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CChheemmiissttrryy WWoorrkksshheeeett aanndd
LLaabboorraattoorryy MMaannuuaall ((TTeerrmm 11))
MMrr.. GGeeiisstt
Per
iodi
c T
able
of E
lem
ents
(A
dditi
onal
Val
ues
and
Con
stan
ts o
n ba
ck p
age)
Yo
u m
ay a
dd
ad
dit
ion
al in
form
atio
n i
n y
ou
r o
wn
han
dw
riti
ng
in
th
is b
ox.
N
ame:
___
____
____
____
____
____
____
__
Exa
m:
MA
ST
ER
CO
PY
P
erio
d: _
____
___
Do
not w
hite
-out
, add
add
ition
al p
aper
, or
tape
. O
nly
writ
e in
box
to th
e le
ft, o
r be
un
able
to
use
this
she
et o
n th
e te
st.
8A
1A
1 H
Hyd
roge
n
1.00
79
2A3A
4A
5A
6A
7A
2 He
Hel
ium
4.00
263 Li
Lith
ium
6.94
1
4 Be
Ber
ylliu
m
9.01
22
5 B
Bor
on
10.8
1
6 C
Car
bon
12.0
11
7 N
Nitr
ogen
14.0
07
8 O
Oxy
gen
15.9
99
9 F
Flu
orin
e
18.9
98
10
Ne
Neo
n
20.1
79
11
Na
Sod
ium
22.9
90
12
Mg
M
agn
esiu
m
24.3
05
13
Al
Alu
min
um
26.9
82
14
Si
Sili
con
28.0
86
15
P
Pho
spho
rus
30.9
74
16
S
Sul
fur
32.0
6
17
Cl
Chl
orin
e
35.4
53
18
Ar
Arg
on
39.9
48
19
K
Pot
assi
um
39.0
98
20
Ca
Cal
cium
40.0
8
21
Sc
Sca
ndiu
m
44.9
56
22
Ti
Tita
nium
47.9
0
23
V
Van
adiu
m
50.9
41
24
Cr
Chr
omiu
m
51.9
96
25
Mn
M
ang
anes
e
54.9
38
26
Fe
Iron
55.8
47
27
Co
C
oba
lt
58.9
33
28
Ni
Nic
kel
58.7
1
29
Cu
C
opp
er
63.5
46
30
Zn
Z
inc
65.3
8
31
Ga
G
alliu
m
69.7
2
32
Ge
G
erm
aniu
m
72.5
9
33
As
Ars
enic
74.9
22
34
Se
Sel
eniu
m
78.9
6
35
Br
Bro
min
e
79.9
04
36
Kr
Kry
pton
83.8
0
37
Rb
R
ubid
ium
85.4
68
38
Sr
Str
ontiu
m
87.6
2
39
Y
Yttr
ium
88.9
06
40
Zr
Zirc
oniu
m
91.2
2
41
Nb
N
iob
ium
92.9
06
42
Mo
M
olyb
den
um
95.9
4
43
Tc
Tec
hnet
ium
(97)
44
Ru
R
uthe
nium
101.
07
45
Rh
R
hod
ium
102.
91
46
Pd
P
alla
dium
106.
4
47
Ag
S
ilver
107.
87
48
Cd
C
adm
ium
112.
41
49
In
Indi
um
114.
82
50
Sn
T
in
118.
69
51
Sb
A
ntim
ony
121.
75
52
Te
Tel
luriu
m
127.
60
53 I
Iodi
ne
126.
90
54
Xe
Xen
on
131.
30
55
Cs
Ces
ium
132.
91
56
Ba
Bar
ium
137.
33
71
L
u
Lute
tium
174.
97
72
Hf
Haf
nium
178.
49
73
Ta
Tan
talu
m
180.
95
74
W
Tun
gste
n
183.
85
75
Re
Rhe
niu
m
186.
21
76
Os
O
smiu
m
190.
2
77
Ir
Irid
ium
192.
22
78
Pt
Pla
tinum
195.
09
79
Au
G
old
196.
97
80
Hg
M
ercu
ry
200.
59
81
Tl
Tha
llium
204.
37
82
Pb
Le
ad
207.
2
83
Bi
Bis
mut
h
208.
98
84
Po
P
olon
ium
(209
)
85
At
Ast
atin
e
(210
)
86
Rn
R
ado
n
(222
) 87
Fr
Fra
nciu
m
(223
)
88
Ra
Rad
ium
(226
)
103
L
rLa
wre
nciu
m
(260
)
104
Rf
Rut
herf
ordi
um
(261
)
105
Db
D
ubn
ium
(262
)
106
Sg
S
eabo
rgiu
m
(263
)
107
Bh
B
ohri
um
(262
)
108
Hs
Has
sium
(265
)
109
Mt
Mei
tner
ium
(266
)
110
Ds
Dar
mst
adtiu
m
(28
1)
111
Rg
R
oent
gen
ium
(282
)
112
Cn
C
oper
nic
ium
(28
5)
113
Nh
N
iho
nium
(286
)
114
Fl
Fle
rovi
um
(289
)
115
Mc
M
osco
vium
(290
)
116
Lv
Live
rmor
ium
(293
)
117
Ts
Ten
ness
ine
(294
)
118
Og
O
gane
sson
(294
)
L
anth
anid
e se
ries
57
La
Lant
han
um
138.
91
58
Ce
Cer
ium
140.
12
59
Pr
Pra
seod
ymiu
m
140.
91
60
Nd
N
eod
ymiu
m
144.
24
61
Pm
P
rom
ethi
um
(145
)
62
Sm
S
amar
ium
150.
4
63
Eu
E
urop
ium
151.
96
64
Gd
G
adol
iniu
m
157.
25
65
Tb
T
erbi
um
158.
93
66
Dy
Dys
pros
ium
162.
50
67
Ho
H
olm
ium
164.
93
68
Er
Erb
ium
167.
26
69
Tm
T
hulli
um
168.
93
70
Yb
Y
tterb
ium
173.
04
A
ctin
ide
serie
s 89
Ac
Act
iniu
m
(227
)
90
Th
T
horiu
m
232.
04
91
Pa
Pro
actin
ium
231.
04
92
U
Ura
nium
238.
03
93
Np
N
eptu
nium
237.
05
94
Pu
P
luto
niu
m
(244
)
95
Am
A
mer
iciu
m
(243
)
96
Cm
C
uriu
m
(247
)
97
Bk
Ber
keliu
m
(247
)
98
Cf
Cal
iforn
ium
(251
)
99
Es
Ein
stei
niu
m
(254
)
100
Fm
F
erm
ium
(257
)
101
Md
M
end
elev
ium
(258
)
102
No
N
obe
lium
(259
)
A
verage E
lectron
egativities o
f the E
lemen
ts
Gen
eral Ph
ysical Co
nstan
ts
Avo
gad
ro’s C
on
stant
6.022 x 102
3 rp/m
ol
Plan
ck’s Co
nstan
t 6.626 x 10
–3
4 Js B
oltzm
ann
’s Co
nstan
t 1.381 x 10
–2
3 J/K
Sp
eed o
f Lig
ht
3.0 x 108 m
/s
Ato
mic M
ass Un
it 1.6605655 x 1
0–
27 kg
R
ydb
erg’s C
on
stant
1.097 x 107 m
–1
Farad
ay’s C
on
stant
96485.309 C/m
ol
Ideal G
as Co
nstan
t 8.31 (LkP
a)/(Km
ol)
0.0821 (L
atm)/(K
mo
l)
62.396 (Lm
m H
g)/(K
mo
l)
1A
2
A
3B
4B
5B
6B
7B
8B
1B
2B
3A
4
A
5A
6
A
7A
8
A
H
2.1
H
e
--
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
N
e
--
Na
0.9
M
g
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
Ar
--
K
0.8
Ca
1.0
S
c 1.3
T
i 1.5
V
1.6
C
r 1.6
M
n
1.5
Fe
1.8
Co
1.9
N
i 1.9
C
u
1.9
Zn
1.6
G
a 1.6
G
e 1.8
A
s
2.0
Se
2.4
Br
2.8
Kr
--
Rb
0.8
S
r 1.0
Y
1.2
Z
r 1.4
N
b
1.6
Mo
1.8
T
c 1.9
R
u
2.2
Rh
2.2
P
d
2.2
Ag
1.9
C
d
1.7
In
1.7
Sn
1.8
S
b
1.9
Te
2.1
I 2.5
X
e --
Cs
0.7
B
a
0.9
Ln
1.0
H
f 1.3
T
a 1.5
W
1.7
R
e
1.9
Os
2.2
Ir 2.2
P
t 2.2
A
u
2.4
Hg
1.9
T
l 1.8
P
b
1.9
Bi
1.9
Po
2.0
A
t 2.2
R
n
--
Fr
0.7
Ra
0.9
A
c
1.0
Th
1.2
P
a 1.5
U
1.7
Other P
olyatomic Ions
HP
O4
2–:
hydro
gen
phosphate
H
2 PO
41–:
dihyd
rogen
phosph
ate
HS
O3
1–:
hydro
gen
sulfite
HS
O4
1–:
hydro
gen
sulfate
HC
O3
1–:
hydro
gen
carbonate
ClO
41
–: perch
lorate
C
lO3
1–:
chlorate
ClO
21
–: chlorite
ClO
1–:
hypochlorite
C2 O
42–:
oxalate
M
ain Po
lyatomic Ions
N
O3
–: nitrate
SO
42
–: sulfate
PO
43
–: phosph
ate
N
O2
–: nitrite
SO
32
–: sulfite
PO
33
–: phosphite
CO
32
–: carbo
nate
C2 H
3 O2
–: acetate
OH
–: h
ydroxide
NH
4+:
amm
onium
S
iO3
2–:
silicate
C
N–:
cyanide
MnO
4–:
perma
nganate
CrO
42
–: chrom
ate C
r2 O7
2–: dichrom
ate
A
ctivity Series of M
etals/Halo
gens (N
OT
E: R
eactivity of the m
etal/halog
en decreases as it gets
low
er on the list.)
S
olubility R
ules in W
ater
Neg
ative Ion
Rule
NO
3–, C
lO3
–, C
lO4
–, C
2 H3 O
2–
All com
poun
ds formed w
ith the N
O3
–, C
lO3
–, ClO
4 –, or C2 H
3 O2 – ion are solu
ble in w
ater.
M
etals H
aloge
ns
Lithium
Potassium
B
arium
Calcium
S
odium
Magn
esium
Alum
inum
M
anganese
Z
inc C
hromium
Iron C
obalt
Nickel
Tin
Lead
H
ydrogen
C
opper
Mercury
Silver
Platin
um
Gold
F
luorine
C
hlorine
B
romin
e
Iodine
I –, Br –, C
l – A
ll comp
ounds form
ed with th
e I –, Br –, or
Cl – ion are solu
ble in water e
xcept Ag
+, P
b2
+, Hg
22
+, and Cu
+.
SO
42
–
Most com
poun
ds formed w
ith the SO
42
– ion are solu
ble in water; e
xcep
tions include S
rSO
4 , BaS
O4 , C
aSO
4 , RaS
O4 ,
Ag
2 SO
4 , and PbS
O4 .
CO
32
–, PO
4 3–,
SO
32
–, C2 O
42–,
CrO
42
–, S2–
All com
poun
ds formed w
ith the C
O3
2–,
PO
43
–, SO
32–, C
2 O4
2–, CrO
42
–, or S2
– ion are insolu
ble in w
ater except those of the
alkali metals a
nd NH
4+.
Monatom
ic Ions
C
u1
+: copper (I) ion
C
u2
+: copper (II) ion
F
e2
+: iron (II) ion
F
e3
+: iron (III) ion
P
b2
+: lead (II) ion
P
b4
+: lead (IV
) ion
Sn
2+:
tin (II) ion
Sn
4+:
tin (IV) ion
C
o2
+: cobalt (II) ion
C
o3
+: cobalt (III) ion
U
seful Conversion F
actors A
nd Co
nversions
Energy:
1 cal = 4.184 J
Length: 1 inch =
2.54 cm
Mass:
1 lb = 0.453
6 kg P
ressure: 1 atm
= 101.3 kP
a
1 atm =
760 mm
Hg
Tem
p.: K
= C
+ 273.1
5
C =
K – 273.15
Volum
e: 1 L =
0.001 m3
1 cm
3 = 1 m
L
OH
–
All com
poun
ds formed w
ith the O
H– ion
are insoluble in
water e
xcept those of the alkali m
etals, NH
4+, S
r 2+, and B
a2
+. (C
a(OH
)2 is slightly soluble.)
Table of Contents
Unit One Worksheet ..................................................................................... 1 Unit Two Worksheet ................................................................................... 15 Unit Three Worksheet ................................................................................ 28 Unit Four Worksheet .................................................................................. 33 Unit One Experiment – 1: Bunsen Burner Operation ................................ 46 Unit One Experiment – 2: Mass and Change ............................................ 49 Unit One Experiment – 3: Density and Relationships ............................... 53 Unit One Experiment – 4: Measurement and Uncertainty ......................... 58 Unit Two Experiment – 1: Average Atomic Mass ...................................... 61 Unit Two Experiment – 2: Flame Test Analysis ........................................ 63 Unit Two Experiment – 3: Cation and Anion Analysis ............................... 65 Unit Four Experiment – 1: Precipitation Reactions ................................... 67 Unit Four Experiment – 2: Pipet Rockets and Synthesis .......................... 70 Unit Four Experiment – 3: Balanced Chemical Equations ........................ 73 Appendix A – Laboratory Equipment, Syllabus, and LPS Safety Contract ......................................................... A-1 Appendix B – SI Units and Conversions ................................................ A-10 Appendix C – Compound Name and Formula Writing ........................... A-13 Appendix D – Chemical Reactions and Quantities ................................ A-15
Appendix E – Practice Tests .................................................................. A-22 Appendix F – Practice Test Keys ........................................................... A-48
page 1 – C – T1 – BOOK
Unit One Worksheet WS – C – U1
Section 1.3 Matching. Match each term with its correct definition. _______1. The use of ones senses to obtain (A) Experiment information directly (B) Hypothesis _______2. A broad and extensively tested (C) Observation explanation of why experiments give (D) Scientific method certain results (E) Scientific law _______3. A logical approach to the solution of (F) Scientific theory scientific problems _______4. A concise statement that summarizes the results of many observation and experiments _______5. A means to test a hypothesis _______6. A proposed explanation for an observation Matching. Match each application with its correct step of the scientific method. _______7. An iron ball falls to Earth when you (A) Experiment drop it. (B) Hypothesis _______8. Earth is a giant magnet. (C) Observation (D) Theory _______9. An iron ball and a piece of wood are dropped at the same time from the same height. _______10. The iron ball and wood fall at the same rate. _______11. The large mass of Earth causes it to exert the same gravitational attraction on any object, regardless of the object’s composition. Short Answer. Answer the following question. 12. Explain the statement “No theory is written in stone.” BOOK PROBLEMS: Section review 1.3, page 17, #8 – 11
page 2 – C – T1 – BOOK
Section 2.1 Matching. Match each term with its correct definition. _______13. A quality or condition of a substance (A) Gas that can be observed or measured (B) Liquid without changing the substance’s (C) Mass composition (D) Matter _______14. Matter that assumes both the shape (E) Physical change and volume of its container (F) Physical property _______15. Matter that has a uniform and (G) Solid
definite composition (H) Substance _______16. Anything that has mass and takes up (I) Vapor
space _______17. Matter that has a definite shape and volume _______18. The amount of matter that an object contains _______19. Matter that has a definite volume and takes the shape of its container _______20. Alteration of a material without changing its chemical composition _______21. Gaseous state of a substance that generally exists as a liquid or solid at room temperature Matching. Match each substance, existing at room temperature, with its state of matter. _______22. Steam (A) Solid (B) Liquid _______23. Gasoline (C) Gas (D) Vapor _______24. Hockey puck _______25. Filtered apple juice _______26. Air Identification. State whether each of the following is a physical or chemical change or property by
writing an ”A” if it is a physical change, “B” if it is a physical property, “C” if it is a chemical change, or “D” if it is a chemical property.
_______27. Melting butter _______34. Boiling water _______28. Flash point _______35. Decomposing flesh _______29. Food spoiling _______36. Freezing liquid iron _______30. Burning gasoline _______37. Density _______31. Breaking a tree twig _______38. Breaking an icicle _______32. Antacid tablet fizzing in water _______39. Vaporizing water _______33. Detonating an explosive _______40. Flammability
page 3 – C – T1 – BOOK
_______41. Dry ice sublimating (turning directly into a gas) BOOK PROBLEMS: Section review 2.1, page 31, #1 – 4 Section 2.2 Matching. Match each term with its correct definition. _______42. A mixture without a completely (A) Distillation uniform composition (B) Heterogeneous mixture _______43. Any part of a system that has uniform (C) Homogenoeus mixture composition and properties (D) Mixture _______44. A mixture with a completely uniform (E) Phase composition (F) Solution _______45. Separation of a liquid solution by
boiling and recondensation _______46. A special name for a homogeneous _______47. A physical blend of two or more
mixture substances Matching. Match each described matter with its type. _______48. Oxygen dissolved in water (A) Compound (B) Element _______49. Carbon mixed with sand (C) Heterogeneous mixture (D) Homogeneous mixture _______50. Hot tea _______51. Sugar (sucrose) _______58. C3H8 _______52. Salt dissolved in water _______59. A classroom _______53. Titanium _______60. Tap water _______54. Salt (sodium chloride) _______61. Hafnium _______55. Air _______62. Carbon dioxide _______56. Vegetable soup _______63. Distilled water _______57. Sterling silver _______64. Cleaning solution Short Answer. Answer the following questions. 65. How might one successfully separate a mixture of salt and water?
page 4 – C – T1 – BOOK
66. Describe a procedure that could be used to separate a mixture consisting of sand and salt. BOOK PROBLEMS: Section review 2.2, page 35, #9, 11, 12 Section 2.3 Matching. Match each substance with its classification. _______67. Plutonium (A) Element (B) Compound _______68. Water _______69. Xenon _______72. Aluminum oxide _______70. Glucose (C6H12O6) _______73. Carbon _______71. Cesium chloride _______74. Sodium Table Completion. Complete the following tables.
Element Symbol Element Symbol 75.
Cr 81.
Pt
Potassium 76.
Fluorine
82.
77.
Rn 83.
Na
Magnesium 78.
Tungsten 84.
79.
Au 85.
Pb
Tin 80.
Nickel 86.
Short Answer. Answer the following questions. 87. How can you distinguish between an element and a compound? 88. A liquid is allowed to evaporate and leaves no residue. Can you determine whether it was an
element, a compound, or a mixture? Explain. BOOK PROBLEMS: Section review 2.3, page 40, #18
page 5 – C – T1 – BOOK
Section 3.1 Calculations. Answer the following questions. 89. Subtract 2.9 x 104 from 5.00 x 105 and express the answer using scientific notation. Show work
or receive no credit. 90. Divide 5.50 x 105 by 2.5 x 104 and express the answer using scientific notation. Show work or
receive no credit. 91. Add 5 x 104 and 6 x 103 and express the answer using scientific notation. Show work or receive
no credit. 92. Multiply 2.5 x 107 by 4.00 x 108 and express the answer using scientific notation. Show work or
receive no credit. BOOK PROBLEMS: Section review 3.1, page 53, #2 – 4 Section 3.2 Short Answer. Answer the following questions. 93. If you measure a line three times with the same ruler, do your measurements become more
accurate? Explain. 94. Give a real-life of example (i.e., from sports, tests, etc.) to illustrate excellent precision but poor
accuracy.
page 6 – C – T1 – BOOK
Significant figures are all the digits you know for sure and one place that is an estimate. Uncertainty is the limit of precision of the reading (based on your ability to estimate the final digit). See examples below.
Rules for zeros: All zeros count except placeholder zeros (the ones that disappear when you write the
number in scientific notation). Examples:
93,000,000 = 9.3 x 107 2 significant figures
0.000372 = 3.72 x 10-4 3 significant figures
0.0200 = 2.00 x 10-2 3 significant figures Readings and Figures. For each of the following, write the scale reading and then the number of
significant figures (SF’s) in the reading. Reading SF’s
95.
page 7 – C – T1 – BOOK
Reading SF’s
96.
97.
98.
99.
100.
101.
102. Identification. Identify the number of significant figures for each of the following measurements. _________103. 0.0000935 m _________108. 500 km _________104. 12.01C _________109. 500.0 km _________105. 0.007000 m _________110. 1.000083 m _________106. 2.350 x 10-5 eV _________111. 40000000000000 kJ _________107. 72 animals _________112. 3000005 V
page 8 – C – T1 – BOOK
Calculations. Answer the following questions. Show work or receive no credit. Include proper units. Michael’s three measurements are 19.0 cm, 20.01 cm, and 24.0 cm. 113. Calculate the average value of his measurements and express the answer with the correct
number of significant figures and/or decimal places. 114. If the actual length of the object is 20.3 cm, what is the error of Michael’s average measurement?
Express the answer with the correct number of significant figures and/or decimal places. 115. If the actual length of the object is 20.3 cm, what is the percent error of Michael’s average
measurement? BOOK PROBLEMS: Section review 3.2, page 62, #14, 16 Section 3.3 Multiple Choice. Select the letter of the equation that is correct. _________116. A) 1 L = 1 cm3 B) 1 mL = 1 cm3 _________117. A) 0C = – 273 K B) 0 K = – 273C _________118. A) 1 kg = 1000 g B) 1 kg = 100 cg _________119. A) 40 cm = 4.0 m B) 500 cm = 5 m Table Completion. Complete the table below by supplying the missing information of what is being
measured, base SI units, and symbols.
Measurement Base unit Symbol
length 123. 126.
120.
kilogram 127.
time 124. 128.
