ib chemistry ii booklet fall...
TRANSCRIPT
Class: Ms. Manakul
Student Name: ____________________
Student Hour: _________
IB Chemistry II
Booklet
Fall Semester
Periodic Table of the Elements
hydrogen
1
H 1.0079
helium
2
He 4.0026
lithium
3
Li 6.941
beryllium
4
Be 9.0122
boron
5
B 10.81
carbon
6
C 12.011
nitrogen
7
N 14.007
oxygen
8
O 15.999
fluorine
9
F 18.998
Neon
10
Ne 20.179
sodium
11
Na 22.990
magnesium
12
Mg 24.305
aluminium
13
Al 26.982
silicon
14
Si 28.086
phosphorus
15
P 30.974
sulphur
16
S 32.06
chlorine
17
Cl 35.453
argon
18
Ar 39.984
potassium
19
K 39.098
calcium
20
Ca 40.08
scandium
21
Sc 44.956
titanium
22
Ti 47.90
vanadium
23
V 50.941
chromium
24
Cr 51.996
manganese
25
Mn 54.938
iron
26
Fe 55.847
cobalt
27
Co 58.933
nickel
28
Ni 58.71
copper
29
Cu 63.546
zinc
30
Zn 65.38
gallium
31
Ga 69.72
germanium
32
Ge 72.59
arsenic
33
As 74.922
selenium
34
Se 78.96
bromine
35
Br 79.904
krypton
36
Kr 83.80
rubidium
37
Rb 85.468
strontium
38
Sr 87.62
yttrium
39
Y 88.906
zirconium
40
Zr 91.22
niobium
41
Nb 92.906
molybdenum
42
Mo 95.94
technetium
43
Tc [98]
ruthenium
44
Ru 101.07
rhodium
45
Rh 102.91
palladium
46
Pd 106.4
silver
47
Ag 107.87
cadmium
48
Cd 112.41
indium
49
In 114.82
tin
50
Sn 118.69
antimony
51
Sb 121.75
tellurium
52
Te 127.60
iodine
53
I 126.90
xenon
54
Xe 131.30
caesium
55
Cs 132.91
barium
56
Ba 137.33
lutetium
71
Lu 174.97
hafnium
72
Hf 178.49
tantalum
73
Ta 180.95
tungsten
74
W 183.85
rhenium
75
Re 186.21
osmium
76
Os 190.2
iridium
77
Ir 192.22
platinum
78
Pt 195.09
gold
79
Au 196.97
mercury
80
Hg 200.59
thallium
81
Tl 204.37
lead
82
Pb 207.2
bismuth
83
Bi 208.98
polonium
84
Po [209]
astatine
85
At [210]
radon
86
Rn [222]
francium
87
Fr [223]
radium
88
Ra [226]
lawrencium
103
Lr [262]
rutherfordium
104
Rf [261]
dubnium
105
Db [262]
seaborgium
106
Sg [263]
bohrium
107
Bh [264]
hassium
108
Hs [265]
meitnerium
109
Mt [268]
darmstadtium
110
Ds [269]
roentgenium
111
Rg [272]
copernicium
112
Cn [277]
ununtrium
113
Uut [284]
flerovium
114
Fl [289]
ununpetium
115
*Uup [288]
livermorium
116
Lv [293]
ununseptium
117
*Uus [294]
ununoctium
118
*Uuo [299]
*Discovery reported by not verified
lanthanum
57
La 138.91
cerium
58
Ce 140.12
praseodymium
59
Pr 140.91
neodymium
60
Nd 144.24
promethium
61
Pm [145]
samarium
62
Sm 150.4
europium
63
Eu 151.96
gadolinium
64
Gd 157.25
terbium
65
Tb 158.93
dysprosium
66
Dy 162.50
holmium
67
Ho 164.93
erbium
68
Er 167.26
thulium
69
Tm 168.93
ytterbium
70
Yb 173.04
actinium
89
Ac [227]
thorium
90
Th 232.04
protactinium
91
Pa 231.04
uranium
92
U 238.03
neptunium
93
Np [237]
plutonium
94
Pu [244]
americium
95
Am [243]
curium
96
Cm [247]
berkelium
97
Bk [247]
californium
98
Cf [251]
einsteinium
99
Es [252]
fermium
100
Fm [257]
mendelevium
101
Md [258]
nobelium
102
No [259]
Relevant Equations
Ideal Gas Law PV = nRT
Speed of Light c = νλ
Heat q = mCΔT
Acid Range
pH = -log10[H3O+]
or
pH = -log10[H+]
Energy E = hν
Spontaneity Δ𝐺° = Δ𝐻° − 𝑇Δ𝑆° Arrhenius’ Equation k = Ae-Ea/RT
Arrhenius’ Point Slope ln k = (-Ea/RT) + ln A
Arrhenius’ Equation
with two points ln
𝑘1
𝑘2=
−𝐸𝑎
𝑅(
1
𝑇1−
1
𝑇2)
Spontaneity Δ𝐺° = −𝑅𝑇𝑙𝑛𝐾
Spontaneity Δ𝐺° = −𝑛𝐹𝐸°
Thermodynamic Quantities for Selected Substances
at 298.15 K (25°C)
Substance ΔH°f (kJ/mol
)
ΔG°f (kJ/mol)
S°f (J/mol K)
C2H2 (g) 227.4 209.9 200.9
NH3 (g) -45.9 -16.4 192.8
C6H6 (l) 49.1 124.5 173.4
CaCO3(s) -1207.6 -1129.1 91.7
CaO(s) -634.9 -603.3 38.1
CO2(g) -393.5 -394.4 213.8
CO(g) -110.5 -137.2 197.7
C(s) graphite 0 0 5.7
C(s) diamond 1.88 2.9 2.4
C2H6 (g) -84.68 -32.0 229.2
C2H5OH(l) -277.6 -174.8 160.7
C2H4 (g) 52.4 68.4 219.3
C6H12O6(s) -1273.3 -910.4 212.1
HBr(g) -36.3 -53.4 198.7
HCl(g) -92.3 -95.3 186.9
HF(g) -273.3 -275.4 173.8
HI(g) 26.5 1.7 206.6
CH4(g) -74.6 -50.5 186.3
CH3OH(l) -238.6 -166.6 126.8
C3H8(g) -103.85 -23.4 270.3
AgCl(s) -127.0 -109.8 96.3
NaHCO3(s) -950.8 -851.0 101.7
Na2CO3(s) -1130.7 -1044.4 135.0
NaCl(s) -411.2 -384.1 72.1
C12H22O11(s) -2226.1 -1544.3 360.24
H2O(l) -285.8 -237.1 70.0
H2O(g) -241.8 -228.6 188.8
O2(g) 0 0 205.2
H2(g) 0 0 130.7
H 2.2
Electron Affinity Table He ---
Li 1.0
Be 1.6
B 2.0
C 2.6
N 3.0
O 3.4
F 4.0
Ne ---
Na 0.9
Mg 1.3
Al 1.6
Si 1.9
P 2.2
S 2.6
Cl 3.2
Ar ---
K 0.8
Ca 1.0
Sc 1.4
Ti 1.5
V 1.6
Cr 1.7
Mn 1.6
Fe 1.8
Co 1.9
Ni 1.9
Cu 1.9
Zn 1.7
Ga 1.8
Ge 2.0
As 2.2
Se 2.6
Br 3.0
Kr ---
Rb 0.8
Sr 1.0
Y 1.2
Zr 1.3
Nb 1.6
Mo 2.2
Tc 2.1
Ru 2.2
Rh 2.3
Pd 2.2
Ag 1.9
Cd 1.7
In 1.8
Sn 2.0
Sb 2.1
Te 2.1
I 2.7
Xe ---
Cs 0.8
Ba 0.9
La 1.1
Hf 1.3
Ta 1.5
W 1.7
Re 1.9
Os 2.2
Ir 2.2
Pt 2.2
Au 2.4
Hg 1.9
Tl 1.8
Pb 1.8
Bi 1.9
Po 2.0
At 2.2
Rn ---
Fr 0.7
Ra 0.9
Ac 1.1
General Physical Constants
Atomic mass unit 1 amu = 1.66 x 10-24 kg
Avogadro’s Number N = 6.02 x 1023 particles/mol
Constant for H2O at 298 K Kw = 1.00 x 10-14 mol2 dm-6
Faraday’s Constant F = 9.65 x 104 C mol-1
Gas constant R = 8.31 L kPa K-1 mol-1
Ideal gas molar volume Vm = 22.4 L mol-1
Planck’s Constant h = 6.626 x 10-34 J s
Standard Temperature and
Pressure (STP) Conditions 273 K and 100 kPa
Speed of Light (in a vacuum) c = 3.00 x 108 m s-1
Specific heat capacity of H2O 4.18 kJ kg-1 K-1 = 4.18 J g-1 °C-1
Useful Conversion Factors
Energy: 1 cal = 4.184 J
Length: 1 angstrom = 0.100 nm
1 inch = 2.54 cm
Mass: 1 lb = 0.4536 kg
Pressure: 1atm = 101.3 kPa
1atm = 760 mm Hg
Temp: K = °C + 273
Volume: 1L = 1 dm3
1cm3 = 1 mL
Bond Enthalpies and average bond enthalpies at 298 K
Bond ΔH/kJ
mol-1
Bond ΔH/kJ
mol-1
Bond ΔH/kJ
mol-1
H – H 436 C – H 413 C – O 358
C – C 347 Si – H 318 C = O 746
C = C 612 N – H 391 C – N 286
C ≡ C 838 P – H 321 C = N 615
C ≈ C
(in benzene)
505 O – H 464
C ≡ N 887
Si – Si 226 S – H 364 C – F 467
N – N 158 F – H 568 C – Cl 346
N = N 410 Cl – H 432 C – Br 290
N ≡ N 945 Br – H 366 C – I 228
Reactivity Series
Metals Halogens
Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Tin Lead Hydrogen Copper Mercury Silver Platinum Gold
Fluorine Chlorine Bromine Iodine
Table of Contents
I. Syllabus page i
II. Classroom Map page iv
III. School Safety Drills page v
IV. Lab Safety Contract page vi
