Differentiated Chemistry Worksheet and Laboratory

Transcription

Differentiated
Chemistry
Worksheet and
Laboratory
Manual (Term 2)
Mr. Geist
21
22
23
24
25
26
27
(260)
(226)
(223)
Lr
Lawrencium
Barium
137.33
88
Cesium
132.91
87
Ra
Lutetium
174.97
103
56
Ba
55
Cs
Radium
71
Lu
87.62
Fr
Yttrium
88.906
Strontium
Rubidium
85.468
Francium
40
Y
Vanadium
(262)
Tungsten
(263)
Seaborgium
Sg
183.85
106
Cerium
140.12
Lanthanum
138.91
90
Th
Thorium
232.04
89
Ac
Actinium
(227)
59
Pr
231.04
Proactinium
Pa
91
140.91
Praseodymium
 Actinide series
58
Ce
57
La
 Lanthanide series
(261)
Dubnium
Db
Tantalum
W
74
Molybdenum
73
Ta
180.95
105
Rf
42
Mo
95.94
Hafnium
Rutherfordium
Chromium
51.996
92.906
Niobium
Nb
41
50.941
178.49
104
Hf
72
91.22
Zirconium
Zr
39
38
Sr
37
Rb
Titanium
47.90
Scandium
44.956
Calcium
40.08
Potassium
39.098
Iron
238.03
Uranium
U
237.05
Neptunium
Np
93
(145)
92
Promethium
144.24
Pm
61
(265)
Hassium
Hs
190.2
108
Osmium
Os
76
101.07
Ruthenium
Ru
44
55.847
Neodymium
Nd
60
(262)
Bohrium
Bh
186.21
107
Rhenium
Re
75
(97)
Technetium
Tc
43
54.938
Manganese
(244)
Plutonium
Pu
94
150.4
Samarium
Sm
62
(266)
Meitnerium
Mt
192.22
109
Iridium
Ir
77
102.91
Rhodium
Rh
45
58.933
Cobalt
Co
28
29
79
Platinum
(243)
Americium
Am
95
151.96
Europium
Eu
63
(269)
Weird
Uum
195.09
110
(247)
Curium
Cm
96
157.25
Gadolinium
Gd
64
(272)
More weird
Uuu
196.97
111
Gold
Au
78
Pt
107.87
Silver
Ag
47
63.546
Copper
Cu
106.4
Palladium
Pd
46
58.71
Nickel
Ni
30
(247)
Berkelium
Bk
97
158.93
Terbium
Tb
65
(277)
Most weird
Uub
200.59
112
Mercury
Hg
80
112.41
Cadmium
Cd
48
65.38
Zinc
Zn
K
Fe
20
Ca
19
Mn
24.305
Cr
Magnesium
Sodium
22.990
V
Aluminum
12
Mg
11
Na
Ti
13
Al
9.0122
6.941
Sc
10.81
Beryllium
Lithium
5
(251)
Californium
Cf
98
162.50
Dysprosium
Dy
66
204.37
Thallium
Tl
81
114.82
Indium
In
49
69.72
Gallium
Ga
31
26.982
Boron
B
4
Be
Li
3A
1.0079
3
2A
51
Tin
(254)
Einsteinium
Es
99
164.93
Holmium
Ho
67
207.2
Lead
Pb
82
118.69
(257)
Fermium
Fm
100
167.26
Erbium
Er
68
208.98
Bismuth
Bi
83
121.75
Antimony
Sb
50
Sn
74.922
Arsenic
As
33
30.974
Phosphorus
P
15
14.007
Nitrogen
N
7
5A
72.59
Germanium
Ge
32
28.086
Silicon
Si
14
12.011
Carbon
C
6
4A
(258)
Mendelevium
Md
101
168.93
Thullium
Tm
69
(209)
Polonium
Po
84
127.60
Tellurium
Te
52
78.96
Selenium
Se
34
32.06
Sulfur
S
16
15.999
Oxygen
O
8
6A
(259)
Nobelium
No
102
173.04
Ytterbium
Yb
70
(210)
Astatine
At
85
126.90
Iodine
I
53
79.904
Bromine
Br
35
35.453
Chlorine
Cl
17
18.998
Fluorine
F
9
7A
(222)
Radon
Rn
86
131.30
Xenon
Xe
54
83.80
Krypton
Kr
36
39.948
Argon
Ar
18
20.179
Neon
Ne
4.0026
10
Helium
2
He
H
Hydrogen
8A
1
Do not white-out, add additional paper, or
tape. Only write in box to the left, or be
unable to use this sheet on the test.
Exam: MASTER COPY Period: ________
Name: _____________________________
1A
You may add additional information in your own handwriting in this box.
Periodic Table of Elements (Additional Values and Constants on back page)
1A
H
2.1
Li
1.0
Na
0.9
K
0.8
Rb
0.8
Cs
0.7
Fr
0.7
3B
Ti
1.5
Zr
1.4
Hf
1.3
Th
1.2
4B
V
1.6
Nb
1.6
Ta
1.5
Pa
1.5
5B
Cr
1.6
Mo
1.8
W
1.7
U
1.7
6B
Mn
1.5
Tc
1.9
Re
1.9
7B
Fe
1.8
Ru
2.2
Os
2.2
8B
Co
1.9
Rh
2.2
Ir
2.2
Ni
1.9
Pd
2.2
Pt
2.2
nitrate
sulfate
phosphate
nitrite
sulfite
phosphite
carbonate
acetate
hydroxide
ammonium
silicate
cyanide
permanganate
chromate
dichromate
1B
Zn
1.6
Cd
1.7
Hg
1.9
2B
B
2.0
Al
1.5
Ga
1.6
In
1.7
Tl
1.8
3A
C
2.5
Si
1.8
Ge
1.8
Sn
1.8
Pb
1.9
4A
N
3.0
P
2.1
As
2.0
Sb
1.9
Bi
1.9
5A
O
3.5
S
2.5
Se
2.4
Te
2.1
Po
2.0
6A
F
4.0
Cl
3.0
Br
2.8
I
2.5
At
2.2
7A
Kb = 0.512C/molal
Kf = –1.86C/molal
8A
He
-Ne
-Ar
-Kr
-Xe
-Rn
--
Fluorine
Chlorine
Bromine
Iodine
Colligative Property Constants
for Water
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Tin
Lead
Hydrogen
Copper
Mercury
Silver
Platinum
Gold
Activity Series of Metals/Halogens
(NOTE: Reactivity of the
metal/halogen decreases as it gets
lower on the list.)
Cu
1.9
Ag
1.9
Au
2.4
Average Electronegativities of the Elements
2A
Sc
1.3
Y
1.2
Ln
1.0
Ac
1.0
–
NO3 :
2–
SO4 :
PO43–:
–
NO2 :
SO32–:
3–
PO3 :
CO32–:
–
C2H3O2 :
OH–:
+
NH :
4
SiO32–:
–
CN :
MnO4–:
2–
CrO4 :
2–
Cr2O7 :
Main Polyatomic Ions
Be
1.5
Mg
1.2
Ca
1.0
Sr
1.0
Ba
0.9
Ra
0.9
Other Polyatomic Ions
HPO42–: hydrogen
phosphate
H2PO41–: dihydrogen
phosphate
HSO31–: hydrogen
sulfite
HSO41–: hydrogen
sulfate
HCO31–: hydrogen
carbonate
ClO41–: perchlorate
1–
ClO3 : chlorate
ClO21–: chlorite
1–
ClO : hypochlorite
2–
C2O4 : oxalate
Monatomic Ions
Cu1+: copper (I) ion
Cu2+: copper (II) ion
2+
Fe : iron (II) ion
Fe3+: iron (III) ion
2+
Pb : lead (II) ion
Pb4+: lead (IV) ion
2+
Sn : tin (II) ion
Sn4+: tin (IV) ion
2+
Co : cobalt (II) ion
Co3+: cobalt (III) ion
Useful Conversion Factors
And Conversions
Energy:
1 cal = 4.184 J
Length:
1 angstrom = 0.100 nm
1 inch = 2.54 cm
Mass:
1 lb = 0.4536 kg
Pressure: 1 atm = 101.3 kPa
1 atm = 760 mm Hg
C = K – 273.15
1 L = 0.001 m3
1 cm3 = 1 mL
Temp.:
Volume:
General Physical Constants
6.022 x 1023 rp/mol
6.626 x 10–34 Js
1.381 x 10–23 J/K
3.0 x 108 m/s
1.6605655 x 10–27 kg
1.097 x 107 m–1
96485.309 C/mol
8.31 (LkPa)/(Kmol)
0.0821 (Latm)/(Kmol)
62.396 (Lmm Hg)/(Kmol)
All compounds formed with the negative
ion are insoluble except those of the
alkali metals and NH4+.
ion are soluble except Ag+, Pb2+, Hg22+,
+
and Cu .
Most compounds formed with the
negative ion are soluble; exceptions
include SrSO4, BaSO4, CaSO4, RaSO4,
Ag2SO4, and PbSO4.
ion are soluble.
Solubility Rules
Rule
Avogadro’s Constant
Planck’s Constant
Boltzmann’s Constant
Speed of Light
Atomic Mass Unit
Rydberg’s Constant
Faraday’s Constant
Ideal Gas Constant
–
Negative Ion
NO3
2–
–
–
–
I , Br , Cl
SO4
CO32–, PO43–,
2–
SO3
OH–
2–
S
alkali metals, NH4+, Sr2+, and Ba2+.
(Ca(OH)2 is slightly soluble.)
alkali metals, alkaline earth metals, and
NH4+.
Table of Contents
Unit Five Worksheet ..................................................................................... 1
Unit Six Worksheet..................................................................................... 16
Unit Seven Worksheet ............................................................................... 21
Unit Eight Worksheet ................................................................................. 38
Unit Five Experiment – 1: Orbital Structures and Identifications ............... 45
Unit Five Experiment – 2: Metal Ions and Flame Tests ............................ 49
Unit Five Experiment – 3: Viewing Spectra ............................................... 50
Unit Six Experiment – 1: Paper Chromatography of Food Dyes ............... 52
Unit Seven Experiment – 1: The Ideal Gas Law ....................................... 54
Unit Seven Experiment – 2: Specific Heat Experiment ............................. 56
Unit Seven Experiment – 3: Solution Preparation and Rate of Reaction .. 59
Unit Seven Experiment – 4: Le Chatelier’s Principle
(Physical and Chemical Changes) ............... 61
Unit Eight Experiment – 1: Using Indicators .............................................. 63
Unit Eight Experiment – 2: Determining Molarity by Neutralization ........... 66
Appendix A – Laboratory Equipment and LPS Safety Contract ............... A-1
Appendix B – SI Units and Conversions .................................................. A-5
Appendix C – Compound Name and Formula Writing ............................. A-8
Appendix D – Chemical Reactions and Quantities ................................ A-10
Appendix E – Study Skills ...................................................................... A-17
Appendix F – Electron Configuration Rules ........................................... A-20
Appendix G – Electron Dot Structure ..................................................... A-22
Appendix H – VSEPR Models ................................................................ A-24
Appendix I – Calorimetry Calculations ................................................... A-26
Appendix J – Water and Solutions ......................................................... A-28
Appendix K – Acid and Base Measurements ......................................... A-32
Appendix L – Acid and Base Notes........................................................ A-33
Appendix M – Neutralization Notes ........................................................ A-36
Appendix N – Practice Tests .................................................................. A-40
Appendix O – Practice Test Keys .......................................................... A-66
Unit Five Worksheet
WS – DC – U5
Section 8.1
Short Answer. Answer the following questions.
1.
What evidence showed that the particles in the beam of Crookes’s tube were negatively charged?
2.
Suppose that two beams pass between a pair of oppositely charged plates. One of the beams is
composed of electrons, and the other is composed of protons. Will the two beams bend in the
same direction or in opposite directions? Explain.
3.
What main feature of Dalton’s atomic model was abandoned after Thomson’s discoveries?
4.
Make a diagram of a lithium atom, based on Thomson’s atomic model.
5.
If a lithium atom lost one electron, forming a positive ion, how would the diagram drawn in question
4 be changed?
page 1 – DC – T2 – BOOK
Section 8.2
Identification. Identify which of the three types of radiation – alpha, beta, and/or gamma – each of the
following describes.
___________6.
Is not deflected by a magnet
___________10.
Consists of ions
___________7.
Has a negative charge
___________11.
Is similar to light rays
___________8.
Moves with the greatest speed
___________12.
Consists of the same particles as
cathode rays
___________9.
Has the highest penetrating
ability
___________13.
Has the lowest penetrating ability
Section 8.3
14.
Identify two differences between protons and electrons.
15.
Which subatomic particle determines the identity of a specific element? What term defines this?
16.
Explain what an isotope is.
17.
What ultimately determines the instability of an isotope?
Table Completion.
Information
Atomic
Number
236
92 U
142
56 Ba
Krypton-91
Iodine-131
27 3+
13 Al
127 
52 I
Fill in the table using the information provided in the left-most column to identify the
following properties of each isotope.
Mass
Number
Number of
Neutrons
Number of
Electrons
Number of
Protons
18.
24.
30.
36.
42.
19.
25.
31.
37.
43.
20.
26.
32.
38.
44.
21.
27.
33.
39.
45.
22.
28.
34.
40.
46.
23.
29.
35.
41.
47.
Calculation. Show work or receive no credit. Express proper units and correct number of significant
figures and decimal places.
48.
The element oxygen contains three naturally occurring isotopes:
16
8O
17
8O
18
8O
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.
49.
The element nitrogen contains two naturally occurring isotopes:
14
7N
15
7N
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.
Chapter 8 General Questions
50.
Elements in the periodic table are ordered according to increasing atomic number rather than
increasing atomic mass. There are several places where the atomic number increases but the
average atomic mass decreases. Identify two of these places, and explain why these exceptions
occur.
51.
Assume the nucleus of a fluorine atom is a sphere with a radius of 5 x 10–13 cm. Calculate the
density of a fluorine nucleus. Compare this density with the atomic density of iridium, whose
density is 22.6 g/cm3. (HINT: Recall how to find the volume of a sphere.)
Section 9.1
Short Answer/Calculation. Answer the following questions and problems.
Write nuclear equations for the following processes.
52.
The alpha decay of polonium-218
53.
The beta decay of lead-210
54.
The alpha decay of americium-241
55.
The beta decay of carbon-14
Write nuclear equations for the following radioactive processes.
56.
alpha decay of francium-208
57.
electron capture by beryllium -7
58.
beta emission by argon-37
59.
positron emission by fluorine-17
60.
After 42 days, a 2.0 g sample of phosphorus-32 contains only 0.25 g of isotope. What is the halflife of phosphorus-32? Show work or receive no credit. Include proper units.
61.
The mass of cobalt-60 in a sample is found to have decreased from 0.800 g to 0.200 g in a period
of 10.5 years. Find the half-life of cobalt-60 and calculate how many years it will take for 0.200 g of
cobalt-60 to decrease in the sample to 0.149 g.
62.
What happens to the mass number and atomic number of an atom that undergoes beta decay?
63.
A radioisotope of an element undergoes alpha particle decay. How do the atomic number and
mass number of the particle change?
64.
Bismuth-211 is a radioisotope. It decays by alpha emission to yield another radioisotope, which
emits beta radiation as it decays to a more stable isotope. Write equations for the nuclear
reactions and name the decay products.
65.
A sample initially contains 70.0 g of an isotope of radon. After 6.6 days, the sample only contains
21.0 g radon. What is the half-life of this isotope of radon, and after how many more days will only
9.5 g radon remain?
Give the composition of the nucleus of the following isotopes.
64
28
136
53
Ni
I
Gold-195
66.
p+: ________
68.
p+: ________
70.
p+: ________
67.
n0: ________
69.
n0: ________
71.
n0: ________
Section 9.2
Short Answer. Complete the equations for the following transmutation reactions.
Li + 01n  42 He + ________
72.
6
3
73.
235
92
74.
27
13
75.
235
92
76.
________ + 01n 
U + 01n 
141
56
Ba + ________ + 3 01n
Al + 42 He  01n + ________
U 
90
38
Sr + ________ + 01n + 4-10e
144
58
Ce +
90
38
Sr + 601n + 2-10e
Section 9.3
77.
Identify two types of nuclear waste produced by nuclear power plants.
78.
Assuming technical problems could be overcome, what are some advantages to producing
electricity in a fusion reactor?
79.
Describe how a nuclear fission power plant operates.
80.
Why are spent fuel rods removed from a reactor core?
81.
What do spent fuel rods contain?
82.
What happens to spent fuel rods after they are removed?
83.
The fission energy of uranium-235 is 2.0 x 107 kcal/g. The heat of combustion of coal is about 8.0
kcal/g. Approximately what mass of coal must be burned to produce the energy released by the
fission of 1.0 gram of uranium-235?
Section 10.1
Calculations.
Solve the following problems. Show work. Include proper units and significant figures.
84.
What is the wavelength of the radiation whose frequency is 5.00 x 1015 s-1?
85.
An inexpensive laser that is available to the public emits light that has a wavelength of 670 nm.
What is the frequency, in hertz, of the radiation?
86.
What is the energy of two moles of photons traveling with a frequency of 2.22 x 1014 s-1?
87.
What is the frequency, in hertz, of a photon whose energy is 6.00 x 10-15 J?
88.
Suppose that an AM radio station broadcasts at a frequency of 1600 kHz. What is the wavelength
in meters of the radiation from the station?
89.
What is the energy of a mole of photons whose wavelength is 658 nm?
90.
Talking on a cell phone is possible because of the electromagnetic signals sent and received.
What is the energy of one photon in a signal that is sent at 850 MHz?
91.
What is the energy, in kilojoules, of a one mole of photons from question 87?
We spend quite a bit of time staring at the red and green lights in a traffic signals. While that image is in
your mind, for questions 92 – 95, decide which would have the following for a red light with a
4.41 x 1014 Hz frequency or a green light with a 6.00 x 1014 Hz frequency. Explain.
92.
Longer wavelength
93.
Greater speed
94.
Greater energy
95.
Greater amplitude
96.
Explain the statement “Water waves transmit energy, not matter.”
Section 10.2
97.
What is a bright-line spectrum?
98.
Explain why the bright-line spectrum of hydrogen is composed of discrete lines and is not a
continuous spectrum.
99.
Explain what happens when the electron of a hydrogen atom changes from a 2s orbital to a 5s
orbital.
100.
Why is it not possible for two different orbitals to have the same first three quantum numbers?
101.
The following four quantum numbers have been supplied for a bound electron. However, one of
the values is incorrect. Explain which value is wrong and correct the error.
n = 4; l = 3; ml = 4; ms = –1/2
102.
The following graphs represent electron probabilities versus average distance from the nucleus.
Shown here are the graphs and figures that represent (not necessarily in order) 1s, 2s and 2p
orbitals. Match the representations with their correct identities:
(a)
(b)
(i)
1s is shown in choice
a
b
c
(ii)
2s is shown in choice
a
b
c
(iii)
2p is shown in choice
a
b
c
(c)
How many orbitals are in each of the following sublevels?
_________103. 4p sublevel
_________105. 4f sublevel
_________104. 3d sublevel
_________106. 2s sublevel
Section 10.3
Write the complete (not abbreviated) electron configurations for the following elements.
107.
Sulfur: ______________________________________________________________________
108.
Potassium: ___________________________________________________________________
109.
Vanadium: ___________________________________________________________________
110.
Argon: ______________________________________________________________________
111.
Iron: ________________________________________________________________________
112.
Sodium: _____________________________________________________________________
113.
Chromium: ___________________________________________________________________
114.
Iodine: _______________________________________________________________________
115.
Calcium: ______________________________________________________________________
116.
Platinum: _____________________________________________________________________
117.
Tellurium: _____________________________________________________________________
118.
Radium: ______________________________________________________________________
119.
Sr2+: _________________________________________________________________________
120.
Br–: __________________________________________________________________________
121.
P3–: __________________________________________________________________________
122.
O2–: __________________________________________________________________________
123.
Rb+: __________________________________________________________________________
124.
Al3+
Identify the elements described below.
125.
Contains a full third energy level
126.
Contains the first p electron
127.
Element that is a noble gas and has its outermost electrons in the fourth energy level
128.
Has an electron configuration of 1s22s22p63s23p4
129.
Has an electron configuration of 1s22s22p63s23p64s23d104p65s14d5
130.
The 3rd period element with a complete outermost energy level
131.
The 5th period element containing only 1 s electron and no d electrons in the outermost energy
level
132.
The 5th period element containing only 3 4d electrons
133.
The first element that fills electrons in the f sublevel
Consider the following individual orbitals; they are drawn to scale and each orbital has n = 3:
s orbital
(1)
p orbital
(2)
d orbital
(3)
d orbital
(4)
134.
Which orbital has the lowest energy? ________
135.
Which orbital(s) could hold a maximum of 2 electrons? ________
136.
Write a reasonable set of quantum numbers for an electron that would be in the orbital pictured in
figure (2) above.
137.
What is the maximum number of electrons that could have n = 3? ________
Sections 11.1 and 11.2
Fill in the Blank. Fill in the blank with the appropriate word or phrases.
138.
Group 7A elements are known as the _____________________________________________.
139.
140.
141.
142.
Group A elements are known as the ______________________________________ elements.
143.
Group B elements are known as the ______________________________________ elements.
Identification. Identify the element based on the provided information.
144.
The 5th period element containing only 3 4d electrons
145.
The first element that fills electrons in the f sublevel
146.
The first element that fills electrons in the d sublevel
147.
Element that has six outer electrons on the second period of the periodic table of elements
148.
Element that has five outer electrons on the fourth period of the periodic table of elements
149.
Element that has an outer electron configuration of 3d54s1
150.
Element that has an outer electron configuration of 3d104s2
151.
Element that has an outer electron configuration of 3s23p6
152.
Element that has an outer electron configuration of 3s1
153.
Element that has an outer electron configuration of 6s26p1
154.
The 6th period Group 5A element
155.
The 2th period Group 7A element
156.
Element that is an alkali metal and in the fifth period
157.
Element that is an alkaline earth metal and is in the fourth period
158.
Element that is a halogen and has its outermost electrons in the fifth energy level
159.
Element that is a noble gas and has its outermost electrons in the fourth energy level
160.
Has an electron configuration of 1s22s22p63s23p4
161.
Has an electron configuration of 1s22s22p63s23p64s23d104p65s14d5
162.
The 3rd period element with a complete outermost energy level
163.
The 5th period element containing only 1 s electron and no d electrons in the outermost energy
level
Short Answer. Write the shorthand electron configurations of the following ions.
164. Cu: __________________________________________________________________________
165. Cr6+: _________________________________________________________________________
166. Cr: __________________________________________________________________________
167. Zn: __________________________________________________________________________
168. Ag: __________________________________________________________________________
169. Sr2+: _________________________________________________________________________
170. P3–: _________________________________________________________________________
Section 11.3
For the following pairs of atoms, circle which one of each pair has the largest ionic radius.
171.
Al
B
172.
S
O
173.
Br
Cl
174.
Na
Al
175.
O
F
For the following pairs of elements, circle which one of each pair has the greater electronegativity.
176.
Ca
Ga
177.
Li
O
178.
S
Cl
179.
As
Br
180.
Br
Cl
For the following pairs of elements, circle which one of each pair has the greater ionization energy.
181.
Ca
Ba
182.
Li
O
183.
S
Cl
184.
As
186.
Would you expect a Cl– ion to be larger or smaller than an Mg2+ ion? Explain.
187.
Which effect on atomic size is more significant: the nuclear charge or the energy level that
electrons are filling? Explain.
Br
185.
Br
Cl
188.
How does the ionic radius of a typical anion compare with the radius for the corresponding neutral
atom? Explain.
189.
How does the ionic radius of a typical cation compare with the radius for the corresponding neutral
atom? Explain.
Unit Six Worksheet
WS – DC – U6
Chapter Thirteen
Drawing. Draw electron dot structures for the following molecules or polyatomic ions.
190.
HBr (hydrobromic acid)
194.
PF3 (phosphorus trifluoride)
191.
C2H2 (ethyne, or acetylene)
195.
NH4+ (ammonium ion)
192.
ClO4– (chlorate ion)
196.
SO3– (sulfite ion)
193.
HCN (hydrogen cyanide)
197.
Br2 (diatomic bromine)
198.
How many resonance structures can be drawn for the carbonate ion? Show the structural
formulas for each.
199.
How many resonance structures can be drawn for the nitrite ion? Show the electron dot structures
for each.
Drawing.
200.
Draw the electron dot structure for each molecule or polyatomic ion. Then identify the
molecular geometry of the following molecule or polyatomic ion.
NO3- (nitrate ion)
202.
Geometry: ___________________
201.
CCl4 (carbon tetrachloride)
Geometry: ___________________
MnO4- (permanganate ion)
(HINT: Assume Mn has 7 valance e–.)
203.
PO43- (phosphate ion)
Geometry: ___________________
Geometry: ___________________
204.
205.
PO33- (phosphite ion)
PI5 (phosphorus pentaiodide)
206.
Geometry: ___________________
Geometry: ___________________
NH3 (ammonia)
207.
Geometry: ___________________
Geometry: ___________________
CO2 (carbon dioxide)
What type of bond – nonpolar covalent, polar covalent, or ionic – will form between each of the following
pairs of atoms?
208.
Na and O: _______________________________________________________
209.
Li and Cl: ________________________________________________________
210.
P and O: ________________________________________________________
211.
N and N: _______________________________________________________
212.
Al and Cl: ________________________________________________________
213.
O and F: ________________________________________________________
214.
Explain why most chemical bonds would be classified as either polar covalent or ionic. (HINT:
Consider why most are NOT nonpolar covalent.)
215.
Would you expect carbon monoxide to be a polar or nonpolar molecule, and is there any difference
in polarity between carbon monoxide and carbon dioxide? Explain.
Draw the structural formulas for each molecule and identify polar covalent bonds by assigning the slightly
positive (+) and the slightly negative ( –) charges to each atom in each bond. Then identify the overall
molecule as polar or nonpolar.
216.
NH3
218.
CF4
217.
CCl4
219.
HF
220.
Which compound would you expect to have the higher melting point: OCl2 or CaCl2? Explain.
Table Completion. Complete the following table.
Symbol
221.
Number of
valence
electrons in
atom/ion
222.
Electron configuration
223.
224.
S2–
Ca2+
Electron dot
formula
225.
8
226.
1s22s22p63s23p6
227.
228.
Na
Rb+
232.
229.
230.
231.
233.
234.
How many electrons will each of the following elements gain or lose in forming an ion? State whether
each is a cation or anion.
Number of e– lost or gained: 235. ________
Anion/Cation? 238. _____________
Phosphorus: Number of e– lost or gained: 236. ________
Anion/Cation? 239. _____________
Number of e– lost or gained: 237. ________
Anion/Cation? 240. _____________
Strontium:
Bromine:
241.
What is the relationship between the group number of the representative elements and the number
of valence electrons?
242.
Why do metals tend to form cations while nonmetals tend to form anions?
Unit Seven Worksheet
WS – DC – U7
Chapter Fourteen
243.
In your own words, explain what a hydrogen bond is.
244.
Depict the hydrogen bonding between three water molecules in a drawing.
245.
How is hydrogen bonding responsible for the high boiling point of water?
246.
Explain how large bodies of water are able to moderate air temperature.
247.
Explain why it gets warmer before it rains.
248.
Explain why the density of ice at 0C is less than the density of liquid water at 0C.
249.
Explain why water has a relatively high boiling point and heat of vaporization.
250.
What is the difference between the structure of liquid water and the structure of ice? How does
this explain why ice floats in water?
251.
Explain why water has a high surface tension.
Chapter Fifteen
Identify the solute and solvent in a dilute aqueous solution of sodium hydroxide.
252.
Solute: _______________________________________________________________
253.
Solvent: ______________________________________________________________
254.
Give an example of a polar molecular compound that dissolves in water and that is a
nonelectrolyte.
Which of the following compounds are soluble in water? Which are insoluble?
255.
CaCl2: ____________________________________________
256.
N2: _______________________________________________
257.
HBr: ______________________________________________
258.
NH4C2H3O2: ________________________________________
Write equations to show how the following compounds dissociate in water.
259.
NH4NO3(s)
260.
K2SO4(s)
Write the formulas for the following hydrates.
261.
Calcium sulfate decahydrate: ________________________________________________
262.
Cobalt (II) chloride hexahydrate: ______________________________________________
263.
Find the percent by mass of water in NiCl2  6H2O.
264.
Why is using water to clean a paintbrush covered with oil-based enamel not an effective cleanup
method?
265.
How can a supersaturated solution be prepared?
266.
You are given a clear aqueous solution containing potassium nitrate (KNO3). How would you
determine experimentally if the solution is unsaturated, saturated, or supersaturated?
Short Answer.
Write complete balanced net ionic equations based on the following chemicals
reacting.
267.
Lead (II) nitrate and sulfuric acid
268.
Sodium phosphate and iron (III) chloride
269.
Ammonium sulfuide and cobalt (II) nitrate
270.
Sulfuric acid and barium chloride
271.
Aluminum sulfate and ammonium hydroxide
272.
Silver nitrate and dihydrogen sulfide
273.
Calcium chloride and lead (II) nitrate
274.
Calcium nitrate and sodium carbonate
275.
Hydrochloric acid and barium hydroxide
276.
Iron (III) nitrate and sodium hydroxide
277.
What are colligative properties of solutions? Give examples of three types of colligative properties.
278.
How many particles in solution are produced by each formula unit of potassium carbonate, K2CO3?
279.
How many moles of particles would 3 mol Na2SO4 give in solution?
280.
What kind of property is vapor-pressure lowering?
An equal number of moles of NaCl and CaCl2 are dissolved in equal volumes of water. Which solution has
the lower
281.
freezing point? _______________________________________________________
282.
vapor pressure? ______________________________________________________
283.
boiling point? _________________________________________________________
284.
Why does a solution have an elevated boiling point and a depressed freezing point compared with
the pure solvent?
Section 6.1
285.
When you talk about the volume of a gas, are you referring to the volume of the molecules
themselves? Explain.
286.
What is the difference between a pressure and a force?
287.
What happens to gas particles when a gas is compressed?
288.
Explain how a mercury barometer works.
289.
If you collect a gas so that it completely fills a 250 cm3 Erlenmeyer flask, the volume of the gas is
actually greater than 250 cm3. Explain why this is true, and explain how you could determine what
the volume of the gas is.
Calculations.
Solve the following problems. Show work and appropriate significant figures and/or
decimal places.
Hydrogen gas is collected by bubbling it through water. Calculate the partial pressure of the hydrogen gas
if:
290.
The total pressure is 94000 Pa and the partial pressure of water is 1200 Pa.
291.
The total pressure is 100.3 kPa and the partial pressure of water is 2600 Pa.
292.
In a flask that has a volume of 273 dm3, you have a sample of two noble gases: neon and xenon.
The partial pressure of the neon is 96950 Pa, and the partial pressure of the xenon is 1.025 atm.
What is the total pressure (in kPa) exerted by these two gases?
Make the following conversions. Show work and appropriate significant figures and/or decimal places.
293.
105 Pa = __________ torr
295.
256.7 mm Hg = __________ atm
294.
12 kPa = __________ psi
296.
285 torr = __________ kPa
Section 6.2
Calculations.
decimal places.
297.
A metal cylinder contains 1 mol of oxygen gas at STP. What will happen to the pressure if another
mole of gas is added to the cylinder, but the temperature and volume do not change?
298.
If a gas is compressed from 10 L to 1 L, and the temperature remains constant, what happens to
the pressure?
299.
A truck driver gets paid according to the quantity of a certain kind of gas he can deliver. The more
gas he delivers, the more he gets paid. He only has time for one trip, and the dimensions of his
tanker are fixed. What properties of gases can he exploit to increase his profit?
300.
The gas in a closed container has a pressure of 3.00 x 102 kPa at 30C. What will the pressure be
if the temperature is lowered to –172C?
301.
Calculate the volume of a gas (in L) at a pressure of 1.00 x 102 kPa if its volume at 1.20 x 102 kPa
is 1.50 x 103 mm3.
302.
A gas with a volume of 3.00 x 102 mL at 150.0C is heated until its volume is 6.00 x 102 mL. What
is the new temperature of the gas if the pressure remains constant during the heating process?
303.
A given mass of air has a volume of 6.00 L at 101 kPa. What volume, in cubic centimeters, will it
occupy at 25.0 kPa if the temperature does not change?
304.
A 3.50-L gas sample at 20C and a pressure of 86.7 kPa expands to a volume of 8.00 L. The final
pressure of the gas is 56.7 kPa. What is the final temperature of the gas, in degrees Celsius?
305.
What is the volume of a sample of carbon dioxide at standard temperature and pressure that has a
volume of 75.0 mL at 30.0C and 91 kPa?
306.
Which of the following samples of gases occupies the largest volume, assuming that each sample
is at the same temperature and pressure – 50.0 g neon, 50.0 g argon, or 50.0 g xenon?
307.
What volume of carbon dioxide gas contains the same number of oxygen atoms as 250.0 cm3 of
carbon monoxide gas, if each gas sample is measured at the same temperature and pressure?
Given the following data,
Volume of Nitrogen Gas (L)
4.28 L
5.79 L
308.
Temperature (K)
303 K
410 K
Draw a graph of the relationship between volume and temperature.
309.
Determine the slope of the line.
310.
Find the slope of a line in relationship to the temperature-volume law expressed as V/T = k.
311.
Calculate the expected volume of the gas when the temperature reaches 1200 K.
Section 7.1
Calculations.
decimal places.
312.
If 4.50 g of methane gas (CH4) is introduced into an evacuated 2.00-L container at 35C, what is
the pressure in the container?
313.
Calculate the number of moles of oxygen gas in a 12.5 L tank if the pressure is 253 kPa and the
temperature is 22C.
314.
Calculate the mass of nitrogen dioxide present in a 275 mL container if the pressure is 240.0 kPa
and the temperature is 28C.
315.
What is the density of nitrogen dioxide given the conditions of Problem 30?
316.
During the metabolic process called respiration, your body obtains energy from the breakdown of
glucose as shown below.
C6H12O6(aq) + 6O2(g)  6H2O(l) + 6CO2(g)
What volume of O2, measured at 37C and 790.0 torr pressure, is required to react with 1.00 g of
glucose (C6H12O6)? Express the volume in cubic centimeters.
317.
Magnesium reacts with oxygen gas to produce magnesium oxide. What volume of oxygen gas, in
cubic centimeters, is required to fully react with 20.2 g of magnesium metal? The reaction is taking
place at 1.02 atm pressure and 24C.
318.
Magnesium reacts with hydrochloric acid to produce hydrogen gas and magnesium chloride. What
volume of hydrogen gas is theoretically produced, in cubic millimeters, if 4.9 g magnesium reacts
with excess hydrochloric acid? The reaction is taking place at 101.5 kPa pressure and 20C.
319.
320.
321.
A truck driver gets paid according to the quantity of a certain kind of gas he can deliver. The more
gas he delivers, the more he gets paid. He only has time for one trip, and the dimensions of his
tanker are fixed. What properties of gases can he exploit to increase his profit?
322.
The gas in a closed container has a pressure of 3.00 x 102 kPa at 30C. What will the pressure be
if the temperature is lowered to –172C?
323.
Calculate the volume of a gas (in L) at a pressure of 1.00 x 102 kPa if its volume at 1.20 x 102 kPa
is 1.50 x 103 mm3.
324.
A gas with a volume of 3.00 x 102 mL at 150.0C is heated until its volume is 6.00 x 102 mL. What
is the new temperature of the gas if the pressure remains constant during the heating process?
325.
A given mass of air has a volume of 6.00 L at 101 kPa. What volume, in cubic centimeters, will it
occupy at 25.0 kPa if the temperature does not change?
326.
327.
Section 7.2
Calculations.
decimal places.
Calculate the number of liters occupied at STP.
328.
2.5 mol N2(g)
329.
0.600 g H2(g)
330.
2.8 g CO2(g)
331.
2.8 x 1021 molecules CO2(g)
332.
Which of the following samples of gases occupies the largest volume, assuming that each sample
is at the same temperature and pressure – 50.0 g neon, 50.0 g argon, or 50.0 g xenon?
333.
What volume of carbon dioxide gas contains the same number of oxygen atoms as 250.0 cm3 of
carbon monoxide gas, if each gas sample is measured at the same temperature and pressure?
Section 7.3
Calculations.
decimal places.
334.
Calculate the ratio of the velocity of helium atoms to the velocity of neon atoms at the same
temperature.
335.
A certain gas effuses four times as fast as oxygen (O2). What is the molar mass of the gas?
336.
During an effusion experiment, it took 75 seconds for a certain number of moles of an unknown
gas to pass through a tiny hole. Under the same conditions, the same number of moles of oxygen
gas passed through the hole in 30 seconds. What is the molar mass of the unknown gas?
337.
At the same temperature and pressure, which gas moves faster: oxygen or neon? How many
times is the speed of the faster gas greater than the slower gas?
338.
In an experiment, it takes an unknown gas 1.5 times longer to diffuse than the same amount of
oxygen gas. Find the molar mass of the unknown gas.
Chapter 16
Calculations.
decimal places.
339.
How many kilojoules of energy are in a donut that contains 205.0 Calories?
340.
What is the specific heat of a substance that has a mass of 25.0 g and requires 525.0 calories to
raise its temperature by 15.0 K?
341.
Suppose 0.20 kg of ice absorbs 125.0 J of heat. What is the corresponding temperature change?
The specific heat capacity of H2O(s) is 2.1 J/(gC).
342.
How many joules of heat energy are required to raise the temperature of 100.0 g of aluminum by
120.0C? The specific heat capacity of aluminum is 0.90 J/(gC).
343.
A student mixed 75.0 mL of water containing 0.75 mol HCl at 25C with 75.0 mL of water
containing 0.75 mol of NaOH at 25C in a foam cup calorimeter. The temperature of the resulting
solution increased to 35C. How much heat in kilojoules was released by this reaction? CH2O =
4.18 J/(gC)
344.
Calculate the amount of heat evolved when 15.0 g of Ca(OH)2 forms from the reaction of CaO(s) +
H2O(l). CaO(s) + H2O(l)  Ca(OH)2(s) H = – 65.2 kJ
345.
Calculate the amount of heat produced when 52.4 g of methane, CH4, burns in an excess of air,
according to the following equation. CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = – 890.2 kJ
346.
When 2 moles of nitric oxide, NO, burn in air to produce 2 moles of nitrogen dioxide, 113.04 kJ of
heat is produced. Write a balanced thermochemical equation for this reaction.
347.
Calculate the amount of heat needed to melt 35.0 g of ice at 0C. Express your answer in
kilojoules.
348.
Calculate the amount of heat needed to convert 1.0 kg of liquid water at 15C to steam at 100C.
349.
Calculate the amount of heat needed to convert 96 g of ice at – 24C to water at 100C. The
specific heat capacity of H2O(s) is 2.1 J/(gC).
350.
How much heat is absorbed when 28.3 g of H2O(s) at 0C is converted to liquid at 0C? The
specific heat capacity of H2O(s) is 2.1 J/(gC).
351.
When 47.5 g of a metal at 425.0C is dropped into 1000.0 g of water at 18.0C, the final
temperature of the metal and water is 21.0C. The specific heat of water is 4.184 J/(gC). What
is the specific heat of the metal in J/(gC)?
352.
In a calorimetry experiment, you collected the following data:
Classification
Mass of metal
Initial temperature of water in cup in degrees
Celsius
Initial temperature of metal in degrees Celsius
(temperature of boiling water)
Maximum temperature of metal and water
Mass of water
Measurement
100.0 g
20.0C
120.0C
30.0C
150.0 g
What is the specific heat of the metal in the experiment in J/(gC)?
Short Answer.
Answer the following questions.
State whether the following physical and chemical changes are endothermic or exothermic.
353.
Melting
357.
Vaporization
354.
Condensation
358.
Fusion
355.
Freezing
359.
Combustion
356.
Sublimation
360.
Evaporation
361.
If a reaction has H < 0, what kind of reaction occurs? Explain?
362.
What is meant by T? Explain two ways of how one can calculate it.
363.
The same quantity of heat is added to an iron nail (3.5 g) and to a metric ton (1000 kg) of iron.
Which would reach the higher temperature? Explain.
364.
Which has more entropy: 1 g of liquid water or 1 g of steam? Explain.
365.
Which has more entropy: 1 g of liquid mercury or 1 g of solid mercury? Explain.
366.
Analogize entropy and enthalpy in terms of the favorability of a reaction.
Unit Eight Worksheet
WS – DC – U8
Chapter Seventeen
Calculation. Answer the following problems. Show work or receive no credit. Show proper units.
367.
An ice machine can produce 120 kg of ice in 24 hours. Express the rate of ice production in kg/hr.
368.
Ethyl acetate (C4H8O2) reacts with a solution of sodium hydroxide (NaOH) in water to form sodium
acetate (C2H3O2Na) and ethyl alcohol (C2H6O). Suppose at 25C two moles of ethyl acetate react
completely in four hours. How would you express the rate of reaction?
369.
A friend tells you that you can recognize a fast reaction because it produces more product than a
slow reaction. What other factors must be included to make this a correct statement?
370.
Ethyl acetate and water are not miscible; thus, the reaction in problem 368 only occurs at the
interface of the two liquids. What would be the effect on the reaction rate by adding a solvent to
make the reaction homogeneous?
371.
What reactant or product would you choose to measure in order to determine the rate of reaction
for the following chemical reaction?
Zn(s) + 2HCl(aq)  H2(g) + ZnCl2(aq)
Explain how you would measure the substance you chose.
372.
What reactant or product would you choose to measure in order to determine the rate of reaction
for the following chemical reaction?
Cu(s) + 2AgNO3(aq)  2Ag(s) + Cu(NO3)2(aq)
Explain how you would measure the substance you chose.
How would the following actions likely change the rate of reaction in problem 368?
373.
the temperature is lowered to 4C
374.
the concentration of sodium hydroxide in water is increased
List three ways that reaction rates can generally be increased.
375.
__________________________________________________________________
376.
__________________________________________________________________
377.
__________________________________________________________________
378.
Explain how a catalyst works as correlated to activation energy and an activated complex.
Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas. Predict which reaction will
occur at a faster rate based on the following circumstances and explain why. Assume all other factors
aside from those specified are the same.
379.
Reaction 1: 1.00 M hydrochloric acid is used
Reaction 2: 6.00 M hydrochloric acid is used
380.
Reaction 1: Chunks of zinc are used
Reaction 2: Powdered zinc is used
381.
Reaction 1: The reaction occurs at 75C.
Reaction 2: The reaction occurs at 25C.
Chapter Eighteen
Calculation. Answer the following problems. Show work or receive no credit. Show proper units.
382.
Write the expression for the equilibrium constant for the following reaction.
2N2O5(g)  4NO2(g) + O2(g)
383.
Calculate the equilibrium constant for the reaction in problem 382 if the equilibrium concentrations
are [N2O5] = 0.50 mol/L, [NO2] = 0.80 mol/L, and [O2] = 0.20 mol/L.
How would the equilibrium position for the equation in problem 382 be affected by the following?
384.
an addition of O2 to the reaction vessel
385.
a decrease in the pressure
Write the equilibrium constant expression for each of the following reactions.
386.
4NO(g) + 2O2(g)  2N2O4(g)
387.
SO2(g) + NO2(g)  SO3(g) + NO(g)
388.
Can a pressure change shift the equilibrium position in every reversible reaction? Explain your
answer.
389.
What two factors determine whether a reaction is spontaneous?
390.
Where does lost free energy typically end up? Does free energy lost as heat ever serve a useful
function? Explain.
391.
What can change a reaction from nonspontaneous to spontaneous?
392.
Suppose the products in a spontaneous process are more ordered than the reactants. Is the
entropy change favorable or unfavorable? Explain.
393.
How can an equilibrium constant be used to determine the favorability of a reaction?
394.
Why are solids not something to be considered when calculating an equilibrium constant?
Classify each of the following acids as monoprotic, diprotic, or triprotic.
395.
HCOOH: _____________________________________________________________
396.
HBr: ________________________________________________________________
397.
H2SO3: ______________________________________________________________
398.
HClO4: _____________________________________________________________
399.
What would you expect to happen when lithium metal is added to water? Show the chemical
reaction.
400.
Identify the hydrogen ion donor(s) and hydrogen ion acceptor(s) for ionization of sulfuric acid in
water. Label the conjugate acid-base pairs.
401.
Identify all of the ions that may be formed when H3PO4 ionizes in water.
Matching. Match each solution with its correct description.
_____402. dilute, weak acid
_____403. dilute, strong base
_____404. concentrated, strong acid
A)
B)
C)
D)
E)
18M H2SO4(aq)
0.5M NaOH(aq)
15M NH3(aq)
0.1M HC2H3O2(aq)
0.1M HCl(aq)
_____405. dilute, strong acid
_____406. concentrated, weak base
407.
Write the expression for the base dissociation constant for hydrazine, N2H4, a weak base.
Hydrazine react with water to form the N2H5+ ion.
408.
Write the base dissociation constant expression for the weak base analine, C6H5NH2.
C6H5NH2(aq) + H2O(l)  C6H5NH3+(aq) + OH–(aq)
Write, complete, and balance the equations for the following acid-base reactions, including states of
matter.
409.
Phosphoric acid + aluminum hydroxide 
410.
Hydroiodic acid + calcium hydroxide 
411.
Nitric acid + sodium hydroxide 
Calculations.
Solve the following problems. Show work with correct significant figures and/or decimal
places. Include proper units.
412.
What is the molarity of a sodium hydroxide solution if 38 mL of the solution is titrated to the end
point with 14 mL of 0.75M sulfuric acid?
413.
If 24.6 mL of a calcium hydroxide solution are needed to neutralize 14.2 mL of 0.0140M acetic
acid, what is the concentration of the calcium hydroxide solution?
414.
A 12.4 mL solution of sulfuric acid is completely neutralized by 19.8 mL of 0.0100M calcium
hydroxide. What is the concentration of the sulfuric acid?
415.
What volume of 0.12M barium hydroxide is needed to neutralize 12.2 mL of 0.25M hydrochloric
acid?
416.
A 55.0 mg sample of aluminum hydroxide is reacted with 0.200M hydrochloric acid. How many
milliliters of the acid are needed to neutralize the aluminum hydroxide?
Unit Five Experiment – 1
Orbital Structures and Identifications
EX – DC – U5 – 1
Introduction:
The purpose of this simulation is to simulate and identify different kinds of atomic orbitals.
Background:
An atomic orbital is a region of space around the nucleus of an atom where there is a high probability of
finding an electron. The means by which the probability of finding an electron was determined is known as
the Schrödinger equation. Although it is a rather complex equation, it can be solved exactly for the
hydrogen atom.
There are four quantum numbers typically associated with an atomic orbital and used by many chemists
and physicists with regard to atomic orbitals, and these are actually used as well in the “Atom in a Box”
software. They are as follows:

Principal Energy Level (n)
This number corresponds to what we have called the principal energy level in our class and also
corresponds with the average distance of electron from the nucleus. We often ascribe the principal
energy level to be synonymous with the period of the element whose electrons we are trying to
convey. Value of n = 1, 2, 3, 4, and so forth.

Sublevels (l)
The general shape of the electron cloud associated with an atomic orbital is called the sublevel. n
and l are related in that l = 0, 1, 2, …, (n – 1). Typically, we can use the following quantum
numbers represented by l to indicate the type of sublevel (s, p, d, and f) we are talking about.
Quantum number (l)
Type of sublevel

0
s
1
p
2
d
3
f
Orbitals (ml)
This quantum number refers to the number of orbitals possible given which type of atomic orbital is
being discussed. We know that an s orbital only has 1 type of orbital whereas a p orbital has 3
different types (x, y, and z). If you know the value of l, indicating the type of orbital, you can
choose specifically what kind of orbital since ml, = l, …, +1, 0, -1, …, -l. Using this, one can
determine that and s orbital has only one kind of orbital, p has three, d has five, and f has seven.
You will need to be familiar with these three quantum numbers as they are in the program you will be
using.
These values can be seen at the top of the screen and can be modified by pressing the plus (+) or minus
(-) sign above the respective values of n, l, and ml. This can be seen in the following screen shot:
Procedure:
1.
Follow the directions provided by your instructor for how to load the “Atom in a Box” software.
2.
For the first atomic orbital, select n = 1. Record in the data table what kind of an orbital this is as
well as its principal energy level.
3.
Now change n = 3. Set the value of l = 1. Record in the data table what kind of orbital this is as
well as its principal energy level.
4.
Change n = 5. Set the value of l = 3 and m = 0. Record in the data table what kind of orbital this is
as well as its principal energy level. Then record whether or not it has any nodes.
5.
Change the value of m to 2. If you observe and noticeable changes, record them accordingly.
6.
Change n = 7. Set the value of l = 3 and m = 0. Record in the data table what kind of orbital this is
as well as its principal energy level. Does this kind of orbital have any observable rings? (Rotate if
necessary to investigate.) If so, record accordingly.
7.
Set the value of l = 6 and m = 0. Record in the data table what kind of orbital this is as well as its
principal energy level. Does this kind of orbital have more or less observable rings than the
previous orbital? Record accordingly.
8.
Experiment with the program to help you answer the questions on the response page.
Unit Five Experiment – 1 – R
Orbital Structures and Identifications
EX – DC – U5 – 1 – R
Procedure
question
n
l
m
Type of
orbital
(i.e., s, p, d, f)
Principal
energy level
Answer to
(i.e, 1, 2, 3) question (if asked)
2
3
Nodes?
4
Noticeable changes?
5
How many rings?
6
More or less rings?
How many?
7
Questions:
Identify the quantum numbers for the following atomic orbitals. You may use the Atom in a Box software
to help you identify them.
1.
2.
3.
4.
5.
What do these orbitals fully represent?
6.
How many different kinds of orbitals are possible in this program given the available information?
***BONUS***
Identify the quantum numbers associated with the following orbital.
Answers:
1. n = ____________ l = ____________
3. n = ____________ l = ____________
2. n = ____________ l = ____________
4. n = ____________ l = ____________
5. _________________________________________________________________________
_________________________________________________________________________
6. _________________________________________________________________________
BONUS: n = ________________ l = ________________
Metal Ions and Flame Tests
EX – DC – U5 – 2
Introduction:
The purpose of this experiment is to determine the cations of a solution based on flame tests.
Background:
When a metal or metal salt is added to a flame, a combustion reaction ensues. This reaction excites an
electron in the metal from its ground state to a higher orbital. In order to return to its ground state, the
electron releases the additional energy in the form of light.
Different metal electrons emit different wavelengths of light to return to their respective ground states, so
the flame colors are varied. These flames can be used to produce atomic emission spectra of the
elements combusted. Using known values of emission spectra, one can perform a flame test on un
unknown substance, gather an emission spectrum from it, and determine which elements are in the
unknown substance.
Procedure:
1.
At each lab station, take a splint and tap off excess fluid from the splint. Hold it over the flame and
observe the color produced. Record your findings in the table. Discard the used splint in the
appropriate place.
2.
Repeat step 1 for each additional station.
Lab
Station
Color Produced
Metal
1
2
3
4
5
6
Your instructor will provide you data about the metal tests following the experiment.
Conclusion/Discussion:
1.
Did you observe certain colors consistent with certain groups/families of metals? Explain.
2.
How do you think scientists use this knowledge today?
3.
How are the results you found related to the photoelectric effect?
Viewing Spectra
EX – DC – U5 – 2
Introduction:
The purpose of this experiment is to examine the spectra produced when you view a variety of light
sources through a diffraction grating in a spectrometer.
Background:
A diffraction grating can break up light from a source into its component colors, much the same as a prism
can break up sunlight into the colors of the rainbow. In the rectangles below, use crayons to sketch the
spectrum you see from each source.
-9
The numbers below the rectangles represent the wavelength of the light in nanometers (x 10 m) After
you have made your observations, answer the Conclusion/Discussion Questions.
Procedure:
1.
View an incandescent light bulb through a diffraction grating. Inside the bulb a tungsten filament is
heated until it glows.
2.
Now view the spectra of three gas discharge tubes. In these tubes, atoms in a low-pressure gas
are excited by being bombarded by a stream of high-energy electrons. Record which gas is in the
discharge tube you observe.
1.
Note how all the colors in the spectrum from the incandescent bulb "run into" one another. We call
this a continuous spectrum. Is there evidence to suggest that the incandescent bulb emits light in
regions of the spectrum other than the visible? Explain.
2.
At which end of the spectrum (blue or red), does light transfer more energy? Explain how you
know in terms of the relationships we have discussed this far.
3.
How do the spectra produced by the excited gases differ from that produced by the hot metal
filament? Do the atoms emit all frequencies of visible light?
4.
Determine, as best you can, the wavelength of the red line in the hydrogen spectrum. Calculate
the frequency of this light. Determine the energy of the photons of light emitted at this frequency.
Show your work clearly.
Unit Six Experiment – 1
Paper Chromatography of Food Dyes
EX – DC – U6 – 1
Purpose:
The purpose of this experiment is to use paper chromatography to separate and identify food dyes in
various samples.
Safety:
Safety goggles will be worn at all times. During this experiment, no open-toe shoes are to be worn.
Procedure:
1.
Cut a 5 cm x 10 cm strip of chromatography paper and label it with a pencil as shown in Figure A.
Food Color Samples
0.1% NaCl solution
Your name
X
X
X
X
Red
Yellow
Green
Blue
2.
Use a different toothpick to place a spot of each of the four food colors on the Xs on the
chromatography paper. Allow the spots to dry for a few minutes.
3.
Fill the 250-mL beaker so its bottom is just covered with the solvent (0.1% NaCl solution). Wrap
the chromatography paper around a pencil. Remove the pencil and place the chromatography
paper, color-spot side down, in the solvent. When the solvent reaches the top of the
chromatography paper, remove the paper and allow it to dry.
Observation/Data Tables:
Record observations from the laboratory experiment below.
Observations:
________________________________________________________________________________
________________________________________________________________________________
________________________________________________________________________________
________________________________________________________________________________
________________________________________________________________________________
________________________________________________________________________________
1.
If a food color yields a single streak or spot, it is usually a pure compound. Which food colors
consist of pure compounds?
2.
Which food colors are mixtures of compounds?
3.
Food colors often consist of a mixture of three colored dyes: Red No. 40, Yellow No. 5, and Blue
No. 1. Read the label on the food color package. Which dyes do your food color samples contain?
4.
Identify each spot or streak on your chromatogram as Red No. 40, Yellow No. 5, or Blue No. 1.
5.
Paper chromatography separates polar covalent compounds on the basis of their relative
polarities. The most polar dyes migrate the fastest and appear at the top of the paper. Which dye
is the most polar? Which dye is the least polar?
Unit Seven Experiment – 1
The Ideal Gas Law
EX – DC – U7 – 1
Purpose:
The purpose of this experiment is to use the ideal gas law, understand the variables involved when
working with gases, and to experimentally find the molar mass of a gas.
Background:
Most gas experiments do not occur at standard temperature and pressure. Therefore, we need to use a
calculation that allows us to account for changes in pressure and temperature. The ideal gas law allows
us to achieve this. The ideal gas law is PV=nRT where P is the pressure of the gas, V is the volume of
the gas, n is the number of moles of the gas, R is the ideal gas constant (found on the back of your
periodic table for different pressure units), and T is the temperature of gas. You will need to know the
atmospheric pressure when you perform this lab. Go to http://www.wunderground.com and place in your
zip code to obtain the atmospheric pressure at your location. Pressure will be given in inches of mercury
(in Hg). This will need to be converted to different units of pressure. This will be done later.
Safety:
Safety goggles will be worn at all times. Never shake or tilt the can of compressed air before or during
usage. Never use the gas or canister around a possible ignition source. Avoid contact with skin. DO
NOT WASTE ANY OF THE COMPRESSED GAS!!! USE ONLY THE GAS REQUIRED!!! Violations of
ANY of these safety provisions will result in reduced credit and/or removal from the laboratory.
Procedure:
1.
Create a data table to record the data for the mass, temperature, pressure and volume of the gas
used in this experiment.
2.
Fill a tray about 2/3 full of water.
3.
Obtain a can of compressed air. Mass the can and record the mass. Also, record the mass
when done with the experiment, meaning do NOT use more air than required, or your
readings will be offset significantly. Subtract the masses to get the mass of the gas used.
4.
Fill a 250 flask with water. Place a watch glass on top of the flask. Turn the flask upside down
making sure that there are no bubbles present. Place the flask in the water and remove the watch
glass.
5.
Holding the flask by the neck, place some glass tubing into the flask. Insert the straw from the
compressed air canister into the tubing and seal with some clay.
6.
Slowly fill the flask with gas using the can of compressed air by spraying it into the glass tubing. To
help insure the clay seal is tight, you may wish to try and hold it in place.
7.
Continue to hold down the lever until the gas level almost matches the water level in the tray.
Remove the glass tubing.
8.
Adjust the flask so that the water level inside the flask is the same level as the outside water
making sure to keep the flask upside down. Using a grease pen or sharpie, mark a line on the
flask to record the water level.
9.
Remove the flask. Fill the flask with water to the pen mark. Mass a 250 mL beaker. Record the
mass. Pour the water into the beaker and mass the water and the beaker. Subtract out the empty
beaker mass to get the mass of the water. Since the density of water is 1 g/mL, the mass of the
water will equal the volume.
10.
Measure and record the temperature of the water bath.
1.
The atmospheric pressure you recorded earlier was in inches of Mercury. This could be converted
to many different units and the units you use for pressure affect what your R value for the ideal gas
law will be. Convert your pressure to mm Hg. The conversion factors is 1 in = 25.4 mm.
2.
Question 1 reflects the total pressure inside the flask. You know this because when your water
levels were equal, the pressure inside and outside the flask were equal. Since your gas was
collected through water, some of the pressure in the flask is water vapor. Use your text and read
about Dalton’s law of partial pressures. Then use the chart below to calculate the pressure exerted
by the gas from the can.
Temperature
(°C)
0.0
5.0
10.0
12.5
15.0
15.5
16.0
16.5
17.0
17.5
18.0
18.5
19.9
Pressure
(mm Hg)
4.6
6.5
9.2
10.9
12.8
13.2
13.6
14.1
14.5
15.0
15.5
16.0
16.5
Water Vapor Pressure Table
Temperature
Pressure Temperature
(°C)
(mm Hg)
(°C)
27.0
19.5
17.0
28.0
20.0
17.5
29.0
20.5
18.1
30.0
21.0
18.6
35.0
21.5
19.2
40.0
22.0
19.8
50.0
22.5
20.4
60.0
23.0
21.1
70.0
23.5
21.7
80.0
24.0
22.4
90.0
24.5
23.1
95.0
25.0
23.8
100.0
26.0
25.2
Pressure
(mm Hg)
26.7
28.3
30.0
31.8
42.2
55.3
92.5
149.4
233.7
355.1
525.8
633.9
760.0
4.
Since you now know the pressure, volume, and temperature of the gas, given that the R value is
62.396 (Lmm Hg)/(Kmol), solve for the number of moles of gas using PV = nRT. (Hints: You
probably measured your volume of gas in milliliter (mL). Make sure and convert this to liters by
dividing by 1000 since 1 L = 1000 mL. Also, your temperature was measured in Celsius, so
convert it to Kelvins. The formula for converting Celsius to Kelvin is K = oC + 273.15.
5.
Now that you know the number of moles of gas collected and the mass of that gas, calculate the
molar mass of the compressed gas.
6.
Compare your answer to question 5 with the actual molar mass of difluoroethane. The formula for
difluoroethane is C2H4F2. Calculate your percent error.
Specific Heat Experiment
EX – DC – U7 – 2
Introduction:
The purpose of this experiment is to determine the specific heat of a substance.
Background:
On a sunny day, the water in a swimming pool may warm up a degree or two while the concrete around
the pool may become too hot to walk on in your bare feet. This may seem strange because both the
concrete and the water are being heated by the same source: the sun. This evidence suggests it takes
more heat to raise the temperature of some substances than others. This, in fact, is true: the amount of
heat that is required to raise the temperature of 1 g of a substance by 1 degree Celsius is called the
specific heat capacity, or simply the specific heat, of that substance. Water, for instance, has a specific
heat of 4.184 J/(gC). This value is high in comparison with the specific heats for other materials, such
as concrete. In this experiment, you will use a simple calorimeter and your knowledge of the specific heat
of water to determine the specific heat of an unknown metal.
Safety:
Safety goggles will be worn at all times. During this experiment, no open-toe shoes are to be worn. Since
Bunsen burners will be used, do not handle hot equipment with only your hands. Use proper protective
equipment. Additionally, never use a thermometer as a stirrer.
Procedure:
As you perform the experiment, record your data in Data Table 1.
1.
Measure the mass of the metal cylinder provided by your instructor to the nearest 0.01 g and
record the measurement. Transfer the cylinder to a large, dry test tube and use a utility clamp to
suspend the test tube in about 400 mL of water in a 600 mL beaker. (If you do not have access to
a 600 mL beaker, that is fine. Use a beaker where the metal in the test tube you will be using can
be below water level, meaning you may even need more or less than 400 mL of water.) Heat the
water until it is boiling gently, and leave the test tube in the boiling water bath for at least 10
minutes.
2.
While the cylinder is heating, measure approximately 100 mL of distilled water in a graduated
cylinder. Assume the density of water is 1 g/mL, meaning that if you get 98.2 mL of distilled water,
you have 98.2 g of water. Record this as the mass of water to one decimal place. Then pour the
water into a plastic-foam cup.
3.
Measure and record the temperature of the water in the plastic-foam cup and of the water in the
boiling bath. This temperature should be recorded to one decimal place.
4.
Remove the test tube from the boiling water and quickly pour the cylinder into the water-filled,
plastic-foam cup. Place a thermometer and a glass stirring rod into the cup. Use a stirring rod to
gently stir the cylinder. Do not stir the cylinder with the thermometer. Note the temperature
frequently and record the maximum temperature reached. This temperature should be recorded to
one decimal place.
5.
Pour the water off, dry the cylinder, and repeat steps 1 – 4. After a second trial, repeat step 5, but
instead of conducting another trial, return the metal cylinder to your teacher.
Observations:
Data Table I: Measurements of Mass and Temperature
Trial 1
Measurement
Classification
Trial 2
Measurement
Mass of cylinder
Initial temperature of water in cup in degrees
Celsius
Initial temperature of metal cylinder in degrees
Celsius (temperature of boiling water)
Maximum temperature of cylinder and water
Mass of water
Data Table II: Specific Heat Capacities of
Common Substances
Substance
Water
Ice
Steam
Stainless steel
Iron
Aluminum
Copper
Brass
Gold
Lead
Carbon (graphite)
Specific heat
capacity
J/(gC)
4.184
2.06
1.87
0.927
0.449
0.897
0.385
0.376
0.129
0.129
0.709
Answer these questions on a separate sheet of paper.
1.
Determine the changes in temperature of the water and of the cylinder for each trial (in other
words,
TH2O (trial 1) , TH2O (trial 2) , TMetal (trial 1) , and TMetal (trial 2) ).
2.
Calculate the heat gained by the water in each trial.
3.
Remembering that the heat gained by the water is equal to the heat lost by the cylinder, calculate
the specific heat of the cylinder.
4.
Calculate the average value for the specific heat of the cylinder in your experiment.
5.
Calculate the percent error in the specific heat value that you determined experimentally. Use the
accepted value given by your teacher based on the identity of the metal.
6.
Identify other possible sources of error in this experiment.
7.
Can specific heat be used to identify substances?
Solution Preparation and Rate of Reaction
EX – DC – U7 – 3
Purpose:
The purpose of this experiment is to learn how to prepare solutions from more concentrated solutions and
how to investigate factors that speed up or slow down chemical reactions.
Safety:
Safety goggles will be worn at all times during the course of the experiment. Also, protect yourself from
any other external contact with the acid and other chemicals used in this experiment.
Procedure:
Part I: Effect of Concentration on Reaction Rate at Constant Temperature
1.
Pour 5 mL of the 6M acid provided by your instructor into a test tube marked “6M HCl”.
2.
Calculate how much 6M hydrochloric acid must be diluted in order to make 3M hydrochloric acid in
a 50-mL beaker using a graduated cylinder to measure the volumes of the water with which you
will dilute and the amount of acid you use. Record the amount of the 6M HCl to be used in Data
Table 1. After making this amount in the 50-mL beaker, pour 5 mL of the 3M acid into a test tube
marked “3M HCl”. Pour the remainder of the 3M hydrochloric acid into a separate 250-mL
Erlenmeyer flask.
3.
Calculate how much 3M hydrochloric acid must be diluted in order to make 1M hydrochloric acid in
a 50-mL beaker using a graduated cylinder to measure the volumes of the water with which you
will dilute and the amount of acid you use. Record the amount of the 3M HCl to be used in Data
Table 1. After making this amount in the 50-mL beaker, pour 5 mL of the 1M acid into a test tube
marked “1M HCl”. Pour the remainder of the 1M hydrochloric acid into the original 250-mL
Erlenmeyer flask.
4.
Take zinc strips or pieces of equal size and put one into each test tube. Record the start time and
end time of each reaction. Record your observations in Data Table 2.
Part II: Effect of Temperature on Reaction Rate
1.
Pour 5.0 mL of 6M hydrochloric acid into a test tube. Pour another 5.0 mL of 6M HCl into a test
tube maintained over a hot water bath in a 250-mL beaker maintained at 50ºC. (This may be done
over a gas burner.)
2.
Drop equal amounts of zinc into each test tube. Record the start time and end time of each
reaction. Record your observations in Data Table 3.
Data Table 1: Preparation of Solutions
Solution to be
prepared
Solution to be
diluted
3M
6M
1M
3M
How much of
solution to be
diluted is needed
How much water is
needed to dilute
solution
Data Table 2: Effect of Concentration on Reaction Rate
Reaction
Condition
Time Reaction
Started
Time Reaction
Ended
Reaction
Duration
Observations
6M HCl
3M HCl
1M HCl
Data Table 3: Effect of Temperature on Reaction Rate
Reaction
Condition
Time Reaction
Started
Time Reaction
Ended
Reaction
Duration
Observations
6M HCl at
room temp.
6M HCl at 50ºC
1.
What equation did you use in order to know by how much you needed to dilute more concentrated
solutions to get the concentration you wanted for a solution?
2.
Write a balanced chemical equation for the reaction between hydrochloric acid and zinc metal.
3.
What happens to the reaction rate as temperature increases? Why does this happen? Explain
this in terms of the collision theory of reactions.
4.
Describe in your own words the effect of concentration on the rate of a reaction. Explain this effect
in terms of the collision theory of reactions.
Le Chatelier’s Principle (Physical and Chemical Changes)
EX – DC – U7 – 4
Purpose:
The purpose of this experiment is to determine how equilibrium systems respond to stress.
Safety:



Safety goggles will be worn at all times during the course of the experiment.
Protect yourself from any other external contact with all chemicals used in this experiment.
Return or dispose of all materials according to the instructions of your teacher.
Procedure:
Part I: Effect of Temperature on a Physical Equilibrium
1.
Add 2 – 3 mL of saturated potassium nitrate solution to a clean test tube. Using a spatula, add one
crystal of potassium nitrate to the solution to act as a seed crystal.
2.
Cool the test tube in a 250-mL of ice water for 10 minutes. Record the results.
3.
Remove the tube from the ice water and place it in the test-tube rack. Record what happens as
the solution warms to room temperature.
Part II: Common Ion Effect on a Chemical Equilibrium
1.
Use a graduated cylinder to add 50 mL of distilled water to a 100-mL beaker. Add 1 mL of 0.1M
iron (III) chloride and 1 mL of 0.1M potassium thiocyanate (KSCN) to the water. Stir the mixture.
The color that appears is due to the presence of ferrothiocyanate ions, FeSCN2+. Your teacher will
write on the overhead projector slide the reaction that is observed. Record your observations.
2.
Label four identical, clean, dry test tubes with the numerals 1 – 4. Pour 5 mL of the mixture from
Step 1 into each. Hold the tubes over a white background and look down into them. The solutions
should appear equally dark.
3.
Tube 1 is the control for this experiment. To tube 2, add 20 drops of 0.1M iron (III) chloride. To
tube 3, add 20 drops of 0.1M potassium thiocyanate. Flick each tube to mix the solutions. To tube
4, add 1 g of potassium chloride crystals. Flick the tube to dissolve the crystals. Compare the
colors of the solution in tubes 2, 3, and 4 with the color of the solution in the control tube (tube 1).
Record your observations.
4.
Discard the solutions as directed by your teacher.
Data Table: Observations
System
Observations
KNO3 (saturated)
(cooled)
KNO3 (saturated)
(warmed)
Fe3+/SCN- reaction
Fe3+/SCN- mixture +
additional Fe3+
Fe3+/SCN- mixture +
additional SCNFe3+/SCN- mixture +
KCl(s)
1.
Write a balanced equation for the equilibrium that exists before the saturated and unsaturated
potassium nitrate was cooled.
2.
Did lowering the temperature (Step 2, Part I) affect the equilibrium? Explain your answer.
3.
Did increasing the temperature (Step 3, Part I) disturb the equilibrium? What evidence do you
have for your answer?
4.
Explain what happened in the potassium nitrate in terms of Le Chatelier’s Principle.
5.
Write a balanced equation for the equilibrium that existed after the ferric (Fe3+) and thiocyanate
(SCN-) ions were combined in the beaker.
6.
What evidence was there that the equilibrium shifted when iron (III) chloride was added? In which
direction did it shift?
7.
What evidence was there that the equilibrium shifted when potassium thiocyanate was added? In
which direction did it shift?
8.
Explain the effect of adding potassium chloride to the system.
9.
Explain the changes observed in the ferrothiocyanate ion system in terms of Le Chatelier’s
Principle.
Unit Eight Experiment – 1
Using Indicators
EX – DC – U8 – 1
Purpose:
The purpose of this experiment is to estimate the pH of solutions by using acid-base indicators.
Safety:




Ammonia is an irritant. Do not inhale ammonia.
Procedure:
1.
Add 1 – 2 mL of the following to five, separate, clean test tubes: 0.1M hydrochloric acid, 0.1M
ethanoic acid, 0.1M ammonia, 0.1M sodium hydroxide, and distilled water. RETURN PIPETS TO
THEIR ORIGINAL CONTAINERS!
2.
Add 2 drops of phenol red indicator solution to each tube. Flick to mix the contents. Record the
final color in Data Table 2 and estimate the pH of the solution by referring to Data Table 1.
RETURN PIPETS TO THEIR ORIGINAL CONTAINERS!
3.
Using fresh samples of the solutions, repeat the procedure for each of the other indicators named
in Data Table 1. If you are using paper indicator strips, use a glass stirring rod to transfer a drop of
the solution to the indicator strip. Record the results of all the tests in Data Table 2. RETURN
PIPETS TO THEIR ORIGINAL CONTAINERS!
4.
Test the common household chemicals that are available to you. Test liquids directly. Solids
should be dissolved or suspended in water before testing. Record your results in Data Table 3.
RETURN PIPETS TO THEIR ORIGINAL CONTAINERS!
5.
Follow your teacher’s instructions for proper disposal of the materials.
6.
Clean all glassware with distilled water and dry accordingly. DO NOT LEAVE WATER BOTTLES
IN SINKS OR HIDDEN AREAS AND DO NOT DISPOSE OF USED PAPER TOWELS IN ANY
OTHER PLACE THAN A TRASH RECEPTACLE. DOING SO WILL RESULT IN A 0% FOR
YOUR LABORATORY GRADE AND DISCUSSION QUESTIONS.
Data Table 1: Common Acid-Base Indicators
Color in Acid (Hin
form)
Indicator
Color in Base (Inform)
pH range
Phenol Red
Red Litmus Paper
Bromthymol Blue
Phenolphthalein
Data Table 2: Indicator Reactions with Standard Solutions
Solution
Methyl Red
Red Litmus
Paper
Bromthymol
Blue
Phenolphthalein
Estimated
pH
0.1M HCl
0.1M CH3COOH
(ethanoic acid)
Distilled water
(from bottles)
0.1M NH3
(ammonia)
0.1M NaOH
Data Table 3: Indicator Reactions with Household Chemicals
Substance
Methyl
Red
Red Litmus
Paper
Bromthymol
Blue
Aspirin
Tea
Baking soda
Cola (diet)
Vinegar
Phenolphthalein
Estimated
pH
1.
Compare the pH of 0.1M ethanoic acid with that of 0.1M hydrochloric acid. Compare the pH of
0.1M ammonia with that of 0.1M sodium hydroxide. Explain any differences.
2.
Which of the following indicators used in this experiment could most accurately identify a neutral
solution? Explain.
3.
Are the household chemicals you tested acidic, basic, or neutral? Explain.
Unit Eight Experiment – 2
Determining Molarity by Neutralization
EX – DC – U8 – 2
Purpose:
The purpose of this experiment is to measure the molarity of hydrochloric acid using a standardized
solution of 0.20 M sodium hydroxide.
Safety:




DO NOT CONTAMINATE SOURCE CONTAINERS OR BURETS WITH OTHER CHEMICALS
THAN THOSE DESIGNATED TO BE IN THEM.
Procedure:
1.
Clean and mount two 25-mL burets. Place a white sheet of paper beneath each buret. Label the
left buret “acid” and the right buret “base.”
2.
Rinse the “acid” with three 5-mL portions of the solution of hydrochloric acid. Let each portion
drain out of the buret before adding the next rinse. Discard these rinses. Fill the buret with the
hydrochloric acid. Before beginning the titration, remove any bubbles trapped in the buret and the
stopcock. Also make sure that the solution is below the “0 mL” mark.
3.
Using the 0.20 M sodium hydroxide solution, rinse and fill the “base” buret. Use a wash bottle of
distilled water to rinse off the tip of each buret; catch the runoff in a sink. Record the initial volume
in each buret to the nearest 0.01 mL.
4.
Add 10-12 mL of the acid solution to a clean 250-mL Erlenmeyer flask. Use the wash bottle to
rinse the last drop of acid from the tip of the buret into the flask. Add at least 50 mL of distilled
water and 5 drops of bromthymol blue to the flask.
5.
Slowly add sodium hydroxide solution from the “base” buret to the flask. As you add the base,
gently swirl the solution in the flask. A blue color will appear and quickly disappear as the solutions
are mixed. As more and more base is added, the blue color will persist for a longer time before
disappearing. This is a sign that you are nearing the equivalence point, also called the end point.
Wash down the sides of the flask and the tip of the buret with distilled water from the wash bottle.
Continue to add sodium hydroxide more slowly, until a single drop of base turns the solution a pale
green color that persists for 15 – 30 seconds.
6.
If you overshoot the end point – that is, if you add too much base so the solution turns bright blue –
simply add a few drops of acid from the acid buret to turn the solution yellow again. Approach the
end point again, adding base drop by drop, until one drop causes the color chance to pale green.
7.
When you are sure that you have achieved the end point, record the final volume reading of each
buret. Note: Do not allow the level of the solution in either buret to go below the 25-mL
mark. If you do, you will have to discard your sample and begin again.
8.
Discard the solution in the Erlenmeyer flask as directed by your teacher, and rinse the flask well
with distilled water. Refill both burets, if necessary. Read the initial volume in each buret and do
another titration, as described in steps 4 – 7.
Data Table 1: Molarity of Hydrochloric Acid
Trial 1
Acid
Trial 2
Base
Acid
Trial 3
Base
Acid
Base
Final volume (mL)
Initial volume (mL)
Volume used (mL)
Molarity of HCl (M)
Average molarity of
HCl (M)
1.
Determine how many moles of sodium hydroxide were used in each trial. (HINT: Remember that
you used 0.20 M NaOH, and that molarity = moles solute  liters solution, meaning you need to
convert the milliliters you used into liters.)
2.
Based on how many moles of sodium hydroxide were used in each trial, calculate how many moles
of hydrochloric acid were used in each trial. (HINT: Use the balanced chemical reaction HCl +
NaOH  H2O + NaCl to determine the molar ratio you will need to use.)
3.
Calculate the molarity of hydrochloric acid for each trial.
4.
Calculate the average molarity of the trials for hydrochloric acid.
5.
Why are the burets rinsed with the acid and base solutions before filling?
Appendix A
Laboratory Equipment and LPS Safety Contract
Triple beam balance
Buret
Graduated
cylinder
Test tube rack
Erlenmeyer
flask
Beaker
Crucible and lid
Bunsen
burner
Ring
stand
Double buret clamp
Funnel
Wire
gauze
page A-1 – DC – T2 – BOOK
Test tube tongs
Distilled water
wash bottle
Clay triangle
Test
tube
Safety goggles
Ring clamp
Scoopula
Test tube brush
Appendix B
SI Units and Conversions
Density
d
m
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?
d
m 0.75 g

 0.25 g mL
v 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?
m
 19.3 g 
 m  dv  
d
10.0 cm3  193 g
3 
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 g
d
v

 0.55 mL
v
d 13.6 g
mL
Specific Gravity

Comparison of densities

Formula: Specific gravity 


density of substance
density 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 50C, what is the temperature in Kelvins?
K = 50 + 273 = 323 K
Example:
If the temperature is 50K, what is the temperature in degrees Celsius?
C = 50 – 273 = – 223 C
Units of Measurement
SI base unit or SI derived
unit
Quantity
*
Length
Volume
Mass
Density
Temperature
Time
Pressure
Energy
Amount of
substance
Luminous
intensity
Electric
current
Symbol
meter
cubic meter
kilogram*
grams per cubic centimeter
or
grams per milliliter
kelvin*
second*
Pascal
m
m3
kg
g/cm3
Joule
mole*
J
mol
candela*
cd
ampere*
A
g/mL
K
s
Pa
Non-SI unit
Symbol
liter
L
degree Celsius
C
atmosphere
millimeter of mercury
calorie
Atm
mm Hg
cal
*
: denotes an SI base unit
Commonly Used Prefixes in the Metric System
Prefix
Symbol
Meaning
mega
M
kilo
k
deci
d
centi
c
milli
m
micro

nano
n
pico
p
1 million times larger than the unit it
precedes
1000 times larger than the unit it
precedes
10 times smaller than the unit it
precedes
precedes
precedes
1 million times smaller than the unit it
precedes
1000 million times smaller than the unit
it precedes
1 trillion times smaller than the unit it
precedes
Important conversions:
1 cm3 = 1 mL
103 mL = 1000 cm3 = 1 L
Scientific
notation
Factor
1 000 000
106
1000
103
1/10
10-1
1/100
10-2
1/1 000
10-3
1/1 000 000
10-6
1/1 000 000 000
10-9
1/1 000 000 000 000
10-12
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
1m
1 kilometer 1000 m 100 cm

 100000 cm
•
•
1
1 km
1m
Appendix C
Compound Name and Formula Writing
Metals/Nonmetals:








The charge of the metal ions in Group 1A is 1+.
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-.
Group 8A has no ions.
Polyatomic ions:

Their charge is always negative except for NH4+.
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.
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
Appendix D
Chemical Reactions and Quantities
Chemical Reaction Classifications:
Synthesis/Combination (Oxidation-Reduction):
A + B  AB
2Na(s) + Cl 2  2NaCl (s)
Decomposition (Oxidation-Reduction):
AB  A + B


 2Hg(l) + O2
2HgO(s) 
Single-Replacement (Oxidation-Reduction):
A + BC  AC + B
2K(s) + 2H2O(l)  2KOH(aq) + H2 (g)
Double-Replacement (Precipitation):
A+B- + C+D-  A+D- + C+BK2CO3(aq) + BaCl2(aq)  2KCl(aq) + BaCO3(s)
Combustion (Oxidation-Reduction):


 
CxHy + x + y O 2  xCO 2 + y H2O
4 

 2 
CH 4 (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
a.
b.
c.
C.
D.
Reduction of iron ore (2Fe2O3 + 3C  4Fe + 3CO2)
2AgNO3 + Cu  2Ag + Cu(NO3)2
Oxidation and reduction always occur simultaneously.
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.
b.
II.
Oxidizing agent: sulfur (gains electrons)
ii.
Reducing agent: magnesium (loses electrons)
Example 2
i.
Oxidizing agent: copper (II) nitrate (gains electrons)
ii.
Reducing agent: magnesium (loses electrons)
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.
2.
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.
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.
The oxidation number of the bromide ion (Br-) is –1 while the oxidation
number of the iron (III) ion (Fe3+) is +3.
The oxidation number of hydrogen in a compound is +1, except in metal hydrides
(i.e., NaH) where it is –1.
The oxidation number of oxygen in a compound is –2. except in peroxides (i.e.,
H2O2) where it is –1.
For any neutral compound, the sum of the oxidation numbers of the atoms in the
compound must equal 0.
For a polyatomic ion, the sum of the oxidation numbers must equal the ionic
charge of the ion.
b.
3.
4.
5.
6.
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
NO3–
I–, Br–, Cl–
SO42–
CO32–, PO43–, SO32–
OH–
S2–
Rule
All compounds formed with the
negative ion are soluble.
negative ion are soluble except Ag+,
Pb2+, Hg22+, and Cu+.
Most compounds formed with the
negative ion are soluble; exceptions
include SrSO4, BaSO4, CaSO4,
RaSO4, Ag2SO4, and PbSO4.
negative ion are insoluble except
those of the alkali metals and NH4+.
those of the alkali metals, NH4+, Sr2+,
and Ba2+. (Ca(OH)2 is slightly
soluble.)
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.
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 mol of representative unit

1
6.02 x 1023 atoms
2.3 x 1026 atoms O
1 mol O

 380 mol O
1
6.02 x 1023 atoms
To go from moles to atoms:
# of moles 6.02 x 1023 molecule # of atoms


mol
molecule
1
23
3.6 mol C6H12O 6 6.02 x 10 molecule 24 atoms


 5.2 x 1025 atoms
1
mol
molecule
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 1.0 g H 1 mol O 16.0 g O



 18.0 g H2O
1 mol H
1
1 mol O
1
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
To go from moles to grams for a compound:
# of moles of substance gam, gfm, or gmm of substance

1
mol
2.85 mol H2O
18.0 g H2O

 51.3 g H2O
1
1 mol H2O
To go from grams to moles for a compound:
mol of substance
# of grams of substance

gam, gfm, or gmm of substance
1
32.5 g H2O 1 mol H2O

 1.81 mol H2O
1
18.0 g H 2O
To go from moles to volume of a gas at STP:
# of moles of gas 22.4 L of gas

1
1 mol of gas
2.8 moles CO2 22.4 L CO 2

 63 L CO2
1
1 mol CO 2
To go from density at STP to molar mass of a gas:
density of gas in
g 22.4 L of gas

L mol of gas
1.43 g O2 22.4 L O2

 32.0 L O2
L O2
mol O2
To calculate percent composition of an element in a compound:
Experimentally:
% mass of Element A =
grams of Element A
 100%
grams of compound
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
 100%
molar mass of the compound
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.
To calculate the mass of an element in a given amount of a compound:
mass of compound mass of element in 1 mol of the compound

molar mass of the compound
1
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 =
12.0 g C
48.3 g CH4

 36.2 g carbon
1
16.0 g CH4
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 mol C

 0.833 mol C
12.0 g C
1
0.80 g H 1 mol H

 0.80 mol H
1
1.0 g H
89.1 g Cl 1 mol Cl

 2.51 mol Cl
35.5 g Cl
1
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
 1.0 mol C
0.80
0.80 mol H
 1.0 mol H
0.80
2.51 mol Cl
 3.1 mol Cl
0.80
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.
Type of Reaction
Synthesis/Combination
Hints / What to Look
For on the Reactant
Side
Two elements, element and a
diatomic gas/liquid/solid
What to Do to
Complete the
Reaction
1.
2.
1.
Decomposition
One compound
2.
1.
Single Replacement
Only one ionic compound; the
other reactant is an element or a
diatomic gas/liquid/solid
2.
3.
1.
Two ionic compounds
Double Replacement
Product cases:
1. One precipitate formed.
2. One gas formed.
3. One liquid formed.
2.
1.
Combustion
A hydrocarbon (something with
carbon and hydrogen) and
oxygen gas
(can be complete or incomplete
combustion)
2.
Combine the elements as you
would if you were forming any
ionic compound.
Balance the equation.
Break down the compound into
its constituent elements and/or
compounds.
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).
If a displaced element exists in
a diatomic state in nature, be
sure to indicate this (i.e. H 
H2).
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).
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).
Appendix E
Study Skills
Reading Skills and Review Skills
Contrary to popular belief, studying and reviewing is not simply a matter of reading through something
trying to memorize the information and occasionally glancing back at notes and supplemental materials.
Reading and reviewing entails a number of different skills, many of which can be close to mastered given
continued practice and direction.
Reading Skills:

Browse through the reading before actually reading and preview each assigned reading
assignment.
If you know what to look for before actually reading as well as how much time you will have to invest
in the reading, you will know what kind of preparation needs to be made for the reading assignment.