121.
kelvin 129.
energy 125. 130.
122.
mole 131.
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Short Answer. Answer the following questions. 132. Between the moon and Earth, where would the same object experience more weight? Explain. 133. Between the moon and Earth, where would the same object experience more mass? Explain. BOOK PROBLEMS: Section review 3.3, page 67, #20, 22 Section 3.4 Calculations. Answer the following questions. Show work or receive no credit. Include proper units. 134. What is the mass of a bar of aluminum measuring 2.0 cm by 1.5 cm by 2.0 cm? (HINT: Refer to
your textbook on page 69 for the density of aluminum.) 135. An object measuring 4.0 cm by 5.0 cm by 5.0 cm has a mass of 220 grams. What is the density
of the object? A fish tank measures 0.40 m long by 200 mm wide by 30 cm high. 136. What is the width of the tank in centimeters? 137. What is the length of the tank in millimeters? 138. What is the volume of the tank in cubic centimeters? (HINT: You may want to get all dimensions
of the same length units.)
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139. What is the mass of water, in grams, that would fill the tank halfway? Show work or receive no credit. (HINT: The density of water is 1.0 g/cm3.)
Short Answer. Answer the following questions. 140. A balloon filled with air is released in a room filled with carbon dioxide. Will the balloon float to
the ceiling or sink to the floor? Explain. (HINT: Refer to page 69 of your textbook.) Modeling. Answer the following questions. Show work or receive no credit. You must also show
proper units and express answers using the correct number of significant figures/decimal places.
Refer to the graph below for questions 141 – 144.
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141. Write the equation of the line for substance A. (No work needs to be shown.) 142. Calculate the mass of a 14.0 cm3 piece of substance A. 143. What would occupy the largest volume: 50 g of substance A, 50 g of substance B, or 50 g of
substance C? Explain. 144. Based on the graph from the previous page and the table at the
right, which element or compound is substance A? 145. What is the volume (in mL) of an object with a density of 7.02 g/cm3
and a mass of 6.00 x 102 g? BOOK PROBLEMS: Section review 3.4, page 72, #26, 28 Section 3.5 Calculations. Answer the following questions. Show work or receive no credit. Include proper
units, decimal places, and/or significant figures. 146. A common temperature used in chemistry is 25.0C at standard pressure. What is this
temperature in Kelvins?
Substance Density (g/mL)
Aluminum Ammonia Carbon dioxide Chlorine gas Corn oil Ethanol Gasoline Nitrogen gas Neon Oxygen gas Sucrose
2.70 0.718
1.83 2.95
0.922 0.789
0.67 1.17 0.84 1.33 1.59
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147. A slurry of dry ice in acetone has a temperature of –78C. What is this temperature in Kelvins? Show work or receive no credit.
148. A typical refrigerator keeps food at 277 K. What is this temperature in degrees Celsius? Show
work or receive no credit. 149. How is absolute zero expressed on the Celsius scale? Show work or receive no credit. BOOK PROBLEMS: Section review 3.5, page 75, #33 – 35 Section 4.2 Calculations. Answer the following questions. Show work or receive no credit. Include proper
units, decimal places, and/or significant figures. 150. How many milliseconds are there in one day? 151. If 30 gits equal 2 erbs, 1 futz equals 12 hews, and 10 erbs equal 1 futz, how many gits equal 16
hews? 152. Express the speed 35 centimeters/minute in kilometers/hour. 153. Teachers in Lincoln Public Schools are contracted for 7.5 hours per day. If Mr. Cooper is really
stressed out and counting how many seconds make up his work day, how many seconds would he accurately count?
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154. Gold has sold for $1500/ounce. Considering that there are 16 ounces (454 grams) in a pound, how many milligrams of gold could a person buy for ten thousand dollars?
155. An automobile can travel 35.0 miles on 1 gallon of gasoline. How many kilometers per liter is
this? (1.61 km = 1 mile; 1 liter = 1.06 quarts; 1 gallon = 4 quarts) 156. The density of water is 1.0 g/mL. What is the density of water in pounds/gallon? (1 quart = 9.46
x 10–1 L; 1 g = 2.20 x 10–3 lb; 1 gallon = 4 quarts) 157. Convert 35 kilomoles to centimoles. Express your answer in scientific notation. 158. In 1976, an airplane was flown at a speed of 2193 miles per hour. What was the speed of the
plane in meters per second? (1 kilometer = 0.621 miles) 159. Light travels at 3.0 x 108 meters/second. How quickly does light travel in kilometers/hour?
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160. Express the density 0.250 g/cm3 in kg/m3. 161. Express the pressure 2.5 g/cm2 in psi, or pounds per square inch (lb/in2). (1 pound = 454 grams;
1 inch = 2.54 cm) Short Answer. Answer the following questions. 162. How does a conversion factor differ from a measurement? Give an example of each. 163. When you use dimensional analysis, identify at least two things you should ALWAYS double-
check. 164. Explain why the value of a conversion factor is always 1 even though both the units and the
numbers in the conversion factor are different from each other. Show an example to illustrate your reasoning.
165. When using dimensional analysis, what determines the number of significant figures in the final
answer? Explain.
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Unit Two Worksheet WS – C – U2
Chapter Five Matching. Match the definition with the term it defines. _____166. Atom A) Number of protons that an atom of an element contains B) Sum of numbers of protons and neutrons for an atom of _____167. Atomic mass an isotope of an element C) Subatomic particle containing no charge _____168. Atomic mass unit D) Negatively-charged subatomic particle E) Positively-charged subatomic particle _____169. Atomic number F) The core of an atom containing protons and neutrons G) Atoms of the same element with different numbers of _____170. Electron neutrons H) The smallest particle of an element that retains the _____171. Isotopes properties of the element I) 1/12 the mass of a carbon-12 atom _____172. Mass number J) Weighted average mass of the atoms in a naturally occurring sample of an element _____173. Neutron _____174. Nucleus _____175. Proton Table Completion. Fill in the table using the information provided in the left-most column to identify
the following numbers corresponding to the term for each isotope.
Information Atomic Number
Mass Number
Number of Neutrons
Number of Electrons
Number of Protons
U23592
176. 177. 178. 179. 180.
P3115
181. 182. 183. 184. 185.
-209F
186. 187. 188. 189. 190.
213756Ba
191. 192. 193. 194. 195.
Oxygen-17 196.
197.
198. 199. 200.
Lanthanum-139
201. 202. 203. 204. 205.
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Short Answer. Answer the following questions. Which person(s) was(were) responsible for the discovery of the following subatomic particles? 206. Neutron: _____________________________________________________________________ 207. Electron: _____________________________________________________________________ 208. Proton: ______________________________________________________________________ 209. Where is the mass of the atom located? Explain. What are the four components of Dalton’s atomic theory? 210. _____________________________________________________________________________ _____________________________________________________________________________ _____________________________________________________________________________ 211. _____________________________________________________________________________ _____________________________________________________________________________ _____________________________________________________________________________ 212. _____________________________________________________________________________ _____________________________________________________________________________ _____________________________________________________________________________ 213. _____________________________________________________________________________ _____________________________________________________________________________ _____________________________________________________________________________ Calculations. Show work or receive no credit. 214. The element oxygen contains three naturally occurring isotopes:
O O O 188
178
168
The relative abundances and atomic masses are 99.759% for oxygen-16 (mass = 15.995 amu), 0.037% for oxygen-17 (mass = 16.995 amu), and 0.204% for oxygen-18 (mass = 17.999 amu). Calculate the average atomic mass of oxygen.
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215. The element nitrogen contains two naturally occurring isotopes:
N N 157
147
The relative abundances and atomic masses are 99.63% for nitrogen-14 (mass = 14.003 amu) and 0.37% for nitrogen-15 (mass = 15.000 amu). Calculate the average atomic mass of nitrogen.
Fill in the Blank. Fill in the blank with the appropriate word or phrases. 216. Group 8A elements are known as the _____________________________________________. 217. Group 7A elements are known as the _____________________________________________. 218. Group 2A elements are known as the _____________________________________________. 219. Group 1A elements are known as the _____________________________________________. 220. Group A elements are known as the ______________________________________ elements. 221. Group B elements are known as the ______________________________________ elements. Identification. Identify each of the following characteristics as being those of a metal, nonmetal, or
metalloid. ________________________222. Elements on the right side of the periodic table of elements ________________________223. Elements along the zig-zag line of the periodic table of elements ________________________224. Elements that are gases ________________________225. The majority of elements ________________________226. Have a luster ________________________227. Do not conduct electricity or conduct it poorly ________________________228. Conduct electricity well ________________________229. The halogens ________________________230. Elements on the left side of the periodic table of elements
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Section 6.1 Matching. Match the definition with the term that best correlates to it. No definition will be used more
than once. _____231. Anion _____234. Ionic compound _____232. Cation _____235. Molecular compound _____233. Ion _____236. Molecule A) Neutral atom that loses or gains electrons B) Smallest electrically neutral unit of a substance that maintains all the properties of the substance C) Compound composed of molecules D) Positively charged ion E) Negatively charged ion F) Compound composed of ions Short Answer. Give the name and symbol of the ion formed when the following neutral atom
experiences the stated behavior. 237. a sulfur atom gains two electrons Name: ____________________________________ Symbol: ___________________ 238. a strontium atom loses two electrons Name: ____________________________________ Symbol: ___________________ 239. a phosphorus atom gains three electrons Name: ____________________________________ Symbol: ___________________ 240. a potassium atom loses one electron Name: ____________________________________ Symbol: ___________________ 241. an aluminum atom loses three electrons Name: ____________________________________ Symbol: ___________________ 242. a chlorine atom gains one electron Name: ____________________________________ Symbol: ___________________
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Table Completion. Fill in the table using the information provided.
Name of ion/atom
Classification (check the correct
one) Symbol/ formula
Number of electrons
lost (if any)
Number of electrons gained (if
any)
Overall charge
(including + or – sign)
Fluoride ion
246. Cation
Anion
Neutral Atom
Molecule
252. 254. 260. 266.
Aluminum ion
247. Cation
Anion
Neutral Atom
Molecule
253. 255. 261. 267.
243. Cation
Anion
Neutral Atom
Molecule
Sr
256. 262. 268.
Carbon tetra-
chloride
248. Cation
Anion
Neutral Atom
Molecule
CCl4 0 0 0
244. 249. Cation
Anion
Neutral Atom
Molecule
Mg
257. 263. 269.
Lithium ion
250. Cation
Anion
Neutral Atom
Molecule
Li+
258. 264. 270.
245. 251. Cation
Anion
Neutral Atom
Molecule
O2-
259. 265. 271.
272. What types of elements tend to form ionic compounds?
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Section 6.2 Matching. Match the definition with the term that best correlates to it. No definition will be used more
than once. _____273. Chemical formula _____276. Law of Multiple Proportions _____274. Formula unit _____277. Molecular formula _____275. Law of Definite Proportions A) Chemical formula for a molecular compound B) Chemical formula for an ionic compound C) In any sample of a compound, the masses of the elements are always in the same proportions D) Shows the kinds and numbers of atoms in the smallest representative unit of a substance E) Whenever two elements form more than one compound, the different masses of one element that
combine with the same mass of the other element are in the ratio of small whole numbers. Short Answer. Answer the following questions. Classify each of the following chemical formulas as a molecular formula or a formula unit by circling the choice that best describes it. 278. Li2O: molecular formula formula unit 279. CO2: molecular formula formula unit 280. H2O: molecular formula formula unit 281. SrCl2: molecular formula formula unit 282. N2O: molecular formula formula unit Identify the number of each kins of atoms present in a molecule of each compound. 283. Ethylenediamine (H2NCH2CH2NH2): 284. Methyl ethyl ketone (CH3COC2H5): 285. Trichloroacetic acid (CCl3COOH):
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286. Mr. Geist’s major vice (C8H10N4O2): 287. One of the products of photosynthesis (C6H12O6): Section 6.3 Short Answer. Answer the following questions. Using only the periodic table, write the formula for the typical ion of each of the following representative elements. 288. Sodium: _________________ 294. Phosphorus: _____________ 289. Oxygen: _________________ 295. Sulfur: __________________ 290. Iodine: __________________ 296. Magnesium: _____________ 291. Lithium: _________________ 297. Calcium: ________________ 292. Cesium: _________________ 298. Potassium: ______________ 293. Aluminum: _______________ 299. Chlorine: ________________ Write the formula (including charge) for each ion. 300. Nitrite: ________________ 303. Permanganate: ______________ 301. Sulfate: ________________ 304. Hydroxide: __________________ 302. Acetate: _______________ 305. Dichromate: _________________ 306. What is a polyatomic ion? Are most of them negative or positive?
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Section 6.4 Short Answer. Answer the following questions. For polyatomic ions you were not asked to memorize,
refer to page 147 in your textbook or the back of your periodic table. Name the following compounds. 307. PbCl2
Name: _______________________________________________________________ 308. Al2(CO3)3
Name: _______________________________________________________________ 309. KMnO4
Name: _______________________________________________________________ 310. ZnCl2
Name: _______________________________________________________________ 311. Fe(NO3)3
Name: _______________________________________________________________ 312. MnO2
Name: _______________________________________________________________ 313. NH4NO2 Name: _______________________________________________________________ 314. Sr3N2
Name: _______________________________________________________________ 315. Li2O
Name: _______________________________________________________________ 316. Cr(PO4)2
Name: _______________________________________________________________ Write the formula of the compound. 317. Aluminum sulfite
Formula: _______________________________________________________________
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318. Tin (IV) fluoride
Formula: _______________________________________________________________ 319. Potassium dichromate
Formula: _______________________________________________________________ 320. Cobalt (III) sulfate
Formula: _______________________________________________________________ 321. Ammonium acetate
Formula: _______________________________________________________________ 322. Lithium nitride
Formula: _______________________________________________________________ 323. Lithium nitrate Formula: _______________________________________________________________ 324. Barium chlorate
Formula: _______________________________________________________________ 325. Manganese (IV) phosphide
Formula: _______________________________________________________________ 326. Manganese (IV) phosphate
Formula: _______________________________________________________________ 327. When are parentheses used in writing a chemical formula?
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Table Completion. Complete the following table by writing in the chemical formulas next to “F:” and names for the compounds formed by combining the indicated positive and negative ions next to “N:”.
NH4
+ Al3+ K+ Cr6+
F–
F: (328)
F: (330)
F: (332)
F: (334)
N: (329)
N: (331)
N: (333)
N: (335)
SO42–
F: (336)
F: (338)
F: (340)
F: (342)
N: (337)
N: (339)
N: (341)
N: (343)
NH4
+ Al3+ K+ Cr6+
PO33–
F: (344)
F: (346)
F: (348)
F: (350)
N: (345)
N: (347)
N: (349)
N: (351)
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NH4
+ Al3+ K+ Cr6+
CN–
F: (352)
F: (354)
F: (356)
F: (358)
N: (353)
N: (355)
N: (357)
N: (359)
Section 6.5 Short Answer. Answer the following questions. Name the following compounds. 360. O2F2
Name: _______________________________________________________________ 361. B2O3
Name: _______________________________________________________________ 362. AsF5
Name: _______________________________________________________________ 363. SF6
Name: _______________________________________________________________ 364. P2O5
Name: _______________________________________________________________ 365. P4O10
Name: _______________________________________________________________ 366. HNO3 (acid) Name: _______________________________________________________________ 367. CO2
Name: _______________________________________________________________
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368. H2SO4 (acid)
Name: _______________________________________________________________ 369. Cl2O
Name: _______________________________________________________________ Write the formula of the compound. 370. Diphosphorus pentasulfide
Formula: _______________________________________________________________ 371. Iodine trichloride
Formula: _______________________________________________________________ 372. Hydrobromic acid
Formula: _______________________________________________________________ 373. Dinitrogen difluoride
Formula: _______________________________________________________________ 374. Xenon difluoride
Formula: _______________________________________________________________ 375. Acetic acid
Formula: _______________________________________________________________ 376. Carbon tetrachloride Formula: _______________________________________________________________ 377. Nitrogen dioxide
Formula: _______________________________________________________________ 378. Dichlorine heptoxide
Formula: _______________________________________________________________ 379. Carbon monoxide
Formula: _______________________________________________________________
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Section 6.6 Short Answer. Answer the following questions. HINT: For polyatomic ions, refer to the back of your
periodic table. Name the following compounds. 380. Ga2O3
Name: _______________________________________________________________ 381. S2Cl2
Name: _______________________________________________________________ 382. Cd(NO3)2
Name: _______________________________________________________________ 383. VF5
Name: _______________________________________________________________ 384. S4N4
Name: _______________________________________________________________ 385. SnO2 Name: _______________________________________________________________ 386. N2O Name: _______________________________________________________________ 387. (NH4)2CO3
Name: _______________________________________________________________ 388. SF4
Name: _______________________________________________________________ 389. Li3PO3
Name: _______________________________________________________________ Write the formula of the compound. 390. Ammonium phosphate
Formula: _______________________________________________________________
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391. Chromium (VI) oxide
Formula: _______________________________________________________________ 392. Potassium hypochlorite
Formula: _______________________________________________________________ 393. Sulfur difluoride
Formula: _______________________________________________________________ 394. Silver phosphate
Formula: _______________________________________________________________ 395. Iron (III) carbonate
Formula: _______________________________________________________________ 396. Phosphoric acid Formula: _______________________________________________________________ 397. Iron (III) hydroxide
Formula: _______________________________________________________________ 398. Hydrochloric acid
Formula: _______________________________________________________________ 399. Sodium fluoride
Formula: _______________________________________________________________ 400. Silicon dioxide
Formula: _______________________________________________________________ 401. Magnesium hydrogen carbonate
Formula: _______________________________________________________________
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Unit Three Worksheet WS – C – U3
Matching. Match the definition with the term that best correlates to it. No definition will be used more
than once. _____402. Gram formula mass (gfm) _____405. Representative particle _____403. Avogadro’s number _____406. Gram molecular mass (gmm) _____404. Gram atomic mass (gam) _____407. Mole A) Amount of a substance that contains 6.022 1023 representative particles of the substance B) Mass, in grams, of one mole of atoms in an element C) Mass, in grams, of one mole of a molecular compound D) Mass, in grams, of one mole of an ionic compound E) An atom, formula unit, ion, or molecule is each an example of this F) 6.022 1023 representative particles of a substance Short Answer. Answer the following questions. Find the gram formula mass or gram molecular mass of each compound. Show work or receive no credit. Include correct significant figures and/or decimal places as well as correct units. 408. Fe(OH)2 (iron (II) hydroxide) 410. C2H5OC2H5 (diethyl ether) 409. (NH4)3PO3 (ammonium phosphite) 411. Li3PO4 (lithium phosphate) How many oxygen atoms are in a representative particle of each substance? _______412. C3H5(NO3)3 (nitroglycerin) _______414. Cr(OH)3 (chromium (III) hydroxide) _______413. (NH4)3PO3 (ammonium phophite) _______415. C8H8O4 (acetylsalicylic acid) How many moles is each of the following? Show work or receive no credit. Include correct significant figures and/or decimal places as well as correct units. 416. 6.022 1027 molecules NO2
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417. 1 trillion (1 1012) molecules Cl2 418. 3.011 1024 molecules O2 419. 9.05 1028 atoms W Matching. Match the definition with the term that best correlates to it. No definition will be used more
than once. _____420. Molar mass _____422. Standard temperature and pressure (STP) _____421. Molar volume A) The mass of one mole of a substance C) 0C and 1 atm B) 22.4 L at STP Calculation. Answer the following questions. Show work or receive no credit. Include proper units
and correct significant figures and/or decimal places. 423. The volume, in liters, of 72.0 g CO2 at standard temperature and pressure 424. The mass, in grams, of 3.28 mol (NH4)2SO4
425. The mass, in grams, of 25.0 L C3H8 at standard temperature and pressure 426. The number of atoms in 125.0 g LiC2H3O2
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427. The number of formula units in 58.3 g Li3PO4 428. The volume, in liters, of 5.00 1028 molecules SO2 at standard temperature and pressure 429. The volume, in liters, of 14.8 mol C2H6, at standard temperature and pressure 430. The mass, in grams, of a molecule of propylene glycol (CH3CHOHCH2OH) 431. The density of CH4 at standard temperature and pressure 432. The density of NO2 at standard temperature and pressure 433. Number of formula units of 202.1 g Cr(C2H3O2)3
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434. Moles of 937 g Al2(CrO4)3 435. The densities of gases A, B, C, and D are 0.17869 g/L, 0.71621 g/L, 1.42848 g/L, and 1.96469
g/L, respectively, at standard temperature and pressure. Calculate the molar mass of each substance (in other words, find the molar mass of gas A, gas B, gas C, and gas D). Then identify each substance as ammonia (NH3), carbon dioxide (CO2), oxygen (O2), helium (He), ethane (C2H6), carbon monoxide (CO), or methane (CH4). Show work or receive no credit.
Short Answer. Answer the following questions. 436. Would four balloons, each containing the same number of molecules of a different gas at
standard temperature and pressure, have the same mass or the same volume? Explain. 437. How can you determine the molar mass of a gaseous compound if you do not know its molecular
formula? 438. Why might the term “molar mass” be used instead of gram molecular mass, gram formula mass,
or gram atomic mass?
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Unit Four Worksheet WS – C – U4
Matching. Match the definition with the term that best correlates to it. _______439. chemical equation _______442. skeleton equation _______440. catalyst _______443. balanced equation _______441. coefficients A) A chemical equation in which mass is conserved and each side of the equation has the same
number of atoms of each element B) A chemical equation that does not indicate the relative amounts of reactants and products C) A substance that causes a reaction to occur or speed up without being used D) A small whole number that appears in front of a formula in a balanced chemical equation E) An expression representing a chemical reaction in which reactants are on the left and products are
on the right Short Answer [Writing]. Balance the following chemical reactions by filling in the coefficients as
needed. If no coefficient is needed, write “1” in the blank. Then identify the general reaction type and specific reaction type by checking the corresponding box.