V. Unit Objectives page viii
VI. Objective Trackers page xi
VII. Homework
Unit 1: The Atom & Periodic Table page 1
Unit 2: Bonding page 14
Unit 3: Moles, Stoichiometry, & Chemical Reactions page 24
Unit 4: Energetics page 35
Unit 5: Entropy & Spontaneity page 43
Unit 6: Kinetics page 48
VIII. Labs
Specialized Tools Lab page 55
Empirical & Molecular Formula Lab page 58
Ideal Gas Lab page 60
Calorimetry Lab page 62
Entropy Lab page 64
Rate of Reaction Lab page 67
IX. Appendix
Stoichiometry Mole Map page A
VSEPR Sheet page B
IUPAC Tables page C
IB Data Tables page D
i
IB Chemistry II Syllabus
Lincoln College Preparatory Academy
Teacher: Melissa Manakul E-mail:[email protected]
Room: 207 Google Classroom: r3ykuj2 *Any communication to me either through e-mail or website will be responded to as soon as humanly possible.
MY WEEKLY SCHEDULE:
Day Monday Tuesday Wednesday Thursday Friday
Morning In classroom by
6:50 am
In classroom
by 6:50 am
In classroom
by 6:50 am
In classroom
by 6:50 am
In classroom
by 6:50 am
Afternoon Meeting Tutoring
2:30 – 3:30 Staff Meeting
Will leave by
3:00 pm
Will leave by
2:30 pm
COURSE DESCRIPTION:
Chemistry is an experimental science that combines academic study with the acquisition of practical and investigational
skills. It is called the central science, as chemical principles underpin both the physical environment in which we live and
all biological systems. Apart from being a subject worthy of study in its own right, chemistry is a prerequisite for many
other courses in higher education, such as medicine, biological science and environmental science, and serves as a useful
preparation for employment. International baccalaureate chemistry is divided in two types of courses, higher level and
standard level. Standard level comprises of core skills and activities, while higher level takes topics into greater depth.
Each are assessed both externally and internally and have prescribed number of hours with to accomplish the curriculum.
GOALS FOR SEMSTER:
1. You will be prepared to obtain a 4 or above on the IB Chemistry Exam
2. You will be prepared for IB Chemistry III or a collegiate General Chemistry course.
3. You will be able to design and execute your own lab.
REQUIRED MATERIALS:
Booklet: This is your bible in Chemistry. Each student receives a booklet at the beginning of the semester, which includes
all unit objectives, homework, labs, and supplemental material. This is to go home every night.
Lab Book: This will be a composition notebook (college ruled or grid paper preferred) which will house all your
completed labs and activities. This will be graded throughout the year, and left in the classroom.
Laptop: This will be used to access the textbook and Google Classroom where any announcements, calendar changes,
copies of activities, additional supplemental information, and extra copies of homework. We will also use it for a
variety of activities to be done at home or in the classroom through a variety of educational websites.
SUGGESTED MATERIALS:
Binder/Folder: This will hold any paper work, including but not limited to tests, quizzes, lab reports, and activities. You
must take this home every night. Any time Tyler or my record is incorrect, it is the student’s job to report this to
Ms. Manakul with a graded copy. This is the only way a grade can be changed for full credit.
Scientific Calculator: TI-30x is a preferred calculator, however any scientific calculator will do.
TOPICS COVERED:
(1) Atomic Structure & Periodicity (2) Bonding (3) Moles, Stoichiometry, & Chemical Reactions (4) Energetics (5)
Entropy & Spontaneity (6) Kinetics (7) Equilibrium (8) Acids & Bases (9) Redox
METHOD OF INSTRUCTION:
We will be moving through material at a steady pace. We will be working through an objective in about a week. There
will be at least 1 quiz per unit and 1 test per unit. There will be around 1 to 2 major labs per unit and activities/discussions
scattered throughout the unit. Participation in labs and activities is a requirement and are done to ensure you properly
comprehend material. You will have homework most nights. All work must be completed for the next class period
unless specified otherwise.
ii
DO NOWS/EXIT TICKETS:
Do Now’s & Exit Tickets are Ms. Manakul’s way to see who does or does not get what is happening in class. I also
believe any work that you do should be worth something. Thus, you will turn your Do Now/Exit Ticket in to Ms. Manakul
BEFORE she goes over the answer. If you answer the question correctly your Do Now/Exit Ticket sheet will go into the
class piggy bank. If you would like to save paper, and show the class that “you got this,” you may go over the answer in
Ms. Manakul’s place. If you choose this route, your name will automatically go into the class piggy bank. Before each
quiz Ms. Manakul will pull out two random Do Now/Exit Ticket Sheets and those two people will gain automatic extra
credit towards the quiz. Before a test Ms. Manakul will pull out four sheets. If you are absent, you may see Ms. Manakul
before or after school (not during school) to make up your Do Now so you do not miss out on the extra credit drawing.
HOMEWORK:
As mentioned above, homework will be assigned at the end of class and will be due in class the following school day.
Homework is a place to better comprehend material, gain proper practice, and solidify knowledge. However, sometimes
homework is confusing or you did not learn material the first time. Should this happen you should ALWAYS ASK
QUESTIONS either with a classmate or Ms. Manakul. Because we all struggle with homework, Ms. Manakul takes
homework as a completion grade. You must have tried the problem, i.e. you do not necessarily need to complete the
problem. Be aware an IDK is not trying and will thus cause your homework to be incomplete. After the class Do Now,
Ms. Manakul will check homework and go through the homework to ensure everyone comprehends the material.
