Download Spring 2014 Chemistry Review

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Due Date:_______________
Chemistry: Spring Final Review
**Worth 2 HW grades
The Basics
1) Review the safety rules for goggles, chemical safety, lab procedure, fire safety, diluting acids, etc. that
were highlighted at the beginning of the year in your lab book/safety contract.
2) Describe the states of matter using the kinetic theory.
a. Solid:
b. Liquid:
c. Gas:
3) What are the 5 signs that a chemical reaction has taken place?
a. ____________________________
b. ____________________________
c. ____________________________
d. ____________________________
e. ____________________________
4) How many sig figs are in:
1.70
0.00550
20,000
10900
8.0370 x10-2
5) Fill in the Chart:
Prefix
Decimal meaning
Saying
Example
Kilo
100 units
Died
Base Unit
decimeter dm
Milli
Chocolate
Milk
6) How many decimeters are in 12.7 hm?
7) Convert 22.5 L to mL.
8) Convert 0.250 kg to mg. Express in scientific notation.
9) What 2 things does the atomic number give you for a neutral atom?
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10) Define atomic mass number.
11) Define valence electrons of an atom?
12) How many valence electrons are in N, Cl and Na?
N: ____
Cl: ____
Na: ____
13) What is the Octet rule?
Chapter 8/9 – Ionic and Covalent Bonding
14) Compare and give examples of the different types of bonds:
Ionic
Covalent
Polar
Nonpolar
15) Define cation and anion.
16) Gaining three electrons gives what overall charge?
17) Draw the Lewis dot structure for N, K, & S2-.
N
K
18) Name the following compounds:
(Ionic or covalent)
S219) Write the formula for the following:
(Ionic or covalent)
___CaCl2 _____________________________
___zinc bromide ____________________
___CuNO3 ____________________________
___dinitrogen monoxide ______________
___CCl4 ______________________________
___lead (IV) oxide __________________
___Ba(C2H3O2)2 ________________________
___copper (I) carbonate _______________
___Si3H8 _____________________________
___calcium fluoride __________________
20) What properties of metals are explained by metallic bonding theory? What major difference in metals
allows this to happen?
21) Define and give an example of an alloy.
22) What does VSEPR stand for? What do we use it for?
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23) What 2 properties about a molecule do you use to determine the Molecular (VSEPR) shape of a molecule?
a. _____________________________
b. _____________________________
24) Draw the Lewis structure and name the molecular shape for each molecule:
BF3
CO2
CH4
H2O
Chapter 10 – Chemical Reactions
25) Define and give an example of the law of conservation of mass.
26) True / False The Law of Conservation of Energy is why we balance equations.
27) A piece of iron was left outside. It reacted with the oxygen in the air to yield iron(III) oxide or rust. Write a
balanced equation for this reaction.
28) Balance the following reactions:
a. ____ SrBr2 + ____ (NH4)2CO3  ____ SrCO3 + ____ NH4Br
b. ____ Al + ____ Pb(NO3)2  ____ Al(NO3)3 + ____ Pb
c. ____ Fe + ____ O2  ____ Fe3O4
d. ____ Li + ____ O2  ____ Li2O
e. ____ ZnCl2 + ____ KOH  ____ Zn(OH)2 + ____ KCl
Identify the Type of Reaction and the category it falls into.
+
O2

CO2 +
Cl2

NaCl
29) _____________________
C2H2
30) _____________________
Na
31) _____________________
FeS
+
Al
32) _____________________
RbOH
+
HNO3
33) _____________________
Na2CO3 +
MgCl2
34) _____________________
MgO
+

H2O
Al2S3 +
Fe

RbNO3 +
H2O

NaCl +
MgCO3
Mg
O2

+
3
For each chemical equation below, circle product is the precipitate using your solubility rules.
35) Pb(NO3)2 + 2KI 
PbI2 + 2 KNO3
36) 3 MgI2 + 2 K3PO4 
Mg3(PO4)2 + 6 KI
Chapter 11...The Mole
Define:
37) Mole:
38) Molar mass:
39) Empirical formula:
40) Molecular formula:
41) Fill in the blanks on the following mole map:
Calculate the following:
42) How many moles of aluminum are in 162 grams of aluminum?
43) What is the molar mass of CaMg(CO3)2?
44) How many molecules of NaOH are in 112 liters of NaOH?
45) How many grams of N2 are present in 67.2 liters of N2?
46) How many molecules of Ca are present in 320 grams of Ca?
47) What is the percent composition of carbon in C6H12O6?
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48) How many moles of oxygen are in three moles of Mg(NO3)2?
49) Determine the empirical formula for the following compound: 22.1% aluminum, 25.4% phosphorus,
52.5% oxygen:
Determine the molecular formula for the following:
50) What is the molecular formula of a compound that has 40.00% carbon, 6.67% hydrogen and 53.33%
oxygen, and has a molar mass of 150 amu’s?
Chapter 12…Stoichiometry
Use the following reaction to answer the questions that follow:
Fe(s) + O2 (g)