Reading once is not reading.
Reading only once is like riding a boat over an ocean: you know what you have gone over, but you
don’t understand what any of it really is. Reread after you have done the initial reading.

Write notes while reading.
Learning is not simply a matter of osmosis for the brain. It helps to have some extra reference than
the book. Writing notes in your own words helps in mastering material.

Focus on important aspects while skipping the fluff.
Writing down everything the book covers, including the “fluff” (the extra things of no importance) is
just like overstudying and can be just as damaging to your learning experience. Pick apart what is
important and focus on recording those in your notes.

Know what parts of the reading are confusing and record them.
Not everything you read will make sense to you. Be sure to record questions you have over the
material as soon as you have them so that you can ask your teacher, friends, family, etc.

Read to understand, not to memorize.
Any field of study requires understanding of the material to perform well, not simply memorizing what
to do and when to do it. Read so that you understand the relationships between different aspects of
what you read, not simply what they are.

Review the material.
The saying “If you don’t use it, you lose it” is all too true. Consistently review material so that you
don’t forget how to do it, whether it is the first week of the semester or the last week of the school
year.

Ask questions regarding what you have read.
Reading and trying to understand without clarifying the confusing points is as futile as simply listening
along and not understanding the relationships between what you have heard. Ask questions so that
everything can come together clearly, and, if things don’t make sense after asking those questions,
ask more questions, whether of your teacher or of your peers.
Study Skills:












Set priorities.
Although work outside of school provides money for a variety of different things, the kind of money
you will make during high school will pale in comparison to the money you can make with a quality
education, as is shown statistically. Prioritize according to long-term goals, not just for the here-andnow. Also realize that, even though sports and athletics allow for competition and personal
development, academics do as well, and you need to prioritize accordingly.
Don’t get behind.
If you get behind in one area, it will impact everything to be learned after that. Do not allow yourself
to slip too far, for catching up may become nearly impossible to achieve later.
Don’t get overcommitted.
Overstudying one subject (aka cramming) can be just as lethal to achievement as not studying at all.
Take adequate breaks and disperse your studying over the range of all subjects being taken.
Absence for any reason has an effect.
Regardless of whether you are gone for sports, clubs, work, family vacation, or another reason, your
absence in class will impact your knowledge and experience with a subject. Make adequate
preparations in advance to know what you are missing in the way of knowledge and avail yourself to
all options that can help you to recover from your absence.
Ask for help.
Teachers do not enter the profession of teaching solely to speak in front of a class and do nothing
else. If you do not understand material covered in class, ask your teacher as soon as the question
arises. If that is not possible, write down the question as soon as you have it and ask your teacher at
an appropriate time. Should that still not be possible, talk with other teachers proficient in the field
you are studying or talk to classmates who can provide assistance with your questions.
Get names and phone numbers of classmates.
Sometimes the individuals who can best assist you are those that think like you – your peers, friends,
and fellow classmates. Know who to call for help in particular subjects.
A study group is not a social group.
Although a friend may enjoy talking with you, the friend may not provide any assistance if they are as
confused about a subject as you are. Involve yourself in study groups that benefit your learning, not
who are simply there for personal discussion.
Study in short segments and then take a break or reward yourself.
Similar to how overstudying can be as disastrous as not studying at all, take breaks and providing
yourself rewards for studying allow your mind to process recently-learned information and help you to
retain it. Give yourself breaks that allow you to retain information.
Review daily, especially in sequence courses.
Continual, methodically study is so much more helpful than trying to cram an entire chapter in your
head during an hour before a test. Sometimes only spending a few minutes a day reviewing material
discussed previously is all that is required to retain information. Analyze how you best retain
information and take appropriate steps to review daily. Also, in sequential courses, material covered
at the very beginning of the semester needs to be reviewed occasionally. Take time to review things
even if they are from “long ago.”
Set definite limits for phone, television, Internet, etc.
Distractions can be all too tempting. Limit distractions until you feel you have mastered material, with
the exception of distracting yourself for occasional breaks.
Select an appropriate place and time for study.
Some people need to study in absolute silence; others have learned that music playing in the
background allows them to focus better. Know when, where, and how to study for material.
Observe proper nutrition.
For some people, “junk food,” loaded with carbohydrates and preservatives along with other
nutrients, helps during studying. For other people, caffeine is a helpful supplement. However, how


you eat throughout the day and what you provide your body to remain nourished is extremely
important. Eat a balanced diet throughout the day to assure yourself that your brain is in top form.
Observe proper sleep habits.
How well can your brain perform when it is more focused on its exhaustion or distraction from lack of
sleep than it can on material you are learning or reviewing? Provide yourself with enough sleep so
that learning can take place more easily for your brain.
Keep daily communication with parents.
Should you feel too frustrated or need feedback, parents are often the greatest source. Keep your
parents apprised of your progress so that they can provide you with feedback, assistance, or
whatever they can do to help you.
Appendix F
Electron Configuration Rules



Aufbau Principle
o “Electrons enter orbitals of lowest energy first.”
o Level 1  Level 2  Level 3 …
o sp…
o Know energy level placements (shown below)
Pauli Exclusion Principle
o “An atomic orbital may describe at most two electrons.”
o Electrons have opposite spin, and can only be up to two electrons per orbital.
Hund’s Rule
o “When electrons occupy orbitals of equal energy, one electron enters each orbital until all
the orbitals contain one electron with parallel spins.”
o Even distribution of electrons; have to spread them out in the orbitals
Aufbau Diagram
Electron Configuration “Cheat” Diagram
s sublevel: can hold 2 electrons
p sublevel: can hold 6 electrons
d sublevel: can hold 10 electrons
f sublevel: can hold 14 electrons
The following chart shows the order in which you should fill the orbitals for an electron
configuration.
Appendix G
Electron Dot Structure
Instructions:
For very simple molecules, Lewis dot structures can often be written by just thinking and
realizing what to write. However, many molecules are more complex. The following are the
steps used to write Lewis dot structures.
1.
Count the number of valence electrons.
For a molecule, add up the valence electrons of the atoms that are present. For a
polyatomic anion, one electron is added for each unit of negative charge. For a
polyatomic cation, a number of electrons equal to the positive charge must be subtracted.
2.
Draw a skeleton structure for the molecule or ion, joining atoms by single bonds.
In some cases, only one arrangement of atoms is possible. In other cases, there may be
two or more ways of drawing the structure. Most of the molecules and polyatomic ions
with which you will be concerned consist of a central atom bonded to two or more
terminal atoms (atoms located on the outer edges of a molecule or ion). NOTE: The
central atoms is usually the atom written first in a formula (i.e., C in CO2; N in NH3); put
this in the center of the molecule or ion. Terminal atoms are most often hydrogen,
oxygen, or a halogen and should be bonded to the central atom.
3.
Determine the number of valence electrons still available to be distributed for
bonds.
To do this, deduct two valence electrons for each single bond written in Step 2.
4.
Determine the number of valence electrons required to fill out an octet for each
atom (except H) in the skeleton. NOTE: Remember that shared electrons are counted
for both atoms.
a.
If the number of electrons available in Step 3 is equal to the number required from
Step 4, distribute the available electrons as unshared pairs so that every atom
(except hydrogen) has an octet.
b.
If the number of electrons available in Step 3 is less than the number required
from Step 4, the skeleton structure must be modified by changing single bonds to
double or triple bonds. If you are two electrons short, convert a single bond to a
double bond; if there is a deficiency of four electrons, convert a single bond to a
triple bond (or converting two single bonds to double bonds).
Example:
Example:
Draw the Lewis structure for SO2.
Solution:
Follow the four steps for creating a Lewis dot structure.
1.
The number of valence electrons is 6 (from S) + 12 (6 from each O) = 18 valence
electrons.
2.
The skeleton structure is
3.
The number of electrons available for distribution is 18 (original number) – 2 (from the
first S – O bond) – 2 (from the second S – O bond) = 14.
4.
Since each bond already provides each oxygen with 2 electrons and sulfur with 4
electrons, each oxygen needs 6 more and sulfur needs 4 more, resulting in a need of 16
electrons. As a result, we are missing two electrons to complete octets for each atom.
This means that a single bond in the skeleton must be converted to a double bond.
or
Therefore, the structure of SO2 is :
or
Appendix H
VSEPR Models
Geometries of Molecules/Ions With Expanded Octets
Species
Type
AX
Molecular
Geometry
Linear
Predicted
bond
angles
180
AX2
Linear
180
CO2
AX3
Trigonal
planar
120
CO32-
AX4
Tetrahedral
109.5
CH4
AX5
Trigonal
bipyramidal
90
120
180
PCl5
AX6
Octahedral
90
180
SF6
Ball and Stick
Representation
-
Example
CO
Example
Representation
-
Geometries of Molecules/Ions With Expanded Octets
Species
Type
AX2E
(E represents a lone pair of electrons on the central atom.)
Species
Species
Specie
Type
Type
Species Type
s Type
Species Type
Bent
NO21–
104.5
AX2E2
Bent
104.5
H2O
AX2E3
Linear
180
XeF2
AX3E
Trigonal
pyramidal
109.5
AX3E2
T-shaped
90
180
ClF3
AX4E
See-saw
90
120
180
SF4
AX4E2
Square
planar
90
180
XeF4
AX5E
Square
pyramidal
90
180
ClF5
-
NH3
-
Appendix I
Calorimetry Calculations
Problem:
A 15.0-g block of ice is heated from -10C to 115C. In order to accomplish this, how many
kilojoules of heat are required to be present?
Solution:
In order to do this, one must realize that the melting point of ice is 0C and the boiling point of
water (melted ice) is 100C. Since the ice is going through two phase changes (ice  water
and water  vapor), one must be able to calculate the heat required as an ice, water, and vapor
using q = mCT as well as the heat generated while passing from one phase to another by
using heats of fusion and vaporization. This means that one needs to know the following
values:
 Cice = 2.1 J/(gC) for when ice is absorbing heat from -10C to 0C.
 Hfus = 6.01 kJ/mol for when ice is becoming water at 0C; this is the same as 6.01 x 103
J/mol since we want all heat units to be the same.
 Cwater = 4.184 J/(gC) for when water is absorbing heat from 0C to 100C.
 Hvap = 40.7 kJ/mol for when water is becoming steam at 100C; this is the same as 40.7
x 103 J/mol since we want all heat units to be the same.
 Csteam = 1.7 J/(gC) for when steam is absorbing heat from 100C to 115C.
 The molar mass of ice and water and steam is 18.015 g/mol; we will use this to convert
moles to grams in the heats of fusion and vaporization.
We have five different heats to add:
 q1  mCice T for when ice is absorbing heat from -10C to 0C.


1
 q 2  m
molar mass
H fus for when ice is becoming water at 0C. We multiply the mass


by the molar mass to convert moles since heat of fusion is expressed in Joules per mole.
 q 3  mCwater T for when water is absorbing heat from 0C to 100C.


1

 q 4  m
molar massHvap for when water is becoming steam at 100C. We multiply the


mass by the molar mass to convert moles since heat of vaporization is expressed in
Joules per mole.
 q 5  mCsteamT for when steam is absorbing heat from 100C to 115C.
This results in q total  q1  q 2  q 3  q4  q 5 .
 2 .1 J 
0C   10C  315 J
q1  mCice T  15.0 g
 g  C 
 mol  6.01 x 10 3 J 

q 2  15.0 g
 5004 J  5000 J

mol
 18.015 g 

 4.184 J 
100C  0C  6276 J  6280 J
q 3  mC water T  15.0 g
 g  C 
 mol 
 40.7 x 10 3 J  33888 J  33900 J
q 4  15.0 J
 18.015 g 


 1. 7 J 
115C  100C  382.5 J  383 J
q5  mC steam T  15.0 g
 g  C 
q total  q1  q 2  q 3  q 4  q 5  315 J + 5000 J + 6280 J + 33900 J + 383 J
 45878 J  45.878 kJ
Thus, to heat ice from -10C to 115C, ice must absorb 45.878 kJ of heat.
HINTS:
1.
Use only the “q”s you need. Don’t think you will have to add up all heats if a certain
thermochemical process is not occurring.
2.
Anytime you are going up to melting point, going between melting and boiling point, or
going beyond boiling point – in other words, are dealing with a temperature difference –
remember that q = mCT.
3.
Anytime you are simply adding heat to melt something or boil something, take your mass
(with appropriate units) divided by your molar mass of your substance times the heat of
fusion or heat of vaporization, respectively.
4.
Draw a picture. If Mr. Geist draws pictures to remember, you might want to as well.
Appendix J
Water and Solutions
I.
Liquid Water and Its Properties
A. The Water Molecule
1. Simple triatomic molecule
2. Polarity
a. Each O-H bond is highly polar
b. O: 2- H: + (partial charges)
c. Region around oxygen has a slight negative charge;
region around the hydrogens has a slight positive charge
d. Hydrogen bonding occurs because of polarity
e. Results of hydrogen bonding
i. High surface tension
ii. Low vapor pressure
iii. High specific heat capacity
iv. High heat of vaporization
v. High boiling point
B. Surface Properties
1. Surface tension
a. Explained by water’s ability to hydrogen bond
b. Water molecules experience uneven attractions at the surface,
pulling inward and minimizing surface area
c. Holds a drop of liquid in a spherical shape while being flattened
by gravity
d. Why water forms a meniscus in a tube
e. Surfactant
i. Surface active agent
ii. Decreases the surface tension of water
iii. Used to interfere with bonding between hydrogen molecules,
causing beads of water to collapse
iv. Example: detergent in water
2. Low vapor pressure
a. Strong hydrogen bonds prevent water from escaping easily
b. Prevents bodies of water from drying up
C. Specific Heat Capacity
1. C = 4.184 J/(gC)
2. Caused by hydrogen bonding
3. Heat released by water during winter; heat absorbed by water in summer
II. Water, Vapor, and Ice
A. Evaporation and Condensation
1. Evaporation
a. Vaporization that occurs at the surface of a liquid that is not boiling
b. Occurs as the result of absorption of heat of vaporization
c. Involves breaking intermolecular hydrogen bonds between water
Molecules
d. Explains high heat of vaporization and high boiling point
2. Condensation
a. Opposite of evaporation
b. Occurs as the result of releasing of heat of vaporization
c. Involves forming intermolecular hydrogen bonds between water
molecules
3. Relates to why temperatures in tropics are cooler than expected
B. Ice
1. Contraction of liquid as cooling occurs
2. Density will increase as temperature decreases because of less volume
3. Below 4C, density begins to decrease because water begins to no
longer behave as a liquid.
4. Insulates heat under the ice
5. Less dense than water, allowing it to float in water
6. Prevents freezing of bodies of water
III. Aqueous Systems
A. Solvents and Solutes
1. Chemically pure water does not exist naturally because of all the
substances it can dissolve.
2. Aqueous solutions
a. Solution in which the solvent is water
b. Solvent: dissolving medium, of which water is for aqueous solutions
c. Solute: particles dissolved in the solvent
d. Homogeneous mixture
B. The Solution Process
1. Solvation: process that occurs when a solute is dissolved
2. Insoluble: attractions between the ions in water in crystals are stronger than the
attractions exerted by the water
C. Electrolytes and Nonelectrolytes
1. Electrolytes: compounds that conduct an electric current in aqueous solution or the
molten state
2. Nonelectrolytes: compounds that do not conduct an electric current in aqueous
solution or the molten state
3. Strong vs. Weak Electrolytes (page 485)
D. Water of Hydration
1. Water in a crystal
2. Also called water of crystallization
3. Example: CuSO4 5H2O implies 5 moles of water to every copper and sulfate pair
IV. Heterogeneous Aqueous Systems
A. Suspensions: mixtures from which particles settle out upon standing
B. Colloids: mixtures containing particles that are intermediate in size between those of
suspensions and true solutions
C. Tyndall effect: the scattering of light in all directions
D. Table 17.6 (KNOW)
Appendix K
Acid and Base Measurements
Part I: The Difference Between Acids and Bases
For any aqueous system, the product of the hydrogen-ion concentration and the hydroxide-ion
concentration is known as the ion-product constant for water (Kw). This is always going to have a value
of 1.0 x 10-14 M2.
  OH   1.0  10
Kw  H
+
14
-
M
2
This is because as hydrogen-ion concentration increases, there has to be less hydroxide-ion
concentration. Since there is only so much space in water for one to occur, as one goes up, the other
concentration goes down.
Take a look at the following reaction when hydrogen chloride dissolves in water to become hydrochloric
acid:
H O
+
-
2
 H(aq)  Cl(aq)
HCl(g) 
More hydrogen ions are being dissociated in water than are hydroxide ions. Since more hydrogen ions
are dissociating in water, the [H+] is greater, meaning that HCl is an acid. This means that this solution is
an acidic solution, a solution in which the concentration of hydrogen ions ([H+]) is greater than the
concentration of hydroxide ions ([OH-]). This implies that [H+] > 1.0 x 10-7 M.
However, take a look at the following reaction when sodium hydroxide dissolves in water:
H O
+
-
2
 Na(aq)  OH(aq)
NaOH(s) 
More hydroxide ions are being dissociated in water than are hydrogen ions. Since more hydroxide ions
are dissociating in water, the [OH-] is greater, meaning that NaOH is a base. This means that this
solution is a basic solution (also known as an alkaline solution), a solution in which the concentration of
hydroxide ions ([OH-]) is greater than the concentration of hydrogen ions ([H+]). This implies that [OH-] >
1.0 x 10-7 M.
Part II: pH and pOH
A person can express hydrogen-ion and hydroxide-ion concentrations in terms of moles per liter, but this
involves working with a lot of scientific notation, which can get really boring and sometimes really difficult
to use. A Danish scientist named Sren Srensen came up with the idea of a pH scale, a means of
rating concentrations of bases and acids on a scale from 0 to 14 with neutral solutions having a pH of 7.
By taking the negative logarithm of a hydrogen-ion concentration, one can easily calculate the pH of a
solution as follows:
 
pH  -log H
+
The pH for acids generally falls between 0 and 7, with 0 being the most acidic and 7 being neutral.
Similarly, bases fall between 7 and 14 with 14 being the most basic and 7 being neutral.
By the same token, if one knows the pH and wishes to find the molar concentration (molarity) of the
solution, one can take 10 to the negative power of the pH as follows:
H  10
+
pH
Sometimes it is helpful, especially when one is more concerned about working with bases, to measure
pOH, the negative logarithm of the concentration of hydroxide ions as follows:
 
-
pOH  -log OH
pOH works opposite of pH since pH + pOH = 14. pOH for acids generally falls between 7 and 14, with
14 being the most acidic and 7 being neutral. Similarly, bases fall between 0 and 7 with 0 being the most
basic and 7 being neutral.
The same way you can work back to find the concentration of hydrogen ions in solution by knowing pH,
you can find the concentration of hydroxide ions in solution by knowing pOH:
OH  10
-
pOH
Realize, especially when double-checking figures, that the following is always true in an aqueous
solution:
pH + pOH = 14
   