Remember if you didn’t get it the first time ALWAYS ASK QUESTIONS!
LAB WORK:
I will go over each lab the day before. You are to prepare your lab book by having you lab book filled out at home based
on my rubric. To prepare you for designing your own lab, each lab is designed to be very similar to your typed lab report.
We will discuss the specifics of the designed lab when we get there. For your safety you will be wearing lab coats,
goggles, and closed toed shoes in the lab room for the majority of labs. If you wear open toed shoes during a lab, you will
wear gloves over your shoes to protect your feet. If this is to change, Ms. Manakul will verbally state it. Ms. Manakul has
high expectations for lab, thus, you will gather, set up, and put away your own lab materials. Furthermore, YOU MUST
BE SAFE at all times during a lab. Ms. Manakul will give ONLY ONE verbal warning. If she must give another, you will
be ejected from lab, and receive an automatic zero for the lab.
LATE WORK:
Ms. Manakul dislikes late work and generally cannot spare the time to regrade items. Thus, any work late will receive
30% off. This means by the 3rd class day, your work will receive an automatic zero. This is to ensure work is turned in a
timely manner, and Ms. Manakul has adequate time to grade your work.
QUIZZES/TESTS:
For every quiz/test you will be given a clean copy of the periodic table. For unit tests only, you will be allowed to have
your own hand written 3” by 5” one sided notecard with any notes you need on them. At the end of the semester you may
take a cumulative test. This will be determined by Ms. Manakul and announced in class when necessary. If you are caught
cheating or talking during a test, Ms. Manakul will take your test and give you an automatic zero. All tests and quizzes are
graded as fast as humanly possible and will be placed in Tyler at the same rate. Any changes will be announced verbally.
ABSENT:
Lesson plans for the week are posted in the classroom. A large calendar is also there which will label future events as
well. You can also check Google Classroom for any updates or changes. If you are absent, it is your responsibility to
make up the lecture notes on your own time. Should you need assistance, you may see Ms. Manakul before or after school
(reference My Weekly Schedule). Remember to turn in work as soon as possible. Any work assigned during your absence
will have a 2 school day grace period. If you are absent on the day an assignment is due, it will be due the first day you
return. If you miss a lab or activity, you are to collect the data from a friend, and finish the rest of your lab report on your
own time. *If Ms. Manakul is unexpectedly absent, you are to refer to the lesson plans posted and watch online for
specific details of any work to be completed in class. If an activity or lab is scheduled for that day, it will be rescheduled.
Ms. Manakul will recap all information and class activities/labs will be done with the class when she returns.
iii
CLASS RULES:
1. Pay close attention to class, and be prepared to work hard.
2. 10/10 Rule. No passes will be given during the first 10 minutes and the last 10 minutes of class. Restroom passes
only in extreme emergencies. We have very little time and will be moving through material quickly. This means
that class time is IMPORTANT and Ms. Manakul will do her best to maximize the time we have together.
3. IF ABSENT: DON’T GET LEFT BEHIND. Remember chemistry a class you must practice to understand,
which means missing a day can be essential to failing. Get help! See My Weekly Schedule.
4. Turn in work on time! Every day your work is late you will lose 30% that means that by the end of day 3 your
work will be an automatic zero. Ms. Manakul will give 2 extra days to make up work for every day you have an
excused absent.
5. No food or drinks will be allowed in the lab room. You may eat in the classroom as long as it does not become a
distraction. If you need to throw away large items or liquids, please toss them in trash cans in the hallways.
6. ALL electronic devices except your computer may not be used in the classroom. There will be occasions
where you may use them. Ms. Manakul will either say this or you may ask if the time is appropriate. If the time is
inappropriate the device will be taken on the second verbal warning and can be retrieved by the end of class.
7. Per district policy: Every 3 tardies = call home. Ms. Manakul considers a tardy as not being in the classroom at
the start of class.
GRADES:
A = 90.0 – 100.0 10% Homework
B = 80.0 – 89.9 20% Activities/Discussions
C = 70.0 – 79.9 20% Lab
D = 60.0 – 69.9 20% Quizzes
F = 0.0 – 59.9 30% Tests
Quarter Grades: Based on breakdown Semester Grade: Average of Quarters Year Grades: Average of Semesters
EXTRA CREDIT:
Extra credit is given throughout the school year and will be collected and graded each quarter. Students have the option to
read, analyze, and explain an approved scientific article following the extra credit rubric. Copies of the articles are located
on Google Classroom in the Class Drive Folder in the Extra Credit Folder. Only one extra credit article can be done per
quarter, and if one is not done for the quarter students do not have an option to make it up later. Due dates for extra credit
will be announced in class and will be in the calendar on Google Classroom.
TIPS TO SURVIVING CHEMISTRY:
1. Come to class prepared. If you bring you’re “A” game, you are more likely to get an “A” for the year.
2. Participate in class. The more you participate, the better you learn.
3. Pay attention in class. Ms. Manakul loves tell you exactly what will be on the test. If you’re listening you will
know what to expect and what to do.
4. Do the practice tests as if it was a real test. This will give you a better idea of what you need to study more and
how long you need for each problem so you can manage your time wisely.
5. Keep all your notecards. They will make it easier to study for the semester tests.
6. If you ever think something is wrong on a test, homework, notes, etc. TELL Ms. Manakul right away. We all get
confused; including Ms. Manakul and it is better to get it clarified then answer a problem wrong.
7. Keep up with any missed notes or work. Chemistry BUILDS! Missing one day of class can put you two to three
days behind your classmate, which will make you struggle more during the test.
8. Do your best on all homework and classwork. Practicing chemistry is like a mental workout. The best workouts
are the ones that require the most mental energy, so DON’T GIVE UP, even if it is tough.
9. Do your own work. Copying others’ homework or labs will only hurt you on the quizzes and tests.
10. Get your own your calculator!
iv
Classroom Map
v
School Safety Drills
On this page you will fill in Manakul’s Classroom procedures for all School Safety Drills.
1. Fire Safety
2. Tornado
3. Lock Down
4. Earthquake
5. Where are the maps located for emergency drills?
vi
vii
viii
Chemistry II Semester I Objectives
Unit 1: The Atom & Periodic Table:
Objective Accomplishment Goal Quiz
Understand and determine the mass
number, atomic number, element,
and number of protons, neutrons,
and electrons
You are able to diagram and label all
the pieces of a mass spectrometer
and emission spectrum.
Explain how the emission spectrum
is produced and its relation to atomic
orbitals
You can explain how electrons jump
from energy levels, then how the
electrons fall to create the emission
spectrum.
Explain electron configuration and
its effects on the periodic table
arrangement
You are able to draw the electron
configuration for every element and
explain how it affects its placement
in the periodic table.
Understand the complexities of
elements when alone and within a
compound.
You are able to determine the
compound based on a physical and
chemical reactive description of the
compound.
Unit 2: Bonding
Objective Accomplishment Goal Quiz
Explain periodic trends in terms of
ionic and covalent bonding
Given any 2 or more elements you
can determine the name and formula
of a compound and explain why
these elements bond together in that
particular manner.
Be able to draw & explain VSEPR
theory and use it to predict bond
length, polarity, shape of compound,
and the physical & chemical
properties of molecule
You are able to draw the Lewis Dot
structure for any compound,
determine the number of σ and π
bonds, and type of hybridization
within the compound.
Identify and explain the
relationships between Lewis dot
structures, molecular shapes, and the
types of hybridization in terms of sp2
and sp3
You are able to explain what it
means when a compound become
hybridized and determine what type
of hybridization a compound has in
terms of sp1, sp2, or sp3
Describe types of intermolecular
forces and atomic orbitals and their
effects on chemical and physical
properties
Be able to explain the three types of
intermolecular forces and how it
affects a molecules chemical and
physical properties
ix
Unit 3: Moles, Stoichiometry, & Chemical Reactions
Objective Accomplishment Goal Quiz
Be able to solve stoichiometric
problems involving chemical
reactions, moles, number of
particles, pressure, temperature, and
volume
You are able to list every given in a
problem and be able to set up any
given stoichiometric problem.
