Download 60. Write the electron configuration for Zn

Spring Semester Final Exam Study Guide- KEY
Honors Chemistry
Name ____Key_________________
Period ________________________
Naming and Formula Writing
AlCl3
1. Write the name or formula for each of the following:
aluminum chloride
HClO2 chlorous acid
Mercury (II) bromide HgBr2
Cu2S copper (I) sulfide
(NH4)2SO4 ammonium sulfate
Phosphorous acid H3PO3
NaCN sodium cyanide
I4O10 tetraiodine decoxide
Nitric acid
HI hydroiodic acid
H3N hydronitric acid
Sodium nitrite NaNO2
PCl3 phosphorus trichloride
NiN nickel (III) nitride
Magnesium hydroxide Mg(OH)2
CrBr3 chromium(III)bromide
H3PO4 phosphoric acid
Carbon trisulfide CS3
HNO3
Chemical Equations
2. State the law of conservation of matter.
Matter is neither created nor destroyed; it is only rearranged.
2H2 + O2 → 2H2O


3. Using the chemical equation above:
a. Underline the subscripts
b. Box the coefficients (bold)
c. Identify the diatomic molecules (arrows)
4. Balance the following equations and identify the type of reaction (synthesis,
decomposition, single replacement, double replacement, combustion)
a. __1_N2 + __3_H2  __2_NH3
Synthesis
b. _2__KClO3  __2_KCl + _3__O2
Decomposition
c. __2_NaCl + _1__F2  __2_NaF + __1_Cl2
Anionic Single Rep.
d. __2_AgNO3 + _1__MgCl2 _2__AgCl + __1_Mg(NO3)2 Double Rep.
e. Aluminum bromide reacts with potassium sulfate to yield potassium
bromide and aluminum sulfate.
__2_ AlBr3 + __3__ K2SO4  __6_ KBr + __1__ Al2(SO4)3 Double Rep.
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Spring Semester Final Exam Study Guide- KEY
f. Sodium and water form sodium hydroxide and hydrogen
__2__ Na + __2__ H2O  __2__ NaOH + __1__ H2 Cationic Single Rep
g. Silver oxide decomposes to silver and oxygen.
__2__ Ag2O  __4__ Ag + __1__ O2 Decomposition
h. Potassium and magnesium bromide yields potassium bromide and
magnesium
__2__ K + __1__ MgBr2  __2__ KBr + __1__ Mg Cationic Single Rep
5. Identify the following symbols:
a. (aq) - Aqueous
b. (s) - Solid
c. (l) - Liquid
d. (g) - Gas
e. H - Change in energy (Heat of Reaction)
f. 
- two way reaction
g. M - Molar
6. What type of reaction (endothermic, exothermic) has more energy in the
products than reactants? What is the sign for H?
Endothermic. + H
7. What type of reaction has more energy in the reactants than products? What
is the sign for H?
Exothermic. - 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
6.44 g S = 0.201 mol S
__2__ S + __3__ O2  __2__ SO3
.201 mol x 791.4 kJ/ 2 mol S = 79.5 kJ released
The Mole
9. How many atoms are in a mole of Ca?
6.02 x 1023 atoms
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Spring Semester Final Exam Study Guide- KEY
10. How many O atoms are on 3.0 x 1015 molecules of CO2?
3.0 x 1015 molecules of CO2 x 2 atoms O/ 1 molecule CO2 = 6.0 x 1015
11. How many atoms are in 2.6 moles of Al?
2.6 mol x 6.02 x 1023 atoms/ 1 mole = 1.6 x 1024
12. Convert 6689 moles of NaCl to molecules of NaCl.
6689 mol x 6.02 x 1023 atoms/ 1 mole = 4.027 x 1027
13. Convert 3.2 x 1028 molecules of hydrofluoric acid to moles.
3.2 x 1028 molecules x I mol/ 6.02 x 1023 atoms = 53000 mol
14. How many moles are in 60,000,000,000 atoms of Zn?
60,000,000,000 atoms x I mol/ 6.02 x 1023 atoms = 1 x 10-13 moles
15. Calculate the molar mass of each of the following compounds:
a. NaOH = 40.00 g
b. MgCl2 = 95.21 g
c. Barium phosphate = Ba3(PO4)2 = 601.93 g
16. How many grams of sulfur are in 5.23 moles of sulfur?
5.23 mol x 32.07 g/ 1 mol = 168 g
17. Convert 192 moles of HCl to g of HCl
192 moles x 36.46g/1 mol = 7000 g
18. Convert 510.2 g of aluminum hydroxide to moles.
510.2 g x 1 mol/ 78.01g = 6.54 mol
19. How many moles are in 62 g of CO2 ?
62g x 1 mol/ 44.01g = 1.4 mol
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Spring Semester Final Exam Study Guide- KEY
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?