Fe2O3 (s)
51) How many grams of Fe2O3 will be produced from 213 grams of Fe?
52)
How many liters of O2 are required to produce 96 grams of Fe2O3?
53) How many grams of iron are required to react with 82 liters of O2?
54) In the following reaction: 3 H2 (g) +
produce 16.8 liters of NH3?
N2(g) 
2 NH3 (g) How many liters of H2 are required to
55) A sample of 23.4 grams of sodium hydroxide is mixed with 18.7 grams of lithium sulfide, producing
lithium hydroxide and sodium sulfide. What is the limiting reagent? How many grams of lithium
hydroxide are produced?
56) I began this reaction with 20 grams of lithium hydroxide.
a. What is my theoretical yield of lithium chloride?
LiOH + KCl  LiCl + KOH
b. I actually produced 6 grams of lithium chloride. What is my percent yield?
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Chapter 13 & 14…Kinetic Theory and Gas Laws
Define:
57) Barometer:
58) Temperature:
59) Gas Pressure:
60) Kinetic Energy:
61) Absolute Zero:
62) What is the relationship between volume and # of collisions: ____________
63) What is the relationship between temperature and # of collisions: ____________
64) What is the relationship between pressure and # of collisions: ____________
65) What is the relationship between volume and pressure: ____________
66) What is the relationship between temperature and volume: ____________
Temperature Conversions:
67) 334 K = ________ °C
Pressure Conversions:
69) 0.96 atm = ________ kPa
68) - 25°C = ________ K
70) 324 kPa = ______ mmHg
71) 956 mmHg = ______ atm
Write the equations for the following and indicate what is held constant:
72) Boyle’s Law:
76) Ideal Gas Law:
73)
Charles’ Law:
77)
Avogadro’s Law:
74)
Gay-Lussac’s Law:
78)
Dalton’s Law of Partial Pressure:
75)
Combined Gas Law:
79) Temperature is in what unit for all gas laws?
80) STP stands for: _________________________________________
81) Standard temperature = __________
82) Standard pressure = __________ atm, __________ kPa or __________mmHg
Calculate the following:
83) A sample of argon gas is cooled and its volume went from 830 mL to 520 mL. If its final temperature
was 35°C, what was its original temperature?
84) A sealed container of nitrogen gas has a pressure of 780 kPa and a temperature of 18°C. The can of gas
is left in a hot car and the temperature of the gas increases to 65°C. What is the new pressure in the can?
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85) A sample of fluorine gas exerts a pressure of 900 mmHg. When the pressure is changed to 1.75 atm, it
occupies 500 mL. What was the original volume of the gas in liters?
86) How many moles of O2 gas will occupy a volume of 435 mL at 267 kPa and 25°C? How many
molecules of O2?
87) A 300 mL sample of CO2 is submitted to conditions of 5.2 atm and 33°C. What would the volume of
this gas be at STP?
Chapter 15…Solutions
Define:
88) Solution:
89) Solute:
90) Solvent:
91) Miscible:
92) Dilute:
93) Concentrated:
94) Saturated solution:
95) Unsaturated solution:
96) Supersaturated solution:
97) Precipitate:
98) In the solvation of solids, solubility rates increase with (high / low) temperatures.
99) In the solvation of gases, solubility rates increase with (high / low) temperatures and (high / low) pressure.
100) Agitation of a solution will (increase / decrease) the dissolving rate of a solid; whereas it will (increase /
decrease) the dissolving rate of a gas.
101) Adding a solute to a solvent will (increase / decrease) the concentration of the solution.
102) The general rule for solvation is “__________ dissolves __________”.
Use the solubility graph to answer questions:
103) How many grams of NaNO3 will dissolve at 30C?
_______
104) Which substance is least soluble at 10C?
________
105) Which two substances have the same solubility at
72C? ___________________
106) 80 grams of KBr placed in 60C creates a
(saturated/unsaturated/supersaturated) solution.
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Chapter 19…Acids and Bases
Define:
107) Acid:
108) Base:
109) Amphoteric:
110) pH scale:
111) Neutralization:
112) Mono-, Di-, and Triprotic acids:
Fill in the following table of acid-base theories:
Acid Definition
Base Definition
Arrhenius Theory
Bronsted-Lowry Theory
Lewis Theory
Name the following:
113) HC2H3O2:________________
114) HF: __________________
115) H2SO3: __________________
Write the formulas for the following:
116) hydrosulfuric acid _______________
117) perchloric acid _______________
118) nitrous acid _______________
119) Label the compounds in the following reaction as conjugate acid (CA), conjugate base (CB), acid and
base:
OH - (aq) + H2CO3 (aq)
↔
HCO3- (aq) + H2O (l)
120) Water ionizes into __________ and __________.
121) The value for Kw is always equal to _______________.
122) If the concentration OH- is 6.3 x 10-3, what is the concentration of H3O+ in water?
123) Is the above solution acidic or basic?
124) What are the products of a neutralization reaction?
Use the following concentration to answer questions: [H+] concentration of 3.62 x 10-12
125) pH =
127) [OH-] concentration =
126) pOH =
128) Is this solution acidic or basic?
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