following molecular/covalent compounds
a. CO2
b. CCl4
c. N2O5
d. H2S
37.
How would you determine if a formula or name is covalent? Ionic?
38.
Write the formulas for the following molecular/covalent compounds:
a.Tetraphosphorus decoxide
b.dinitrogen tetrafluoride
c.nitrogen monoxide
39.
Complete the following table concerning covalent, ionic and metallic bonds.
Covalent
Ionic
Metallic
Electrons are
Made up of
Give an example
List 2 properties
NUCLEAR
40. Write a complete nuclear equation for each of the following:
a. The decay of
by beta emission.
b. The decay of
by alpha emission.
c. The decay of
by gamma emission.
41. Provide the missing term in each of the following equations:
a.
+ →
+ ________
b.
→
+ ________
c.
+
→
+ 3
+ ________
d.
+ _______ →
+
42. How many milligrams of 15.0 mg sample of radium 226 remain after 6396 years? The half-life of
radium-226 is 1599 years.
43. The half-life of plutonium-239 is 24,110 years. Of an original mass of 100. g, how much
plutonium-239 remains after 96, 440 years?
44. The half-life of isotope X is 2.0 years. How many years would it take for a 4.0 mg sample of X to
decay and have only 0.50 mg of it remain?
ENERGY
45. Draw a phase diagram of water.... label melting, freezing, sublimation, critical point, triple point,
vaporization, deposition, condensation. Label the x and y axes, pressure or temperature. Show an
arrow indicating exothermic and endothermic.
46. Is the thermochemical equation below exothermic or endothermic? How do you know?
N2 (g) + O2 (g) → 2NO(g) ΔH = 180.6 kJ/mol
47. Determine whether the following are endothermic or exothermic processes:
a. 4 Fe(s) + 3O2(g)  2Fe2O3(s) + 1625 kJ
b. NH4NO3(s) + 27 kJ  NH4+(aq) + NO3-(aq)
48. How is heat transferred within liquids?
49.
How is heat transferred between solids?
50.
Heat transferred through waves is called ___________?
51.
When a hot metal is placed in cold water the heat _____by the metal equals the heat _______ by
the water.
52.
Energy of position is called ___________. Show where this energy changes on a heating curve?
53.
Energy of motion is called ____________. Show where this energy changes on a heating curve?
54.
A fruit and oatmeal bar contains 142 nutritional Calories. Convert the energy to calories. 1 cal =
4.184 J 1000cal=1kcal = 1 Cal
A 95.0 g sample of a metal at 98.0º C is dropped into a container of 249 g of water at 15.0ºC, The
final temperature is 21.3 ºC. What is the specific heat of the metal?
55.
56.
A 4.50 g of pure silver was heated from 25.0°C to 500.0°C. How much heat was absorbed by the
gold? The specific heat of gold is 0.235J/g °C.
57.
Given the following reaction diagram, calculate the ΔH and
indicate whether the reaction is endothermic or exothermic.
GASES
58. Convert 555 mmHg to kPa.
59. What is the total pressure in a canister that contains oxygen gas at a partial pressure of 804 mmHg,
nitrogen at a partial pressure of 220 mmHg and hydrogen at a partial pressure of 445 mmHg?
60. The pressure of air in a 2.25 L container is 1.20 atm. What is the new pressure if the sample is
transferred to a 65.0 L container? Temperature is constant.
61. A gas at 55.0ºC occupies a volume of 3.60 L. What volume will it occupy at 30.0ºC? Pressure is
constant.
62. The volume of a gas is 0.668 L at 66.8ºC. At what Celsius temperature will the gas have a volume
of 0.942 L, assuming the pressure remains constant?
63. Air in a tightly sealed bottle has a pressure of 0.978 atm at 25.0 ºC. What will the pressure be if
the temperature is raised to 46.0ºC?
64. Calculate the number of moles of gas that occupy a 3.45 L container at a pressure of 150 kPa and a
temperature of 45.6 ºC.
65. A sample of neon gas at a pressure of 1.08 atm fills a flask with a volume of 250 mL at a
temperature of 24.0 ºC. If the gas is released into a 10.0 L reaction vessel at a temperature of 33.6
ºC, what will the pressure be inside the reaction vessel?
66. Calculate the volume of chlorine gas at STP that is required to completely react with 3.50 g of
silver, using the following equation:
2Ag(s) + Cl2(g)  2AgCl(s)
67. How many moles of air are in a 6.06 L tire at STP?
68. Determine the volume of methane (CH4) gas needed to react completely with 0.660 L of oxygen
gas to form methanol (CH3OH).
69. If volume stays the same, and pressure increases, temperature increases, decreases or stays the
same.
70. If volume increases and pressure decreases, temperature increases decreases or stays the same?