444. Fe + O2 FeO Balanced equation: _______Fe + _______O2 _______FeO General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 445. N2O5 + H2O HNO3 Balanced equation: _______N2O5 + _______H2O _______HNO3 General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 446. P + O2 P2O5 Balanced equation: _______P + _______O2 _______P2O5 General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base
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447. Fe(OH)3 Fe2O3 + H2O Balanced equation: _______Fe(OH)3 _______Fe2O3 + _______H2O General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 448. NaOH + HCl NaCl + H2O Balanced equation: _______NaOH + _______HCl _______NaCl + _______H2O General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 449. NiCO3 NiO + CO2 Balanced equation: _______NiCO3 _______NiO + _______CO2 General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 450. NH4NO3 N2O + H2O Balanced equation: _______ NH4NO3 _______N2O + _______H2O General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 451. Al + H2SO4 Al2(SO4)3 +H2
Balanced equation: _______Al + _______ H2SO4 _______ Al2(SO4)3 + _______H2 General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base
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452. Zn + AgNO3 Zn(NO3)2 + Ag Balanced equation: _______Zn + _______ AgNO3 _______ Zn(NO3)2 + _______Ag General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 453. H2 + Fe3O4 Fe + H2O Balanced equation: _______H2 + _______Fe3O4 _______Fe + _______H2O General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base Short Answer [Writing]. Write skeleton equations representing the following reactions and then
balance them. Include all needed symbols for states of matter and catalysts. Then identify the general reaction type and specific reaction type by checking the corresponding box.
454. Pure copper metal can be produced by heating solid copper (II) sulfide in the presence of oxygen
gas from the air. Sulfur dioxide gas is also produced in this reaction. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 455. Solid iron (III) oxide and hydrogen gas react to produce iron metal and liquid water. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base
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456. When nitric acid is poured into magnesium hydroxide solution, a reaction occurs in which magnesium nitrate and water are produced.
Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 457. When barium metal is dropped into hydrochloric acid, barium chloride is created with hydrogen
gas being given off. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 458. Sulfur dioxide gas reacts with oxygen gas to produce sulfur trioxide gas. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 459. Aluminum metal will react with copper (II) sulfate solution to produce copper metal and aluminum
sulfate solution. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base
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460. Solid silver oxide is heated to produce solid silver and oxygen gas. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 461. Aqueous sodium chloride and sulfuric acid react to yield sodium sulfate and hydrochloric acid,
both of which are aqueous. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 462. Solid phosphorus reacts with oxygen gas to produce solid diphosphorus pentoxide. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base 463. Sulfuric acid reacts with aqueous sodium hydroxide to produce sodium sulfate and water. Skeleton equation: ____________________________________________________ Balanced equation: ____________________________________________________ General reaction type: Acid-base Oxidation-reduction Precipitation Specific reaction type: Synthesis Single-replacement Combustion Decomposition Double-replacement Acid-base
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Short Answer [Writing]. Write a balanced chemical equation for the following problems by predicting the correct products, writing the equation with proper formulas and symbols for each of the following reactions, and including states of matter. If the reaction is not possible, circle the “Not Possible” phrase below the blank.
464. Potassium metal reacts with chlorine gas to produce ... Balanced equation: ___________________________________________________________ Not possible 465. Aqueous solutions of aluminum chloride and sodium carbonate react to produce ... Balanced equation: ___________________________________________________________ Not possible 466. Metallic magnesium reacts with aqueous zinc sulfate to produce … Balanced equation: ___________________________________________________________ Not possible 467. Aqueous solutions of ammonium sulfate and barium chloride react to produce … Balanced equation: ___________________________________________________________ Not possible 468. Metallic silver reacts with aqueous sodium nitrate to produce … Balanced equation: ___________________________________________________________ Not possible
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Short Answer. Answer the following questions. 469. How is the law of conservation of mass related to the balancing of a chemical equation? 470. What is the purpose of a catalyst? 471. The equation for the formation of water from its elements, H2 (g) + O2 (g) H2O (l), can easily be
“balanced” by changing the formula of the product to H2O2. Explain why this is incorrect. 472. How do you predict the correct formula for the combination reaction between a Group A metal
and a nonmetal? 473. Explain why the following is true: 2Na + 2HCl 2NaCl + H2 Ag + HCl No reaction For the following chemical reactions, identify A) the element being oxidized, B) the element being reduced, C) the oxidizing agent, and D) the reducing agent. 474. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) A) Oxidized element: _________ C) Oxidizing agent: ____________________ B) Reduced element: _________ D) Reducing agent: ____________________ 475. N2(g) + 3H2(g) 2NH3(g) A) Oxidized element: _________ C) Oxidizing agent: ____________________ B) Reduced element: _________ D) Reducing agent: ____________________
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476. Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g) A) Oxidized element: _________ C) Oxidizing agent: ____________________ B) Reduced element: _________ D) Reducing agent: ____________________ 477. 2AgNO3(aq) + Cu(s) Cu(NO3)2(aq) + 2Ag(s) A) Oxidized element: _________ C) Oxidizing agent: ____________________ B) Reduced element: _________ D) Reducing agent: ____________________ Short Answer/Calculation. Answer the following questions. Show work or receive no credit.
Include proper units and correct significant figures and/or decimal places.
4NH3(g ) 5O2 (g ) 4NO(g ) 6H2O(g ) Refer to the above equation for questions 478 – 482. 478. How many moles of oxygen gas are required to react with 8 moles of ammonia (NH3)? 479. How many moles of water result from 7.5 moles of oxygen gas reacting with an excess amount of
ammonia? 480. If 3.5 L of oxygen gas at STP react with an excess amount of ammonia, how many liters of
nitrogen monoxide will be produced? 481. If 3.5 L of oxygen gas at STP react with an excess amount of ammonia, how many grams of
nitrogen monoxide will be produced?
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482. If 15 grams of ammonia (NH3) react with an excess amount of oxygen gas, how many grams of nitrogen monoxide will result?
16HCl 2KMnO4 2KCl 2MnCl2 5Cl2 8H2O Refer to the equation above for questions 483 – 486. 483. If 110 moles of hydrochloric acid (HCl) are used in the reaction, how many moles of chlorine gas
will be produced? 484. If 110 grams of hydrochloric acid (HCl) are used in the reaction, how many grams of chlorine gas
will be produced? 485. If 110 grams of hydrochloric acid (HCl) are used in the reaction, how many moles of water will be
produced? 486. How many moles of hydrochloric acid (HCl) are needed to react with 4 moles of potassium
permanganate (KMnO4)?
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487. In a chemical reaction, what two things are conserved? BOOK PROBLEMS: Section review 9.2, page 250, #19 – 22 (all parts) Short Answer/Calculation. Answer the following questions. Show work or receive no credit for ALL
problems involving calculations. Include proper units and correct significant figures and/or decimal places.
For questions 488 – 491, refer to the following scenario. Scenario: A 500.0 g sample of aluminum sulfate is made to react with 450.0 g of calcium hydroxide. A
total of 596 grams of calcium sulfate is produced. 488. Write a balanced equation for this reaction. HINT: You will need to predict the other product. 489. What is the limiting reagent in this reaction? 490. How many grams of excess reagent are unreacted? 491. What is the percent yield of calcium sulfate?
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For questions 492 – 495, refer to the following scenario. Scenario: Plants often consume more water than they actually need to react with the carbon dioxide
for photosynthesis or attain more carbon dioxide than they have to react with water. In a particular plant, 4.50 grams of water reacts with 10.0 grams of carbon dioxide.
Chemical equation: 2612622 O6OHCOH6CO6 492. What is the limiting reagent? Justify your answer.
493. How much excess reagent is there? 494. How many grams of glucose are produced from the reaction?
495. After carrying out the previously stated reaction, the plant actually only produces 1.00 g of
glucose. What is the percent yield of the glucose? 496. Merck is a leading pharmacological company that produces many different drugs. When
producing a certain medication, they find that the actual yield of a certain drug is 6.55 kg. After doing some calculations, Merck discovers that it only has an 85% yield of the drug. What was the theoretical mass of the drug Merck produced?
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For questions 497 – 501, refer to the following scenario. Scenario: A 3.1 mol sample of sulfur dioxide is made to react with a 2.7 mol sample of oxygen gas to
produce sulfur trioxide. 497. Write a balanced equation for this reaction. 498. What is the limiting reagent? Justify your answer. 499. How much excess reagent remains? 500. How many grams of sulfur trioxide are produced? 501. If the reaction takes place at standard temperature and pressure, how many liters of sulfur
trioxide are produced? For questions 502 – 506, refer to the following scenario. Scenario: 75.00 grams of zinc react with 120.0 grams of sulfuric acid. 502. Write a balanced equation for this reaction. HINT: You will need to predict the products.
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503. What is the limiting reagent? Justify your answer. 504. How much excess reagent remains? 505. How many grams of hydrogen gas are produced? 506. If only 1.05 grams of hydrogen gas are produced, what is the percent yield? BOOK PROBLEMS: Section review 9.3, page 259, #30, 32
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Unit One Experiment – 1 Bunsen Burner Operation
EX – C – U1 – 1
Introduction: The purpose of this experiment is to learn to operate a Bunsen burner and make careful observations regarding combustion. Background: One of the most useful and exciting aspects of chemistry is the use of the Bunsen burner. Based on the way it is operated, you can attain different temperatures, gas and air mixtures, and other necessary combinations for use in a laboratory setting. Safety: Safety goggles will be worn at all times. Procedure: 1. Light the Bunsen burner by holding a spark igniter next to the barrel of the burner and turning on the gas.
NOTES: The barrel should be opened half way. The gas is flowing to the burner where the handle is parallel to the spigot or nozzle of the gas outlet. There are two points of control on the burner: the barrel itself and a screw adjustment knob, known as
the “wheel”, at the base of the burner. Be careful in making adjustments of either the barrel or the wheel since turning these pieces too far will disassemble the burner. Additionally, be especially careful since making such adjustments while the burner is in operation can result in more critical accidents.
Turning the barrel adjusts the mixture of air with the gas. Turning the wheel adjusts the amount of gas that flows through the burner. Note that you should never
leave a burner turned off when the wheel is adjusted for a high flame as it poses a significant danger to the next person utilizing the burner.
2. Turn the barrel of the burner while it in operation. Observe the barrel openings near the base of the tube.
Record how turning the barrel controls the oxygen availability, and note the relationship between hole size and air availability.
3. Turn the barrel to decrease the flow of air. (Turn the barrel down, in other words.) Observe the color of the
flame. Record a description of the color and structure of the flame, and compare how this appears compared to a candle flame.
4. Use the wheel to adjust the flame to about 4 inches or 10 centimeters. Describe how adjustments of the
wheel raises or lowers the flame height. 5. The proportion of air to gas in the mixture determines the flame’s temperature. Lean mixtures (high air/low
gas) burn hot. Rich mixtures (low air/high gas) burn cool. The structure of the flame is a function of both the rate of flow and the richness of burning mixture.
When placing a glass item, such as a beaker, over a flame, soot will often form. If soot forms on the
beaker, then the combustion is incomplete, giving off carbon in the form of soot. If no soot forms on the beaker, then the combustion is complete, giving off carbon dioxide instead from carbon reacting with oxygen gas as well as also producing water.
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Using a low air supply, hold a 250-mL beaker of water with beaker tongs over the flame and observe the beaker’s bottom. Record any observations regarding the appearance of the bottom of the beaker. Which element, if any, is in your observation? Clean the outside of the beaker following your observations.
6. Turn the barrel upward to increase the air supply. You may also need to adjust the gas supply, using the
wheel, if the flame is either too low or too high. Adjust the air supply so that the flame has an inner blue cone and the burner makes a low roaring sound. Record your observations of the flame structure as well as a labeled diagram of what you observe.
7. Fill the cleaned 250-mL beaker with ¾ cold tap water. Using the beaker tongs, hold the beaker over the
burner and watch for condensation (dew). Turn the burner off. Record your speculated source of the condensed water, recalling that lowering the temperature usually precedes dew formation on the exterior of the beaker wall. Cite evidence related to your speculation.
8. Turn off the burner and put away all equipment.
Figure 1 Observation/Data Tables:
Procedure Step Recorded observations
Step 2
Step 3
Step 4
Step 5
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Procedure
Step Recorded observations
Step 6
Step 6 Diagram
Step 7
Conclusion/Discussion: Compare AND contrast the burning of a candle and the burning of a Bunsen burner. Do so in three different ways per box.
Compare
Contrast
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Unit One Experiment – 2 Mass and Change
EX – C – U1 – 2
Introduction: The purpose of this experiment is to learn about changes in mass based on physical and chemical changes that occur. Background: Although the appearance of matter can be different, it is often questionable as to whether the mass of the matter changes, particularly when burning something or reacting something under a set of circumstances. This experiment will help you to determine how mass can change, if at all. Safety: Safety goggles will be worn at all times. Procedure: Part 1 – Mass of Steel Wool 1. Determine the mass of the wad of steel wool provided to you using the triple-beam balance. Record the
mass using the correct number of decimal places. 2. Carefully pull the wad apart so that it occupies a volume roughly twice as great as before. Then determine
the mass of the expanded wad of steel wool. Record the mass using the correct number of decimal places. 3. Calculate the change in mass by subtracting the expanded wool mass by the contracted wool mass.
Record this mass using the correct number of decimal places. Part 2 – Mass of Ice and Water 1. Find the mass of the vial and a small piece of ice using the triple-beam balance. Record the mass using
the correct number of decimal places. 2. Because the ice takes a while to melt, set the vial aside and go on to part 3 rather than wait for the process.
Periodically warm the vial in your hands to speed up the process. Once melted, find the mass of the vial and the water using the triple-beam balance. Record the mass using the correct number of decimal places.
3. Calculate the change in mass by subtracting the mass of the vial and water by the mass of the vial and ice. Record the mass using the correct number of decimal places.
4. Wash the vials in soapy water, rinse with tap water, and then rinse with distilled water. Part 3 – Mass of a Precipitate 1. Fill a clean, empty vial with no more than 1/3 full of the calcium nitrate (Ca(NO3)2) solution. Then fill
another clean, empty vial with no more than 1/3 full of the sodium carbonate (Na2CO3) solution. Cap the vials and find the mass of both vials together using the triple-beam balance. Record the mass using the correct number of decimal places.
2. Carefully pour the contents of one vial into the other, and cap the non-empty vial. Put both vials and caps back on the balance pan. Find the mass of both vials and caps together using the triple-beam balance. Record the mass using the correct number of decimal places.
3. Calculate the change in mass by subtracting the mass of the combined solutions (vials and caps) by the mass of the individual solutions (vials and caps). Record the mass using the correct number of decimal places.
4. Pour the solution and precipitate into the waste bottle provided. Wash the vials in soapy water, rinse with tap water, and then rinse with distilled water.
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Part 4 – Mass of Burning Steel Wool 1. Find the mass of the steel wool provided using the triple-beam balance. Record the mass using the correct
number of decimal places. 2. Set up a Bunsen burner and light the burner. Hold the steel wool by using crucible tongs over an
evaporating dish, and heat the steel wool until it glows. Turn the steel wool around in the flame so that all sides are exposed. Any pieces of the steel wool that break free during heating should fall into the dish and then be transferred to the triple-beam balance. After the wool has been heated for awhile, allow the wool to cool. Find the mass of the steel wool provided using the triple-beam balance. Record the mass using the correct number of decimal places, and record any observations about the heated wool compared to the original wool.
3. Calculate the change in mass by subtracting the mass of the heated wool by the mass of the original wool. Record the mass using the correct number of decimal places.
4. Discard the cooled wool per your teacher’s instructions. Part 5 – Mass of Dissolved Alka-Seltzer 1. Fill a vial about 1/2 full of water. Put about 1/4 of a tablet of Alka-Seltzer in the cap of the vial. Place the
vial, water, cap and Alka-Seltzer on the pan of the triple-beam balance. Record the mass using the correct number of decimal places.
2. Put the piece of Alka-Seltzer into the vial, and loosely cap the vial. When the piece of tablet has completely dissolved, find the mass of the vial and contents again. Record the mass using the correct number of decimal places.
3. Calculate the change in mass by subtracting the mass of the dissolved Alka-Seltzer solution by the mass of the original water and separate Alka-Seltzer. Record the mass using the correct number of decimal places.
4. Pour the solution and precipitate into the waste bottle provided. Wash the vials in soapy water, rinse with tap water, and then rinse with distilled water.
Observation/Data Tables: Part 1 – Mass of Steel Wool
Expanded steel mass
– Contracted steel mass
Difference in steel mass
Part 2 – Mass of Ice and Water
Mass of vial and water
– Mass of vial and ice
Difference in mass
Part 3 – Mass of a Precipitate
Mass of combined solutions (vials and caps)
Mass of individual solutions (vials and
caps)
Difference in mass
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Part 4 – Mass of Burning Steel Wool
Mass of heated wool
– Mass of original wool
Difference in mass
Observations of heated wool: ________________________________________________________________ ________________________________________________________________ ________________________________________________________________ Part 5 – Mass of Dissolved Alka-Seltzer
Mass of dissolved solutions (vial and cap)
Mass of water and separate Alka-Seltzer (vial and cap)
Difference in mass
Discussion/Conclusion Questions: The histogram is a way to represent the class results. The only real difficulty with the use of this tool is in introducing the idea of “bins” to store the results. Histograms are often used by teachers in a program they use (Easy Grade Pro) to see the number of students who received grades in certain ranges, and shown below.
Let us now create a classroom set of histograms regarding the data. Part 1 – Mass of Steel Wool
-0.05 -0.03 -0.01 +0.01 +0.03 +0.050
Change in mass (g)
60 70 80 90
5
15
10
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Part 2 – Mass of Ice and Water
-0.05 -0.03 -0.01 +0.01 +0.03 +0.050
Change in mass (g) Part 3 – Mass of a Precipitate
-0.05 -0.03 -0.01 +0.01 +0.03 +0.050
Change in mass (g) Part 4 – Mass of Burning Steel Wool
-0.05 -0.03 -0.01 +0.01 +0.03 +0.050
Change in mass (g) Part 5 – Mass of Dissolved Alka-Seltzer
-0.05 -0.03 -0.01 +0.01 +0.03 +0.050
Change in mass (g)
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Unit One Experiment – 3 Density and Relationships
EX – C – U1 – 3
Introduction: The purpose of this experiment is to determine the densities of unknown metals. Background: Have you ever noticed that some pipes tend to weigh a lot more than others, even if they are smaller? This aspect is strongly related to a relationship expressed by the physical property called density. Density is defined as the ratio of a substance’s mass to the volume it occupies. The formula for density is as follows:
(mL) Volume
(g) MassDensity
In this experiment, you will measure the mass and volume of three unknown metals. You will then use your data to explore the relationship between the mass and volume of the metals to determine their respective densities. Safety: Safety goggles will be worn at all times. During this experiment, no open-toe shoes are to be worn. Procedure: As you perform this experiment, record your data in Data Tables 1 and 2. 1. Determine the mass of the three different unknown metal samples to the nearest 0.01 gram using the
centigram balance. Record the masses of the metal samples into Data Tables 1 and 2. 2. Find the volume of each metal sample by water displacement. Fill one of the provided graduated cylinders
about half-full with water, measure the volume, and record as “volume of water alone” in Data Table 1. Tilt the graduated cylinder and carefully slide one of the metal samples down the side. Make sure the metal sample is completely submerged in the water. Measure the volume and record the measurement as “volume of water + metal” in Data Table 1.
3. Repeat Step 2 using the other metal samples. Dry all samples and then return them to their respective
places. 4. Compute the volume of each metal sample using data from Data Table 1. Compare the density of each
metal sample, showing your work (including units), in Data Table 1. 5. Complete Data Table 2 by recording the mass and volume data collected by your classmates and yourself.
Using the class data, plot a graph of mass versus volume. Represent the plotted points for each metals with a different symbol. Draw a “best fit” straight line through each group of plotted points. Determine the slope of each of the lines in the graph. Record the slope of each line and your method of calculation in Data Table 3. NOTE: The general equation for a line is y = mx + b where m is the slope and b is the value of the y-intercept. Be sure to use correct units for the slope.
Calculations:
(mL) Volume
(g) MassDensity
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100% value accepted
value alexperiment - value accepted errorpercent
Table of Reagents and Products:
Table of Metals Densities of Common Metals
Metal
Density (g/mL)
Aluminum 2.699 Brass 8.44
Chromium 7.19 Cobalt 8.90 Copper 8.96
Gold 19.32 Iron 7.87 Lead 11.34 Nickel 8.88
Platinum 21.4 Silver 10.491
Stainless steel 7.75 Tin 7.29
Titanium 4.5 Tungsten 19.3 Vanadium 6.11
Zinc 7.1 Observation/Data Tables:
Data Table 1: Individual Data and Calculations Metal A Metal B Metal C
Mass (in grams)
Volume of water alone (in milliliters)
Volume of water + metal (in milliliters)
Volume of metal (in milliliters)
Density of metal (in grams/milliliter)
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Data Table 2: Class Data – Mass and Volume of Metal Samples Metal A Metal B Metal C
Lab Pair
Mass (g)
Volume (mL)
Mass (g)
Volume (mL)
Mass (g)
Volume (mL)
1
2
3
4
5
6
7
8
9
10
11
12
Data Table 3: Density Calculations from Class Data (Slopes)
Metal A Metal B Metal C
mL
g
mL
g
x
y
mL
g
mL
g
x
y
mL
g
mL
g
x
y
Discussion/Conclusions: 1. What does the slope of each metal represent? HINT: Refer to Data Table 1.