Be able to determine the empirical
and molecular formula.
You are able to set up the process to
determine the empirical formula and
molecular formula given
experimental data.
Solve problems involving
concentration, amount of solute, and
volume of solution
You are able to list every given in a
problem and be able to set up any
given solution problem.
Determine the limiting reactant and
theoretical yield
You are able to determine which of
the reactants is limiting and which is
in excess.
Unit 4: Energetics Packet
Objective Accomplishment Goal Quiz Test
Be able to define the terms:
endothermic, exothermic,
standard enthalpy change of
reaction (ΔH°), average bond
enthalpy, standard enthalpy
change of formation (ΔHf°), and
standard change of combustion
You can define every term
related to energetics.
Be able to relate temperature
change and enthalpy change to
endothermic and exothermic.
You are able to determine if a
reaction is endothermic or
exothermic given a chemical
reaction.
Be able to calculate enthalpy
change
You can calculate enthalpy
change using Hess’s Law or
Bond Energy Potentials.
x
Unit 5: Entropy & Spontaneity
Objective Accomplishment Goal Quiz Test
You are able to describe entropy
and be able to calculate the
change in entropy
You can list the factors which
affect entropy and can use the
values on the back of your
periodic table to determine the
change in entropy for a reaction.
You are able to describe
spontaneity and be able to
calculate Gibbs Free Energy.
You can list the factors which
affect spontaneity and can use the
values on the back of your
periodic table to determine Gibbs
Free Energy for a reaction.
Unit 6: Kinetics
Objective Accomplishment Goal Quiz Test
You will be able to describe
rates of reactions and collision
theory.
You can explain how gas
molecules move and react.
You will be able to determine
the rate expression, rate
constant, overall order of a
reaction, and order of a reaction
with respect to a particular
reactant.
You can explain in words rate
expression, rate constant, overall
order of a reaction, and order of a
reaction.
You will be able to draw,
explain, and determine
components to a reaction
mechanism.
You can draw a 0 order, 1st order,
or a 2nd order reaction.
xi
Objective Graphs
On this page you will create a bar graph for each unit of your goal, quiz and test scores per
objective (these are obtained from the trackers done after each quiz and test). Use test data for
best accuracy & remember to put units and labels. SUGGESTION: Use different colors.
Unit 1: Atom & Periodic Table Unit 2: Bonding
Unit 3: Moles, Stoichiometry, & Gas Laws Unit 4: Energetics
xii
Unit 5: Entropy & Spontaneity Unit 6: Kinetics
Spare Graphs
1
Unit 1: The Atom & Periodic Table Homework
SECTION 1.1: The Atom
1. Use the periodic table to identify the sub-atomic particles present in the following
species.
Species Number of Protons Number of Neutrons
Number of
Electrons
A 7Li
B 1H
C 14C
D 19F-
E 56Fe3+
2. Isoelectronic species have the same number of electrons. Identify the following
isoelectronic species by giving the correct symbol and charge. You will need a periodic
table. The first one has been done as an example.
Species Number of Protons Number of Neutrons
Number of
Electrons
40Ca2+ 20 20 18
A 18 22 18
B 19 20 18
C 17 18 18
3. Which of the following species contain more electrons than neutrons?
a. 𝐻11 b. 𝐵5
11 c. 𝑂2−8
16 d. 𝐹−9
19
4. What is the same for an atom of phosphorus-26 and an atom of phosphorus-27?
a. Atomic number and mass number
b. Number of protons and electrons
c. Number of neutrons and electrons
d. Number of protons and neutrons
2
5. The relative abundance of the two isotopes of chlorine are shown in this table:
Isotope Relative Abundance 35Cl 75% 37Cl 25%
Determine the average atomic mass for chlorine.
6. Magnesium has three stable isotopes – 24Mg, 25Mg, and 26Mg. The lightest isotope has an
abundance of 78.90%. Calculate the percent abundances of the other isotopes.
7. The Geiger-Marsden experiment, supervised by Ernest Rutherford, gave important
evidence for the structure of the atom. Positively charged alpha particles were fired at a
piece of gold foil. Most the particles passed through with only minor deflections but a
smaller number rebounded from the foil.
How did this experiment change our knowledge of the atom?
a. It provided evidence for the existence of discrete atomic energy levels.
b. It provided evidence for a positively charged dense nucleus.
c. It provided evidence that electrons move in unpredictable paths around the
nucleus.
d. It proved evidence for the existence of an uncharged particle in the nucleus.
3
SECTION 1.2: Electron Configuration
1. Emission and absorption spectra both provide evidence for:
a. The existence of neutrons
b. The existence of isotopes
c. The existence of atomic energy levels
d. The nuclear model of the atom
2. The diagram shows the lowest five energy levels in the hydrogen atom.
n = 5
n = 4
n = 3
n = 2
n = 1
Deduce how many different frequencies in the emission spectrum of the atomic hydrogen
would arise as a result of electron transitions between these levels.
a. 3 b. 4 c. 6 d. 10
3. Identify which of the following provide evidence to support the Bohr model of the
hydrogen atom?
I. The energy lines in the emission spectra of atomic hydrogen
II. The energy of the missing lines in the absorption spectra of helium as seen
from the sun.
III. The relative intensity of the different spectral lines in the emission spectrum
of atomic hydrogen.
a. I only b. II only c. I and II d. I and III
4. List the 4d, 4f, 4p, and 4s atomic orbitals in order of increasing energy.
5. Which statement is correct for the emission spectrum of the hydrogen atom?
a. The lines converge at lower energies
b. The lines are produced when electrons move from lower to higher energy levels.
c. The lines in the visible region involve electron transitions into the energy level
closest to the nucleus
d. The line corresponding to the greatest emission of energy is in the ultraviolet
region.
4
6. Identify the sub-level which does not exist.
a. 5d b. 4d c. 3f d. 2p
7. State the full ground-state electron configuration of the following elements.
a. V
b. K
8. Identify the excited state (i.e. not the ground state) in the following electron
configuration.
a. [Ne]3s23p3 c. [Ne]3s23p64s1
b. [Ne]3s13p64s1 d. [Ne]3s23p63d14s2
9. Deduce the number of unpaired electrons present in the ground state of a titanium atom.
a. 1 b. 2 c. 3 d. 4
10. State the full ground-state electron configuration of the following ions.
a. O2-
b. Cl-
c. Ti3+
d. Cu2+
11. State the electron configuration of the following transition mental ions by filling in the
boxes below. Use arrows to represent the electron spin.
Ion 3d 4s
A Ti2+
B Fe2+
C Ni2+
D Zn2+
5
SECTION 1.3: Electrons in atoms
1. The first four ionization energies for a particular element are 738, 1450, 7730, and 10550
kJ mol-1 respectively. Deduce the group number of the element.
a. 1 b. 2 c. 3 d. 4
2. Successive ionization energies for an unknown element are given in the table below.
First Ionization
energy (kJ mol-1)
Second ionization
energy (kJ mol-1)
Third ionization
energy (kJ mol-1)
Fourth ionization
energy (kJ mol-1)
590 1145 4912 6491
Identify the element.
a. K b. Ca c. S d. Cl
3. The successive ionization energies (in kJ mol-1) for carbon are tabulated below.
1st 2nd 3rd 4th 5th 6th
1086 2352 4619 6620 37820 47280
a. Explain why there is a large increase between the fourth and fifth values.
b. Explain why there is an increase between the second and third values.
4. Sketch a graph to show the expected pattern for the first seven ionization energies of
fluorine.