__4__ Al + __3__ O2  __2__ Al2O3
2.3 mol Al2O3 x 4 Al/2 Al2O3 = 4.6 mol Al
b. How many moles of oxygen are required to react completely with 8.4
moles of aluminum?
8.4 mol Al x 3 mol O2/ 4 mol Al = 6.3 mol O2
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?
__1__ CaCl2 + __1_ K2S  __2__ KCl + __1__ CaS
50.0 g CaCl2 x 1 mol/ 110.98g = 0.451 mol CaCl2
0.451 mol CaCl2 x 2 mol KCl/ 1 mol CaCl2 = 0.901 mol KCl
0.901 mol KCl x 74.55 g/ 1 mol = 67.2 g KCl
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.
__2__ KClO3  __2__ KCl + __3__ O2
25.0g KClO3 x 1 mol/ 122.55g = 0.204 mol KClO3
0.204 mol KClO3 x 3 mol O2/ 2 mol KClO3 = 0.304 mol O2
0.304 mol O2 x 22.4 L/1 mol = 6.85 L
23. Predict the products and write a balanced equation for the following: Silver
nitrate reacts with barium chloride to produce ___ AgCl ___ and __ Ba(NO3)2__?
Balanced equation: __2_ AgNO3 + __1_ BaCl2  _2__ AgCl + _1__ Ba(NO3)2
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?
5.0 g AgNO3 x 1 mol/ 169.88g = 0.029 mol AgNO3 – LIMITING REACTANT
4.0 g BaCl2 x 1 mol/ 208.23g = 0.019 mol BaCl2 – EXCESS REAGENT
0.029 mol AgNO3 x 2 mol AgCl/ 2 mol AgNO3 = 0.029 mol AgCl produced
0.029 mol AgNO3 x 1 mol BaCl2/ 2 mol AgNO3 = 0.015 mol BaCl2 used
0.004 mol BaCl2 excess x 208.23g/1 mol = 0.83 g BaCl2 remains
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Spring Semester Final Exam Study Guide- KEY
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.
7.20g C = .600 mol C
Empirical Formula: CH2O
Mass EF: 30.03g
1.20g H = 1.19 mol H
9.60g O = .600 mol O
Molar Mass: 180g/ EF: 30.03g = multiplier of 6
Molecular Formula: C6H12O6
26. Define:
a. Actual yield – Amount actually produced through the lab
b. Theoretical yield – Amount that could have been produced, ideally
c. % yield – Percent of product possible that was retained in the lab
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?
__4__ K + __1__ O2  __2__ K2O
4.28g O2 x1 mol/32g = 0.134 mol O2 x2 mol K2O/1 mol O2= 0.268 mol K2O
0.268 mol K2O x 94.2g/1 mol = 25.2 g K2O Expected
17.36g actual/ 25.2g theoretical x 100% = 68.9% yield
Gases (again)
28. Name four variables that affect gas behavior and identify the relationship.
a. V T – Volume, Temperature, Direct
b. P T – Pressure, Temperature, Direct
c. P V – Pressure, Volume, Inverse
d. P n – Pressure, quantity, Direct
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?
PV = nRT
atm, L, mol, K
Law can be used when one set of conditions are to be compared back to
standard conditions
30. A gas has a pressure of 50.0mmHg at 540K. What will be the pressure at
56C, assuming volume does not change
50.0mmHg x (329K/540K) = 30.5mmHg
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Spring Semester Final Exam Study Guide- KEY
31. An underground cavern contains 98 moles of methane gas at a pressure of
15atm and a temperature of 22C. How many liters of methane does this natural
gas deposit contain?
V = nRT/P
L  atm
98mol x 0.0821
x 295K/ 15atm = 158 L
mole  K
32. What volume does 16.0 g of O2 occupy at STP?
16.0g x 1 mol/32g x 22.4 L/ 1 mol = 11.2 L
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?