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2. Looking at your graph, what does this experiment demonstrate about the density of a metal? What does it demonstrate about the densities of different metals?
3. Calculate the percent error in the density calculations for the two samples. Your teacher will provide you
with the accepted value for the density of each metal. (Show all relevant calculations.) 4. Calculate the percent error in the values of density obtained from the slopes of the lines in your graph.
(Show all relevant calculations.) 5. Look back at the percent errors calculated in problems 3 and 4. Generally, the slope of the line will give a
more accurate value of density than a single sample. Explain why this is true. 6. Can you identify a metal if you know its density? Explain your answer.
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7. Do you think that determining the volumes of your metal samples by measuring their dimensions and calculating would be more or less accurate than determining these volumes by water displacement? Explain. Would measuring the dimensions of a solid always be possible? Explain.
8. You originally want to use 20 grams of iron as a mass while fishing but decide that copper may be better
since it will not rust. What volume of copper will provide the same mass as the 20 grams of iron? (HINT: Use the density table provided in the laboratory exercise.)
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Unit One Experiment – 4 Measurement and Uncertainty
EX – C – U1 – 4
Introduction: The purpose of this experiment is to study the nature of measurement and gather data. Background: Everyone deals with measurements every day. We hear statements such as "The time at the tone is 10 p.m.", "It is currently 79 degrees and sunny," and "7.8 gallons of gas - That will be $23.24." The measured values in these three statements are printed in boldface type. Are these and other measurements always exact? An exact measurement is a perfectly correct value containing no error. Right now, before you begin this experiment, select the one statement below you think is most correct. A. Measurements are exact if correctly done. B. Measurements may or may not be exact. It depends who did them and how they were done. C. There is some inexactness in every measurement. Safety: Safety goggles will be worn at all times. During this experiment, no open-toe shoes are to be worn. Avoid skin contact with solids and solutions. Exercise caution and use proper techniques in handling hot materials safely. Dispose of all solutions in the container designated by your teacher. Wash your hands before leaving the laboratory. Procedure: Perform this laboratory activity with a partner. You and your partner will determine each mass value independently on two different balances. After you exchange data you will have four values for each mass. On a separate sheet of paper, prepare a data table to organize and record the 48 data values you will obtain. Read the following steps to decide how to organize and label your data table. NOTE: Your teacher will collect this table on a separate sheet of paper. 1. Label a clean, dry 250 mL beaker with your name and the letter A. Your partner should label a beaker
similarly with his/her name and the letter B. 2. Measure the mass of your beaker to the nearest 0.01 g on two different balances. Your partner should do
the same with the other beaker. Record each mass value and the balance number in your data table. 3. Exchange beakers. Repeat Step 2. Record your partner's data in your data table. You should have eight
mass values and eight balance numbers. Good data values for a measurement are consistent (very close to each other). Repeat measurements for any data values which are greatly different from the others.
4. Using a spatula, add four scoops of solid potassium sulfate, K2SO4(s), to Beaker A. One partner should use two different balances to measure the mass of Beaker A and its contents. The other partner should measure the mass of a piece of filter paper on two different balances. Record masses and balance numbers as before.
5. Exchange Beaker A and filter paper. Complete two mass measurements. Record and share data. You should now have 16 mass values and 16 balance numbers.
6. Use a graduated cylinder to add 18 mL (± 1 mL) of distilled water to Beaker A. Swirl gently for approximately three minutes.
7. Set up a funnel with filter paper above Beaker B as shown in the diagram below.
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8. Decant (pour) the liquid from Beaker A into the filter paper, collecting the filtered liquid in Beaker B. Do not
be concerned if some undissolved solid collects on the filter paper. 9. After filtering is complete, use forceps to remove the filter paper from the funnel. Place the filter paper in
Beaker A. Rinse the empty funnel with 5 mL (± 1 mL) of distilled water. 10. Place Beakers A and B on a hot plate adjusted to medium heat to evaporate water from both beakers.
Keep the heat setting low enough to prevent splashing, crackling, or charring. (Caution: Do not handle hot beakers with bare hands.) When all drops of moisture have evaporated from the sides of the beakers, remove them from the heat with beaker tongs. Allow each beaker to cool.
11. Measure the mass of each beaker on two different balances. Record data and exchange beakers as before.
12. Save Beakers A and B until you have completed your data analysis. Then dispose of the contents of the beakers as directed by your teacher.
13. Wash hands thoroughly before leaving the laboratory. Observation/Data Tables:
Data Table 1: Data Comparison
Measurement (g) Range
Best estimate of mass Uncertainty
Best estimate with
uncertainty
(1) Mass of empty Beaker A
(2) Mass of empty Beaker B
(3) Mass of filter paper
(4) Mass of Beaker A and dry potassium sulfate
(5) Mass of Beaker B and contents after heating
(6) Mass of Beaker A and contents after heating
Calculations: Complete the following calculations on a separate sheet of paper. 1. Calculate the best estimate of the mass of dry potassium sulfate in Beaker A. (Use (1) and (4) from the
table.) 2. Calculate the uncertainty in the mass of dry potassium sulfate in Beaker A. 3. Calculate the best estimate of the mass of potassium sulfate in Beaker B after heating. What is the
uncertainty in this value? 4. Calculate the best estimate and uncertainty of the mass of potassium sulfate in Beaker A after heating.
Remember to consider the mass of the filter paper.
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5. Calculate the best estimate and uncertainty of the total mass of potassium sulfate in Beakers A and B after
heating. 6. Use best estimate and uncertainty values in Calculations 2 through 5 to determine the range in mass of
each of the following:
Data Table 2: Data Comparison Calculated value Mass range
Mass of dry potassium sulfate (use calc. 2 and 3)
Total mass, in grams, of potassium sulfate in beakers A and B after heating (use calc. 6)
Discussion/Conclusions: Complete the following questions on a separate sheet of paper. 1. After all your work, do you know the exact mass of potassium sulfate you used in this activity? Explain. 2. Do the ranges of mass of potassium sulfate before and after heating overlap? 3. Is it possible that the mass of potassium sulfate after heating equals the mass of potassium sulfate before
heating? Explain. 4. The Law of Conservation of Matter states that matter is neither created nor destroyed during any physical
or chemical change. Are experimental results obtained by you and others consistent with this law? Explain.
5. Sally uses Balance 4 to find the mass of Beaker A and records a value of 67.15 g. Would Sally be correct
in accepting this as an exact value? Explain. 6. If the procedure had asked you to measure each mass only once instead of four times, would you know
more or less about the best estimate and uncertainty of your data? Explain. 7. Should the accepted value for the mass of potassium sulfate be expressed as a single value or as a best
estimate with uncertainty? Explain with your experimental results. 8. Based on experience and knowledge gained in this activity, would you still select the same statement you
chose in the Introduction? Explain. 9. What would you do differently if you repeated this laboratory activity? 10. Based on your experimental data, do 250 mL beakers have equal mass? Explain. 11. Each student in Mr. Geist's chemistry class measures the mass of the same pencil. Many different mass
values are reported by the students. Suggest three reasons why. (Hint: Consider procedure, equipment, and techniques.)
12. What mass of potassium sulfate dissolved in the water during your laboratory activity? Include best
estimate and uncertainty. 13. If tap water were used in place of distilled water, how would this affect your results?
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Unit Two Experiment – 1 Average Atomic Mass
EX – C – U2 – 1
Introduction: The purpose of this experiment is to determine the average atomic mass of a sample of an “element”. Background: If you look at the periodic table of elements, you will see decimals underneath the symbol of each element. These are referred to as average atomic masses, most of which you notice are not whole numbers. These are weighted averages of all isotopes of an element based on their relative abundance and individual atomic masses. For this experiment, you will learn how to determine the mass of each isotope, find the percent abundance of each isotope, and then calculate the average atomic mass. NOTE: Beanium is the name for the isotope (comprised of beans). Procedure: 1. Sort the Beanium sample into the different isotopes (by color). Diagram each isotope. Isotope #1 Isotope #3 Isotope #2 2. Pick one of the isotopes to be #1. Record the mass of all isotopes #1. 3. Count the number of atoms of isotope #1 and record in the data table. Verify this number by having your
lab partner count again. If you do not agree on the number, count them again together. 4. Calculate the average mass of one isotope#1 using the following formula and record: Average mass of isotope #1: Total mass of all isotope #1 / Number of atoms of isotope #1 When you are through with isotope #1, put it back into the zip-lock baggie. Be careful not to spill any
atoms on the floor! 5. Repeat steps 2 to 4 for isotopes #2 and #3. Be sure to record the mass of each isotope and the exact
number of each isotope. Record the average mass of each isotope. Be sure to return all isotopes to the zip-lock baggie while making sure not to spill any.
Observation/Data Tables: Isotope Total Mass Number of Atoms Average Mass
1
2
3
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TOTAL NUMBER OF ATOMS: _________________________ Conclusion/Discussion: 1. Calculate the percent abundance of each isotope.
Percent Abundance of Isotope #1: __________ % isotope #1 = (count of #1 isotope / count of ALL isotopes) X 100%
Percent Abundance of Isotope #2: __________
% isotope #2 = (count of #2 isotope / count of ALL isotopes) X 100% Percent Abundance of Isotope #3: __________
% isotope #3 = (count of #3 isotope / count of ALL isotopes) X 100% 2. Calculate the average atomic mass of Beanium.
AVERAGE ATOMIC MASS OF BEANIUM: __________ amu Average atomic mass = (% abundance of isotope #1 average mass of isotope #1)
+ (% abundance of isotope #2 average mass of isotope #2) + (% abundance of isotope #3 average mass of isotope #3)
3. Why isn’t the atomic mass of most of the elements on the periodic table of element an integer (not a
decimal)? 4. If the heaviest isotope was more abundant, and the other two isotopes were less abundant, what would
happen to the atomic weight of Beanium? Why?
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Unit Two Experiment – 2 Flame Test Analysis
EX – C – U2 – 2 Introduction: The purpose of this simulation is to identify cations based on flame tests. Background: Detectives in mystery novels often rush evidence from the crime scene to the lab for analysis. One common type of analysis is referred to as a qualitative analysis, a form of analysis where identification of an unknown substance is done and where quantity is not necessarily important. By conducting a qualitative analysis whereby cations (positively-charged ions) are exposed to chemical tests and results are compared to the results given by known cations, lab technicians can specifically identify the presence of cations. Another practical application of flame tests deals with fireworks. If you are a firework aficionado (or simply a pyromaniac), you will notice when lighting off fireworks that many have different colors. Many of those will be seen in this experiment. In this laboratory experiment, you will identify colors of flames for specific ions and infer relationships between ions. Data Table/Observations:
Cations and Associated Colors Ion Chemical Observations of Color
Na+ (sodium ion)
Na2CO3 (sodium carbonate)
K+ (potassium ion)
KCl (potassium chloride)
Cu2+ (copper (II) ion)
CuSO4 (copper (II) sulfate)
Na+ (sodium ion)
Na2B4O7 (sodium borate)
Li+ (lithium ion)
LiCl (lithium chloride)
Sr2+ (strontium ion)
SrCl2 (strontium chloride)
Cu2+ (copper (II) ion)
CuCl2 (copper (II) chloride)
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Cations and Associated Colors
Ion Chemical Observations of Color
Fe3+ (iron (III) ion)
Fe(NO3)3 (iron (III) nitrate)
Fe3+ (iron (III) ion)
FeCl3 (iron (III) chloride)
Ba2+ (barium ion)
BaCl2 (barium chloride)
Conclusion/Discussion: 1. Did any differences exist between the copper (II) ion (Cu2+) tests? Explain. 2. Did any differences exist between the sodium ion (Na+) tests? Explain. 3. It is possible to get a false-positive or a false-negative result when testing for ions. Propose a situation that
could lead to a false positive for a particular ion. Choose a different ion and show how a false negative could result. Which do you think is more likely to happen – a false-positive or a false-negative result? Explain your reasoning.
4. Why is it necessary to know the color of the burning splint alone before testing the burning splint with the
“soaked” chemicals?
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Unit Two Experiment – 3 Cation and Anion Analysis
EX – C – U2 – 3
Introduction: The purpose of this experiment is to determine what cations and anions are present in a sample. Background: When bomb experts and forensic scientists test a site that has been bombed, often they will do residue analyses in order to determine what kind of bomb was detonated. They will analyze chemicals in soil, observe heat and temperature impacts, and other assorted components of the site in order to determine the chemical makeup of the bomb. The nitrate (NO3
-) ion is common in most C-4 explosives such as RDX, TNT, and others, but, as was the case with the Oklahoma City bombing, the ammonium (NH4
+) ion was also present. Your goal for this experiment will be to determine what ions are present in the sample you have been given. Safety: Safety goggles will be worn at all times. During this experiment, no open-toe shoes are to be worn. Procedure: Test 1: For CO3
2- ion Fit a flask with a one hole stopper and a bent glass tube as shown on below.
Add 2 mL of the sample solution. Stopper immediately after pouring in 2 mL of 6M HCl. Be sure the bent tube is below the surface of the limewater (saturated Ca(OH)2 solution) which should be clear (one may have to filter if it is cloudy). Heat the flask gently to the boiling point to drive CO2 over. A white precipitate of calcium carbonate (CaCO3) in the limewater indicated the carbonate ion in the sample.
CO32 2H H2OCO2(g)
CO2 Ca2 2OH CaCO3 (s) H2O
The only other gas that could be produced would be H2S, but this will not interfere as CaS is soluble.
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Test 2: For NO3- ion and NH4
+ ion Add 1 mL of the sample solution to a test tube. Add 2 mL of 6M NaOH. Warm and test for ammonia (NH3) vapor.
NH4 OH H2ONH3(g )
Ammonia vapor will turn moist red litmus paper to blue. Be sure to hold the paper above the test tube so that it does not come into contact with the NaOH. If ammonia is present, then the ammonium ion is also present. Also, if ammonia is present, carefully boil the solution until no change is produced with a new piece of moist litmus paper. The boiling will remove the ammonia which prevents testing for NO3
-.
Add about a quarter-inch of metallic zinc to the test tube and place a loose cotton wad (to prevent splattering of the NaOH solution) into the test tube. Warm and test for the presence of ammonia vapor while being careful not to get NaOH solution on the test paper. The zinc reduces the NO3
- to NH3 and reacts with NaOH to produce hydrogen gas. However, the hydrogen gas will not bother the ammonia vapor test. It takes several minutes of heating before this test appears.
NO3 7OH 4Zn 6H2O 4Zn(OH)4
2 NH3(g)
Zn 2OH 2H2O Zn(OH)42 H2 (g )
Observation/Data Tables:
Test
Cation/Anion
Observations
Is the anion/cation present?
Test 1 Carbonate (CO32-)
Test 2 Ammonium (NH4
+)
Nitrate (NO3-)
Conclusion/Discussion: 1. What anions/cations were present in the sample solution? 2. The nitrate ion is present in most explosives, including ammonium nitrate, TNT, RDX, PETN, nitroglycerine,
and a number of others. What property might one assume to be associated with the nitrate ion? 3. This laboratory experiment was one that involved a qualitative analysis. How does this differ from a
quantitative analysis? (You may need to take a look in a dictionary or on the Internet to answer this.) 4. Why do the equations and reactions listed in this experiment have charges associated inside them?
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Unit Four Experiment – 1 Precipitation Reactions
EX – C – U4 – 1 Introduction: The purpose of this experiment is to observe, identify, and write balanced equations for precipitation reactions. Background: The majority of ionic solid are soluble in water. Those that are not account for the formation of an insoluble salt called a precipitate. The formation of a precipitate is predicted by using the general rules for solubility of ionic compounds. Ionic compounds are made up of positive ions (cations) and negative ions (anions) held together by the attractive, electrostatic forces between the oppositely-charged particles. When soluble ionic compounds are placed in water, they break apart to give separate ions, a process known as dissociation. When two ionic solutions are combined, the resulting mixture contains positive and negative ions from each solution. The mixing allow new combinations of ions and, if one or more of these new ion combinations happens to be insoluble in water, it falls out of solution as a solid compound. The insoluble product formed in this way is called a precipitate. As an example of this, think of the addition of sodium chloride to a solution of silver nitrate.
Sodium chloride: NaCl Na+(aq) + Cl–(aq) Silver nitrate: AgNO3 Ag+(aq) + NO3
– (aq)
Na+(aq) + Cl–(aq) + Ag+(aq) + NO3– (aq) Na+(aq) + NO3
– (aq) + AgCl(s) The equation above is called a complete ionic equation. It shows the ions that are present, even if some are not involved in the formation of the precipitate. The ions that do not create the precipitate are referred to as spectator ions. An equation that doesn’t show those spectator ions is called a net ionic equation. The net ionic equation for the reaction, then, is as follows:
Ag+(aq) + Cl– (aq) AgCl(s) In this experiment, you will mix six different ionic solutions in all possible combinations of two to determine which combinations result in precipitates being formed. Based on your results, you will write complete equations for each reaction. Safety: Safety goggles will be worn at all times. During this experiment, no open-toe shoes are to be worn. Procedure: 1. Mix each pair of solutions within each set on a separate spot plate depression using no more than two
drops of each solution. Be careful not to contaminate the dropper from one bottle with a different solution. Simply shake the plate lightly to mix the solutions.
2. Create the following charts for your data section to keep track of your precipitates. Place a “ppt” in the box
if a precipitate forms. Write “no rxn” in the box if no reaction is observed.
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Solution Set One
Chemical Al2(SO4)3 MgCl2 Na2SO4 Mg(NO3)2 AlCl3 Ba(NO3)2
Al2(SO4)3
MgCl2
Na2SO4
Mg(NO3)2
AlCl3
Ba(NO3)2
Solution Set Two
Chemical KCl MgCl2 Na2SO4 NaOH BaCl2 MgSO4
KCl
MgCl2
Na2SO4
NaOH
BaCl2
MgSO4
3. Observe each mixture carefully for signs of a precipitate Since many precipitates are light in color and
difficult to notice, you may wish to vary the color of the background behind the plate with different colors of paper. Record evidence of any precipitate in your data tables.
4. Use distilled water to rinse the plate into the spent chemical container. 5. After completely reacting the chemicals in set one and set two, clean your work area and wash your hands
thoroughly before leaving the laboratory. Conclusion/Discussion: 1. For each combination of solutions that gave a precipitate, write correct formulas, not equations, for the two
new compounds that could form with the ions present. (HINT: Remember to balanced charges.) Do so in the following tables. IF NO NEW PRODUCTS ARE FORMED, PUT AN “X” IN THE BOX. ALSO INCLUDE STATE OF MATTER (i.e., (s), (aq), etc.)
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Solution Set One
Chemical Al2(SO4)3 MgCl2 Na2SO4 Mg(NO3)2 AlCl3 Ba(NO3)2
Al2(SO4)3
MgCl2
Na2SO4
Mg(NO3)2
AlCl3
Ba(NO3)2
Solution Set Two
Chemical KCl MgCl2 Na2SO4 NaOH BaCl2 MgSO4
KCl
MgCl2
Na2SO4
NaOH
BaCl2
MgSO4
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Unit Four Experiment – 2 Pipet Rockets and Synthesis
EX – C – U4 – 2
Introduction: The purpose of this experiment is to examine the relationship between hydrogen and oxygen gases generated in a synthesis reaction and correlate the optimum proportion for launching a pipet rocket. Background: Hydrogen and oxygen gases react with each other in a very quick, exothermic manner. The explosiveness of this reaction is greatest when the hydrogen and oxygen gases are mixed in just the right proportion. In this experiment, you are generating hydrogen and oxygen gases and testing their explosiveness. Your final goal will be to find the most powerful mixture and then use it to launch a pipet rocket as far as you can. Safety:
Safety goggles will be worn at all times. No open-toe shoes are to be worn. This experiment involves mini-explosions. FOLLOW ALL TEACHER GUIDELINES OF TESTING AND
PREPARATION SAFETY!!! Dispose of all materials according to the instructions of your teacher. DO NOT DISPOSE OF IN THE SINK
OR TRASH!!! Procedure: 1. Fill a Petri dish ¾ full of tap water. This will serve as your water source and recycling supply. Calibrating the Bulb of Pipet: 2. Fill the bulb completely with water. Squeeze the bulb and dip the mouth in the Petri dish of water and
release the squeeze. Then, with the bulb mouth held upward, squeeze a second time, just to the point where the water inside the pipet starts to come out. Then, still squeezing, dip the mouth into the Petri dish again and draw up the remaining water needed to fill the bulb.
3. Squeeze the water out into a 10 mL graduated cylinder. Leave enough water in the bulb to fill the launch
nozzle up to the line already on the bulb. The amount of water in the graduated cylinder is the volume of your gas chamber. Record this volume. Divide this volume by 6 and record this volume. Refill your gas chamber (bulb of pipet). Squeeze out your calculated number of the water into the graduated cylinder. Use a permanent marker and mark the water level. Make sure your bulb is dry first. Squeeze out another 1/6 of water and mark again. This should serve to increment the bulb into 6 equal volumes. (Note that there will still be some water left in the bulb.) Now refill your bulb with water.
Setting Up the Gas Generators: 4. The generators are plastic film canisters with nozzle-fitted caps containing the proper chemicals. Obtain a
canister for hydrogen and oxygen. Do not mix chemicals until ready for gas production. 5. To generate hydrogen, place enough 1.0M HCl (hydrochloric acid) in the canister labeled “hydrogen” to rise
up to 1 cm. When you are ready for gas production, add a few pieces of magnesium and put the cap back on. Then set the generator in your Petri dish.
6. To generate oxygen, place enough H2O2 (hydrogen peroxide) in the canister labeled “oxygen” to rise up to
1 cm. When you are ready for gas production, add 5 drops of KI (potassium iodide, the catalyst), put the cap back on, and swirl. Then set the generator in your Petri dish.