6
5. The first ionization energies of the Period 3 elements Na to Ar are below.
Na Mg Al Si P S Cl Ar
496 738 578 787 1012 1000 1251 1520
a. Explain the general increase in ionization energy across the period.
b. Explain why the first ionization energy of magnesium is greater than that of
aluminum.
c. Explain why the first ionization energy of sulfur is less than that of phosphorous.
SECTION 1.4: The Periodic Table
1. Use the periodic table to identify the position of the following elements:
Element Period Group
A Helium
B Chlorine
C Barium
D Francium
2. Phosphorus is in Period 3 and Group 15 on the periodic table.
a. Distinguish between the terms period and group.
b. State the electron arrangement of phosphorus and relate it to its position in the
periodic table.
7
3. How many valence (outer shell) electrons are present in the atoms of the element with
atomic number 51?
4. Which of the following elements is a metalloid?
a. Calcium b. manganese c. germanium d. magnesium
5. Which of the following materials is the best conductor of electricity in the solid state?
a. Silicon b. graphite c. phosphorus d. antimony
6. Which of the following properties is used to arrange the elements in the modern Periodic
Table?
a. Relative atomic number
b. Number of valence electrons
c. Atomic number
d. Effective nuclear charge
SECTION 1.5: Periodic Trends
1. Explain what is meant by the atomic radius of an element.
2. The atomic radii of the elements are found in the slides.
a. Explain why no values for ionic radii are given for the noble gases.
b. Describe and explain the trend in atomic radii across the Period 3 elements.
3. Si4+ has an ionic radius of 4.2 x 10-11 m and Si4- has an ionic radius of 2.71 x 10-10 m,
explain the large difference in size between the Si4+ and Si4- ions.
4. Which of the following is a property of gaseous atoms?
I. ionization energy
II. electron affinity
III. electronegativity
a. I and II b. I and III c. II and III d. I, II, and III
8
5. Identify the element which is likely to have an electronegativity value most similar to that
lithium?
a. Beryllium b. sodium c. magnesium d. hydrogen
6. The graph represents the variation of a property of the Group 2 elements.
Identify the property:
a. Ionic radius
b. Atomic radius
c. Neutron/proton ratio
d. First ionization energy
7. Atomic radii and ionic radii are found in the slides. Explain why:
a. The potassium ion is much smaller than the potassium ion.
b. There is a large increase in ionic radius from silicon (Si4+) to phosphorus (P3-).
c. The ionic radius of Na+ is less than that of F-.
8. Explain why sulfur has a higher melting point than phosphorus.
9. Which physical property generally increases down a group but decreases from left to
right across a period?
a. Melting point b. Electronegativity c. Ionization Energy d. Atomic Radius
10. The elements in the Periodic Table are arranged in order of increasing:
a. Relative atomic mass
b. Ionic radii
c. Nuclear charge
d. Ionization Energy
9
11. What is the order of decreasing radii for the species Cl, Cl+, and Cl-?
12. The following graph shows the variation of a physical property, X, of the first 20
elements in the Periodic Table with the atomic number.
Identify the property X.
a. atomic radius c. ionic radius energy
b. first ionization d. melting point
13. How do the reactivity’s of the alkali metals and the halogens vary down the group?
14. Chlorine is a greenish-yellow gas, bromine is a dark red liquid, and iodine is a dark grey
solid. Identify the property which most directly causes these differences in volatility.
a. The halogen-halogen bond energy
b. The number of neutrons in the nucleus of the halogen atom
c. The number of outer electrons in the halogen atom
d. The number of electrons in the halogen molecule
15. A paper published in April 2010 by Yu. Ts. Oganessian and others claim the synthesis of
isotopes of a new element with atomic number 117. One of the isotopes is 𝑈𝑢𝑠117293 . Which
of the following statements is correct?
a. The nucleus of the atom has a relative charge of +117
b. 𝑈𝑢𝑠117293 has a mass number of 117
c. There are 262 neutrons in 𝑈𝑢𝑠117293
d. The atomic number is 293 – 117
10
16. Use the data table below to:
Oxides Melting Point (K) Boiling Point (K)
MgO 3125 3873
SiO2 (quartz) 1883 2503
P4O10 297 448
SO2 200 263
a. Identify the state of the four oxides listed under standard conditions.
b. Explain the difference in melting points by referring to the bonding and structure
in each case.
SECTION 1.6: First Row d-Block Elements
1. Identify the property/properties which are characteristic of an element found in the d
block of the Periodic Table.
a. All the compounds of the element are ionic
b. The element exhibits a variety of oxidation states and color in its compounds
c. The element has a low melting point
d. The element is a good conductor of heat and electricity
2. Identify the oxidation number which is the most common among the first-row transition
elements.
a. +1 b. +2 c. +4 d. +6
3. An element has the electron configuration of 1s22s22p63s23p63d34s2. Which oxidation
state(s) would this element show?
a. +2 and +3 only c. +3 and +5 only
b. +2 and+5 only d. +2, +3, +4, and +5
4. State the full electron configuration of zinc (Zn).
5. State the full electron configuration of Zn2+.
6. Explain why zinc is not classed as a transition metal.
7. State the oxidation states/charges shown by calcium and chromium and explain the
difference in their behavior.
8. Why do some metals require a complex with ligands?
11
9. Identify the species which cannot act as a ligand?
a. H2O b. CO c. CH4 d. Cl-
10. Consider the following reaction below:
[Cu(H2O)6]2+(aq) + 4HCl (aq) [CuCl4]
2- (aq) + 6H2O (aq) + 4H+(aq)
Which of the following is acting as a ligand?
a. H+ only c. H2O and Cl- only
b. H+ and Cl- only d. H+, H2O, and Cl-
11. The color and formulas of some coordination compound of hydrated forms of chromium
chloride are listed in this table.
I II III
Formula [Cr(H2O)6]Cl3 [CrCl(H2O)5]Cl2•H2O [CrCl2(H2O)4]Cl•2H2O
Color Purple Blue-green Green
What are the charges on each of the complex ions?
I II III
A 0 0 0
B + 2+ 3+
C 2+ 3+ +
D 3+ 2+ +
12. Identify the feature which is an essential characteristic of all ligands.
a. A negative charge
b. An electronegative atom
c. The presence of a non-bonding pair of electrons
d. The presence of two or more atoms
13. Which of the following elements would be expected to be paramagnetic?
a. Ca b. Zn c. He d. Mn
14. Which of the elements is the most paramagnetic?
b. Sc b. Ti c. V d. Cr
15. Explain why chromium is the most paramagnetic element in the first transition series and
why diamagnetic.
12
16. Study the structure of the heme shown below.
a. What is the oxidation state of the central iron ion?
b. What is the geometry of the nitrogen atoms around the central iron ion?
c. Explain why the complex is ideally suited to carry oxygen around the body.
17. Distinguish between homogenous catalysis and heterogeneous catalysis.
18. Explain why the transition metals make effective heterogeneous catalysts.
19. Explain why heterogeneous catalysts are generally used in industrial processes.
13
SECTION 1.7: Colored Complexes
1. The color of transition metal complexes depends on several factors.
a. Suggest why the color of [Cr(H2O)6]3+ is different from [Fe(H2O)6]
3+.
b. Suggest why the color of [Fe(H2O)6]2+ is different from [Fe(H2O)6]
3+.
c. Suggest why the color of [Fe(NH3)6]2+is different from [Fe(H2O)6]
3+.
2. Explain why Fe2+(aq) is colored and can behave as a reducing agent, whereas Zn2+ (aq) is
not colored and does not behave as a reducing agent.
3. The absorption spectrum of [Ti(H2O)6]3+ is shown below. Use the color wheel to suggest
a color for the complex.
Purple Red
4. Predict whether the splitting of the d orbitals in [Fe(CN)6]4- would be less than or greater
than the splitting of [Fe(H2O)6]2+.