Kw  H+  OH-  1.0  1014 M 2
Appendix L
Acid and Base Notes
I.
Acid-Base Theories
A.
___________________________ Acids and Bases
1.
Acids contain hydronium ions (H3O+) commonly referred to as hydrogen ions
(H+) that dissociate in water
a.
Different acids release different numbers of H+, known as protons since
the hydrogen loses its electron, resulting in only one proton (positive
charge)
Acid
HNO3 (nitric acid)
HC2H3O2 (acetic acid)
HCl (hydrochloric acid)
HBr (hydrobromic acid)
HF (hydrofluoric acid)
HI (hydroiodic acid)
H2SO4 (sulfuric acid)
H2CO3 (carbonic acid)
H3PO4 (phosphoric acid)
H3PO3 (phosphorous acid)
Common Acids and Types
Number of H+ ions per mole
Type of acid
b.
c.
d.
2.
B.
Not all compounds containing hydrogen are acids
Not all hydrogens in an acid will necessarily dissociate in water
Dissociation only occurs when very polar bonds are present because the
hydrogen ions are stabilized by dissolving in solution (i.e., forming
hydronium ions in solution)
Bases contain hydroxide ions (OH-) that dissociate in water
a.
Differences in solubility in water (page 227, “Solubility Rules for Ionic
Compounds”)
1.
High solubility: KOH, NaOH, hydroxides with Group 1 elements
2.
Low solubility: Ca(OH)2, Mg(OH)2, hydroxides with Group 2
elements
b.
React with acids to produce salt and water via double-replacement reaction
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
______________________________________________ Acids and Bases
1.
Some bases do not give off hydroxide ions but are still basic (i.e., NH3, Na2CO3)
2.
_________: hydrogen-ion (H+) donor; __________: hydrogen-ion (H+) acceptor
3.
__________________________________: what makes the solution acidic;
4.
5.
__________________________________: what makes the solution basic
Conjugate acid-base pair related by the loss or gain of a single hydrogen ion
Examples
NH3 (aq )  H2O(aq )  NH+4 (aq )
Ammonia
hydrogen ion
acceptor; BronstedLowry base
Water
hydrogen ion
donor; BronstedLowry acid

Ammonium ion
Conjugate acid
OH- (aq )
Hydroxide ion
Conjugate base
since this makes
the solution basic
HCl(aq )  H2O(l )  H3O+ (aq )  Cl- (aq )
Hydrochloric acid
Water
hydrogen ion
hydrogen ion
donor; Bronstedacceptor; BronstedLowry acid
Lowry base
6.
C.
Hydronium ion
Conjugate acid
since this makes the
solution acidic
Chloride ion
Conjugate base
__________________________________ substance
a.
Substance that can act as either an acid or a base
b.
Example: _______________________________
Lewis Acids and Bases
1.
Lewis _________: can accept a pair of electrons to form a covalent bond
2.
Lewis _________: can donate a pair of electrons to form a covalent bond
D.
Summary of acids and bases
Type
Arrhenius
Bronsted-Lowry
Lewis
II.
Acid-Base Definitions
Acid
H+ producer
H+ donor
Electron-pair acceptor
Base
OH- producer
H+ acceptor
Electron-pair donor
Strengths of Acids and Bases
A.
Strong and Weak Acids and Bases
1.
Strong acids completely ionize in water
2.
Weak acids only slightly ionize in water
B.
Acid dissociation constant
HCl(aq )  H2O(l )  H3O+ (aq ) + Cl- (aq )
[acid]
+
[H ]
[conjugate base]
Acid dissociation constant: _____________________________________________________
1.
Calculation done at equilibrium
2.
The smaller the constant, the less likely the acid will ionize in water
3.
The smaller the constant, the weaker the acid
4.
Each hydrogen ionizing in water has a different ionization constant
C.
Base dissociation constant
NH3 (aq )  H2O(aq )  NH+4 (aq ) + OH- (aq )
[base]
[conjugate acid]
[OH ]
Base dissociation constant: ____________________________________________________
1.
Calculation done at equilibrium
2.
The smaller the constant, the less likely the base will ionize in water
3.
The smaller the constant, the weaker the base
Relative Strengths of Common Acids and Bases
Substance
Formula
Relative Strength
HCl
Hydrochloric acid
HNO3
Strong acids
Nitric acid
Sulfuric acid
H2SO4
Phosphoric acid
H3PO4

Ethanoic acid
CH3COOH
|
Carbonic acid
H2CO3
Increasing strength of acid
Hydrosulfuric acid
H2S
Hypochlorous acid
HclO
Boric acid
H3BO3
Neutral solution
Sodium cyanide
Ammonia
Methylamine
Sodium silicate
Calcium hydroxide
Sodium hydroxide
Potassium hydroxide
NaCN
NH3
CH3NH2
Na2SiO3
Ca(OH)2
NaOH
KOH
Increasing strength of base
|

