Spring Semester Final Exam Study Guide
Honors Chemistry Name ______
Period ______
Naming and Formula Writing
1. Write the name or formula for each of the following:
AlCl3 HClO2 Mercury (II) bromide
Cu2S (NH4)2SO4 Phosphorous acid
NaCN I4O10 Nitric acid
HI H3N Sodium nitrite
PCl3 NiN Magnesium hydroxide
CrBr3 H3PO4 Carbon trisulfide
Chemical Equations
2. State the law of conservation of matter.
2H2 + O2 → 2H2O
3. Using the chemical equation above: a. Underline the subscripts b. Box the coefficients c. Identify the diatomic molecules
4. Balance the following equations and identify the type of reaction (synthesis, decomposition, single replacement, double replacement, combustion)
a. ___N2 + ___H2 ___NH3
b. ___KClO3 ___KCl + ___O2
c. ___NaCl + ___F2 ___NaF + ___Cl2
d. ___AgNO3 + ___MgCl2 ___AgCl + ___Mg(NO3)2
e. Aluminum bromide reacts with potassium sulfate to yield potassium bromide and aluminum sulfate.
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f. Sodium and water form sodium hydroxide and hydrogen
g. Silver oxide decomposes to silver and oxygen.
h. Potassium and magnesium bromide yields potassium bromide and magnesium
5. Identify the following symbols: a. (aq)
b. (s)
c. (l)
d. (g)
e. △H
f.
g. M
6. What type of reaction (endothermic, exothermic) has more energy in the products than reactants? What is the sign for △H?
7. What type of reaction has more energy in the reactants than products? What is the sign for △H?
8. How many kJ of heat would you expect to be transferred when 6.44g of sulfur react with excess oxygen gas to produce sulfur trioxide? △H = -791.4kJ
The Mole
9. How many atoms are in a mole of Ca?
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15 10. How many O atoms are on 3.0 x 10 molecules of CO2?
11. How many atoms are in 2.6 moles of Al?
12. Convert 6689 moles of NaCl to molecules of NaCl.
13. Convert 3.2 x 1028 molecules of hydrofluoric acid to moles.
14. How many moles are in 60,000,000,000 atoms of Zn?
15. Calculate the molar mass of each of the following compounds: a. NaOH
b. MgCl2
c. Barium phosphate
16. How many grams of sulfur are in 5.23 moles of sulfur?
17. Convert 192 moles of HCl to g of HCl
18. Convert 510.2 g of aluminum hydroxide to moles.
19. How many moles are in 62 g of CO2 ?
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Stoichiometry
20. Aluminum reacts with oxygen to form aluminum oxide. a. How many moles of aluminum are needed to form 2.3moles of aluminum oxide?
b. How many moles of oxygen are required to react completely with 8.4 moles of aluminum?
21. Calcium chloride reacts with potassium sulfide to produce potassium chloride and calcium sulfide. How many grams of potassium chloride are produced when 50.0 grams of calcium chloride react with xs potassium sulfide?
22. Potassium chlorate decomposes into potassium chloride and oxygen gas. What volume of oxygen gas will be produced is 25.0 grams of potassium chlorate decompose at STP.
23. Predict the products and write a balanced equation for the following: Silver nitrate reacts with barium chloride to produce ______and ______?
Balanced equation: ______
24. Using the equation above if 5.0 g of silver nitrate react with 4.0 g of barium chloride how much silver chloride will be produced? What is the limiting reactant? How many grams of excess reagent remains?
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Empirical Formulas / % yield
25. A compound is composed of 7.20g of C, 1/20g of hydrogen and 9.60g of oxygen. The molar mass of the compound is 180g/mole. Determine the empirical and the molecular formula of the compound.
26. Define: a. Actual yield
b. Theoretical yield
c. % yield
27. In the reaction between excess K(s) and 4.28 g of O2(g), potassium oxide is formed . What mass would you expect to form (theoretical yield)? If 17.36 g of K2O is actually produced, what is the percent yield?
Gases (again)
28. Name four variables that affect gas behavior and identify the relationship. a. V T b. P T c. P V d. P n
29. Write the equation for the ideal gas law. When is the ideal gas law used? What units do the variables need to be in to use the Ideal Gas Law?
30. A gas has a pressure of 50.0mmHg at 540K. What will be the pressure at 56C, assuming volume does not change
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31. An underground cavern contains 98 moles of methane gas at a pressure of 15atm and a temperature of 22C. How many liters of methane does this natural gas deposit contain?
32. What volume does 16.0 g of O2 occupy at STP?
33. When calcium carbonate is heated strongly, carbon dioxide gas is evolved. CaCO3(s) –––> CaO(s) + CO2 (g) If 4.74 g of calcium carbonate is heated, what volume of CO2 (g) would be produced when collected at STP?