14
Unit 2: Bonding Homework
SECTION 2.1: Ionic Bonding
1. Write the formula for each of the following compounds:
a. Potassium bromide d. copper (II) bromide
b. Zinc oxide e. chromium (III) sulfate
c. Sodium sulfate f. aluminum hydride
2. Name the following compounds:
a. Sn3(PO4)2 d. BaSO4
b. Ti(SO4)2 e. Hg2S
c. Mn(HCO3)2
3. Explain what happens to the electron configuration of Mg and Br when they react to form
the compound magnesium bromide.
4. You are given two white solids and told that only one of them is an ionic compound.
Describe three tests you could carry out to determine which it is.
15
SECTION 2.2: Covalent Bonding
1. Explain how some covalent compounds have ionic bonding.
2. Which of the following molecules contains the shortest bond between carbon and
oxygen?
a. CO2 b. H3COCH3 c. CO d. CH3COOH
3. For each of these molecules, identify any polar bonds and label them using δ+ and δ-
appropriately.
a. HBr d. O2
b. CO2 e. NH3
c. ClF
4. Use the electronegativity values from the IB Data Tables to predict which bond in each
of the following pairs is more polar.
a. C-H or C-Cl
b. Si-Li or Si-Cl
c. N-Cl or N-Mg
5. Put the following species in order of increasing carbon-oxygen bond length:
CO CO2 CO32- CH3OH
6. By reference to their resonance structures, compare the nitrogen-oxygen bond lengths in
nitrate (V) (NO3-) and nitric (V)acid (HNO3).
16
SECTION 2.3: Covalent Structures
1. Draw the Lewis structures of:
a. HF b. CF3Cl c. C2H6
d. PCl3 e. C2H4 f. C2H2
2. Use Lewis structures to show the formation of a coordinate bond between H2O and H+ to
form H3O+.
3. Draw the Lewis Structures of:
a. NO3- b. NO+ c. NO2
-
4. Predict the shape and bond angles of the following molecules and ions:
a. H2S b. CF4 c. HCN
d. NF3 e. BCl3 f. NH2Cl
g. OF2
17
5. Predict the shape of each of the following ions:
a. CO32- b. NO3
‾ c. NO2+
d. NO2‾ e. ClF2
+ f. SnCl3-
6. Predict whether the following will be polar or non-polar molecules overall.
a. PH3 e. C2H4
b. CF4 f. ClF
c. HCN g. F2
d. BeCl2 h. BF3
7. The molecule C2H2Cl2 can exist as two forms known as cis- and trans- isomers, which
are shown below:
Determine whether either of these has a net dipole movement.
18
SECTION 2.4: Intermolecular Forces
1. The physical properties of five solids labelled A, B, C, D, and E are summarized below.
The substances are an ionic compound, a non-polar solid, a metal, a polar molecular
solid, and a giant molecular substance. Classify each correctly.
Sample Solubility in
H2O
Conductivity
of solution
Conductivity
of solid
Relative
melting point
A In soluble - Yes Third to melt
B Insoluble - No Highest
C Soluble No No Second to melt
D Insoluble - No Lowest
E Soluble Yes No Fourth to melt
2. Which substance is the most soluble in water?
a. CH3OH c. C2H6
b. CH4 d. C2H5OH
3. State the intermolecular forces that exits between molecules of each of the following:
a. Dry ice, CO2(s) c. N2(l)
b. NH3(l) d. CH3OCH3
4. Which of each pair has the lower boiling point?
a. C2H6 and C3H8 c. Cl2 and Br2
b. H2O and H2S d. HF and HCl
19
SECTION 2.5: Metallic Bonding
1. Which is the best definition of metallic bonding?
a. The attraction between cations and anions
b. The attraction between cations and delocalized electrons
c. The attraction between nuclei and electron pairs
d. The attraction between nuclei and anions
2. Aluminum is a widely used metal. What properties make it suitable for the following
applications?
a. Baking foil
b. Aircraft bodywork
c. Cooking pans
d. Tent frames
SECTION 2.6: Unique Covalent Bonding
1. What bond angles do you expect for each of the following?
a. F-Kr-F angle in KrF4 b. Cl-P-Cl angle in PCl3 c. F-S-F angle in SF6
2. Use the concept of formal charge to explain why BF3 is an exception to the octet rule.
20
3. Draw two different Lewis (electron dot) structures for SO42-, one of which obeys the
octet rule for all its atoms, the other which has an octet for S expanded to 12 electrons.
Use formal charges to determine which is the preferred structure?
4. Explain why ozone can be dissociated by light with a longer wavelength than that
required to decompose oxygen.
5. Outline ways in which ozone levels are decreased by human activities, using equations
to support your answer.
SPECTION 2.7: Hybridization
1. What is the difference in spatial distribution between electrons in a pi bond and electrons
in a sigma bond?
2. What hybridization would you expect for the bolded atom in each of the following?
a. H2C=O b. BH4- c. SO3 d. BeCl2 e. CH3COOH
21
3. Cyclohexane, C6H12, has a puckered, non-planar shape. Benzene C6H6 is planar.
Explain the difference by making reference to the C – C – C bond angles and the type
of hybridization of carbon in each molecule.
SECTION 2.8: Fundamentals of Organic Chemistry
1. Name the following molecules:
a. CH3CH2CH2COOH d. CH3CONH(CH3)
b. CHCl2CH2CH3 e. C3H7CHO
c. CH3CH2COCH3 f. CH3CH2CH2CH2COOCH2CH3
22
2. Give the structural formulas for the following molecules (Condensed form is accepted):
a. Hexanoic acid e. Ethyl methanoate
b. Butanal f. Hexanamide
c. Pent-1-ene g. N,N - dimethylhexanamine
d. 1-bromo-2methylbutane
3. Which of the following is an amine?
a. CH3CH2NH2 c. CH3CH2CN
b. CH3CONH2 d. C2H5CONH(CH3)
4. Which compound is a member of the same homologous series as 1-bromopropane?
a. 1-iodopropane c. 1-bromopropene
b. 1,2 – dibromopropane d. 1-bromopentane
5. Draw and name all the structural isomers of C3H3Cl5.
23
6. Which formula is that of a secondary halogen alkane?
a. CH3CH2CH2CH2Br c. (CH3)2CHCH2Br
b. CH3CHBrCH2CH3 d. (CH3)3CBr
7. Describe the bonding in a benzene molecule and use it to explain benzene’s energetic
stability.
8. When comparing the boiling points of different classes of compound, why is it
important to choose molecules that have similar molar mass?
9. Explain how you would expect the solubility of alcohols in hexane to change with
increasing chain length.
24
Unit 3: Moles, Stoichiometry, & Chemical Reactions Homework
Section 3.1: Chemical Reactions
1. Write balanced chemical equations for the following reactions:
a. The decomposition of copper carbonate into copper (II) oxide and carbon
dioxide?
b. The combustion of magnesium in oxygen to form magnesium oxide.
c. The neutralization of sulfuric acid (H2SO4) with sodium hydroxide to form
sodium sulfate and water.
d. The synthesis of ammonia from nitrogen and hydrogen.
e. The combustion of methane to produce carbon dioxide and water.
2. Write balanced chemical equations for the following reactions:
a. K + H2O KOH + H2
b. C2H5OH + O2 CO2 + H2O
c. Cl2 + KI KCl + I2
25
d. Fe2O3 + C CO + Fe
e. C4H10 + O2 CO2 + H2O
Section 3.2: Kinetic Movement
1. Which of the following occurs at the melting point when solid sulfur is converted to its
liquid from?
I. Movement of the particles increases
II. Distance between the particles increases
a. I only
b. II only
c. Both I and II
d. Neither I nor II
2. Gasses deviate from ideal gas behavior because their particles:
a. Have negligible volume
b. Have forces of attraction between them
c. Are polyatomic
d. Are not attracted to one another
3. A mixture of two gases, X and Y, which both have strong but distinct smells, is released.
From across the room the smell of X is detected more quickly than the smell of Y. What
can you deduce about X and Y.