Strong bases
Appendix M
Neutralization Notes
I.
Neutralization Reactions
A.
Acid-Base Reactions
1.
An acid reacts with a base to produce water and salt (generally)
2.
Examples (strong acids reacting with strong bases)
a.
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
b.
H2SO4(aq) + 2KOH(aq)  K2SO4(aq) + H2O(l)
3.
Solubilities of salts dictated by the Solubility Rules for Ionic Compounds (p. 227)
4.
Neutralization reaction
a.
Reaction in which an acid and a base react in an aqueous solution to
produce a salt and water
b.
A double-replacement reaction
c.
Neutralization will not necessarily occur between weak acids and/or weak
bases
d.
Type of reaction to prepare pure samples of salt (i.e., NaCl from the
reaction shown above)
B.
Titration
1.
Process of adding a known amount of solution of known concentration to
determine the concentration of another solution
2.
Steps
a.
A measured volume of an acid solution of unknown concentration is added
to a flask.
b.
Several drops of an indicator are added to the solution.
c.
Measured volumes of a base of known concentration are mixed into the
acid until the indicator just barely changes color and maintains that color.
This occurs at the “end point”.
3.
Examples for strong acids and strong bases (refer to Chapter 20 Notes)
4.
When titration is complete at the end point, the contents of the flask are only salt
and water.
Example 1:
How many moles of sulfuric acid are required to neutralize 0.75 mol of potassium
hydroxide?
Solution 1:
It helps to first know the equation of neutralization. A reaction between sulfuric acid (a
strong acid) and potassium hydroxide (a strong base) is a double-replacement reaction.
Example 2:
What is the molarity of sodium hydroxide if 20.0 mL of the solution is neutralized by 17.4
mL of 1.00M phosphoric acid?
It helps to first know the equation of neutralization. A reaction between phosphoric acid (a
strong acid) and sodium hydroxide (a strong base) is a double replacement reaction.
C.
Titration Curves
1.
Strong-Acid; Strong-Base Titration
a.
The end point (equivalence point) has a pH at or very close to when the pH
= 7.
b.
Both acid and base completely ionize and therefore do not create
hydrolyzing salts (salts produced that remove hydrogen ions from water or
donate hydrogen ions to water)
c.
2.
c.
Titration curve for a strong-acid, strong-base titration
Strong-Acid; Weak-Base Titration
a.
The end point (equivalence point) has a pH < 7.
b.
Hydrolyzing salts that are produced from the titration will donate hydrogen
ions to the water, increasing the concentration of hydrogen ions ([H+]) and
therefore decrease the pH of the solution comparatively.
Titration curve for a strong-acid, weak-base titration
3.
Weak-Acid; Strong-Base Titration
a.
The end point (equivalence point) has a pH > 7.
b.
Hydrolyzing salts that are produced from the titration will remove hydrogen
ions to the water, decreasing the concentration of hydrogen ions ([H+]) and
therefore increase the pH of the solution comparatively.
c.
Titration curve for a weak-acid, strong-base titration
4.
Indicator selection for titration
a.
The indicator’s point of color change must be taken into consideration.
i.
Bromythol blue works for strong-acid, strong-base titrations
because it will change to green right at or close to pH = 7.
Phenolphthalein has an end point at pH = 8 but is often used
because it works even if one is colorblind. Also, the pH changes so
rapidly near the end point (4 to 10 with one drop of base) that it will
work well.
ii.
Methyl red works for strong-acid, weak-base titrations because it
will change color at pH = 5. Other indicators would not show a
change or would show it too prematurely.
iii.
Phenolphthalein works well for weak-acid, strong-base titrations
because of its end point in the upper pH range. Alizarin yellow R
works also as its color will change in the range of 10 < pH < 12.
b.
pH meters
i.
Often used in industry for more precise measurement of endpoint
ii.
Helpful in creating a titration curve
Summary of titrations and solutions
a.
Strong acid + Strong base = Neutral solution
b.
Strong acid + Weak base = Acidic solution
c.
Weak acid + Strong base = Basic solution
5.
Appendix N
Practice Tests
Unit Five Practice Test
PT – DC – U5
Multiple Choice.
On the answer sheet for each question, write the upper-case letter of the answer
that best completes or answers the statement or question in the adjacent
corresponding blank.
LPS Standard(s): 12.2.4a, 12.2.5a, 12.2.5b
State Standard(s): 12.8.3
1.
Which scientist developed a model with a positive region that encapsulated negatively-charged
particles he discovered?
(A) Dalton (B) Bohr (C) Rutherford (D) Thomson (E) Schrodinger
2.
Which scientist was responsible for discovering the nucleus and developed a model of the atom
that incorporated this in his model of the atom?
3.
Which scientist developed an indivisible model of the atom?
4.
Which scientist developed the current model of the atom?
(A) Dalton (B) Bohr (C) Rutherford (D) Thomson
(E) Schrodinger
5.
Which scientist described the model of the atom based on electrons in orbits that could be
quantized?
6.
Which scientist discovered the electron?
(A) Dalton (B) Bohr (C) Rutherford
(D) Thomson
(E) Schrodinger
7.
Which scientist developed the earliest model of the atom based on experimentation?
8.
In the Bohr model of the atom, an electron in an orbit has a fixed ___.
(A) position
(B) color
(C) energy
9.
The quantum mechanical model exactly predicts which characteristic of electrons in an atom?
(A) position (B) energy (C) orbit (D) charge (E) None of the choices listed
10.
Which of the following is an accurate description of Thomson’s model of the atom?
(A) Electrons occupy fixed positions around the protons, which are at the center of the atom.
(B) The electrons orbit in specified energy levels around the protons, which are at the center of
the atom.
(C) The electrons, like “raisins,” are stuck into a lump of protons, like “dough,” in a “plum
pudding” atom.
(D) The electrons and protons move throughout the atom.
LPS Standard(s): 12.2.4a, 12.2.5a, 12.2.5c
11.
State Standard(s): 12.3.1a
In which of the following is the number of neutrons correctly represented?
(A) 167 N has 9 neutrons.
(D) 239
94 Pu has 333 neutrons.
(B) 146 C has 14 neutrons.
(E) 42 He has 6 neutrons.
(C) 188 O has 8 neutrons.
12.
Which of the following sets of symbols represents isotopes of the same element?
90
90
51
52
(C) 50
(A) 90
42 J 43 J 44 J
19 L 19 L 19 L
(B) 168 O 179 O 18
10 O
(D) None of those listed
13.
How many electrons are in a neutral atom of the isotope plutonium-239?
(A) 94
(B) 145
(C) 239
(D) 333
(E) None of those listed
14.
How many protons are in a neutral atom of the isotope plutonium-239?
(A) 94
(B) 145
(C) 239
(D) 333
15.
How many neutrons are in a neutral atom of the isotope plutonium-239?
(A) 94
(B) 145
(C) 239
(D) 333
16.
What is the mass number of the isotope plutonium-239?
(A) 94
(B) 145
(C) 239
(D) 333
17.
What is the atomic number of the isotope plutonium-239?
(A) 94
(B) 145
(C) 239
(D) 333
Refer to the following isotope for questions 18 – 20.
36 
17 Cl
18.
How many neutrons does the ion of this isotope contain?
(A) 17
(B) 18
(C) 19
(D) 36
19.
How many electrons does the ion of this isotope contain?
(A) 17
(B) 18
(C) 19
(D) 36
20.
How many protons does the ion of this isotope contain?
(A) 17
(B) 18
(C) 19
(D) 36
LPS Standard(s): ---
State Standard(s): 12.3.1, 12.3.6
21.
In the first principal energy level (n = 1), what orbitals can exist in that energy level?
(A) s and p (B) s, p, and d (C) Only s (D) Only p (E) Only d
22.
If only two electrons occupy two p orbitals, what is the direction of the spins of these two
electrons?
(A) Both are always clockwise.
(B) Both are always counterclockwise.
(C) They are either both clockwise or both counterclockwise.
(D) One is clockwise and the other is counterclockwise.
Refer to the following diagram for questions 23 – 26.
23.
What type of orbital does the picture represent?
(B) p
(C) d
(D) f
(E) I
(A) s
24.
What is the maximum number of this kind of orbital in any principal energy level?
(A) 1
(B) 2
(C) 3
(D) 5
(E) 7
25.
How many electrons can all the orbitals of this sublevel hold collectively in any energy level?
(A) 2
(B) 6
(C) 10
(D) 14
(E) infinitely many
26.
What does the black spaces near the center in the picture represent?
(A) areas of high electron concentration
(C) modes of movement for electrons
(B) a node
(D) areas where protons will move
27.
Which electron configuration of the 4d energy sublevel is the most stable?
(A) 4d1
(B) 4d2
(C) 4d3
(D) 4d4
(E) 4d5
28.
Which electron configuration of the 4d energy sublevel is the most stable?
(B) 4d7
(C) 4d8
(D) 4d9
(E) 4d10
(A) 4d6
29.
In the second principal energy level (n = 2), what orbitals can exist in that energy level?
(A) s and p (B) s, p, and d (C) Only s (D) Only p (E) Only d
30.
If only two electrons occupy one p orbital, what is the direction of the spins of these two
electrons?
(A) Both are always clockwise.
(B) Both are always counterclockwise.
(C) They are either both clockwise or both counterclockwise.
(D) One is clockwise and the other is counterclockwise.
31.
What type of orbital has the greatest number of orientations?
(B) p
(C) d
(D) f
(E) I
(A) s
32.
What type of orbital is the only one contained in a hydrogen atom?
(B) p
(C) d
(D) f
(E) I
(A) s
33.
Which of the following sublevels exists in boron but not beryllium?
(B) p
(C) d
(D) f
(E) I
(A) s
34.
How many p orbitals can exist in any energy level?
(A) 1
(B) 2
(C) 3
(D) 5
(E) 7
35.
What is the maximum number of electrons that all the p orbitals of an energy level can hold?
(A) 2
(B) 4
(C) 6
(D) 8
(E) 10
36.
The electron configuration for rhodium (Rh) is ___.
(C) 1s22s22p63s23p64s24d104p65s25d7
(A) 1s22s22p63s23p64s23d104p65s24d7
2
2
6
2
6
2
10
6
1
8
(B) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d
(D) 1s22s22p63s23p64s24d104p65s15d8
37.
The electron configuration for copper (Cu) is ___.
(C) 1s22s22p63s23p64s24d9
(A) 1s22s22p63s23p64s23d9
(B) 1s22s22p63s23p64s13d10
(D) 1s22s22p63s23p64s14d10
38.
The electron configuration for the rubidium ion (Rb+) is ___.
(C) 1s22s22p63s23p64s23d104p65s1
(A) 1s22s22p63s23p64s23d104p6
2
2
6
2
6
2
10
6
2
10
6
(B) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p
39.
The electron configuration for calcium (Ca) is ___.
(C) 1s22s22p103s23p4
(A) 1s22s22p23s23p24s2
2
2
2
6
2
(B) 1s 2s 3s 3p 3d
(D) 1s22s22p63s23p64s2
40.
The electron configuration for vanadium (V) is ___.
(C) 1s22s22p63s23p64s24d3
(A) 1s22s22p63s23p64s23d3
(B) 1s22s22p63s23p64s13d4
(D) 1s22s22p63s23p64s14d4
41.
The electron configuration for chromium (Cr) is ___.
(C) 1s22s22p63s23p64s24d4
(A) 1s22s22p63s23p64s23d4
2
2
6
2
6
1
5
(B) 1s 2s 2p 3s 3p 4s 3d
(D) 1s22s22p63s23p64s14d5
42.
The electron configuration for the selenium ion (Se2–) is ___.
(C) 1s22s22p63s23p64s23d104p4
(A) 1s22s22p63s23p64s23d104p6
2
2
6
2
6
2
10
5
(B) 1s 2s 2p 3s 3p 4s 3d 4p
43.
The electron configuration for the sodium ion (Na+) is ___.
(B) 1s22s22p6
(C) 1s22s22p63s2
(A) 1s2
44.
The electron configuration for the fluoride ion (F–) is ___.
(B) 1s22s22p4
(C) 1s22s22p6
(A) 1s22s22p2
45.
The oxide ion (O2–) has the same electron configuration as ___.
(A) F
(B) C
(C) N
(D) Ne
46.
The calcium ion (Ca2+) is isoelectronic to ___.
(A) K
(B) Ar
(C) Br
(D) Kr
47.
How many half-filled orbitals are there in a nitrogen atom?
(A) 0
(B) 1
(C) 2
(D) 3
(E) 4
48.
How many unpaired electrons are there in a sulfur ion?
(A) 0
(B) 1
(C) 2
(D) 3
(E) 4
49.
How many unpaired electrons are there in a chloride ion?
(A) 0
(B) 1
(C) 2
(D) 3
(E) 4
50.
What is the number of electrons in the outermost energy level of a boron atom?
(A) 1
(B) 2
(C) 3
(D) 4
(E) 5
LPS Standard(s): 12.2.5d
State Standard(s): 12.3.2b
51.
What is a general classification for the elements of neon, calcium, and oxygen?
(A) Group A elements
(D) noble gases
(B) Group B elements
(E) halogens
(C) Group C elements
52.
What is a general classification for the gold, silver, and zinc?
(A) Group A elements
(D) noble gases
(B) Group B elements
(E) halogens
(C) Group C elements
53.
Which of the following elements is a representative element?
(A) cesium
(B) chromium
(C) californium
(D) cerium
54.
Which of the following groupings contains only transition metals?
(A) Magnesium, chromium, silver
(C) Copper, cobalt, cadmium
(B) Nickel, iron, polonium
(D) Aluminum, magnesium, lithium
55.
Of the elements Pt, As, V, Li, and Kr, which is a metalloid?
(A) Pt
(B) As
(C) V
(D) Li
(E) Kr
56.
What is another name for a family of elements in the periodic table called?
(A) period
(B) transition
(C) list
(D) group
57.
Who first arranged the elements according to atomic number and is responsible for our current
periodic table of elements?
(A) John Dalton
(D) Antoine Lavoisier
(B) Dmitri Mendeleev
(E) Henry Moseley
(C) Louis Pasteur
58.
What is a horizontal row of elements in the periodic table called?
(A) period
(B) transition
(C) list
(D) group
59.
The periodic law states that there is a periodic repetition of the physical and chemical properties
of elements ___.
(A) when they are arranged in the order Dmitri Mendeleev ordered them
(B) when they are arranged in the order Henry Moseley ordered them
(C) when they are arranged in the order Eugene Kirionov ordered them
(D) if only metalloids are considered
60.
The symbol for the 4th period Group 4A element is ___.
(A) As
(B) Ge
(C) In
(D) Sn
(E) none of those listed
61.
State Standard(s): 12.3.1, 12.3.2b
What is true of the electron configurations of the inner transition metals?
(A) The outermost s and f sublevels are very close in energy and have electrons in them.
(B) The outermost s and p sublevels are partially filled.
(C) The outermost s and d sublevels are very close in energy and have electrons in them.
(D) The outermost s and p sublevels are filled.
62.
What is true of the electron configurations of cobalt, molybdenum, and titanium?
63.
What is true of the electron configurations of calcium, phosphorus, and fluorine?
64.
The symbol of the first element that fills electrons in the s sublevel is ___.
(A) B
(B) H
(C) He
(D) Sc
65.
The symbol of the second element that fills electrons in the d sublevel is ___.
(A) Ca
(B) Sc
(C) Ti
(D) Zn
66.
The symbol for the 4th period element containing only 6 3d electrons is ___.
(A) Fe
(B) Mn
(C) Ru
(D) Tc
67.
The category of elements that is characterized by the filling of d orbitals is the ___.
(A) alkali metals
(C) inner transition metals
(B) alkaline earth metals
(D) transition metals
68.
On the periodic table, every period correlates to ___.
(A) a principal energy level
(C) an energy sublevel
(B) an atomic number
(D) an atomic mass
69.
The category of elements that would end with an s1 electron configuration would be the ___.
(A) alkali metals
(C) halogens
(D) noble gases
70.
The category of elements that would end with an p6 electron configuration would be the ___.
(A) alkali metals
(C) halogens
(D) noble gases
71.
Which of the following factors contributes to the greater ionization energy of the lower-atomicnumber elements in a family in the periodic table?
(A) Smaller distance from the nucleus
(C) Smaller number of protons in nuclei
(B) Smaller nuclei
(D) Greater number of valence electrons
72.
What term is used to describe the energy required to remove an electron from a gaseous atom?
(A) excitation energy
(D) heat of vaporization
(B) ionization energy
(E) electrolytic energy
(C) polarization energy
73.
Which of the following factors contributes to the lower ionization energy of the elements on the
left side of a period in the periodic table?
(A) Less shielding by inner electrons
(C) Smaller number of protons in nuclei
(B) Smaller nuclei
(D) Greater number of valence electrons
74.
Which group of the periodic table has the lowest electronegativity?
(A) 1A
(B) 6A
(C) 3A
(D) 7A
(E) 2A
75.
Of the following, which element has the greatest first ionization energy?
(A) strontium (B) phosphorus (C) fluorine (D) carbon
76.
Of the following, which element’s atoms have the largest ionic radius?
(A) lithium (B) potassium (C) rubidium (D) sodium
77.
Of the following, which element’s atoms have the largest atomic radius?
(A) iodine (B) rubidium (C) strontium (D) tellurium
78.
Of the following, what is the most electronegative element?
(A) iodine (B) rubidium (C) strontium (D) tellurium
79.
As you move from top to bottom down the first group of the periodic table, ___.
(A) the ionization energy decreases
(C) the electronegativity increases
(B) the atomic radii decrease
(D) the atomic mass decreases
80.
As you move from left to right across the third period of the periodic table, ___.
(A) the ionization energy increases
(C) the electronegativity decreases
(B) the atomic radii increase
(D) the atomic mass decreases
81.
What does an unstable nucleus NOT do to become more stable?
(A) lose electrons (B) lose neutrons (C) lose protons (D) gain protons
82.
Why do radioactive isotopes emit radiation?
(A) To achieve a proper protons to neutron ratio
(B) To ionize and feel like a noble gas
(C) To have enough protons to bond to other atoms
(D) To gain energy for more radiation
Choices for this section are as follows:
(A) alpha emission
(B) beta emission
(C) gamma emission
83.
Emission definitely involved when an element loses 4 in its mass number and 2 in its atomic
number
84.
Has no mass
85.
Least penetrating emission
86.
Consists of the same particles J.J. Thomson discovered
87.
Emission definitely involved when a mass number does not change but the atomic number does
88.
E = mc2 was an equation developed by Einstein to show the relationship between mass and
energy with ___ as a constant of proportionality between the two.
(A) specific heat (B) Planck’s constant (C) the speed of light (D) gamma radiation
89.
In the equation E = mc2, m represents ___.
(A) meters (B) mass (C) momentum
(D) Mr. Geist
State Standard(s): 12.1.2a; 12.3.1
90.
Which of the following half-lives indicates the most stable element?
(A) 2 seconds (B) 2 hours (C) 2 days (D) 2 weeks (E) 2 years
91.
If the half-life of sodium-24 is 15 hours, how much remains from a 10.0 g sample after 30 hours?
(A) 0.63 g (B) 1.3 g (C) 2.5 g (D) 5.0 g (E) None of the answers listed
92.
What is the half-life of iodine-131 if, after 24 days, 0.125 g remains from one 1.00 g starting
sample?
(A) 3 days (B) 6 days (C) 8 days (D) 12 days (E) None of the answers listed
Choices for this section are as follows:
(A) nuclear fission
(B) nuclear fusion
(C) None of those listed
93.
Occurs in most, if not all, stars of the universe
94.
Produces radioactive waste that must be contained in a secure facility
95.
The combination of atoms
State Standard(s): 12.3.6d
96.
What is the frequency of 6.502 x 10–7 m wavelength light?
(A) 4.614 x 105 Hz (B) 4.614 x 1014 Hz (C) 1.951 x 102 Hz
97.
What is the approximate energy of a photon having a frequency of 4 x 106 Hz?
(A) 3 x 10–27 J (B) 3 x 10–28 J (C) 2 x 10–42 J (D) 3 x 1041 J (E) 1 x 10–19 J
(D) 1.951 x 1011 Hz
For questions 98 – 100, refer to the following choices:
(A) Wave A ( = 5.1 x 105 Hz)
(B) Wave B ( = 5.1 x 1010 Hz)
(C) Waves A and B are equal for this.
(D) Not enough information provided or
not a valid question
98.
Which wave would possess the greater energy?
99.
Which wave would possess the longer wavelength?
100.
Which wave would possess the greater speed?
Unit Six Practice Test (Multiple Choice)
PT – DC – U6
Multiple Choice.
On the scantron for each question, fill in the rectangle of the corresponding letter of
the answer that best completes or answers the statement or question in the adjacent
LPS Standard(s): 12.2.5f
State Standard(s): 12.3.3c
1.
To form aluminum chloride, ___.
(A) one aluminum atom gains five electrons from five chlorine atoms
(B) one chlorine atom gains one electron from one aluminum atom
(C) three aluminum atoms gain three electrons from one chlorine atom
(D) three chlorine atoms gain three electrons from one aluminum atom
2.
To form potassium iodide, ___.
(A) one potassium atom gains seven electrons from an iodine atom
(B) one iodine atom gains seven electrons from a potassium atom
(C) one potassium atom gains one electron from an iodine atom
(D) one iodine atom gains one electron from a potassium atom
3.
To form strontium chloride, ___.
(A) one strontium atom gains two electrons from two chlorine atoms
(B) one chlorine atom gains two electrons from two strontium atoms
(C) two strontium atoms gain two electrons from one chlorine atom
(D) two chlorine atoms gain two electrons from one strontium atom
4.
To form sodium oxide, ___.
(A) one sodium atom gains six electrons from two oxygen atoms
(B) one oxygen atom gains six electrons from two sodium atoms
(C) one sodium atom gains two electrons from two oxygen atoms
(D) one oxygen atom gains two electrons from two sodium atoms
5.
How does aluminum obey the octet rule when reacting to form ionic compounds?
(A) It gains electrons. (B) It loses electrons. (C) It neither gains nor loses electrons.
6.
How does sulfur obey the octet rule when reacting to form ionic compounds?
(A) It gains electrons. (B) It loses electrons. (C) It neither gains nor loses electrons.
7.
What is the formula of the ion formed when sulfur achieves a noble-gas electron configuration?
(B) S+
(C) S–
(D) S2–
(E) S3–
(A) S6+
8.
What is the formula of the ion formed when lithium achieves a noble-gas electron configuration?
(B) Li+
(C) Li–
(D) Li2–
(E) Li3–
(A) Li2+
9.
How many electrons does barium have to lose/gain up to achieve a noble-gas electron
configuration?
(A) lose 1
(B) lose 2
(C) gain 1
(D) gain 2
10.
How many electrons does bromine have to lose/gain up to achieve a noble-gas electron
configuration?
(A) lose 1
(B) lose 2
(C) gain 1
(D) gain 2
11.
State Standard(s): 12.1.2a; 12.3.3c
Which of the following covalent bonds is NOT polar?
(A) C – C
(B) H – Br
(C) C – Cl
(D) C – Br
(E) C – S
12.
Which of the following pairs of elements can be joined by an ionic bond if the atoms ionize?
(A) sodium and carbon
(C) carbon and carbon
(B) nitrogen and carbon
(D) lithium and oxygen
13.
The polarity of the bond between a carbon atom and chlorine atom would best be identified as
a(n) ___ bond.
(A) ionic
(B) nonpolar covalent
(C) polar covalent
14.
In the bond between sodium and fluorine to make sodium fluoride, the sodium atom would have
___.
(D) a complete charge of –1
(A) a partial charge shown with +
(B) a partial charge shown with –
(E) no partial charge
(C) a complete charge of +1
15.
In the bond between hydrogen and chlorine in H – Cl, the hydrogen atom would have ___.
(D) no partial charge
(A) a partial charge shown with +
(E) a partial charge shown with +/–
(B) a partial charge shown with –
(C) a complete charge of –1
State Standard(s): 12.1.2a; 12.3.3c
16.
Which of the following elements can form diatomic molecules held together by single covalent
bonds and adhering to the octet rule?
(A) hydrogen
(B) nitrogen
(C) oxygen
(D) sodium
17.
___ covalent bonds possess the least energy.
(A) Double
(B) Single
(C) Triple
18.
Which of the following compounds is not covalently bonded?
(A) nitrogen
(B) oxygen
(C) strontium
(D) sulfur
19.
___ electrons are shared in a triple covalent bond.
(A) 0
(B) 2
(C) 4
(D) 6
20.
When reacting with atoms of their own element, nitrogen atoms form ___ covalent bonds to
create nitrogen molecules.
(A) single
(B) double
(C) triple
(D) no
21.
Which of the following elements can form diatomic molecules held together by single covalent
bonds, not following the octet rule, and not achieving 8 valence electrons per atom?
(A) hydrogen
(B) nitrogen
(C) oxygen
(D) sodium
22.
What is the total number of covalent bonds normally associated with a single carbon atom as the
central atom in a compound?
(A) 1
(B) 2
(C) 3
(D) 4
23.
How many covalent bonds are there in a covalently bonded molecule containing one phosphorus
atom and five chlorine atoms?
(A) 3
(B) 4
(C) 5
(D) 6
24.
Which of the following elements do not exist as diatomic molecules?
(A) bromine
(B) iodine
(C) oxygen
(D) phosphorus
25.
A ___ covalent bond is the only bond contained in carbon monoxide.
(A) double
(B) single
(C) triple
26.
A molecule or polyatomic ion with only double covalent bonds is ___.
(B) SO3
(C) CO2
(D) SO32–
(A) CH4
27.
A molecule or polyatomic ion with only single covalent bonds is ___.
(C) CO2
(D) CO32–
(A) HCCH
(B) CH4
28.
A molecule or polyatomic ion with a single covalent bond and a triple covalent bond is ___.
(B) HCN
(C) SO3
(D) N2
(A) H2O2
29.
How do atoms achieve noble-gas electron configurations in single covalent bonds?
(A) One atom completely loses two electrons to the other atom in the bond.
(B) Two atoms share two electrons.
(C) Two atoms share four electrons.
(D) Two atoms share six electrons.
30.
___ covalent bonds are the shortest in length.
(A) Double
(B) Single
(C) Triple
31.
Which of the following molecules has an electron dot structure that does NOT obey the octet
rule?
(C) PF3
(D) HCN
(E) CCl4
(A) NO
(B) CS2
32.
Which of the following violates the octet rule?
(B) IF3
(C) PF3
(D) SbF3
(A) NF3
(E) AsF3
33.
In which of the following compounds/ions is the octet expanded to include 14 electrons?
(B) BF3
(C) PCl5
(D) IF7
(E) SBr6
(A) H2O
34.