4.74g CaCO3 x 1 mol CaCO3/ 100.09g x 1 CO2/1 CaCO3 = 0.0474 mol CO2
0.0474 mol CO2 x 22.4 L/ 1 mol = 1.06 L CO2
Atomic Structure/Periodic Trends
34. What are the five postulates of Dalton’s atomic theory?
1. Elements are made of tiny particles called atoms.
2. All atoms of a given element are identical
3. The atoms of a given element are different from those of any other
element.
4. Atoms of one element can combine with atoms of other elements to
form compounds. A given compound always has the same relative
numbers and types of atoms.
5. Atoms are indivisible in chemical processes. That is, atoms are not
created or destroyed in chemical reactions. A chemical reaction
simply changes the way the atoms are grouped together.
35. Atomic number is the same as the number of _PROTONS___
36. 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
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Spring Semester Final Exam Study Guide- KEY
37. The atomic number and mass of chlorine are:
a. 17, 35
b. 33, 75
c. 16, 32
d. 29, 64
38. The neutrons of an atom are found _IN THE NUCLEUS__
39. List the three subatomic particles. Name their location. Identify their
charges.
Proton
Nucleus
+1
Neutron
Nucleus
0
Electron
Quantum Levels -1
40. An element loses three electrons to form an ion that has a total of 18
electrons. What is the positive charge in the nucleus? What is the element?
Positive Charge in Nucleus: 21
Element: Scandium (Sc)
41. Which subatomic particle has the smallest mass? Electron
42. An ion is formed when an atom _Gains or loses Electrons______
43. The number of protons always equals the number of ___Electrons____ in a
neutral atom.
44. The mass number of an element tells you the # of __Protons_ and
__Neutrons__.
45. Who is credited with discovering…
The nucleus _Rutherford_
The electron _Thomson_
The Planetary Model of the Atom __Bohr__
46. Isotopes have the same number of ___Protons___ but different numbers of
_Neutrons__ and different ____Mass___ numbers.
47. Vertical columns on the periodic table are called _Groups/Families_.
48. Horizontal rows on the periodic table are called __Periods_____.
49.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
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Spring Semester Final Exam Study Guide- KEY
50. KNOW ALL PERIODIC TRENDS for atomic radius and ionization energy
51. As you go across a period from left to right ionization energy:
a. Increases
b. Decreases
c. Doesn’t change
WHY?
As you move left to right on the periodic table, the elements are all in the
same quantum level but have increasing numbers of protons in the
nucleus. That means that the farther right the element, the stronger the
pull from the nucleus.
52. Define: Atomic radius – The distance from the nucleus to the farthest
electron
Ionization energy – Amount of energy required to remove one
electron
53. Why does atomic radius decrease as you go across a period?
Atomic radius decreases because as you move left to right on the periodic
table, the elements are all in the same quantum level but have increasing
numbers of protons in the nucleus. That means that the farther right the
element, the stronger the pull from the nucleus and the electrons would be
pulled closer.
54.Which groups of elements are the most reactive _Alkali Metals, Halogens_?
55. Which group of elements is the least reactive __Noble Gases____?
56. Which group (family) do the following elements belong to?
Li- Alkali Metals
C- Carbon Group
Be- Alkaline Earth Metals
N- Nitrogen Group
V- Transition Metals
O- Oxygen Group
La- Inner Transition Metals
F- Halogens
B- Boron Group
Ne- Noble Gases
57. What ions will the following elements form?
Li- +1
Be- +2
Al- +3
Ne- None
N- -3
58. Classify the following as metals, semimetals, or non-metals.
As- Semimetal
Si- Semimetal
Na- Metal
B- Semimetal
Ne- Nonmetal
Ar- Nonmetal
Cl- Nonmetal
H- Hydrogen!
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Spring Semester Final Exam Study Guide- KEY
ELECTRONS:
59. Draw the orbital diagram for Phosphorus.
60. Write the electron configuration for Zn
1s22s22p63s23p64s23d10
61. 1s22s22p63s1 is the electron configuration of _Sodium (Na) __.