4. Ice floats on water. Comment on why this is not what you would expect from kinetic
theory of matter.
26
5. A closed flask contains a pure substance, a brown liquid that is at its boiling point.
Explain what you are likely to observe in the flask, and distinguish between the inter-
particle distance and the average speeds of the particles in the two states present.
6. During very cold weather, snow often gradually disappears without melting. Explain
how this is possible.
7. Explain why a burn to the skin caused by steam is more serious that a burn caused by
the same amount of boiling water at the same temperature.
8. You are given a liquid substance at 80°C and told that it has a melting point of 35°C.
You are asked to take its temperature at regular time intervals while it cools to room
temperature (25°C). Sketch the cooling curve that you would expect to obtain.
27
9. A road cyclist pumps his tires up very hard before a trip over a high mountain pass at
high altitude. Near the summit one of his tires explodes. Suggest why this may have
occurred.
10. List the main features of the kinetic theory for ideal gases
11. Explain the reason for the difference in behavior between real and ideal gases at low
temperatures.
12. Ammonia forms a relatively strong type of intermolecular attraction known as hydrogen
bonding whereas methane does not. Explain the relative deviation from ideal behavior
that each gas is likely to show.
28
SECTION 3.3: Reacting Gases
1. How many moles are present in each of the following at STP?
a. 54.5 dm3 CH4
b. 250.0 cm3CO
2. Which sample contains more molecules 3.14 dm3 of bromine Br2 or 11.07 g of chlorine
Cl2, when measured at the same temperature and pressure?
3. A 2.50 dm3 container of helium at a pressure of 85 kPa was heated from 25°C to 75°C.
The volume of the container expanded to 2.75 dm3. What was the final pressure of the
helium?
4. After a sample of nitrogen with a volume of 675 cm3 and a pressure of 1.00 x 105 Pa was
compressed to a volume of 350 cm3 and a pressure of 2.00 x 105 Pa, its temperature was
27.0°C. Determine its initial temperature.
29
5. The absolute temperature of 4.0 dm3 of hydrogen gas is increased by a factor of three and
the pressure is increased by a factor of four. Deduce the final volume of the gas.
6. To find the volume of a flask, it was first evacuated so that it contained no gas at all.
When 4.40 g of carbon dioxide was introduced, it exerted 90 kPa at 27°C. Determine the
volume of the flask.
7. An unknown noble gas has a density of 5.84 g dm3 at STP. Calculate its molar mass and
its identity.
8. A 12.1 mg sample of gas has a volume of 255 cm3 at a temperature of 25.0 °C and a
pressure of 1300 Pa. Determine its molar mass.
9. Which has a greater density at STP, hydrogen or helium?
30
SECTION 3.4: The Mole Concept
1. Calculate how many hydrogen atoms are present in:
a. 0.020 moles of C2H5OH
b. 2.50 moles of H2O
c. 0.10 moles of C(HCO3)2
2. Propane has the formula C3H8. If a sample of propane contains 0.20 moles of C, how
many moles of H are present?
3. Calculate the amount of sulfuric acid, H2SO4, which contains 6.02 x 1023 atoms of
oxygen.
4. Calculate the molar mass of the following compounds:
a. Magnesium phosphate c. Calcium nitrate
b. Ascorbic acid C6H8O6 d. Hydrated sodium thiosulfate, Na2S2O3·5H2O
31
5. Calcium arsenate, Ca3(AsO4)2, is a poison which was widely used as an insecticide.
What is the mass of 0.475 mol of calcium arsenate?
6. Which contains the greater number of particles, 10.0 g water or 10.0 g of mercury?
7. Put the following in descending order of mass?
a. 1.0 mol N2H4 c. 3.0 mol NH3
b. 2.0 mol N2 d. 25.0 mol H2
8. Give the empirical formulas of the following compounds:
a. Glucose C6H12O6
b. SucroseC12H22O11
c. Ethanoic acid CH3COOH
9. A street drug has the following composition: 83.89% C, 10.35% H, 5.76%N. Determine
its empirical formula.
32
SECTION 3.5: Stoichiometry
1. Calculate the mass of potassium hydroxide required to prepare 250 cm3 of a 0.200 mol
dm-3 solution
2. Calculate the mass of magnesium sulfate heptahydrate MgSO4·7H2O required to prepare
0.100 dm3 of a 0.200 mol dm-3 solution.
3. Calculate the number of moles of chloride ions in 0.250 dm3 of 0.0200 mol dm-3 zinc
chloride solution.
4. 250 cm3 solution contains 5.85 g of sodium chloride. Calculate the concentration of
sodium chloride in mol dm-3.
5. Concentrated nitric acid HNO3 is 16.0 mol dm-3. What volume would you need to prepare
100 cm3 of 0.50 mol dm-3 HNO3?
33
6. In a titration a 15.00 cm3 sample of H2SO4 required 36.42 cm3 of 0.147 mol dm-3 NaOH
solution for complete neutralization. What is the concentration of the H2SO4?
SECTION 3.6: Limiting and Excess 1. Limestone is mostly calcium carbonate, CaCO3, but also contains other minerals. When
heated the CaCO3 decomposes into Ca and CO2. A 1.605 g sample of limestone was
heated and gave off 0.657 g of CO2.
a. Formulate the equation for the thermal decomposition of calcium carbonate.
b. Determine the percentage mass of CaCO3 in limestone.
c. State the assumption that you are making in this calculation.
2. Methanol, CH3OH, is useful fuel that can be made as follows:
CO(g) + 2H2(g) CH3OH
A reaction mixture used 12.0 g of H2 and 74.5 g of CO.
a. Determine the theoretical yield of CH3OH
b. Calculate the amount of excess reactant that remains unchanged at the end of the
reaction.
34
3. Calcium carbonate CaCO3 is able to remove sulfur dioxide from the waste gases by a
reaction in which they react in a 1:1 stoichiometric ratio to form equimolar amounts of
CaSO3. When 255g of CaCO3 reacts with 135g of SO2, 198 g of CaSO3 were formed.
Determine the percentage yield of CaSO3.
4. Pentyl ethanoate CH3COOC5H11 which smells like bananas, is produced from the
esterification reaction:
CH3COOH(aq) + C5H11OH CH3COOC5H11 (aq) + H2O (l)
A reaction uses 3.58 g of CH3COOH and 4.75 C5H11OH and has a yield of 45.00%.
Determine the mass of ester that forms.
5. A chemist has to make a 100 g sample of Chlorobenzene C6H5Cl, from the following
reaction:
C6H6 + Cl2 C6H5Cl + HCl
Determine the minimum quantity of benzene that can be used to achieve this with a yield
of 65%.
35
Unit 4: Energetics Homework
SECTION 4.1: Measuring Energy Changes
1. Explain the meaning of the term ΔH and describe how it is measured.
2. When a sample of NH4SCN is mixed with solid Ba(OH)2 · 8H2O in a glass beaker, the
mixture changes to a liquid and the temperature drops sufficiently to freeze the beaker to
the table.
Which statement is true about the reaction?
a. The process is endothermic and ∆H is –
b. The process is endothermic and ∆H is +
c. The process is exothermic and ∆H is –
d. The process is exothermic and ∆ H is +
3. Which one of the following statements is true of all exothermic reactions?
a. They produce gases.
b. They give out heat
c. They occur quickly
d. They involve combustion
4. If 500 J of heat is added to 100.0g samples of each of the substances below, which will
have the largest temperature increase?