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7. If gas production becomes slow, check your generator. For example, you may need to add more magnesium to the hydrochloric acid if the magnesium is gone. If there is still magnesium, you may need to add more hydrochloric acid. You may also need to do the same for the oxygen generator by adding more hydrogen peroxide or potassium iodide.
Collecting Gas by Water Displacement: 8. Once your gas generators are producing and the lids are on, take your water-filled graduated bulb and
place it with the mouth of the bulb downward over the nozzle of the generator. The fit should be loose, enabling water to leak out as the bulb collects the gas. As soon as you have the amount of gas required, remove the nozzle.
Launching the Pipet Bulb: 9. Fill your graduated bulb with the required ratio of gases. (There should still be a little water left.) Take your
bulb over to the launch area. Place the wires on the launch pad into the bulb and push the launch button. Record the distance traveled.
Collecting and Testing Different Ratios: 10. Begin collecting oxygen gas in your bulb. Once it is 1/6 full, move the bulb from the oxygen generator to
the hydrogen generator and continue collecting until full, remembering some water is left. This gives you a 1:5 mixture of oxygen gas to hydrogen gas. Launch it and record its distance. Then repeat this procedure making the switch-over at various points so as to create mixtures of various proportions as listed in Data Table 1. Record the distances.
Observation/Data Tables:
Data Table 1: Volumes and Distances Volume of bulb
(in mL):
Volume of 1/6 of bulb (in mL):
Parts H2 6 5 4 3 2 1 0 Parts O2 0 1 2 3 4 5 6
Distance (cm)
Conclusion/Discussion: 1. Write balanced equations for the reactions taking place inside the two generator vials. 2. Explain the launch test for pure hydrogen gas and oxygen gas. 3. Write the balanced equations for the reaction of hydrogen gas and oxygen gas. 4. According to the balanced equation, what is the correct ratio for maximum explosiveness?
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5. According to your data, which ratio was the most explosive? Does this match your answer to question 4? If not, explain.
6. How many gas molecules would the pipet hold at standard temperature and pressure? 7. In the perfect ratio of hydrogen gas to oxygen gas, how many molecules of each gas are in the bulb? 8. What would be the mass, in grams, of each of the gases in question 7?
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Unit Four Experiment – 3 Balanced Chemical Equations
EX – C – U4 – 3
Introduction: The purpose of this experiment is to examine the relationship between amounts of reactants and products in a chemical reaction. Background: Although every chemical reaction involves using chemicals, it is common for individuals to overuse chemicals and therefore waste them. In this experiment, you will examine the stoichiometry – the relationship between amounts of materials – of the reaction between calcium chloride and sodium phosphate. You will also explore the concepts of limiting and excess reagents. A limiting reagent is completely used up in a reaction. The amount of product that can be formed depends on the quantity of limiting reagent present. An excess reagent is so called because, after the reaction is complete, there is still an amount left unreacted. Finally, you will estimate the actual yield of one product, calcium phosphate, and compare it with the theoretical yield predicted from the balanced equation. Safety:
Safety goggles will be worn at all times. No open-toe shoes are to be worn. Dispose of all materials according to the instructions of your teacher. DO NOT DISPOSE OF IN THE SINK
OR TRASH WITHOUT TEACHER APPROVAL!!! Procedure: Part 1: 1. Number, with the numerals 1 – 6, six large, clean, and dry test tubes. (NOTE: Test tubes NEED to be the
same length and diameter for accurate measuring during Part 2 of the experiment.) 2. Mount two 50-mL burets on a ring stand, using a double buret clamp. Identify the left buret as “Na3PO4”
and the right buret as “CaCl2”. 3. Obtain 50 mL of 0.5M Na3PO4 and 50 mL of 0.5M CaCl2. Fill the labeled burets with these solutions,
making sure that the solutions are below the 0 mL mark of each buret. 4. Using the filled burets, add the solutions to the six test tubes according to the following table. Place each
test tube in a test-tube rack.
Chemical Tube 1 Tube 2 Tube 3 Tube 4 Tube 5 Tube 6 0.5M Na3PO4 (mL) 1.00 2.00 3.00 4.00 5.00 6.00
0.5M CaCl2 (mL) 7.00 6.00 5.00 4.00 3.00 2.00 Keep the remaining solutions in the burets for later use. Record all volumes used as accurately as possible
in Data Table 1. NOTE: Use the correct number of decimal places. 5. Seal each tube with a rubber stopper. Mix the contents by inverting each test tube three times. Do not
shake. 6. Leave the test tubes undisturbed in the test-tube rack for at least ten minutes.
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7. Measure the height of the precipitate in each tube to the nearest 0.1 cm and record the measurement in Data Table 1.
Part 2: 1. The liquid above a settled precipitate is referred to as a supernatant. You can test the supernatant for
excess reagent. With a dropper pipet, remove a sample of supernatant from tube 1. Add three drops of supernatant to each of two adjacent depressions on a reaction plate. Rinse the dropper with distilled water and repeat this procedure for tubes 2 – 6.
2. Add three drops of 0.5M Na3PO4 to one set of samples from tubes 1 – 6. 3. Add three drops of 0.5M CaCl2 to the other set of samples. 4. Record the results of these spot-plate tests in Data Table 2. 5. Follow your teacher’s instructions for proper disposal of the materials. Observation/Data Tables:
Data Table 1: Data for Reaction Mixtures Test tube number
Na3PO4 CaCl2 Height Ppt. (cm)
Maximum theoretical yield of ppt. (mol) (mL) (mol) (mL) (mol)
1
2
3
4
5
6
Data Table 2: Spot Tests of Supernatant Samples from Test Tubes Substance
added Tube 1
precipitate? Tube 2
precipitate? Tube 3
precipitate? Tube 4
precipitate? Tube 5
precipitate? Tube 6
precipitate? Na3PO4
(yes or no?)
CaCI2 (yes or no?)
Reagent present in
excess
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Conclusion/Discussion: 1. What is the balanced equation for the reaction that occurred between calcium chloride and sodium
phosphate? Include states of matter for each chemical. 2. Calculate the number of moles of calcium chloride and sodium phosphate added to each tube. (HINT:
Calculate how many moles would be in 1 mL and then multiply by the number of milliliters in the sample.) Enter those results in Data Table 1.
3. For the tube with the greatest amount of precipitate, calculate the mole ratio between calcium chloride and
sodium phosphate. 4. Plot two separate bar graphs showing the height of calcium phosphate (cm) (on the y-axis) versus tube
number (on the x-axis). One graph should show the height of calcium phosphate actually obtained in each tube, and the other should show the maximum theoretical number of moles (also on the y-axis) of calcium phosphate in each tube. Number the tubes 1 – 6 from left to right.
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5. In which tube was there little or no reaction of the supernatant with either calcium chloride or sodium phosphate? What is the mole ratio of the reactants in this tube?
6. How are the two graphs similar? 7. Examining class data from conferring with other students, explain any inconsistency you observe in the
results. 8. Based on the results of this experiment, develop a hypothesis to explain how the stoichiometry could be
determined for a reaction which forms a gas (for example, calcium carbonate and hydrochloric acid reacting) instead of a precipitate.
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Appendix A Laboratory Equipment, Syllabus, and LPS Safety Contract
Triple beam balance Beaker Crucible and lid Test tube tongs
Buret Graduated Bunsen Ring Distilled water cylinder burner stand wash bottle
Test tube rack Double buret clamp Clay triangle
Erlenmeyer Funnel Wire Test flask gauze tube
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Safety goggles Ring clamp
Scoopula
Test tube brush
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Chemistry Syllabus
Mr. Michael Geist E-mail: [email protected] Office: B117 http://isite.lps.org/mgeist Availability: Before Block 1 and after Block 3 Course Description: Chemistry is the study of the structure, properties, and composition of substances and the changes that substances undergo. Laboratory experiences are used to reinforce classroom presentations. Chemistry is an essential class for students considering nursing, engineering, medical, or scientific areas of study. Algebra and geometry are prerequisites for this course. Course Objectives: Course objectives for all science courses can be accessed here: https://docushare.lps.org/docushare/dsweb/View/Collection-316260. District Common Assessments (DCAs): The data provided by the district common assessments are used to gauge the extent to which students are meeting state standards, to provide students and parents with information about student progress, to enhance school improvement planning, and to improve instruction. The two district common assessments provided this terms are as follows: Measurement and Matter Reactions and Stoichiometry District Grading Policy/Criteria: Quality, thoughtful assessment facilitates improvement in learning for students and teachers. Formative assessment provides ongoing information and guidance about student needs to facilitate
student learning towards objectives and guide future actions taken by both student and teacher. Summative assessment provides a measurement of student learning of objectives at the conclusion
of a specific instructional period. Report card grade reflects the achievement level of learning objectives at the end of a specific
grading period. A combination of formative and summative assessments comprise the final grade, with the vast majority of the grade determined by summative assessment.
Course Materials: Book: Wilbraham, Staley, Matta, and Waterman. Pearson Chemistry. 2012. Pearson. Lab manual: Geist, Michael. Chemistry Worksheet and Laboratory Manual (Term 1). 2017. LPS. Website: http://isite.lps.org/mgeist A calculator, writing utensil, and notes Safety: The safety manual, located at http://wp.lps.org/science/safety , was created to provide the classroom teacher, parents, and students with an additional resource of information pertaining to safety in the science classroom as outlined for the Lincoln Public Schools District.
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Class Outline: You will be expected to participate in all class discussions and will read any and all material assigned
in order to prepare for these discussions. Even though the reading in chemistry is not like those of other subjects, reading and practice is expected and required.
You will participate in and document laboratory experiments (generally once a week on average). You, as an individual, will submit a lab report from what you performed and concluded in certain experiments. Although you will be working in a group, you as an individual will document your findings. Additionally, a specified format for laboratory experiment reports will be handed out to you and adhered to throughout the course of the term.
You will do assignments from each chapter of the book as indicated by the instructor to reinforce important material covered in the course of that chapter.
Any project in the course of the term related to concepts you have learned or will learn during the course of the term will be entered as a grade in the Summative category.
The final exam will be developed by your instructor and administered at the end of the term. Classroom Rules/Expectations: EVERYONE is to be treated with respect at all times. People of different gender, race, religion,
and sexual orientation contribute to a better experience to be shared by all in and out of class. It is in your best interests as well as enjoyment of the class to cooperate, share, and contribute to all people in class as best you can. Anything to disrupt respect in the classroom will be addressed immediately according to school and district policy.
It is important for all classroom participants, teachers and students, to be punctual and always prepared. Although we cannot always avoid being late or unprepared, habitual tardiness and lack of preparation will not be allowed and will be reflected in your grade. You are considered tardy if you enter class following the bell ringing to signal the beginning of class. During third block classes, you may still be counted tardy if returning from lunch late.
Time will not be permitted in class for doing homework. Albeit guided practice and independent practice will be done in class, time does not permit doing homework (hence the name, homework). Additionally, as is also done in college, should one section end in class, you may be reasonably expected to move on to a new section. In other words, there will be no idle time in class.
You will receive a copy of the lab safety rules in addition to this syllabus. Both must be signed and adhered to throughout the course of the semester. NOTE: If a laboratory experiment requires you to wear any personal protective equipment (i.e., goggles, apron, etc.), they must be worn at all times. Failure to do so will result in your dismissal from the laboratory experiment and consequently a lower laboratory grade.
No electronic devices (i.e., mobile phones, headphones, etc.) will be allowed in the classroom or computer lab. If you have a question about this or related guidelines, please see me before bringing an item into class. Should you bring such an item, it will be confiscated and referred to administration.
No material should be opened, written, or applied in class that will interfere with the maintenance of classroom management and continuity. This includes, but is not limited to, magazines, separate reading material, personal notes or notes related to other classes, application of makeup, styling of hair, etc. Any materials which are not related to class that are used or created during class will be confiscated. Personal notes will not be returned at any time.
Books are to be covered at all times. Books should be covered with papers, not with adhesive plastic. Also, at the end of the semester, you will be required to vacate the book of all papers and the cover you have put on the book. Any damage incurred to the book by not covering or mistreatment and negligence of the book will be charged to you at the end of the semester.
As this class demands participation and preparedness, it is in your best interests to come to class fully rested. Sleeping/napping is not acceptable behavior in class and will result in dismissal from the
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classroom. Such activity logically suggests possible drug usage or suspicious behavior and will be investigated accordingly by the school nurse or administration.
You are expected to bring a pencil, notebook, calculator, and textbook with you to class every day. You will be issued a textbook during the first week of class. It is your responsibility to report any damage that may already be present with the textbook you receive so you are not fined for that damage at the end of the course.
Classroom Preparation and Grading: Full preparation for classroom discussion and participation will most likely require anywhere from thirty minutes to an hour or more a day outside of class. This includes reviewing notes, reading the material from the book, and working on problems. Some days, including those before and after more intense concepts, will require more time. It is your responsibility to be adequately prepared and make sufficient time reservations outside of class to come to class prepared. My means of grading the same as those of the Lincoln Southwest High School Science Department and are as follows:
A = 90.0 – 100.0 B+ = 85.0 – 89.9 B = 80.0 – 84.9 C+ = 75.0 – 79.9 C = 70.0 – 74.9 D+ = 65.0 – 69.9 D = 60.0 – 64.9 F = 0.0 – 59.9
Although the points a test or quiz is worth might change, as follows is the normal scheme for percentages of your final grade in this course: Tests/Lab Activities (Summative): Approximately 70% You may retake any test objective on the condition that you showed relevant work on that test
objective throughout the entirety of the test objective. Retakes will be provided on specified dates. Partial credit will be provided on all tests as long as you show RELEVANT work. On tests, work
MUST be shown to receive full credit. Otherwise, the logical conclusion to draw from seeing a right answer is that a student has cheated or guessed, neither of which demonstrates that the student has mastered the associated objectives. Failure to show relevant work on an objective will prohibit the student taking the exam from retesting that exam objective.
Quizzes (Formative): Approximately 20% Quizzes cannot be retaken, and you may expect a quiz at least twice a week. Partial credit will be provided on all quizzes as long as you show RELEVANT work. On quizzes,
work MUST be shown to receive full credit. Otherwise, the logical conclusion to draw from seeing a right answer is that a student has cheated or guesses, neither of which demonstrates that the student has mastered the associated objective. Quizzes will all be based on homework problems you have done, and you may use the homework you have done directly on the quizzes.
Homework (Formative): 0% Homework will be assigned every day save possibly on test days. Homework is designed to help you gauge your understanding of material. Credit is given for
homework indirectly in the quizzes. Homework will be reviewed and discussed the following day for understanding, but a grade will not be assigned for it. Statistically, there is a direct correlation between students who do their own homework and ask questions to higher test grades.
If you are not doing well in homework, seek help immediately from me. Homework will directly correlate with in-class quizzes. Final exam(s) (Summative): Approximately 10% There are NO retests for a final exam(s).
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Grades will be made available as often as is possible and feasible to that you are able to track your progress in this class. If the grades made available are not as current as the most recently graded assignment, feel free to ask me what your current grade is and I will be more than happy to share it with you. Classroom procedures (other reasonable procedures may be created as the semester progresses): Raised hand. When I raise my hand in class, that signal indicates that conversation with any other
individual should stop until I ask for a response from you. Talking when I raise my hand is disrespectful and will be addressed accordingly. Similarly, when another student has his/her hand raised or is talking to me during classroom discussion, the same policy is in effect.
Visitors. When a visitor enters the room, that is someone who is not normally in class (i.e., the principal, another teacher, etc.), it is expected that you will respect those visitors and myself as is normally expected.
Tests and quizzes. When a test or quiz is being given, at no time should anyone be talking. If you have a question, silently raise your hand and I will assist you as much as I can. NOTE: Any talking or suspicious communication (including any nonverbal communication) made during or after a test or quiz prior to teacher approval will result in the removal of your test or quiz, an automatic disciplinary referral, a replacement test without the use of written aids, and no ability to retest. If you continue to talk or otherwise communicate with others after this has occurred, you will be removed from the classroom and appropriate consequences will follow. Tests and quizzes are some of the greatest assessment and grading tools at your disposal, and interrupting another individual's self-assessment is disrespectful and, more importantly, damaging to the students still taking their tests or quizzes.
Substitute teachers. If I am unable to be in class and have a substitute teacher take my place for any amount of time, you will be expected and required to provide the utmost respect for the substitute teacher and extend the teacher every courtesy and fulfill every request. Failure to do so will result in administrative action upon my return. Should I receive no negative comments from the substitute teacher upon my return and you wish to have the teacher return should I be absent again, I will personally request that substitute return. However, should I receive no negative comments from the substitute teacher upon my return and you do not wish to have the teacher return should I be absent again, I will personally request that the substitute will not return and I will make other arrangements.
Retests. A student will only be allowed to retest an objective or objectives if the student has done relevant work for all problems on the objective(s) of the test. Failure to do so will preclude the student from being able to retest that exam. Also, once a student has begun a retest, the student may only have the time they are in there for that part of the day to do the retest. As in a normal testing period, the time given during that day of the retest being taken is the only time given – once you begin the test on that day, it must be completed during the allotted time. Failure to retest before the deadline for retesting will result in the inability to retest over that unit. The grade received from the retest will be the grade entered, whether higher or lower than the original test objective grade.
Other important notes: GET HELP AS SOON AS POSSIBLE. Students who put off asking questions or getting help when
they need it get further and further behind. E-mail me ([email protected] - the best way to get a hold of me outside of school), try calling me, hunt me down between classes, during class, or whenever is convenient.
Check your grades often. From previous experience, students who do not keep accurate recollections of their grades and find out at the last minute have difficult times trying to get the grades they desire (i.e., overstudying for final exams, etc.).
Let your parents know how to contact me. If you happen to get behind at all, it helps tremendously for your parents and me to collaborate on an effective way of helping you.
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Keep a positive mental attitude. Chemistry will undoubtedly be a much different experience for you, but it is a very fun subject.
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Appendix B SI Units and Conversions
Density
md
v
Example: If a substance has a mass of 0.75 g and a volume of 3.0 mL, what is the substance’s density?
m 0.75 g gd 0.25 mLv 3.0 mL
Example: Gold has a density of 19.3 g/cm3. If one has 10.0 cm3 of gold, what mass of gold is present?
33
m 19.3 gd m dv 10.0 cm 193 g
v cm
Example: Mercury has a density of 13.6 g/mL. If there are 7.48 g of mercury present, how many milliliters of mercury are there?
m m 7.48 gd v 0.55 mL
gv d 13.6 mL
Specific Gravity Comparison of densities
Formula: density of substance
Specific gravitydensity of water
Same units must be used in numerator and denominator Used to diagnoses certain illnesses, such as diabetes; used to check the condition of
the antifreeze in a vehicle; used for car batteries
Temperature Ways to convert: K = C + 273 C = K – 273 Example: If the temperature is 50C, what is the temperature in Kelvins?
K = 50 + 273 = 323 K
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Example: If the temperature is 50K, what is the temperature in degrees Celsius?
C = 50 – 273 = – 223 C
Units of Measurement
Quantity SI base unit or SI derived
unit Symbol Non-SI unit Symbol
Length meter* m Volume cubic meter m3 liter L Mass kilogram* kg Density grams per cubic centimeter
or grams per milliliter
g/cm3 g/mL
Temperature kelvin* K degree Celsius C Time second* s Pressure Pascal Pa atmosphere
millimeter of mercury Atm mm Hg
Energy Joule J calorie cal Amount of substance
mole* mol
Luminous intensity
candela* cd
Electric current
ampere* A
*: denotes an SI base unit
Commonly Used Prefixes in the Metric System
Prefix Symbol Meaning Factor Scientific notation
mega M 1 million times larger than the unit it precedes
1 000 000 106
kilo k 1000 times larger than the unit it precedes
1000 103
deci d 10 times smaller than the unit it precedes
1/10 10-1
centi c 100 times smaller than the unit it precedes
1/100 10-2
milli m 1000 times smaller than the unit it precedes
1/1 000 10-3
micro 1 million times smaller than the unit it precedes
1/1 000 000 10-6
nano n 1000 million times smaller than the unit it precedes
1/1 000 000 000 10-9
pico p 1 trillion times smaller than the unit it precedes
1/1 000 000 000 000 10-12
Important conversions: 1 cm3 = 1 mL 103 mL = 1000 cm3 = 1 L
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Weight and Mass
Mass: amount of matter an object has Weight: force that measures the pull on a given mass by gravity Mass does not change based on location; weight does.
Conversions (prelude to Chapter Four) Example: How many centimeters are in a kilometer? Solution: Since there are 100 centimeters in a meter and 1000 meters in a kilometer, find a way that will cancel out units. 1 km 1000 m 100 cm
• •1 1 km 1 m1 kilometer 1000 m 100 cm
• • 100000 cm1 1 km 1 m
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Appendix C Compound Name and Formula Writing
Metals/Nonmetals:
The charge of the metal ions in Group 1A is 1+. The charge of the metal ions in Group 2A is 2+. The charge of the metal ions in Group 3A is 3+. The charge of the transition metals and such elements as Sn, Pb, Hg, and Sb may have
more than one charge. The charge of the nonmetal ions in Group 5A is 3-. The charge of the nonmetal ions in Group 6A is 2-. The charge of the nonmetal ions in Group 7A is 1-. Group 8A has no ions.
Polyatomic ions:
Their charge is always negative except for NH4+.
Memorize or look at the table on page 147 for the name, formula, and charge of the polyatomic ions.
Forming ionic compounds:
Compounds have electrical neutrality. Na+ and S2- must be written as Na2S since you need two positive charges to balance the 2- charge on the S. Fe3+ and O2- must be written Fe2O3 since you need two 3+ charges to balance three 2- charges (6 + -6 = 0).