Solutions and Acids / Bases
34. In a solution the substance being dissolved is called the ______and the substance it is dissolved in is called the ______.
35. To what volume should 25mL of 15M nitric acid be diluted to prepare a 3.0M solution?
36. What is the molarity of a solution in which 10.0g of AgNO3 is dissolved in 500mL of solution?
37. What is the pH range of acids? Bases?
38. A solution with more hydroxide ions than hydrogen ions is called a ______.
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39. Litmus paper turns what color in the presence of an acid? Base?
40. Which of the following pH’s if for the weakest acid? pH 3 , pH 1, pH 6
41. What is the [H+] of a solution with a pH of 4?
42. What differentiates a weak base from a strong base? What differentiates a weak acid from a strong acid?
43. What signals the end of most lab titrations?
44. What volume of 0.175 M HCl is needed to exactly neutralize 25.0 mL of 0.150 M KOH solution?
a. Neutralization Reaction:
b. Solution:
Atomic Structure/Periodic Trends
45. What are the five postulates of Dalton’s atomic theory?
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46. Atomic number is the same as the number of ______
47. An atom is defined as the smallest part of an element that a. Has protons, neutrons and electrons b. Has protons and neutrons c. Retains the chemical identity of that element d. Can form an ion
48. The atomic number and mass of chlorine are: a. 17, 35 b. 33, 75 c. 16, 32 d. 29, 64
49. The neutrons of an atom are found ______
50. List the three subatomic particles. Name their location. Identify their charges.
51. An element loses three electrons to form and ion that has a total of 18 electrons. What is the positive charge in the nucleus? What is the element?
52. Which subatomic particle has the smallest mass?
53. An ion is formed when an atom ______
54. The number of protons always equals the number of ______in a neutral atom.
55. The mass number of an element tells you the # of ______and ______.
56. Who is credited with discovering… The nucleus ______The electron ______The Planetary Model of the Atom______
57. Isotopes have the same number of ______but different numbers of ______and different ______numbers.
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58. Vertical columns on the periodic table are called ______.
59. Horizontal rows on the periodic table are called ______.
60.Which of the following is true about groups of elements: a. Elements in a group have similar chemical properties b. Elements in a group have approximately the same atomic mass c. Elements in a group have the same number of total electrons d. Elements in a group have the same number of valence electrons e. Both a and d
61. KNOW ALL PERIODIC TRENDS for atomic radius and ionization energy
62. As you go across a period from left to right ionization energy: a. Increases b. Decreases c. Doesn’t change WHY?
63. Define: Atomic radius
Ionization energy
64. Why does atomic radius decrease as you go across a period?
65.Which groups of elements are the most reactive ______?
66. Which group of elements is the lease reactive______?
67. Which group (family) do the following elements belong to? Li ______C ______Be ______N ______V ______O ______La ______F ______B ______Ne ______
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68. What ions will the following elements form? Li Be Al Ne N 69. Classify the following as metals, semimetals, or non-metals. As Si Na B Ne Ar Cl H
ELECTRONS: 70. Draw the orbital diagram for Phosphorus.
71. Write the electron configuration for Zn
72. 1s22s22p63s1 is the electron configuration of ______.
73. What rule states that electrons enter the lowest energy orbitals until all electrons are accounted for? ______
74. How many orbitals are there in the: s group p group d group f group
75. How many electrons can each of the sublevels above hold?
VSPER Shapes and Polarity
76. Compare and contrast ionic and covalent bonds in terms of:
a. Shape
b. Melting and Boiling Point
c. Conductivity
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77. What happens to electrons in ionic bonds?______78. What happens to electrons in covalent bonds? ______
79. What type of bond do the following Lewis Dot Structures Represent?
Na S Na
O = C = O
80. Draw Lewis Dot Structures to show what happens to electrons when the following elements form ionic compounds (Hint: Write the formula first)
Ca + Br
Na + F
Mg + O
Ba + S
81. Draw Lewis Dot Structures to show what happens to electrons in the following covalent bonds:
F2
N2
H2O
CO2
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82. Draw, name the VSPER shapes and identify as polar or nonpolar the following compounds:
H2O
NH3
CHCl3
CO2
Equilibrium
83. Describe a system at equilibrium in terms of the rates of reaction.
84. How is Keq is determined?
85. Determine Keq of the following equation: 2H2 (g) + O2 (g) ⇄ 2H2O (l)
Use the equation below to answer questions 86-88
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2NBr3 (g) ⇄ N2 (g) + 3Br2 (g)
-3 [NBr3 (g)] = 2.07 x 10 M -2 [N2 (g)] = 4.11 x 10 M -3 [Br2 (g)] = 1.06 x 10 M
86. Set up the equilibrium expression.
87. Find the value of Keq at the conditions shown by the molarities above.
88. What compound or element(s) does the reaction favor?
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