In which of the following compounds/ions is the octet expanded to include 12 electrons?
(B) SO42–
(C) SBr6
(D) SO32–
(E) PCl5
(A) SO3
35.
Which of the following compounds/ions violates the octet rule?
(B) CH4
(C) SCl6
(D) CO
(E) SO2
(A) CO2
For the following questions, identify the molecular geometry of the specified compounds using the
following choices:
(A) bent
(B) tetrahedral
(C) linear
(D) trigonal pyramidal
(E) none of the choices listed
36.
H2O (water)
39.
SO32– (sulfite ion)
37.
SO2 (sulfur dioxide)
40.
CH4 (methane)
38.
HCl (hydrochloric acid)
For the following questions, identify the molecular geometry of the specified compounds using the
following choices:
(A) trigonal planar
(D) linear
(B) octahedral
(C) trigonal bipyramidal
(E) none of the choices listed
41.
NH3 (ammonia)
44.
SiO2 (silicon dioxide)
42.
ClO (hypochlorite ion)
45.
PCl5 (phosphorus pentachloride)
43.
PO43– (phosphate ion)
46.
Which of the following compounds is the most polar?
(B) CH4
(C) CO2
(D) HBr
(A) Br2
47.
Diatomic molecules are ___ compounds.
(A) nonpolar covalent
(B) polar covalent
(C) ionic
48.
What would best describe sodium chloride as a compound?
(B) polar covalent
(C) ionic
49.
What would best describe oxygen gas as a compound?
(B) polar covalent
(C) ionic
50.
What would best describe sulfur dioxide as a compound?
(B) polar covalent
(C) ionic
51.
Which of the following causes the boiling point of HF to be much higher than that of HCl or HBr?
(A) hydrogen bonds
(C) covalent bonds
(B) van der Waals forces
(D) coordinate covalent bonds
52.
The strongest forms of intermolecular attractions are ___.
(A) dipole interactions
(B) dispersion forces
(C) hydrogen bonds
(D) ionic bonds
The weakest forms of intermolecular attractions are ___.
(A) dipole interactions
(C) hydrogen bonds
(D) ionic bonds
53.
54.
Why is hydrogen-bonding only possible with hydrogen?
(A) because hydrogen is the only atoms whose nucleus is not shielded by electrons when it is
involved in a covalent bond
(B) because hydrogen is the only atom that is the same size as an oxygen atom
(C) because hydrogen has the highest electronegativity of any element in the periodic table
55.
In an electric field, which region of the water molecule is attracted to the positive pole?
(A) the oxygen region of the molecule
(B) the hydrogen region of the molecule
(C) No part of the water molecule is attracted to the positive pole.
For the following questions, identify the following as ionic or molecular compounds using the following
(B) molecular compound
choices: (A) ionic compound
56.
The representative unit of these compounds is the molecule.
57.
These compounds tend to have higher conductivity.
58.
Li3PO4
59.
Have melting points usually below 300C
60.
Generally involve entirely nonmetallic elements
61.
Involve single, double, and/or triple bonds
62.
C2H5OH
63.
NH4Cl
64.
SCl2
65.
Li+ and S2– would be involved in this kind of compound.
Unit Six Practice Test (Short Answer)
PT – DC – U6
Name____________________________________________ Period_______________
LPS Standard(s): --Structures and Geometry.
For each of the following molecules or anions, draw the correct electron
dot structure for the molecule or anion.
1. PCl5 (phosphorus pentachloride)
3. NO21– (nitrite ion)
2. XeF2 (xenon difluoride)
4. ClO4– (perchlorate ion)
Short Answer.
Answer the following questions.
5.
Draw a structural formula of carbon dioxide. Also use + and – in the picture. Make sure the
picture is understandable and clear as to which atom is which. Then identify whether the
molecule is polar or nonpolar and explain why.
6.
Draw a structural formula of carbon monoxide. Also use + and – in the picture. Make sure the
picture is understandable and clear as to which atom is which. Then identify whether the
molecule is polar or nonpolar and explain why.
Unit Seven Practice Test (Multiple Choice)
PT – DC – U7
NOTE:
For all calculations on this test, use the constants and values from the periodic table of
elements provided to you.
Multiple Choice.
On the scantron for each question, fill in the rectangle of the corresponding letter of
the answer that best completes or answers the statement or question in the adjacent
LPS Standard(s): 12.2.4c
State Standard(s): 12.1.3e; 12.1.2d
1.
A gas has a pressure of 555 kPa at 227C. What will its pressure be at 53C if the volume does
not change?
(A) 59 kPa
(B) 130 kPa
(C) 362 kPa
(D) 578 kPa
(E) 2380 kPa
2.
A 30-g mass of carbon dioxide occupies 17.0 L at a pressure of 156 kPa. Find the volume of
carbon dioxide when the pressure is increased to 215 kPa at the same temperature.
(A) 0.000507 L
(B) 0.0811 L
(C) 12.3 L
(D) 23.4 L
(E) 1970 L
3.
At standard pressure and temperature, a gas occupies 22.4 L. If the volume and pressure of the
gas are changed to 3.50 L and 70.1 kPa, respectively, what will the new temperature of the gas
be in degrees Celsius?
(B) 0C
(C) 29.5C
(D) 2257C
(E) 2530C
(A) –243.6C
4.
A gas occupies a volume of 0.70 L at 50.C. What volume will the gas occupy at 100.C if the
pressure does not change?
(A) 0.35 L
(B) 0.61 L
(C) 0.81 L
(D) 1.4 L
(E) 2.8 L
5.
If a gas has a volume of 24.0 L at 20.C, at what temperature, in degrees Celsius, will the gas
have if the volume of the gas is increased to 55.0 L and the pressure does not change?
(A) –227°C
(B) 8.7°C
(C) 46°C
(D) 398°C
(E) 671°C
6.
Why does the pressure inside a container of gas increase if more gas is added to the container?
(A) because there is a corresponding increase in the number of particles striking an area of
the wall of the container per unit time
(B) because there is a corresponding increase in the temperature
(C) because there is a corresponding decrease in volume
(D) because there is a corresponding increase in the force of the collisions between the
particles and the walls of the container
7.
If the volume of a container holding a gas is reduced, the pressure within the container ___.
(A) decreases
(B) increases
(C) remains the same
8.
The temperature of a gas ___ when the gas is compressed.
(A) decreases
(B) increases
9.
If a balloon is rubbed vigorously and the volume remains constant, the pressure of the air inside
the balloon ___.
(A) decreases
(B) increases
10.
If there is no change in pressure for a sample of gas at 40C, which temperature will cause a
decrease in the volume of this gas?
(A) 260 K
(B) 280 K
(C) 300 K
(D) All of the temperatures listed.
11.
If a gas’s temperature is decreased and volume is held constant, its pressure will ___.
(A) decrease
(B) increase
(C) remain the same
12.
If the volume of a gas is increased and its temperature remains constant, what will happen to its
pressure?
(A) decrease
(B) increase
(C) remain the same
13.
If a capped syringe is heated, in which direction will the syringe plunger move?
(A) In
(B) Out
(C) No sliding will occur.
14.
How do gas particles respond to an increase in pressure?
(A) An increase in kinetic energy and an increase in temperature
(B) A decrease in kinetic energy and a decrease in volume
(C) An increase in temperature and an increase in volume
(D) A decrease in kinetic energy and an increase in temperature
15.
One way to increase the pressure of a gas is to ___.
(A) lower the kinetic energy of the gas particles
(B) decrease the number of gas particles
(C) increase the temperature of the gas
(D) increase the volume of the gas
16.
A sealed vessel contains 0.200 mol of oxygen gas, 0.100 mol of nitrogen gas, and 0.200 mol of
argon gas. The total pressure of the gas mixture is 5.00 atm. The partial pressure of the argon is
___ atm.
(A) 0.200
(B) 0.500
(C) 1.00
(D) 2.00
(E) 5.00
17.
Atmospheric pressure on the surface of Mars is 6.0 torr. The partial pressure of carbon dioxide is
5.7 torr. What percent of the Martian atmosphere is carbon dioxide?
(A) 5.0% (B) 6.0% (C) 95% (D) 96% (E) 98%
18.
A breathing mixture used by deep-sea divers contains helium, oxygen, and carbon dioxide. What
is the total pressure of the air if PHe = 84.5 kPa, PO2 = 2.8 kPa,and PCO2 = 0.1 kPa?
(A) 87.4 kPa
(B) 23.7 kPa
(C) 81.6 kPa
(D) None of the choices are correct.
19.
What happens to the partial pressure of nitrogen in the air if the air temperature is decreased?
(A) It decreases.
(B) It increases.
(C) It remains the same.
20.
If oxygen is added to a scuba tank, what happens to the total pressure of the air in the tank?
(A) It decreases.
(B) It increases.
LPS Standard(s): 12.2.4c
21.
State Standard(s): 12.1.3e; 12.1.2d
At standard temperature and pressure, 22.4 L of a gas is discovered to have a mass of 64.06
grams. What is the gas?
(B) CO2
(C) NO2
(D) P2O5
(E) Cl2
(A) SO2
22.
Which of the following gases would have the largest volume at 10C and 760 mm Hg?
(B) 10 g He
(C) 64 g O2
(D) 56 g N2
(A) 88 g CO2
23.
If 0.214 mol of argon occupies 652 mL at a given temperature and pressure, what is the volume
of 0.214 mol of butane at the same temperature and pressure?
(A) 652 mL
(B) 1304 mL
(C) 3047 mL
(D) 6094 mL
24.
Calculate the approximate temperature of a 0.500 mol sample of gas at 750. mm Hg and a
volume of 12.0 L.
(B) 11C
(C) 15C
(D) 288C
(A) –7C
25.
What is the approximate volume of gas in a 1.50 mol sample that exerts a pressure of 0.922 atm
and has a temperature of 10.0ºC?
(A) 13.0 L
(B) 14.2 L
(C) 37.8 L
(D) 378 L
26.
In collisions between ideal gas molecules, the total energy of the gas ___.
(A) decreases (B) increases significantly (C) increases slightly (D) remains the same
27.
The ideal gas law will be least likely to predict the behavior of ___ gas.
(A) helium
(B) hydrogen
(C) neon
(D) water vapor
28.
What is the pressure exerted by 1.2 mol of a gas with a temperature of 20.ºC and a volume of
9.5 L?
(A) 0.030 atm
(B) 1.0 atm
(C) 3.0 atm
(D) 30. atm
29.
What does the ideal gas law lead a scientist to calculate that the other gas laws cannot lead to
calculating?
(A) mass
(B) temperature
(C) volume
(D) energy
(E) pressure
30.
A sample of gas at 25ºC has a volume of 11 L and exerts a pressure of 660 mm Hg. How many
moles of gas are in the sample?
(A) 0.39 mol
(B) 3.9 mol
(C) 9.3 mol
(D) 87 mol
LPS Standard(s): 12.2.4b
State Standard(s): 12.1.2a; 12.1.2b
31.
At low temperatures and pressures, how does the volume of a real gas compare with the volume
that would be predicted for an ideal gas under the same conditions?
(A) It is much greater.
(B) It is much less.
(C) There is no difference.
32.
Another way of describing temperature is as ___ kinetic energy.
(A) adequate
(B) average
(C) static
(D) total
33.
An ideal gas cannot be ___.
(A) expanded
(B) compressed
(C) heated
(D) frozen
34.
Which term is used to describe changing a gas into a liquid?
(A) condensation (B) freezing (C) melting (D) vaporization
35.
Which of the following is not one of the assumptions of kinetic theory?
(A) Particles in a gas are assumed to have a significant volume.
(B) All gas particles move in constant random motion.
(C) Gases consist of hard spherical particles.
(D) No attractive and repulsive forces exist between gas particles.
State Standard(s): 12.3.5; 12.3.2e
36.
A 5.00-g sample of liquid water at 25.0C is heated by the addition of 84.0 J of energy. What is
the final temperature of the water if the specific heat capacity of liquid water is 4.184 J/(gC)?
(A) 4.02C
(B) 21.0C
(C) 29.0C
(D) 95.2C
37.
If the heat involved in a chemical reaction has a negative sign, ___.
(A) heat is lost to the surroundings
(B) heat is gained from the surroundings
(C) no heat is exchanged in the process
38.
If you want to cool a hot drink, it is best to use a spoon with a relatively ___ specific heat.
(A) low
(B) high
(C) The specific heat of the spoon does not matter.
39.
When heat is added to melting ice, its temperature ___.
(A) decreases
(B) increases
40.
If the heat of a substance decreases, its temperature ___.
(A) decreases
(B) increases
(C) stays the same
41.
Which of the following would you measure using a calorimeter?
(A) specific heat
(B) weight
(C) specific gravity
(D) density
42.
Compared to 100 grams of iron, a 10-gram sample of iron has ___.
(A) a higher specific heat
(B) a lower specific heat
(C) the same specific heat
43.
How many calories are there in 148 Joules?
(A) 6.61 J
(B) 35.4 J
(C) 148 J
(D) 619 J
(E) 3320 J
44.
In an exothermic reaction, the energy stored in the chemical bonds of the reactants is ___.
(A) equal to the energy stored in the bonds of the products
(B) greater than the energy stored in the bonds of the products
(C) less than the energy stored in the bonds of the products
(D) less than the heat released
(E) less than the heat absorbed
45.
If a substance gets colder, what happens to the average kinetic energy of the particles of the
substance?
(A) It decreases.
(B) It increases.
46.
How much energy would be released as the temperature of 150. grams of copper (specific heat =
0.0924 cal/(gC) drops from 79.0°C to 50.0°C?
(A) 3.00 cal
(B) 402 cal
(C) 3650 cal
(D) 3950 cal
47.
What is the energy required to melt one mole of a solid at its melting point?
(A) molar heat of melting
(C) molar heat of vaporization
(B) molar heat of fusion
(D) molar heat of solution
48.
A process that releases heat is a(n) ___ process.
(A) ectothermic
(B) endothermic
(C) exothermic
(D) polythermic
49.
Materials with a very high specific heat capacity can absorb little energy and show ___ change in
temperature.
(A) great
(B) little
(C) no
50.
The freezing of a liquid is a(n) ___ process.
(A) ectothermic
(B) endothermic
(C) exothermic
(D) polythermic
51.
The amount of heat absorbed by a vaporizing liquid ___ the amount of heat lost by the same
vapor if it is condensing.
(A) is the same as
(B) is less than
(C) is greater than
52.
Which of the following equations correctly represents an exothermic reaction?
(B) A + B + heat  C + D
(A) A + B  C + D + heat
53.
If Mr. Geist jumps into a cold pool, the pool will be experiencing a(n) ___ process.
(A) ectothermic
(B) endothermic
(C) exothermic
(D) polythermic
54.
When heat is released from steam, its temperature ___.
(A) increases
(C) depends on the amount of water
(B) decreases
(D) remains constant
55.
If you were to touch the flask in which an exothermic reaction were occurring, ___.
(A) the flask would feel the same as before the reaction started
(B) the flask would probably feel cooler than before the reaction started
(C) the flask would probably feel warmer than before the reaction started
56.
Which of the following is NOT a result of water’s high specific heat capacity?
(A) The temperature of water goes up rapidly as it absorbs solar energy.
(B) The temperature of cities near large bodies of water are moderated.
(C) For the same increase in temperature, iron needs to absorb only about one-tenth as
much energy as water.
(D) Water is an excellent medium for the storage of solar energy.
57.
How does the vapor pressure of water compare with the vapor pressures of other molecules of
similar size?
(A) It is higher.
(B) It is lower.
(C) It is about the same.
58.
Which of the following is primarily responsible for the high heat of vaporization of water?
(A) polar covalent bonds
(C) ionic bonds
(D) hydrogen bonds
59.
Which of the following is primarily responsible for the low vapor pressure of water?
(A) polar covalent bonds
(C) ionic bonds
(D) hydrogen bonds
60.
At what temperature does liquid water freeze?
(B) 4C
(C) 37C
(D) 100C
(A) 0C
State Standard(s): 12.3.2; 12.1.2b
61.
Which of the following substances is the most insoluble in water?
(A) HCl
(B) CH4
(C) CaCO3
(D) NaOH
62.
Which of the following compounds would be the most soluble in water?
(B) PCl5
(C) NaCl
(D) SF6
(A) CO2
63.
Which of the following substances is the most soluble in water?
(A) sodium fluoride
(B) oxygen gas
(C) carbon
(D) liquid bromine
64.
What is the term for the substance being dissolved into a solution?
(A) solvator
(B) solute
(C) emulsifier
(D) solvent
65.
Why are two polar substances able to dissolve in each other?
(A) Polar substances cannot dissolve in each other.
(B) They combine to produce a polar substance.
(C) There is no attractive or repulsive force between them.
(D) There are attractive and repulsive forces between them.
State Standard(s): 12.1.3e
66.
If the temperature of a liquid decreases, the solubility of a solid in the liquid generally ___.
(A) decreases
(B) increases
(C) stays the same
67.
If the pressure above a liquid decreases, the solubility of a gas in the liquid ___.
(A) decreases
(B) increases
(C) stays the same
68.
If the temperature of a liquid decreases, the solubility of a gas in the liquid generally ___.
(A) decreases
(B) increases
(C) stays the same
69.
The solubility of a gas at constant temperature and 2.0 atm pressure is 3.28 g/mL. What would
the solubility of the gas be at 4.0 atm pressure and without a temperature change?
(A) 0.410 g/mL
(B) 0.610 g/mL
(C) 2.44 g/mL
(D) 6.56 g/mL
70.
Which of the following substances is more soluble in hot water than in cold water?
(A) CO
(B) O2
(C) NH3
(D) LiCl
Unit Seven Practice Test (Short Answer/Calculation)
PT – DC – U7
Name____________________________________________ Period_______________
Short answer/Calculation. Answer the following questions. For questions involving mathematics,
show work or receive no credit, and include correct significant figures,
decimal places, and proper units.
State Standard(s): 12.1.3e
1.
How could you experimentally determine if a solution is unsaturated, saturated, or
supersaturated?
2.
Among the following gases, which of the following, if they possess the same temperature, would
have the slowest velocity (rate of effusion)? Circle your answer AND explain your answer.
Cl2
3.
C2H6
CO2
CO
How much heat (in kJ) does it take to convert 1550 g of water at 10.0C to steam at 105.0C?
Cice = 2.1 J/(gC); Cwater = 4.184 J/(gC); Csteam = 1.7 J/(gC); Hfus = 6.01 kJ/mol; Hvap = 40.7
kJ/mol
4.
During the metabolic process called respiration, your body obtains energy from the breakdown of
glucose as shown below.
C6H12O6(aq) + 6O2(g)  6H2O(l) + 6CO2(g)
What volume of O2, measured at 37C and 790.0 torr pressure, is required to react with 1.00 g of
glucose (C6H12O6)? Express the volume in milliliters.
5.
Why is ice less dense than water?
6.
Explain why water has a relatively high heat of vaporization.
Unit Eight Practice Test (Multiple Choice)
PT – DC – U8
LPS Standard(s): --Multiple Choice.
Identify the letter of the choice that best completes the statement or answers the
question.
1.
When an acid reacts with a base, what is one compound that is always formed?
(A) salt
(B) sugar
(C) carbohydrate
(D) protein
2.
Which of the following is a property of an acid?
(A) reactive with bases (B) nonreactive with bases (C) nonelectrolyte (D) nonconductive
3.
What is the charge on the hydronium ion?
(A) 2–
(B) 1–
(C) 0
(D) 1+
(E) 2+
4.
Which of the following is a property of an acid?
(A) sour taste
(B) strong color
(C) nonelectrolyte
(D) unreactive
5.
What is the name of the compound H3PO3?
(A) hydrophosphoric acid
(C) phosphoric acid
(B) hydrophosphorus acid
(D) phosphorous acid
6.
What is the name of the compound KOH?
(A) potassium oxygen hydride
(B) potassium (I) hydroxide
(C) potassium (II) hydroxide
(D) potassium hydroxide
7.
What is the formula of lithium hydroxide?
(D) Li(OH)2
(A) LiH
(B) LiOH
(C) LiH2
8.
What is the formula of chloric acid?
(A) HCl
(B) HClO
(C) HClO2
9.
10.
What is the formula of carbonic acid?
(C) H2CO3
(A) HC
(B) H4C
(D) HClO3
(D) H3CO3
What is the formula of hydrobromic acid?
(C) HBrO3
(D) HBrO4
(A) HBr
(B) HBrO2
11.
An Arrhenius acid ___.
(B) produces OH–
(A) produces H+
(C) accepts H+
(D) accepts OH–
12.
An Arrhenius base ___.
(B) produces OH–
(A) produces H+
(C) accepts H+
(D) accepts OH–
13.
A Bronsted-Lowry acid ___.
(B) donates OH–
(A) donates H+
(C) accepts H+
(D) accepts OH–
14.
A Bronsted-Lowry base ___.
(B) donates OH–
(A) donates H+
(C) accepts H+
(D) accepts OH–
15.
A substance that can behave as an acid or a base is ___.
(A) andropic
(B) amphoteric
(C) analgesic
(D) analytic
16.
Which of the following is a Bronsted-Lowry base, but not an Arrhenius base?
(B) NaOH
(C) LiOH
(D) HCl
(A) NH3
17.
In the chemical equation LiOH + H2O  OH– + LiOH2+, what is the conjugate acid?
(A) LiOH
(B) H2O
(C) OH–
(D) LiOH2+
18.
In the chemical equation NH3 + H2O  NH4+ + OH–, what is the conjugate base?
(B) H2O
(C) NH4+
(D) OH–
(A) NH3
19.
In the chemical equation HCl + H2O  Cl– + H3O+, what is the conjugate acid?
(C) Cl–
(D) H3O+
(A) HCl
(B) H2O
20.
In the chemical equation HCl + H2O  Cl– + H3O+, what is the conjugate base?
(A) HCl
(B) H2O
(C) Cl–
(D) H3O+
LPS Standard(s): --Identification.
Identify the following properties, definitions, or compounds as being those of
(A) acidic, (B) basic, or (C) neutral.
21.
Solution where pOH = 2
26.
Solution where pH = 3.8
22.
HNO3
27.
Solution where pOH = 7
23.
H2O
28.
Produce hydronium ions in solution
24.
H2SO4
29.
Solution where [H+] = 3.1 x 10–10 M
25.
KOH
30.
Solution where [OH–] = 4.8 x 10–4 M
31.
Distilled water containing only HC2H3O2 dissolved in it
32.
Distilled water containing only NaOH dissolved in it
33.
In HNO3 + H2O  NO3– + H3O+, what H2O would be
34.
A solution containing substantially more hydronium ions than hydroxide ions
35.
A solution containing substantially more hydrogen ion acceptors than hydrogen ion donors
36.
Hydrogen ion concentration of 1  10–7 M
37.
Hydroxide ion concentration of 3  10–4 M
38.
Hydronium ion concentration of 3  10–4 M
39.
pH = 6.0
40.
pOH = 6.0
Unit Eight Practice Test (Short Answer)
PT – DC – U8
Name_________________________________________________ Period__________________
LPS Standard(s): --Calculations.
Solve the following problems. Show work or receive no credit. Include proper units
and significant figures and/or decimal places.
41.
What volume (in milliliters) of 0.275M sulfuric acid is needed to neutralize 17.2 mL of 0.550M
sodium hydroxide?
42.
A 100.0 mL solution of sodium hydroxide is completely neutralized by 100.0 mL of 0.5000M
phosphoric acid. What is the concentration of the sodium hydroxide?
A lab assistant tests poolwater near a park and determines it to have a pOH of 5.9.
43.
What is the hydrogen-ion concentration [H+] of the solution?
44.
What is the hydroxide-ion concentration [OH-] of the solution?
45.
What is the pH of the solution?
46.
What is the ion-product constant of the solution?
47.
Is this solution acidic, basic, or neutral?
Identification. Identify the following piece of equipment involved in the titration process.
48.
49.
50.
Equipment:
48.
_________________________________________________________________________
49.
_________________________________________________________________________
50.
_________________________________________________________________________
Appendix O
Practice Test Keys
Unit Five Practice Test Key
PTK – DC – U5
1. D
2. C
3. A
4. E
5. B
6. D
7. A
8. C
9. E
10. C
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
A
C
A
A
B
C
A
C
B
A
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
C
C
C
D
C
B
E
E
A
D
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
D
A
B
C
C
A
B
A
D
A
41.
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
B
A
B
C
D
B
D
C
A
C
A
B
A
C
B
D
E
A
A
B
61.
62.
63.
64.
65.
66.
67.
68.
69.
70.
A
C
B
B
C
A
D
A
A
D
71.
72.
73.
74.
75.
76.
77.
78.
79.
80.
A
B
C
A
C
C
B
A
A
A
81.
82.
83.
84.
85.
86.
87.
88.
89.
90.
A
A
A
C
A
B
B
C
B
E
91. C
92. C
93. B
94. A
95. B
96. B
97. A
98. B
99. A
100. C
Unit Six Practice Test Key (Multiple Choice)
PTK – DC – U6
1.
2.
3.
4.
5.
6.
7.
D
D
D
D
B
A
D
8. B
9. B
10. C
11. A
12. D
13. C
14. C
15.
16.
17.
18.
19.
20.
21.
A
A
B
C
D
C
C
22.
23.
24.
25.
26.
27.
28.
D
C
D
C
C
B
B
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
B
C
A
B
D
C
C
A
A
C
D
B
E
D
43.
44.
45.
46.
47.
48.
49.
E
D
C
D
A
C
A
50.
51.
52.
53.
54.
55.
56.
B
A
C
B
A
A
B
57.
58.
59.
60.
61.
62.
63.
A
A
B
B
B
B
A
64. B
65. A
Unit Six Practice Test Key (Short Answer)
PTK – DC – U6
1.
2.
3.
4.
5.
Nonpolar; partial charges cancel out because of symmetry
6.
Polar; partial charges do not cancel out because of creation of
poles on each side
Unit Seven Practice Test Key (Multiple Choice)
PTK – DC – U7
1.
2.
3.
4.
5.
6.
7.
C
C
A
C
D
A
B
8. B
9. B
10. D
11. A
12. A
13. B
14. A
15.
16.
17.
18.
19.
20.
21.
C
D
C
A
A
B
A
22.
23.
24.
25.
26.
27.
28.
B
A
C
C
D
D
C
29.
30.
31.
32.
33.
34.
35.
A
A
B
B
D
A
A
36.
37.
38.
39.
40.
41.
42.
C
A
A
C
A
A
C
43.
44.
45.
46.
47.
48.
49.
B
B
A
B
B
C
B
50.
51.
52.
53.
54.
55.
56.
C
A
A
B
B
C
A
57.
58.
59.
60.
61.
62.
63.
B
D
D
A
B
C
A
64.
65.
66.
67.
68.
69.
70.
B
D
A
A
B
D
D
Unit Seven Practice Test Key (Short Answer/Calculation)
PTK – DC – U7
1.
Add more solute. If the solute dissolves in the solution, the solution is unsaturated. If the
solute will not dissolve in and goes to the bottom, the solution is saturated. If the the
solute causes the solution to crystallize, the solution is supersaturated.
2.
Cl2: 2(35.453) = 70.906 g/mol
C2H6: 2(12.011) + 6(1.0079) = 24.022 + 6.0474 = 30.069 g/mol
CO2: 1(12.011) + 2(15.999) = 12.011 + 31.998 = 44.009 g/mol
CO: 1(12.011) + 1(15.999) = 12.011 + 15.999 = 28.010 g/mol
Cl2 would have the slowest velocity (rate of effusion) because it has the greatest molar
mass.
3.
q3: Heat needed for water to elevate from 10.0C to 100.0C
q4: Heat needed to vaporize water into steam
q5: Heat needed for water to elevate from 100.0C to 105.0C
1550 g 4.184 J 100.0C  10.0C  1550 g 4.184 J 90.0C 