The positive ion is always written before the negative ion. If two or more polyatomic ions are used in the formula, enclose the polyatomic ion in
parentheses and put the number of polyatomic ions you need on the outside of the parentheses as a subscript. For example, Mg2+ and OH- must be written Mg(OH)2 since you need two negative charges of the OH- ion to balance the 2+ charge on the Mg.
Do not write the charge of the ion in the formula. For example, sodium sulfide is Na2S, not Na2+S2-, 2Na+S2-, or Na2
+S2-. Naming ionic compounds:
When a metal is involved, the name of the metal is used. For example, magnesium becomes “magnesium ion” when it becomes a cation.
When the metal ion can have two different charges, the charge of the ion is indicated by writing it in Roman numerals in parentheses after the name of the metal. For example, Cu+ is written as the Copper (I) ion. Cu2+ is written as the Copper (II) ion.
When a nonmetal is involved, ide is added as a suffix to the root word of the nonmetal (usually the first syllable). For example, phosphorus become the “phosphide ion” as oxygen becomes the “oxide ion.”
Polyatomic ions retain their names. To name a metal and a nonmetal together, combine the ion names. For example, when
Copper (II) ion is together with the nitride ion, the compound is Copper (II) nitride.
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Naming binary molecular compounds:
The first nonmetal gets its full name. The second nonmetal gets its root word + ide. Both nonmetals get a prefix denoting how many atoms are used to make the compound. However, when only one atom is used in the first nonmetal, the prefix mono is not attached. Examples:
o CO is carbon monoxide, not monocarbon monoxide. o N2O5 is dinitrogen pentaoxide.
Prefixes: o 1 atom – mono (or mon if it begins with an “o”) o 2 atoms – di o 3 atoms – tri o 4 atoms – tetra o 5 atoms – penta o 6 atoms – hexa o 7 atoms – hepta o 8 atoms – octa o 9 atoms – nona o 10 atoms – deca
Naming acids:
Use the list of acids to name them. Examples:
o HC2H3O2: acetic acid o H2CO3: carbonic acid o HNO3: nitric acid o H2SO4: sulfuric acid o H3PO4: phosphoric acid o HCl: hydrochloric acid o HBr: hydrobromic acid o HI: hydroiodic acid o HF: hydrofluoric acid
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Appendix D Chemical Reactions and Quantities
Chemical Reaction Classifications:
Synthesis/Combination (Oxidation-Reduction):
A + B AB
2Na(s) + Cl2 2NaCl (s)
Decomposition (Oxidation-Reduction):
AB A + B
2HgO(s) 2Hg(l) + O2
Single-Replacement (Oxidation-Reduction):
)(H + (aq)MgCl 2HCl(aq) + Mg(s)
B + AC BC + A
22 g
Double-Replacement (Precipitation):
A+B- + C+D- A+D- + C+B-
K2CO3(aq) + BaCl2(aq) 2KCl(aq) + BaCO3(s)
Combustion (Oxidation-Reduction):
CxHy + x + y4
O2 xCO2 + y
2
H2O
CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
Redox reactions:
I. The Meaning of Oxidation and Reduction
A. Oxidation 1. Classical definition: combination of an element with oxygen to produce oxides 2. Modern definition: complete or partial loss of electrons or gain of oxygen 3. Examples
a. Rusting (2Fe + 3O2 2Fe2O3) b. Methane oxidation (CH4 + 2O2 CO2 + 2H2O) c.
B. Reduction 1. Classical definition: loss of oxygen from a compound 2. Modern definition: complete or partial gain of electrons or loss of oxygen 3. Examples
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a. Reduction of iron ore (2Fe2O3 + 3C 4Fe + 3CO2) b. 2AgNO3 + Cu 2Ag + Cu(NO3)2
c.
C. Oxidation and reduction always occur simultaneously. D. Oxidation-reduction reactions
1. Reactions that involve oxidation and reduction occurring 2. Often called “redox reactions” 3. Electrons of one side must equal electrons of other side
a. Example 1
Mg(s) S(s) MgS(s)
i. Oxidizing agent: sulfur (gains electrons)
ii. Reducing agent: magnesium (loses electrons) b. Example 2
i. Oxidizing agent: copper (II) nitrate (gains electrons)
ii. Reducing agent: magnesium (loses electrons)
II. Oxidation Numbers
A. A positive or negative number assigned to a combined atom according to a set of arbitrary rules
B. Generally the charge an atom would have if the electrons in each bond were assigned to the atoms of the more electronegative element
C. Rules for assigning oxidation numbers
1. The oxidation number of an element in an elementary substance is 0. a. The oxidation number of chlorine in Cl2 or of phosphorus in P4 is 0. b. The oxidation number of Fe by itself is 0.
2. The oxidation number of an element in a monatomic ion is equal to the charge of that ion. a. In the ionic compound NaCl, sodium has an oxidation number of +1 and
chlorine has an oxidation number of –1.
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b. The oxidation number of the bromide ion (Br-) is –1 while the oxidation number of the iron (III) ion (Fe3+) is +3.
3. The oxidation number of hydrogen in a compound is +1, except in metal hydrides (i.e., NaH) where it is –1.
4. The oxidation number of oxygen in a compound is –2. except in peroxides (i.e., H2O2) where it is –1.
5. For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 0.
6. For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.
Solubility Rules
If a salt is said to be soluble, then it will not be a precipitate of the solution. Salts that are said to be insoluble will precipitate out of the solution.
Negative ion Rule
NO3–
All compounds formed with the negative ion are soluble.
I–, Br–, Cl– All compounds formed with the
negative ion are soluble except Ag+, Pb2+, Hg2
2+, and Cu+.
SO42–
Most compounds formed with the negative ion are soluble; exceptions
include SrSO4, BaSO4, CaSO4, RaSO4, Ag2SO4, and PbSO4.
CO32–, PO4
3–, SO32–
All compounds formed with the negative ion are insoluble except
those of the alkali metals and NH4+.
OH–
All compounds formed with the negative ion are insoluble except
those of the alkali metals, NH4+, Sr2+,
and Ba2+. (Ca(OH)2 is slightly soluble.)
S2–
All compounds formed with the negative ion are insoluble except those of the alkali metals, alkaline
earth metals, and NH4+.
Rules for Balancing Equations:
1. Be sure to write all the correct formulas for all the reactants and products in the reaction. In some cases, you may also need to write in parentheses the state of matter they are in. (i.e., Fe(s), Br2(l), etc.)
2. Write the formulas for the reactants on the left and the formulas for the products on the right with a yield sign () in between. If two of more reactants are involved, separate their formulas with a plus sign (+). When finished, you will have a skeleton equation.
3. Count the number of atoms of each element in the reactants and products. To be as easy as possible, a polyatomic ion appearing the exact same on both sides of the equation can be counted as a single unit.
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4. Balance the elements one at a time by using coefficients (the numbers out in front of the formulas). When no coefficient is written, it is assumed to be 1. It is best to begin the balancing operation with elements that appear only once of each side of the equation. You must not attempt to balance an equation by changing the subscripts in the chemical formula of a substance.
5. Check each atom or polyatomic ion to be sure that the equation is balanced. 6. Make sure all the coefficients are in the lowest possible ratio that balances.
Stoichiometric/Molar Conversions and Calculations:
To go from atoms to moles:
# of atoms
1
1 mol of representative unit
6.02 x 1023 atoms2.3 x 1026 atoms O
1
1 mol O
6.02 x 1023 atoms 380 mol O
To go from moles to atoms:
# of moles
1 6.02 x 1023 molecule
mol # of atoms
molecule
3.6 mol C6H12O6
1 6.02 x 1023 molecule
mol 24 atoms
molecule 5.2 x 1025 atoms
What is gram atomic mass (gam)? Gram atomic mass is the average mass of an element per mole. This is shown on the Periodic Table of Elements underneath the symbol of the element. What is the gram molecular mass (gmm) and how is it calculated? The gram molecular mass of any molecular compound is the mass of one mole of that compound. To calculate it, add the gram molecular masses of the atoms that make it up. For example, the mass of water would be calculated by doing the following (since there are two hydrogen atoms and one oxygen atom in each mole of water):
2 mol H
1
1.0 g H
1 mol H
1 mol O
1
16.0 g O
1 mol O 18.0 g H2O
What is the gram formula mass (gfm) and how is it calculated? The gram formula mass of any ionic compound is the mass of one mole of the formula unit of that ionic compound. It is calculated the exact same way as the gram molecular mass of a molecule except that it is done for an ionic compound. To calculate, simply add up the atomic masses of the ions in the formula of the compound. For example, in magnesium hydroxide (Mg(OH2)) where the gmm of Mg is 24.3 g, H is 1.0 g, and O is 16.0 g, the gfm for magnesium hydroxide would be calculated as follows:
1 x 24.3 g Mg + 2 x 1.0 g H + 2 x 16.0 g O = 58.3 g Mg(OH)2
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To go from moles to grams for a compound:
# of moles of substance
1
gam, gfm, or gmm of substance
mol
2.85 mol H2O
1
18.0 g H2O
1 mol H2O 51.3 g H2O
To go from grams to moles for a compound:
# of grams of substance
1 mol of substance
gam, gfm, or gmm of substance
32.5 g H2O
1 1 mol H2O
18.0 g H2O 1.81 mol H2O
To go from moles to volume of a gas at STP:
# of moles of gas
1 22.4 L of gas
1 mol of gas
2.8 moles CO2
1 22.4 L CO2
1 mol CO2
63 L CO2
To go from density at STP to molar mass of a gas:
density of gas in g
L 22.4 L of gas
mol of gas
1.43 g O2
L O2
22.4 L O2
mol O2
32.0 L O2
To calculate percent composition of an element in a compound:
Experimentally: % mass of Element A =
grams of Element A
grams of compound100%
For example, if a compound is made up of 7.65 g hydrogen and 5.25 g carbon, the total mass of the compound is 12.90 g. To calculate the percent mass of hydrogen in the compound, you would divide 7.65 g by 12.90 g and multiply by 100% to get a percent composition of 59.3% hydrogen. Theoretically:
% mass of Element A =
grams of Element A in 1 mol of the compound
molar mass of the compound100%
For example, the molar mass of hydrogen peroxide (H2O2) is 2 x 1.01 g + 2 x 16.00 g = 34.02 g. Out of that 34.02 g, the mass of hydrogen that is in that mole of hydrogen peroxide is 2 x 1.01 g = 2.02 g. To calculate the percent composition of hydrogen, you would divide 2.02 g by 34.02 g and multiply by 100% to get a percent composition of 5.94% hydrogen.
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To calculate the mass of an element in a given amount of a compound:
mass of compound
1
mass of element in 1 mol of the compound
molar mass of the compound
For example, if you were asked to calculate the mass of carbon in 48.3 g of methane (CH4), you would know that for every molar mass of methane, which is approximately 16.0 g, 12.0 g of that mole of methane is made up of carbon. Therefore, to calculate the mass present in 48.3 g of methane,
g of carbon =
48.3 g CH4
1
12.0 g C
16.0 g CH4
36.2 g carbon
To calculate the empirical formula of a compound: Example: What is the empirical formula of a compound that is 10.0% carbon, 0.80% hydrogen, and 89.1%
chlorine. 1. Realize that in a 100.0 g sample of this compound, 10.0 g would be carbon, 0.80 g would be hydrogen, and
89.1 g would be chlorine. 2. Convert the grams of each of the elements to moles.
10.0 g C
1
1 mol C
12.0 g C 0.833 mol C
0.80 g H
1
1 mol H
1.0 g H 0.80 mol H
89.1 g Cl1
1 mol Cl35.5 g Cl
2.51 mol Cl
3. The mole ratio is C0.833H0.80Cl2.51. This is not the correct empirical formula though because it is not the lowest
whole-number ratio. To do this, we need to divide all the molar quantities by the smallest number of moles. This will give a 1 for the element with the smallest number of moles.
0.833 mol C
0.80 1.0 mol C
0.80 mol H
0.80 1.0 mol H
2.51 mol Cl
0.80 3.1 mol Cl
4. The mole ratio is now CHCl3.1. Given how close the 3.1 is to 3, the empirical formula for this is CHCl3. If the
mole ration was something like CHCl0.5, we would need to multiply each molar quantity by a value such as 2 to get all whole numbers, resulting in C2H2Cl.
To calculate the molecular formula of a compound given molar mass: Example: What is the molecular formula of the compound whose molar mass is 180.0 g and the empirical formula
is CH2O? 1. Calculate the empirical formula mass. In this case, the molar mass of CH2O would be 30.0 g. 2. Divide the compound’s molar mass by the empirical formula mass. In this case, you would divide 180.0 g by
30.0 g to get a value of approximately 6. 3. Multiply the subscripts in the empirical formula by the value you calculated in step 2 to get the molecular
formula. Multiplying the example empirical formula subscripts by 6, the answer would be C6H12O6.
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Type of Reaction
Hints / What to Look For on the Reactant
Side
What to Do to Complete the
Reaction
Synthesis/Combination Two elements, element and a diatomic gas/liquid/solid
1. Combine the elements as you would if you were forming any ionic compound.
2. Balance the equation.
Decomposition One compound
1. Break down the compound into its constituent elements and/or compounds.
2. Balance the equation.
Single Replacement Only one ionic compound; the other reactant is an element or a diatomic gas/liquid/solid
1. Switch around the two anions or the two cations that need to be replaced with each other. Remember the Activity Series of Metals and of Halogens when it comes to displacing a metal. Also be sure to balance charges in the new compound formed (i.e. Ca replacing Ag in AgCl has a 2+ charge, resulting in CaCl2 for charges to balance).
2. If a displaced element exists in a diatomic state in nature, be sure to indicate this (i.e. H H2).
3. Balance the equation.
Double Replacement
Two ionic compounds Product cases: 1. One precipitate formed. 2. One gas formed. 3. One liquid formed.
1. Switch around the two cations that need to be replaced with each other. Also be sure to balance charges in the new compound formed (i.e. Ca replacing Ag in AgCl has a 2+ charge, resulting in CaCl2 for charges to balance).
2. Balance the equation.
Combustion
A hydrocarbon (something with carbon and hydrogen) and oxygen gas (can be complete or incomplete combustion)
1. If there is a sufficient amount of oxygen, carbon dioxide and water will be the products (complete combustion). If there is an insufficient amount of oxygen, carbon monoxide and water will be the products (incomplete combustion).
2. Balance the equation.
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Appendix E Practice Tests
Unit One Practice Test (Multiple Choice) PT – C – U1
LPS Standard(s): 12.2.4b State Standard(s): 12.1.2a Identification. Identify the following as (A) a physical property or (B) a chemical property for ethanol (C2H5OH). 1. Ethanol is a liquid at room temperature. 2. Ethanol can chemically react with another chemical to become an aldehyde. 3. Ethanol has a density of 0.789 g/cm3 4. Ethanol is colorless. 5. Ethanol is flammable. 6. Ethanol is a central nervous system depressant in the human body.
LPS Standard(s): 12.2.4e, 12.2.6b State Standard(s): 12.1.2a Identification. Identify the following as (A) a physical change or (B) a chemical change.
7. Burning paper 11. Decomposing animal remains 8. Melting ice 12. Rusting iron 9. Cutting paper 13. Freezing liquid mercury 10. Grilling a steak 14. Igniting an explosive
LPS Standard(s): 12.2.4b State Standard(s): 12.1.2a Classification. Classify each of the examples as one of the following. NOTE: A classification may be
used more than once.
Choices: (A) compound (B) element (C) heterogeneous mixture (D) homogeneous mixture 15. Salad 20. Titanium 16. Sterling silver 21. Carbon dioxide 17. C12H22O11 22. Orange juice 18. Mud 23. Brass 19. (NH4)3PO4
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Multiple Choice. Identify the letter of the choice that best completes the statement or answers the question.
24. If 2.5 g of nitrogen gas reacts with hydrogen gas to produce 38.5 g of ammonia, how much
hydrogen gas reacted with the nitrogen gas? (A) 15.4 g (B) 36.0 g (C) 41.0 g (D) Not enough information to determine LPS Standard(s): --- State Standard(s): --- Significant Figures. Identify how many significant figures are in each reading (typed or visual) using
the following choices. Choices: (A) 1 (B) 2 (C) 3 (D) 4 (E) 5
25. 0.00385 mL 31. 32. 26. 2.01 x 103 g 27. 17.30 kg 28. 254.25 cm3 29. 0.02 eV 30. 0.05820 L Calculation. Perform the following operations, expressing your answers to the proper number of
significant figures and/or decimal places. Then select the correct choice corresponding to that calculation.
33. 2.00 x 4.0
(A) 8 (B) 8.0 (C) 8.00 (D) 8.000 34. 5.10 x 2.391 (A) 12 (B) 12.2 (C) 12.19 (D) 12.194 35. 0.0218
0.2419 (A) 0.09 (B) 0.090 (C) 0.0901 (D) 0.09012 36. 8.3 + 1.05 (A) 9 (B) 9.3 (C) 9.35 (D) 9.4 37. 3.2 – 3.192 (A) 0 (B) 0.01 (C) 0.008 (D) 0.0080 LPS Standard(s): --- State Standard(s): 12.3.3a,d Scientific Notation. Express each of the following results in scientific notation with appropriate
decimal places or significant figures by selecting the corresponding choice.
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38. 4800000 (A) 4.8 ×10-6 (B) 48 ×10-5 (C) 4.8 ×10-23 (D) 48 ×105 (E) 4.8 ×106 39. 0.00006351 (A) 0.6351 × 10-4 (B) 63.51 × 10-6 (C) 6351 × 10-8 (D) 6.351 × 10-5 (E) 6351.×10-8
40. Fifty-two (A) 52 × 100 (B) 5.2 × 101 (C) 0.52 × 102 (D) 520 × 10-1 (E) 52 Calculation. Answer the following questions with appropriate decimal places and/or significant figures
by selecting the corresponding choice. The following mass measurements were taken by Mr. Geist using several different scales during an experiment: 9.95 g, 10.102 g, 9.89 g, and 10.316 g. 41. What is the average of their results? (A) 10.06 g (B) 10.064 g (C) 10.07 g (D) 10.1 g 42. How many centimeters are there in 3 kilometers? (A) 0.00003 cm (B) 0.003 cm (C) 3000 cm (D) 300000 cm 43. What is the density 5.6 g/cm3 in kg/m3? (A) 5600000 kg/m3 (B) 5600 kg/m3 (C) 560 kg/m3 (D) 0.0056 kg/m3 (E) 0.0000056 kg/m3 LPS Standard(s): --- State Standard(s): 12.1.2 Modeling. Select the answer that best answers or completes the specified questions and statements.
For questions involving calculations, choose answers with proper units and the correct number of significant figures.
Refer to the graph below for questions 44 – 49.
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44. What is the equation of the line for substance B? (A) m = – 0.05000V + 2.708 (C) y = 2.708x – 0.05000 (B) y = – 0.05000x + 2.708 (D) m = 2.708V – 0.05000 45. What is the equation of the line for substance A? (A) m = 4.540V + 1.000 (C) y = 1.000x + 4.540 (B) y = 4.540x + 1.000 (D) m = 1.000V + 4.540 46. What is the mass of a 75.0 cm3 piece of substance B? (A) 0.0361 g (B) 27.7 g (C) 203 g (D) 341 g (E) Not enough information available 47. Which of the following would occupy the largest volume? (A) 50 g of substance A (C) (A) and (B) would occupy the same space. (B) 50 g of substance B (D) Not enough information available 48. Based on the graph from the previous page and the table at the right, which element or compound is
substance B? (A) Aluminum (B) Carbon dioxide (C) Ethanol (D) Titanium (E) Not enough information available 49. Based on the graph from the previous page and the table at the right, which element or compound is
substance A? (A) Corn oil (B) Nitrogen gas (C) Ammonia (D) Titanium (E) Not enough information available 50. What is the volume (in mL) of 344 grams of ice if it has a density of 0.92 g/mL? (A) 0.0027 mL (B) 37 mL (C) 320 mL (D) 370 mL
Substance Density(g/mL) Aluminum Ammonia Carbon dioxide Chlorine gas Corn oil Ethanol Gasoline Nitrogen gas Neon Oxygen gas Sucrose Titanium
2.70 0.718
1.83 2.95
0.922 0.789
0.67 1.17 0.84 1.33 1.59
4.5
page A-26 – C – T1 – BOOK
Unit One Practice Test (Short Answer/Calculation) PT – C – U1
LPS Standard(s): --- State Standard(s): 12.3.3a,d Calculations. Answer the following questions. Show work or receive no credit. Use the correct
number of significant figures and/or decimal places. Show proper units, and express answers using scientific notation.
51. Convert 1.0 g/cm3 to lb/ft3. (1 lb. = 454 g; 1 in. = 2.54 cm; 12 in. = 1 ft.) 52. In the United States, we measure our car speeds in miles per hour, but scientists often measure
speed in meters per second. If a rocket is traveling at 1550 miles per hour, what is its speed in meters per second? (1 km = 0.621 miles)
page A-27 – C – T1 – BOOK
Unit Two Practice Test (Multiple Choice) PT – C – U2
LPS Standard(s): 12.2.4a, 12.2.5a, 12.2.5c State Standard(s): 12.3.1a Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question.