62. What rule states that electrons enter the lowest energy orbitals until all
electrons are accounted for? _Aufbau Principle _
63. How many orbitals are there in the:
s group- 1
p group- 3
d group- 5
f group- 7
64. How many electrons can each of the sublevels above hold?
s group- 2
p group- 6
d group- 10
f group- 14
VSPER Shapes and Polarity
65. Compare and contrast ionic and covalent bonds in terms of:
a. Shape: Ionic is Lattice, Covalent is VSEPR
b. Melting and Boiling Point: Ionic High, Covalent Low
c. Conductivity: Ionic is when in solution, Covalent is not
66. What happens to electrons in ionic bonds?__Transfered__
67. What happens to electrons in covalent bonds? __Shared____
68. What type of bond do the following Lewis Dot Structures Represent?
Na  S  Na
Ionic
O=C=O
Covalent
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Spring Semester Final Exam Study Guide- KEY
69. Draw Lewis Dot Structures to show what happens to electrons when the
following elements form ionic compounds (Hint: Write the formula first)
Ca + Br
CaBr2
Na + F
NaF
Mg + O
MgO
Ba + S
BaS
70. Draw Lewis Dot Structures to show what happens to electrons in the
following covalent bonds:
F2
F–F
N2
NN
H2O
H–O–H
CO2
O=C=O
71. Draw, name the VSPER shapes and identify as polar or nonpolar the
following compounds:
H2O
Bent
Polar
NH3
Pyramidal
Polar
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Spring Semester Final Exam Study Guide- KEY
CHCl3
Tetrahedral
Polar
CO2
Linear
Nonpolar
Equilibrium
72. Describe a system at equilibrium in terms of the rates of reaction.
Rate of forward reaction = rate of reverse reaction
73. How is Keq is determined?
Concentration of the products/Concentration of reactants
74. Determine Keq of the following equation: 2H2 (g) + O2 (g)  2H2O (l)
1 / [H2]2 [O2]
Use the equation below to answer questions 86-88
2NBr3 (g)  N2 (g) + 3Br2 (g)
[NBr3 (g)] = 2.07 x 10-3 M
[N2 (g)] = 4.11 x 10-2 M
[Br2 (g)] = 1.06 x 10-3 M
75. Set up the equilibrium expression.
[Br2]3 [N2]
[NBr3]2
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Spring Semester Final Exam Study Guide- KEY
76. Find the value of Keq at the conditions shown by the molarities above.
(1.06 x 10-3 M)3 x 4.11 x 10-2 M / (2.07 x 10-3 M)2 = 1.14 x 10-5
77.
What compound or element(s) does the reaction favor?
Reactants: NBr3
Solutions and Acids / Bases
78. In a solution the substance being dissolved is called the ___SOLUTE___
and the substance it is dissolved in is called the ___SOLVENT__.
79. To what volume should 25mL of 15M nitric acid be diluted to prepare a 3.0M
solution?
0.025 L x 15 mol/ 1 L = 0.375 mol
0.375 mol x 1 L/ 3.0 mol = 0.125 L or 125 mL
80. What is the molarity of a solution in which 10.0g of AgNO3 is dissolved in
500mL of solution?
10.0g x 1 mol/ 169.88g = .0588mols AgNO3 .0588mols / .500L = .118M
81. What is the pH range of acids? Bases?
Acids: 0-7; Bases: 7-14
82. A solution with more hydroxide ions than hydrogen ions is called a
___BASE__________.
83. Litmus paper turns what color in the presence of an acid? Base?
Acid: Red
Base: Blue
84. Which of the following pH’s if for the weakest acid? pH 3 , pH 1, pH 6
85. What is the [H+] of a solution with a pH of 4?
1x10-4M
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Spring Semester Final Exam Study Guide- KEY
86. What differentiates a weak base from a strong base? What differentiates a
weak acid from a strong acid?
Strong acids and bases completely dissociate in water.
Weak acids and bases do not completely dissociate in water.
87. What signals the end of most lab titrations?
Color Change of Indicator
88. What volume of 0.175 M HCl is needed to exactly neutralize 25.0 mL of
0.150 M KOH solution?
a. Neutralization Reaction:
HCl + KOH  H2O + KCl
b. Solution: Find moles of KOH .025L x .150mols/1L = .00375 mols KOH
1:1 Ratio between KOH and HCl so .00375 mols HCl
.00375 mols HCl x 1L/.175 mols = .0214L or 21.4mL
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