1
g  C
1
1
g  C
1
= 584000 J = 584 kJ
1550 g
1 mol
40.7 kJ
q4 


 3500 kJ
1
18.015 g 1 mol
1550 g 1.7 J 105.0C  100.0C 1550 g 1.7 J 5.0C
q 5  mCT 





1
g  C
1
1
g  C
1
= 13000 J = 13 kJ
q 3  mCT 
Total heat = 584 kJ + 3500 kJ + 13 kJ = 4097 kJ
4.
1.00 g C 6H12 O 6
1 mol C 6H12 O 6
6 mol O 2


 0.0333 mol O 2
1
180.155 g C 6H12 O 6 1 mol C 6H12 O 6
P: 790.0 torr (mm Hg)
n: 0.0333 mol O2
R: 62.396 (Ltorr)/(Kmol)
T: 37C + 273.15 = 310. K
PV = nRT
(790.0 torr)V = (0.0333 mol O2)(62.396 (Ltorr)/(Kmol))(310. K)
V 
0.0333 mol62.396 (L  torr)/(K  mol)310. K   0.815 L  815 mL O
2
790.0 torr 
5.
The structure of ice is a very regular, open framework in which the water molecules are
farther apart from each other than they are in liquid water, in great part due to hydrogen
bonds. When ice melts, this open framework collapses and the water molecules move
closer together. As a result, the water is denser than the ice.
6.
Water has a high heat of vaporization as a result of its hydrogen bonding. Because of its
extensive network of hydrogen bonds, the molecules of water are held together more
tightly than are the molecules of many other liquids. The attractive force of these
hydrogen bonds must be overcome in order for water to vaporize.
Unit Eight Practice Test Key (Multiple Choice)
PTK – DC – U8
1.
2.
3.
4.
A
A
D
A
5.
6.
7.
8.
D
D
B
D
9. C
10. A
11. A
12. B
13.
14.
15.
16.
A
C
B
A
17.
18.
19.
20.
D
D
D
C
21.
22.
23.
24.
B
A
C
A
25.
26.
27.
28.
B
A
C
A
29.
30.
31.
32.
B
B
A
B
33.
34.
35.
36.
B
A
B
C
Unit Eight Practice Test Key (Short Answer)
PTK – DC – U8
41.
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l)
1 L H2 SO 4
0.550 mol NaOH 0.0172 L NaOH 1 mol H2 SO 4



1L NaOH
1
2 mol NaOH 0.275 mol H2 SO 4
= 0.0172 L H2SO4 = 17.2 mL H2SO4
42.
H3PO4(aq) + 3NaOH(aq)  Na3PO4(aq) + 3H2O(l)
0.5000 mol H3PO 4 0.1000 L H3PO 4 3 mol NaOH
1



1 L H3PO 4
1
1 mol H3PO 4 0.1000 L NaOH
= 1.500 M NaOH
43.
[H+] = 10–8.1 = 7.94  10–9 M
(8.1 from answer to question 45)
44.
[OH-] = 10–5.9 = 1.26  10–6 M
45.
pH + pOH = 14
pH + 5.9 = 14
pH = 8.1
48.
Funnel
49.
Ring stand
46.
1  10–14 M2
50.
Beaker
47.
Basic
37.
38.
39.
40.
B
A
A
B

Differentiated Chemistry Worksheet and Laboratory

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