1. The number of neutrons in an atom of the isotope depicted as C146 is ___.
(A) 6 (B) 8 (C) 12.011 (D) 14 (E) 20
2. The number of protons in an atom of the isotope depicted as C146 is ___.
(A) 6 (B) 8 (C) 12.011 (D) 14 (E) 20
3. The mass number of the isotope depicted as C146 is ___.
(A) 6 (B) 8 (C) 12.011 (D) 14 (E) 20
4. The atomic number of the isotope depicted as C146 is ___.
(A) 6 (B) 8 (C) 12.011 (D) 14 (E) 20
5. The number of electrons in a neutral atom of the isotope depicted as C146 is ___.
(A) 6 (B) 8 (C) 12.011 (D) 14 (E) 20 6. The number of neutrons in an atom of the isotope oxygen-18 is ___. (A) 8 (B) 10 (C) 15.999 (D) 18 (E) 26 7. The number of protons in an atom of the isotope oxygen-18 is ___. (A) 8 (B) 10 (C) 15.999 (D) 18 (E) 26 8. The mass number of the isotope oxygen-18 is ___. (A) 8 (B) 10 (C) 15.999 (D) 18 (E) 26 9. The atomic number of the isotope oxygen-18 is ___. (A) 8 (B) 10 (C) 15.999 (D) 18 (E) 26 10. The number of electrons in a neutral atom of the isotope oxygen-18 is ___. (A) 8 (B) 10 (C) 15.999 (D) 18 (E) 26 LPS Standard(s): 12.2.5a, 12.2.5d State Standard(s): 12.3.2d Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question. 11. ___ is a halogen. (A) Argon (B) Chlorine (C) Magnesium (D) Sodium 12. ___ is an alkali metal. (A) Beryllium (B) Fluorine (C) Lithium (D) Neon
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13. Elements without luster and that cannot conduct electricity are ___. (A) metals (B) nonmetals (C) metalloids 14. Beryllium, magnesium, and strontium are a family of elements that are most accurately called
the ___. (A) alkali metals (B) alkaline earth metals (C) halogens (D) noble gases 15. Xe, Ar, and He are a family of elements that are most accurately called the (A) alkali metals (B) alkaline earth metals (C) halogens (D) noble gases 16. ___ is in the same group as beryllium. (A) Carbon (B) Magnesium (C) Nitrogen (D) Oxygen 17. The elements radium, radon, and rubidium are best classified as being ___. (A) representative elements (C) periodic elements (B) transition elements (D) inner transition elements 18. The elements uranium, plutonium, and einsteinium are best classified as being ___. (A) representative elements (C) periodic elements (B) transition elements (D) inner transition elements 19. The elements copper, cadmium, and zinc are best classified as being ___. (A) representative elements (C) periodic elements (B) transition elements (D) inner transition elements 20. ___ has the same chemical and physical properties as sulfur. (A) Fluorine (B) Magnesium (C) Oxygen (D) Xenon LPS Standard(s): --- State Standard(s): 12.3.2 Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question. 21. The name for the compound with the formula N2O4 is ___. (A) nitrogen oxide (D) nitrogen tetroxide (B) nitrogen (II) oxide (E) dinitrogen tetraoxide (C) nitrogen (IV) oxide 22. The name for the compound with the formula SnO is ___. (A) tin oxide (D) tin (II) oxide (B) monotin dioxide (E) tin (IV) oxide (C) tin (I) oxide 23. The name for the compound with the formula P2O is ___. (A) phosphorus oxide (D) phosphorus (I) oxide (B) diphosphorus oxide (E) phosphorus (II) oxide (C) diphosphorus monoxide 24. The name for the compound with the formula Cu2O is ___. (A) copper oxide (D) copper (I) oxide (B) copper monoxide (E) copper (II) oxide (C) dicopper monoxide
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25. The name for the compound with the formula (NH4)2Cr2O7 is ___. (A) ammonium chromide (D) dinitrogen octahydrogen dichromate (B) ammonium chromate (E) nitrogen tetrahydrogen monochromate (C) ammonium dichromate 26. The name for the compound with the formula SO2 is ___. (A) sulfur oxide (D) sulfur (I) oxide (B) sulfur dioxide (E) sulfur (IV) oxide (C) monosulfur dioxide 27. The name for the compound with the formula ZnCl2 is ___. (A) zinc chloride (D) zinc (I) chloride (B) zinc dichloride (E) zinc (II) chloride (C) monozinc dichloride 28. The name for the compound with the formula Cr(ClO2)6 is ___. (A) chromium chlorite (D) chromium (I) chloride (B) chromium (I) chlorite (E) chromium (VI) chloride (C) chromium (VI) chlorite 29. The name for the compound with the formula NO is ___. (A) nitrogen oxide (D) nitrogen monoxide (B) nitrogen (I) oxide (E) mononitrogen monoxide (C) nitrogen (II) oxide 30. The name for the compound with the formula Fe(CN)3 is ___. (A) iron cyanide (D) iron (II) cyanide (B) iron tricyanide (E) iron (III) cyanide (C) monoiron tricyanide 31. The name for the compound with the formula Mg3N2 is ___. (A) magnesium nitrate (D) magnesium (II) nitride (B) magnesium nitride (E) magnesium (II) nitrite (C) magnesium nitrite 32. The name for the compound with the formula Sn(CrO4)2 is ___. (A) tin chromate (D) tin (IV) chromate (B) tin dichromate (E) tin (IV) dichromate (C) tin (II) dichromate 33. The name for the compound with the formula Ca(NO3)2 is ___. (A) calcium dinitrate (D) calcium nitrate (B) calcium dinitrogen hexoxide (E) calcium (II) nitrate (C) calcium mononitrogen trioxide 34. The name for the compound with the formula PF5 is ___. (A) monophosphorus pentafluoride (D) phosphorus (V) fluoride (B) phosphorus pentafluoride (E) monophosphorus (V) fluoride (C) phosphorus fluoride 35. The name for the compound with the formula Cu3(PO4)2 is ___. (A) tricopper diphosphate (D) copper diphosphate (B) copper (II) phosphate (E) tricopper diphosphide (C) copper (III) phosphate
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LPS Standard(s): --- State Standard(s): 12.3.2 Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question. 36. The formula for ammonium sulfate is ___. (A) (NH4)2S (B) (NH4)2SO3 (C) (NH4)2SO4 (D) (NH4)3SO3 (E) (NH4)3SO4 37. The formula for barium chlorate is ___. (A) Ba(ClO)2 (B) Ba(ClO2)2 (C) Ba(ClO3)2 (D) Ba(ClO4)2 (E) BaCl2 38. The formula for potassium sulfite is ___. (A) KHSO3 (B) KHSO4 (C) K2SO3 (D) K2SO4 (E) K2S 39. The formula for calcium dihydrogen phosphate is ___. (A) CaH2PO4 (B) Ca2H2PO4 (C) Ca(H2PO4)2 (D) Ca(H2HPO4)2 40. The formula for tin (II) chloride is ___. (A) SnCl (B) Sn4Cl (C) SnCl4 (D) SnCl2 (E) Sn2Cl 41. The formula for silver oxide is ___. (A) AgO (B) AgO2 (C) Ag2O (D) Ag3O2 (E) Ag2O3 42. The formula for lithium phosphide is ___. (A) Li3PO4 (B) Li3PO3 (C) Li3PO2 (D) Li3PO (E) Li3P 43. The formula for dioxygen dibromide is ___. (A) OBr (B) O2Br2 (C) OBr2 (D) O2Br (E) (O)2(Br)2 44. The formula for copper (II) oxide is ___. (A) CuO (B) Cu2O (C) CuO2 (D) Co2O (E) CoO2 45. The formula for zinc phosphate is ___. (A) ZnP (B) ZnPO4 (C) Zn3PO4 (D) Zn3(PO4)2 (E) Zn2(PO4)3 46. The formula for sulfur hexachloride is ___. (A) SCl (B) S5Cl (C) SCl5 (D) S6Cl (E) SCl6 47. The formula for cadmium chloride is ___. (A) CdCl (B) Cd2Cl (C) CdCl2 (D) Cd3Cl (E) CdCl3 48. The formula for carbon monoxide is ___. (A) CO (B) CO2 (C) C2O (D) C3O2 (E) C2O3 49. The formula for dinitrogen tetroxide is ___. (A) N2O5 (B) N5O2 (C) NO (D) N2O4 (E) N4O2 50. The formula for lithium sulfite is ___. (A) Li2S (B) Li2SO3 (C) Li2SO4 (D) LiS (E) LiSO3
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Unit Two Practice Test (Short Answer/Calculation) PT – C – U2
Calculation. Answer the following questions. Show work or receive no credit. You must also show
proper units. 51. The element chromium contains four naturally occurring isotopes:
Cr Cr Cr Cr 5424
5324
5224
5024
The relative abundances and atomic masses are as follows: 4.31% for chromium-50 (mass = 50.000 amu) 83.76% for chromium-52 (mass = 52.000 amu) 9.55% for chromium-53 (mass = 53.000 amu) 2.38% for chromium-54 (mass = 54.000 amu)
Calculate the average atomic mass of chromium.
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Unit Three Practice Test (Multiple Choice) PT – C – U3
LPS Standard(s): --- State Standard(s): --- Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question. 1. Which of the following is the closest to the molar mass of Ag2O? (A) 124 g (B) 140 g (C) 232 g (D) 248 g (E) 340 g 2. Which of the following is the closest to the molar mass of Cu2CrO4? (A) 131.541 g (B) 179.538 g (C) 231.534 g (D) 243.084 g (E) 399.072 g 3. Which of the following is the closest to the molar mass of lithium oxide? (A) 22.940 g (B) 29.881 g (C) 36.822 g (D) 38.939 g (E) 59.762 g 4. Which of the following is the closest to the molar mass of sodium sulfate? (A) 71.05 g (B) 94.04 g (C) 119.05 g (D) 142.04 g (E) 284.20 g 5. Which of the following is the closest to the molar mass of ammonium phosphate? (A) 113.009 g (B) 121.072 g (C) 141.023 g (D) 149.086 g (E) 174.957 g LPS Standard(s): --- State Standard(s): 12.1.3a Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question. 6. Which of the following is equal to Avogadro’s number? (A) the number of molecules of lithium oxide in 1 mol Li2O (B) the number of atoms of chlorine in 1 mol Cl2 (C) the number of formula units of carbon dioxide in 1 mol CO2
(D) the number of molecules of sulfur hexafluoride in 1 mole SF6 7. Which of the following contains the most formula units: 10.0 mol AlCl3, 10.0 mol Ba(NO3)2, or
10.0 mol (NH4)3PO4? (A) 10.0 mol AlCl3 (C) 10.0 mol (NH4)3PO4 (B) 10.0 mol Ba(NO3)2 (D) They all contain the same number of formula units. 8. Which of the following contains the most atoms: 10.0 mol AlCl3, 10.0 mol Ba(NO3)2, or 10.0 mol
(NH4)3PO4? (A) 10.0 mol AlCl3 (C) 10.0 mol (NH4)3PO4 (B) 10.0 mol Ba(NO3)2 (D) They all contain the same number of atoms. 9. Which of the following is NOT a representative particle? (A) atom (B) cation (C) formula unit (D) molecule (E) neutron 10. Avogadro’s number is ___. (A) 22.4 L (D) 0C (B) 1 atm (E) 6.022 1022 (C) the number of representative particles in a mole of a substance
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LPS Standard(s): --- State Standard(s): 12.1.2d Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question. 11. 0.075 mol of titanium are equivalent to ___ atoms of titanium. (A) 1.2 10–25 (B) 3.6 (C) 6.4 102 (D) 4.5 1022 (E) 2.2 1024 12. 9.0 1023 formula units of SrCl2 is equivalent to ___ grams of SrCl2. (A) 0.0094 g (B) 110 g (C) 160 g (D) 240 g (E) 3.4 1045 g 13. 1.8 1020 atoms of silver are equivalent to ___ moles of silver atoms. (A) 3.0 10–4 (B) 3.3 10–3 (C) 0.30 (D) 3.0 102 (E) 1.1 1044 14. The formula of sucrose is C12H22O11. What mass will 1.50 moles of sucrose have? (A) 0.00438 g (B) 33.6 g (C) 228 g (D) 342 g (E) 513 g 15. To convert moles into mass for a substance, you must ___. (A) divide the moles by Avogadro’s number (B) multiply the moles by Avogadro’s number (C) divide the moles by the molar mass (D) multiply the moles by the molar mass (E) divide the moles by the molar volume 16. To convert the number of representative particles into the number of moles for a substance, you
must ___. (A) divide the number of representative particles by Avogadro’s number (B) multiply the number of representative particles by Avogadro’s number (C) divide the number of representative particles by the molar mass (D) multiply the number of representative particles by the molar mass (E) divide the number of representative particles by the molar volume 17. 4.0 moles of sodium are equivalent to ___ grams of sodium. (A) 0.174 g (B) 0.179 g (C) 5.75 g (D) 89.6 g (E) 92.0 g 18. How many formula units are in 5.00 grams of lithium chloride? (A) 3.52 10–22 (B) 2.64 100 (C) 7.10 1022 (D) 1.34 1023 (E) 5.11 1024
19. ___ moles of calcium bromide are in 5.0 grams of calcium bromide. (A) 2.5 10–2 (B) 4.2 10–2 (C) 4.0 101 (D) 1.0 103 (E) 3.0 1024 20. How many molecules are in 25.0 moles of propane (C3H8)? (A) 4.15 10–23 (B) 5.67 10–1 (C) 1.10 102 (D) 2.41 1022 (E) 1.51 1025 LPS Standard(s): --- State Standard(s): 12.1.2d Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question. 21. What is the number of moles in 5.0 L of SO3 gas at STP? (A) 0.062 mol (B) 0.22 mol (C) 4.5 mol (D) 16 mol (E) 110 mol
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22. How many atoms of chlorine are there in 75.0 liters of chlorine gas at STP? (A) 2.79 10–21 atoms of chlorine (D) 4.03 1024 atoms of chlorine (B) 6.37 1023 atoms of chlorine (E) 1.01 1027 atoms of chlorine (C) 2.02 1024 atoms of chlorine 23. At STP, how many liters of nitrogen dioxide are occupied by 6.1 1022 molecules of nitrogen
dioxide? (A) 2.3 L (B) 4.7 L (C) 220 L (D) 1.3 1023 L (E) 1.6 1045 L 24. What is the density of oxygen gas at standard temperature and pressure? (A) 5.31 10–23 g/L (B) 0.714 g/L (C) 1.43 g/L (D) 358 g/L (E) 717 g/L 25. At standard temperature and pressure, what is the volume, in liters, of 3.2 moles of argon gas? (A) 0.080 L (B) 0.14 L (C) 7.0 L (D) 72 L (E) 130 L 26. At STP, what is the volume, in liters, of 5.00 grams of nitrogen gas? (A) 0.00797 L (B) 4.00 L (C) 6.25 L (D) 8.00 L (E) 3140 L 27. If the density of a gas is 0.902 g/L at standard temperature and pressure, that gas is ___. (A) H2 (B) He (C) Ne (D) F2 (E) SO3
28. Standard temperature and pressure is equivalent to ___. (A) 22.4 L (B) 0C and 1 atm (C) 6.022 1023 particles (D) the molar mass 29. The volume of one mole of a substance is 22.4 L at STP for all ___. (A) compounds (B) elements (C) gases (D) liquids (E) solids 30. What is the volume, in liters at standard temperature and pressure, of 0.500 mol of propane
(C3H8)? (A) 0.0335 L (B) 5.60 L (C) 11.2 L (D) 16.8 L (E) 22.4 L
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Unit Three Practice Test (Short Answer/Calculation) PT – C – U3
Calculation. Answer the following questions. Show work or receive no credit. You must also show
proper units. 31. How many molecules are present in 1.0 grams of aspirin, C9H8O4? 32. There are 7.85 x 1025 molecules of a gas in a chamber at 0C and one atmosphere of pressure.
How many liters of the gas are in the chamber?
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Unit Four Practice Test (Multiple Choice) PT – C – U4
LPS Standard(s): 12.2.6a State Standard(s): 12.3.3a Multiple Choice. Identify the letter of the choice that best completes the statement or answers the
question. 1. In the chemical equation 2H2 + O2 2H2O, the H2O is a ___. (A) catalyst (B) coefficient (C) inhibitor (D) product (E) reactant 2. When the chemical equation Mg + HCl → MgCl2 + H2 is balanced, the coefficient of H2 is ___. (A) 1 (B) 2 (C) 3 (D) 6 3. When the chemical equation N2 + H2 → NH3 is balanced, the coefficient of H2 is ___. (A) 1 (B) 2 (C) 3 (D) 4 4. When iron reacts with oxygen gas, iron (III) oxide is produced. The coefficient of iron (III) oxide in
the balanced chemical equation for this reaction is ___. (A) 1 (B) 2 (C) 3 (D) 4 5. Aluminum chloride and bubbles of hydrogen gas are produced when metallic aluminum is placed
in hydrochloric acid. What is the balanced chemical equation for this reaction? (A) H + AlCl → Al + HCl (D) Al + 2HCl → AlCl2 + H2
(B) 2Al + 6HCl → 2AlCl3 + 3H2 (E) H2 + AlCl3 → Al + 2HCl (C) Al + HCl3 → AlCl3 + H 6. If you rewrite the following word equation as a balanced chemical equation, what will the
coefficient and symbol for fluorine be? nitrogen trifluoride nitrogen gas + fluorine gas (A) 3F (B) 6F2 (C) F3 (D) 6F (E) 3F2 7. Which of the following is NOT a true statement concerning what happens in all chemical
reactions? (A) The ways in which atoms are joined together are changed. (B) New atoms are formed as products. (C) The starting materials are referred to as reactants. (D) The bonds of the reactants are broken and new bonds of the products are formed. 8. What are the missing coefficients for the skeleton equation: Cr(s) + Fe(NO3)2(aq) Fe(s) + Cr(NO3)3(aq) (A) 4, 6, 6, 2 (B) 2, 3, 2, 3 (C) 2, 3, 3, 2 (D) 1, 3, 3, 1 (E) 2, 3, 1, 2 9. What are the missing coefficients for the skeleton equation: NH3(g) + O2(g) N2(g) + H2O(l) (A) 4, 3, 2, 6 (B) 2, 1, 2, 3 (C) 1, 3, 1, 3 (D) 2, 3, 2, 3 (E) 3, 4, 6, 2 10. Chemical equations must be balanced to satisfy the ___. (A) law of definite proportions (C) law of multiple proportions (B) law of conservation of mass (D) principle of Avogadro
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LPS Standard(s): --- State Standard(s): 12.3.3a Classification. Classify each of the examples or reactions as one of the following. NOTE: A
classification may be used more than once. Choices: (A) combustion reaction (D) single-replacement reaction (B) decomposition reaction (E) synthesis reaction (C) double-replacement reaction 11. Mg + 2HCl MgCl2 + H2 12. N2O5 + H2O 2HNO3
13. 2H2 + O2 2H2O 14. 2Fe(OH)3 Fe2O3 + 3H2O 15. 3KSCN + FeCl3 3KCl + Fe(SCN)3 16. CH4 + 2O2 CO2 + 2H2O 17. BaCl2 + K2CO3 BaCO3 + 2KCl 18. CaCO3 CaO + CO2 19. Cl2 + 2KI 2KCl + I2 20. 2NaCN + H2SO4 2HCN + Na2SO4
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Unit Four Practice Test (Short Answer/Calculation) PT – C – U4
(a) LPS Standard(s): --- State Standard(s): 12.3.3a,d (b) LPS Standard(s): --- State Standard(s): --- (c) LPS Standard(s): --- State Standard(s): 12.1.3d (d) LPS Standard(s): --- State Standard(s): 12.1.2d Calculations. Solve the following problems. Show work or receive no credit. Show proper units
and express all answers using correct significant digits and/or decimal places. For any molar masses or constants, use the values from the table following this problem or receive reduced credit.
21. As you notice the importance of steel in your life, you may also note that the production of steel
depends on available iron. The following balanced equation shows one of the overall reactions for the production of iron.
2 Fe2O3 + 3 C → 4 Fe + 3 CO2
(a) If a manufacturer began with 1.00 kg of Fe2O3, how many kg of carbon would be required
to fully react with the Fe2O3? (b) If a manufacturer began with 1.00 kg each of Fe2O3 and C, what would be the limiting
reagent? Explain. Also calculate the amount of excess reagent there would be. Limiting reagent: ________________ Amount of excess reagent: ___________ kg
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(c) How many kilograms of iron would be produced based on information in part (b)? (In other words, what is the theoretical yield of iron?)
Amount of Fe produced: _____________________ kg (d) If the actual amount of iron produced from the reaction was 0.500 kg Fe, what is the
percent yield of the iron? % yield of Fe: ______________%
Constants to Use
Molar volume of gas at STP
Avogadro’s Constant Compound/Element
Molar mass of compound/element
(in g/mol)
22.4 L/mol 6.022 1023
representative particles/mol
Carbon 12.011 Carbon dioxide (CO2) 44.009
Iron (Fe) 55.847 Iron (III) oxide (Fe2O3) 159.691
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LPS Standard(s): 12.2.6a State Standard(s): 12.1.1a, 12.3.3a, 12.1.2d Reactions. Write a balanced chemical equation for the following problems by predicting the correct
products, writing the equation with proper formulas and symbols for each of the following reactions, and including states of matter. If the reaction is not possible, circle the “Not Possible” phrase below the blank. FAILURE TO FOLLOW THESE INSTRUCTIONS WILL RESULT IN NO CREDIT.
22. Potassium metal reacts with chlorine gas to produce ... Balanced equation: ___________________________________________________________ Not possible 23. Aqueous solutions of aluminum chloride and sodium carbonate react to produce ... Balanced equation: ___________________________________________________________ Not possible 24. Metallic magnesium reacts with aqueous zinc sulfate to produce … Balanced equation: ___________________________________________________________ Not possible 25. Metallic silver reacts with aqueous sodium nitrate to produce … Balanced equation: ___________________________________________________________ Not possible
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Chemistry Term One Practice Test PT – C – T1
DO NOT WRITE ON THIS TEST. Use the scratch paper provided for any work. Multiple Choice. On the scantron sheet for each question, fill in the rectangle corresponding with
the upper-case letter of the answer that best completes or answers the statement or question. NOTE: If the rectangle is not completely filled in or not otherwise done correctly (i.e., done in pen, etc.), the answer may be considered incorrect and will not be checked by the teacher personally (all grading is done via the scantron machine).
1. Titanium is a(n) ___. (A) compound (B) element (C) heterogeneous mixture (D) homogeneous mixture 2. The air in a scuba tank is a ___ solution. (A) gas-gas (B) liquid-liquid (C) solid-solid (D) solid-gas
3. Which of the following cannot be classified as a substance? (A) carbon dioxide (B) stainless steel (C) hydrogen gas (D) iron 4. Which of the following is a chemical property of acetone? (A) colorless (C) flammable (B) liquid at room temperature (D) low melting point 5. Which of the following is a chemical change? (A) melting mercury (C) evaporating alcohol (B) detonating dynamite (D) freezing bromine 6. The ___ scale is the SI temperature scale that is used. (A) Celsius (B) Fahrenheit (C) Joule (D) Kelvin
7. Density is calculated by dividing ___. (A) weight by volume (C) volume by weight (B) mass by volume (D) volume by mass 8. If a liter of water increases in temperature from 20C to 60C, its density ___. (A) increases (B) decreases (C) stays the same 9. If a temperature changes by 53C, by how much does it change in K? (A) –273 K (B) 0 K (C) 53 K (D) 273.15 K 10. How many significant figures are there in the measurement 0.00540 kg? (A) 2 (B) 3 (C) 5 (D) 6 11. How many significant figures are there in the measurement 501000 mg? (A) 2 (B) 3 (C) 5 (D) 6 12. How many significant figures are there in the measurement 40500.0 mg? (A) 2 (B) 3 (C) 5 (D) 6 13. What is the measurement 222.0095 mm rounded off to five significant digits? (A) 222 mm (B) 222.0 mm (C) 222.00 mm (D) 222.01 mm (E) 222.001 mm
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14. Which are used when rounding for the answer to 86.6 + 85.43? (A) significant figures (B) decimal places (C) neither 15. Which are used when rounding for the answer to 2.56 8.982? (A) significant figures (B) decimal places (C) neither 16. Values of 52, 53, and 53 compared to an accepted value of 53 best demonstrate ___. (A) precision (B) accuracy (C) precision and accuracy 17. Values of 65, 65, and 66 compared to an accepted value of 21 best demonstrate ___. (A) precision (B) accuracy (C) precision and accuracy 18. Express the density of 5.6 g/cm3 in kg/m3. (A) 5.6 x 106 kg/m3 (C) 5.6 x 10–3 kg/m3 (B) 5.6 x 103 kg/m3 (D) 5.6 x 10–6 kg/m3
19. Express 60 m/s in km/hr. (A) 21,600 km/hr (C) 0.216 km/hr (B) 216 km/hr (D) 0.00216 km/hr
20. Which of the following ratios is a correct conversion factor to multiply to change meters to
centimeters?
(A) cm 100
m 1 (B)
cm 10
m 1 (C)
m 1
cm 10 (D)
m 1
cm 100
21. A conversion factor ___. (A) is never equal to one (B) is a ratio of equivalent measurements (C) changes the value of a measurement (D) can never be used to change one unit to another type of unit 22. Five kilometers is equal to ___ centimeters. (A) 5.0 x 10–3 (B) 5.0 x 105 (C) 5.0 x 10–5 (D) 5.0 x 103 23. 18 g/mol is equal to ___ kg/kmol. (A) 0.018 (B) 18 (C) 18000 (D) 18000000 24. The density of aluminum is 2.70 g/cm3. The mass of a cube of aluminum with a 1.0 cm3 volume
is ___ g. (A) 2.7 (B) 5.4 (C) 27 (D) 81 25. The lightest subatomic particle is the ___. (A) electron (B) neutron (C) proton 26. Robert Millikan discovered the charge of a(n) ___. (A) electron (B) nucleus (C) neutron (D) proton 27. A(n) ___ has a positive charge. (A) electron (B) neutron (C) proton 28. A(n) ___ has a negative charge. (A) electron (B) neutron (C) proton
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29. A(n) ___ has no charge. (A) electron (B) neutron (C) proton 30. Rutherford discovered the ___. (A) electron (B) nucleus (C) neutron (D) proton 31. Chadwick discovered the ___. (A) electron (B) nucleus (C) neutron (D) proton 32. Average atomic mass is based on a(n) ___ average. (A) normal (B) skewed (C) transitional (D) weighted Refer to the following isotope for questions 33 – 38.
Silver-108
33. How many neutrons does an atom of this isotope contain? (A) 47 (B) 61 (C) 107.87 (D) 108 (E) 155 34. How many electrons does a neutral atom of this isotope contain? (A) 47 (B) 61 (C) 107.87 (D) 108 (E) 155 35. How many protons does an atom of this isotope contain? (A) 47 (B) 61 (C) 107.87 (D) 108 (E) 155 36. What is the mass number of this isotope? (A) 47 (B) 61 (C) 107.87 (D) 108 (E) 155 37. What is the atomic number of this isotope? (A) 47 (B) 61 (C) 107.87 (D) 108 (E) 155 38. How can this isotope be expressed?
(A) Ag107.8747 (B) Ag108
47 (C) Ag6147 (D) Ag47
107.87
39. Isotopes of the same element have different ___. (A) numbers of protons (C) atomic numbers (B) numbers of neutrons (D) chemical behavior 40. Who was responsible for creating the first periodic table of elements? (A) Louis Pasteur (C) Dmitri Mendeleev (B) Henry Moseley (D) John Dalton 41. Who was responsible for creating the current periodic table of elements? (A) Louis Pasteur (C) Dmitri Mendeleev (B) Henry Moseley (D) John Dalton 42. Group 1A of the periodic table of elements contains the ___. (A) alkali metals (B) alkaline earth metals (C) halogens (D) noble gases 43. Group 8A of the periodic table of elements contains the ___. (A) alkali metals (B) alkaline earth metals (C) halogens (D) noble gases
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44. Which element is a transition metal? (A) silicon (B) silver (C) sodium (D) sulfur 45. Which element is a representative element? (A) cadmium (B) cerium (C) cesium (D) chromium 46. Which of the following formulas represents an ionic compound? (A) N2O (B) NH4Cl (C) NH3 (D) NCl3
47. Of the following choices, the elements that would most likely make an ionic compound when
combined are ___. (A) nitrogen and oxygen (C) sodium and calcium (B) sulfur and chlorine (D) lithium and sulfur 48. How many valence electrons does a neutral atom of sodium have? (A) 0 (B) 1 (C) 7 (D) 8
49. How many valence electrons does a sodium ion have? (A) 0 (B) 1 (C) 7 (D) 8 50. How many valence electrons does an atom of any alkaline earth metal have? (A) 0 (B) 1 (C) 2 (D) 7 51. What is the charge of an ion having 10 protons and 12 electrons? (A) 1+ (B) 2+ (C) 2– (D) 1– 52. What is the formula of the sulfide ion? (A) S1+ (B) S2+ (C) S1– (D) S2– 53. What is the formula of the strontium ion? (A) Sr1+ (B) Sr2+ (C) Sr1– (D) Sr2– 54. Which of the following elements does not exist as a diatomic molecule? (A) Hydrogen (B) Helium (C) Fluorine (D) Oxygen 55. What is the net charge of the compound with formula CaCl2? (A) 2+ (B) 1+ (C) 0 (D) 1– (E) 2–
56. Ions do NOT form when atoms ___. (A) have a charge (B) lose or gain electrons (C) lose or gain protons 57. What is the formula for nitric acid? (A) HNO2 (B) HNO3 (C) H3N (D) HN 58. What is the formula for lead (IV) chloride? (A) PbCl (B) PbCl2 (C) Pb4Cl (D) PbCl4 59. What is the formula for chromium (III) oxide? (A) CrO (B) Cr3O2 (C) Cr2O3 (D) Cr3O 60. Which of the following is the name for the compound with formula CrO3? (A) chromium (II) oxide (C) chromium (VI) oxide (B) chromium (III) oxide (D) dichromium trioxide
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61. Which of the following is the name for the compound with formula H2SO4? (A) Hydrosulfic acid (C) Sulfuric acid (B) Hydrosulfuric acid (D) Sulfurous acid 62. What of the following is the name for the compound with formula Sn3(PO4)4? (A) tritin diphosphate (C) tin (III) phosphate (B) tin (II) phosphate (D) tin (IV) phosphate 63. Which set of chemical name and chemical formula for the same compound is correct? (A) diiron trioxide, Fe2O3 (C) tin bromide, SnBr4 (B) magnesium chloride, MgCl2 (D) potassium chloride, PCl
64. Which set of chemical name and chemical formula for the same compound is correct? (A) ammonium sulfide, (NH4)2SO4 (C) copper chloride, CuCl (B) iron (III) phosphate, Fe3(PO4)2 (D) chromium (III) sulfide, Cr2S3 65. What is the formula for ammonium phosphate? (A) NH4PO4 (B) NH4P (C) (NH4)3P (D) (NH4)3PO4 66. The Roman numeral in chromium (VI) nitrate indicates the ___. (A) group number on the periodic table (B) positive charge on the chromium ion (C) number of chromium ions in the formula (D) number of nitrate ions in the formula 67. When the chemical equation H2O2 → H2O + O2 is balanced, the coefficient of H2O2 is ___. (A) 1 (B) 2 (C) 3 (D) 4 For questions 68 – 77, classify the reaction as one of the following: (A) synthesis/combination (D) double replacement (B) decomposition (E) combustion (C) single replacement 68. Cl2 + 2KBr 2KCl + Br2
69. C10H8 + 12O2 10CO2 + 4H2O 70. NH4NO3 N2O + 2H2O
71. 4Fe + 3O2 2Fe2O3 72. Hg(NO3)2 + 2NH4SCN Hg(SCN)2 + 2NH4NO3 73. 2Al + 3CuSO4 Al2(SO4)3 + 3Cu 74. 2Al(OH)3 Al2O3 + 3H2O 75. C4H8 + 6O2 4CO2 + 4H2O
76. 2Mg + O2 2MgO 77. 3Ag2SO4 + 2AlCl3 6AgCl + Al2(SO4)3
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78. All of the following are formulas for ionic compounds except ___. (A) CO2 (B) NH4Cl (C) Al(NO3)3 (D) CaCl2 (E) Al2O3 79. All of the following are formulas for molecular compounds except ___. (A) N2O (B) CH4 (C) SO3 (D) NH4OH (E) NO 80. If the density of a gas is 1.696 g/L at standard temperature and pressure, that gas is ___. (A) H2 (B) He (C) Ne (D) F2 (E) SO3 81. How many moles of helium atoms are there in 500 L of helium gas at standard temperature and
pressure? (A) 0.05 mol (B) 0.2 mol (C) 20 mol (D) 90 mol (E) 10000 mol 82. Which of the following gas samples would have the largest number of representative particles at
standard temperature and pressure? (A) 4.8 L H2 (B) 3.7 L N2 (C) 5.6 L Ne (D) 0.78 L SO2 (E) 0.64 L XeF2 83. Which of the following gas samples would have the smallest number of representative particles at
standard temperature and pressure? (A) 0.5 L Cl2 (B) 1.0 L O2 (C) 2.0 L CO2 (D) 3.0 L F2 (E) 4.0 L C2H6 84. Standard pressure is ___. (A) 0 K (B) 0C (C) 0 atm (D) 1 atm 85. Which of the following reactions will not take place spontaneously in the direction written? (A) Li + H2SO4 (D) Zn + HCl (B) Mg + Zn(NO3)2 (E) Fe + H3PO4 (C) Au + Li2SO4 86. Use the activity series of metals to write a balanced chemical equation for the reaction of solid calcium and silver nitrate solution. (A) Ca(s) + AgNO3(aq) CaNO3(aq) + Ag(s) (B) Ca(s) + 2AgNO3(aq) Ca(NO3)2(aq) + 2Ag(s) (C) Ca(s) + Ag(NO3)2(aq) Ca(NO3)2(aq) + Ag(s) (D) Ca(s) + 3AgNO3(aq) Ca(NO3)3(aq) + 3Ag(s) (E) No reaction takes place because silver is less reactive than potassium. 87. How many liters of chlorine gas can be produced when 0.98 L of HCl react with excess oxygen
gas, at STP? 4HCl(g) + O2(g) 2Cl2(g) + 2H2O(g) (A) 0.98 L (B) 0.49 L (C) 3.9 L (D) 2.0 L (E) 0.25 L 88. The equation below shows the decomposition of lead (II) nitrate. How many grams of oxygen gas
are produced when 11.5 grams of nitrogen dioxide are formed? 2Pb(NO3)2(s) 2PbO(s) + 4NO2(g) + O2(g) (A) 1.00 g (B) 2.00 g (C) 2.88 g (D) 32.0 g (E) 46.0 g 89. Glucose, C6H12O6, is a good source of food energy. When it reacts with oxygen gas, carbon
dioxide and water are formed. How many liters of carbon dioxide are produced when 126 grams of glucose completely reacts with oxygen? Reaction: C6H12O6 + 6O2 6CO2 + 6H2O
(A) 4.21 L (B) 5.33 L (C) 15.7 L (D) 94.1 L (E) 185 L
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90. Phosphorous trichloride, PCl3, is a commercially important compound used in the manufacture of pesticides, gasoline additives, and a number of other products. It is made by the direct combination of phosphorous and chlorine. How many total moles of PCl3 can be produced from the reaction of 125 g of Cl2 with 125 g of P4? The reaction is P4 + 6Cl2 4PCl3.
(A) 1.01 mol (B) 1.18 mol (C) 1.76 mol (D) 2.64 mol (E) 4.04 mol 91. How many liters of hydrogen gas are needed to react with CS2 to produce 2.50 L of CH4 at
standard temperature and pressure? The reaction is 4H2(g) + CS2(l) CH4(g) + 2H2S(g). (A) 2.50 L (B) 0.625 L (C) 5.00 L (D) 7.50 L (E) 10.0 L 92. Identify the limiting reagent and the volume of product formed when 11 L CS2 reacts with 11 L O2
to produce CO2 and SO2 at standard temperature and pressure. The reaction is CS2 + 3O2 CO2 + 2SO2 (A) Limiting reagent: CS2; 11 L CO2 (D) Limiting reagent: O2; 3.7 L CO2
(B) Limiting reagent: O2; 11 L CO2 (E) Limiting reagent: O2; 7.3 L CO2 (C) Limiting reagent: CS2; 3.7 L CO2 93. Methane and hydrogen sulfide form when hydrogen reacts with carbon disulfide. Identify the
excess reagent and calculate how much remains after 36 L of H2 reacts with 12 L of CS2 at standard temperature and pressure. The reaction is 4H2 + CS2 CH4 + 2H2S.
(A) 3 L CS2 (B) 6 L CS2 (C) 9 L CS2 (D) 12 L H2 (E) 24 L H2 94. The thermite reaction has been used for welding railroad rails, in incendiary bombs, and to ignite
solid-fuel rocket motors. If 80.0 grams of Fe2O3 react with 60.0 grams of Al, how much iron will be produced, in grams? The reaction is Fe2O3 + 2Al 2Fe + Al2O3.
(A) 28.0 g (B) 51.1 g (C) 56.0 g (D) 112 g (E) 166 g 95. For a given chemical reaction, the actual yield is ___ greater than the theoretical yield. (A) sometimes (B) always (C) never 96. The reagent present in the largest amount is ___ the limiting reagent. (A) sometimes (B) always (C) never 97. 6.022 1023 representative particles is equal to one ___. (A) kilogram (B) gram (C) liter (D) Kelvin (E) mole 98. How many grams are in 5.90 mol C8H18? (A) 0.0512 g (B) 19.4 g (C) 389 g (D) 673 g (E) 3.55 1024 g 99. What is the volume, in liters, of 6.8 mol of Kr gas at STP? (A) 0.30 L (B) 3.3 L (C) 25 L (D) 150 L (E) 13000 L 100. What is the molar mass of Cr2(SO4)3? (A) 148.1 g (B) 200.0 g (C) 288.0 g (D) 344.2 g (E) 392.2 g
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Appendix F Practice Tests
Unit One Practice Test Key (Multiple Choice) PTK – C – U1
1. A 11. B 21. A 31. C 41. C 2. B 12. B 22. C 32. C 42. D 3. A 13. A 23. D 33. B 43. B 4. A 14. B 24. B 34. B 44. D 5. B 15. C 25. C 35. C 45. A 6. B 16. D 26. C 36. D 46. C 7. B 17. A 27. D 37. A 47. B 8. A 18. C 28. E 38. E 48. A 9. A 19. A 29. A 39. D 49. D 10. B 20. B 30. D 40. B 50. D
Unit One Practice Test Key (Short Answer/Calculation)
PTK – C – U1 51. 52.
Unit Two Practice Test Key (Multiple Choice) PTK – C – U2
1. B 6. B 11. B 16. B 21. E 26. B 31. B 36. C 41. C 46. E 2. A 7. A 12. C 17. A 22. D 27. A 32. D 37. C 42. E 47. C 3. D 8. D 13. B 18. D 23. C 28. C 33. D 38. C 43. B 48. A 4. A 9. A 14. B 19. B 24. D 29. D 34. B 39. C 44. A 49. D 5. A 10. A 15. D 20. C 25. C 30. E 35. B 40. D 45. D 50. B
Unit Two Practice Test Key (Short Answer/Calculation) PTK – C – U2
51. (0.0431)(50.000) + (0.8376)(52.000) + (0.0955)(53.000) + (0.0238)(54.000) = 2.16 + 43.56 + 5.06 + 1.29 = 52.07 amu (NOTE: Different than periodic table because of slight changes in values provided.)
333
33
33
33
3 lb/ft 62ft. 1
in. 12
in. 1
cm 54.2
g 454
lb. 1
cm 1
g 0.1
m/s 693s 60
min 1
min 60
hr 1
km 1
m 1000
miles 0.621
km 1
hr 1
miles 1550
page A-49 – C – T1 – BOOK
Unit Three Practice Test Key (Multiple Choice) PTK – C – U3
1. C 6. D 11. D 16. A 21. B 26. B 2. D 7. D 12. D 17. A 22. D 27. C 3. B 8. C 13. A 18. C 23. A 28. B 4. D 9. E 14. E 19. A 24. C 29. C 5. D 10. C 15. D 20. E 25. D 30. C
Unit Three Practice Test Key (Short Answer/Calculation) PTK – C – U3
31. Since representative particles are involved (molecules), use 6.022 x 1023 molecules/mol. Since mass is involved, use molar mass. Molar mass C9H8O4 = 9C + 8H + 4O = 9(12.011) + 8(1.0079) + 4(15.999) = 180.158 g/mol
48921
489
48923
489
489489 OHC molecules 10 3.3OHC mol 1
OHC molecules 10 022.6
OHC g 158.180
OHC mol 1
1
OHC g 0.1
32. Since representative particles are involved (molecules), use 6.022 x 1023 molecules/mol. Since volume at STP is involved (1 atmosphere and 0C), use 22.4 L/mol.
25
323
7.85 10 molecules 1 mol 22.4 L2.92 10 L
1 1 mol6.022 10 molecules
Unit Four Practice Test Key (Multiple Choice) PTK – C – U4
1. D 5. B 9. A 13. E 17. C 2. A 6. E 10. B 14. B 18. B 3. C 7. B 11. D 15. C 19. D 4. B 8. C 12. E 16. A 20. C
Unit Four Practice Test Key (Short Answer/Calculation) PTK – C – U4
21. (a)
(b) Since you have 1.00 kg of carbon but only need 0.113 kg of carbon, carbon is the
excess reagent, meaning Fe2O3 is the limiting reagent. Excess reagent = 1.00 kg C – 0.113 kg C = 0.887 kg C = 0.89 kg C
C kg 113.0C g 0001
C kg 1
C mol 1C g 011.12
OFe mol 2C mol 3
OFe g 691.159OFe mol 1
OFe kg 1OFe g 0001
1OFe kg .001
3232
32
32
3232
page A-50 – C – T1 – BOOK
(c) (d) 22. 2K(s) + Cl2(g) 2KCl(s) 23. 2AlCl3(aq) + 3Na2CO3(aq) 6NaCl(aq) + Al2(CO3)3(s) 24. Mg(s) + ZnSO4(aq) Zn(s) + MgSO4(aq)
25. Not possible
Chemistry Term One Practice Test PT – C – T1
1. B 11. B 21. B 31. C 41. B 51. C 61. C 71. A 81. C 91. E 2. A 12. D 22. B 32. D 42. A 52. D 62. D 72. D 82. A 92. D 3. B 13. D 23. B 33. B 43. D 53. B 63. B 73. C 83. A 93. A 4. C 14. B 24. A 34. A 44. B 54. B 64. D 74. B 84. D 94. C 5. B 15. A 25. A 35. A 45. C 55. C 65. D 75. E 85. C 95. C 6. D 16. C 26. A 36. D 46. B 56. C 66. B 76. A 86. B 96. A 7. B 17. A 27. C 37. A 47. D 57. B 67. B 77. D 87. B 97. E 8. B 18. B 28. A 38. B 48. B 58. D 68. C 78. A 88. B 98. D 9. C 19. B 29. B 39. B 49. D 59. C 69. E 79. D 89. D 99. D 10. B 20. D 30. B 40. C 50. C 60. C 70. B 80. D 90. B 100. E
Fe kg 699.0Fe g 0001
Fe kg 1
Fe mol 1
Fe g 847.55
OFe mol 2
Fe mol 4
OFe g 691.159
OFe mol 1
OFe kg 1
OFe g 0001
1
OFe kg .001
3232
32
32
3232
yield%5.71%100kg 699.0
kg 0.500%100
yieldlTheoretica
yieldActual yieldPercent