Substance Specific heat capacity (J g-1 K-1)
A Gold 0.129
B Silver 0.237
C Copper 0.385
D Water 4.18
5. The temperature of a 5.0 g sample of copper increases from 27°C to 29°C. Calculate how
much heat has been added to the system. (Specific heat capacity of Cu = 0.385 J g-1 K-1).
a. 0.770 J b. 1.50 J c. 3.00 J d. 3.85J
6. Consider the specific heat capacity of the following metals.
Metal Specific heat capacity (J g-1 K-1)
Al 0.897
Be 1.82
Cd 0.231
Cr 0.449
1 kg samples of metals at room temperature are heated by the same electrical heater for
10 minutes. Identify the metal which has the highest final temperature.
a. Al b. Be c. Cd d. Cr
36
7. The specific heat of metallic mercury is 0.138 J g-1 °C-1. If 100.0 J of heat is added to a
100.0 g sample of mercury at 25.0°C, what is the final temperature of the mercury?
8. The mass of the burner and its content is measured before and after the experiment. The
thermometer is read before and after the experiment. What are the expected results?
Mass of the burner
and contents
Reading on
thermometer
A Decrease Increase
B Decrease Stays the same
C Increase Increase
D Increase Stays the same
9. The experimental arrangement below is used to determine the enthalpy of combustion of
an alcohol. Which of the following would lead to an experimental result which is less
exothermic than the literature value?
I. Heat loss from the sides of the copper calorimeter
II. Evaporation of alcohol during the experiment
III. The thermometer touches the bottom of the calorimeter.
a. I and II only b. I and III only c. II and III only d. I, II, and III
10. A copper calorimeter was used to determine the enthalpy of combustion of butan-1-ol.
The experiment value obtained was -2100 ± 200 kJ mol-1 and the data booklet value is
-2676 kJ mol-1. Which of the following accounts for the difference between the two
values?
I. random measurement errors
II. incomplete combustion
III. heat loss to the surroundings
a. I and II only b. I and III only c. II and III only d. I, II, and III
37
11. The heat released from the combustion of 0.0500 g of white phosphorus increases the
temperature of 150.0 g of water 25.0°C to 31.5°C. Calculate a value for the enthalpy
change of combustion of phosphorus. Discuss possible sources of error in the experiment.
SECTION 4.2: Hess’s Law
1. The diagram illustrates changes of a set of reactions.
Which of the following statements are correct?
I. P S ΔH = -10 kJ
II. R Q ΔH = +90 kJ
III. P R ΔH + 20 kJ
a. I and II only b. I and III only c. II and III only d. I, II, and III
2. Calculate the standard enthalpy change, ∆H ̊, for the reaction:
C(graphite) + ½ O2(g) CO(g)
From the information below:
C(graphite) + O2(g) CO2(g) ∆H ̊ = - 394 kJ
CO(g) + ½ O2(g) CO2(g) ∆H ̊ = - 283 kJ
38
3. Calculate the standard enthalpy change, ∆H ̊, for the reaction :
2NO(g) + O2 (g) 2NO2 (g)
From the information below:
N2 (g) + O2 (g) 2NO(g) ∆H ̊ = + 180.5 kJ
N2(g) + 2O2 (g) 2NO2 (g) ∆H ̊ =+66.4 kJ
4. Calculate the enthalpy change for the dimerization of nitrogen dioxide:
2NO2(g) N2O4 (g)
From the following data:
½ N2(g) + O2(g) NO2(g) ∆H ̊ = +33.2 kJ mol-1
N2(g) + 2O2 (g) N2O4 (g) ∆H ̊ = +9.16 kJ mol-1
5. The thermochemical equations for three related reactions are shown.
CO(g) + ½ O2(g) CO2 (g) ΔH1 = -283 kJ mol-1
2H2(g) + O2(g) 2H2O (l) ΔH2 = -572 kJ mol-1
CO2(g) + H2(g) CO(g) + H2O(l) ΔH3 = ? Determine ΔH3.
a. +289 kJ mol-1 c. -298 kJ mol-1
b. b. -3 kJ mol-1 d. -855 kJ mol-1
6. Which of the following does not have a standard heat of formation value of zero at 25°C
and 1.00 x 105 Pa.
a. Cl2(g) b. I2(s) c. Br2(g) d. Na(s)
7. Which of the following does not have a standard heat of formation value of zero at 25°C
and 1.00 x 105 Pa.
a. H(g) b. Hg(s) c. C(diamond) d. Si(s)
8. For which equation is the enthalpy changes described as an enthalpy change of
formation?
a. CuSO4(aq) + Zn(s) ZnSO4(aq) + Cu(s)
b. Cu(s) + S(s) + 2O2(g) CuSO4(aq)
c. 5H2O(l) + CuSO4(s) CuSO4·5H2O(s)
d. Cu(s) + S(s) + 2O2(g) CuSO4(s)
39
SECTION 4.3: Enthalpy of Formation and Combustion
1. (a) Write the thermochemical equation for the standard enthalpy of formation of
propanone.
(b) State the conditions under which the standard enthalpy changes are measured.
2. Calculate the ΔH° (in kJ mol-1) for the reaction
Fe3O4(s) + 2C(graphite) 2Fe(s) + 2CO2(g)
From the data below:
ΔH° (in kJ mol-1)
Fe3O4(s) -1118
CO2(g) -394
3. Calculate the ΔH° (in kJ mol-1) for the reaction
2NO2(g) N2O4(g)
From the data below:
ΔH° (in kJ mol-1)
NO2(g) +33.2
N2O4(g) +9.2
40
4. Hydrogen peroxide slowly decomposes into water and oxygen.
2H2O2(l) 2H2O(l) + O2(g)
Calculate the enthalpy change of this reaction from the data below:
ΔH° (in kJ mol-1)
H2O2(l) -188
H2O(l) -286
a. +98 kJ mol-1 c. +196 kJ mol-1
b. -98 kJ mol-1 d. -196 kJ mol-1
5. Calculate the enthalpy change for the hypothetical reduction of magnesium oxide by
carbon, according to the equation below from the data given. Comment on its feasibility.
2MgO(s) + C(s) CO2(g) + 2Mg(s)
From the data below:
ΔH° (in kJ mol-1)
CO2(g) -394
MgO(s) -602
SECTION 4.4: Bond Enthalpies
1. Which of the following processes are endothermic?
I. H2O(s) H2O (g)
II. CO2 (g) CO2 (s)
III. O2(g) 2O(g)
a. I and II only b. I and III only c. II and III only d. I, II, and III
2. Identify the equation which represents the bond enthalpy for H-Cl bond.
a. HCl(g) H(g) + Cl(g)
b. HCl(g) ½ H2 (g) + ½ Cl2 (g)
c. HCl(g) H+ (g) + Cl- (g)
d. HCl(aq) H+ (aq) + Cl- (aq)
41
3. Which of the following processes are endothermic?
I. CO2(g) CO2(g)
II. H2O (s) H2O (g)
III. O2(g) 2O (g)
a. I and II only b. I and III only c. II and III only d. I, II, and III
4. Identify bonds which are broken in the following process.
C2H6 (g) 2C(g) + 5H(g)
5. Which of the following is equivalent to the bond enthalpy of carbon – oxygen bond in
carbon monoxide?
a. CO(g) C(s) + O(g) c. CO(g) C(s) + ½ O2(g)
b. CO(g) C(g) + O(g) d. CO(g) C(g) + ½ O2(g)
6. Use the bond enthalpies from your IB Data Tables to calculate the ΔH for the reaction:
CH2CH2 + H2 CH3CH3
7. Use the bond enthalpies from your IB Data Tables to calculate the ΔH for the reaction:
2H2 (g) + O2(g) 2H2O(g)
8. Hydrogenation of the alkene double bond in unsaturated oils is an important reaction in
margarine production. Calculate the enthalpy change when one mold of C=C bond is
hydrogenated from the bond energy data shown below.
9. Use the bond enthalpy data from your IB Data Tables to calculate the enthalpy change of
reaction between methane and fluorine.
C2H4(g) + F2(g) CH2FCH2F(g)
a. +776 b. +164 c. -511 d. -776
42
10. Use the bond enthalpies from your IB Data Tables to estimate the enthalpy of combustion
of ethanol and comment on the reliability of your result.
11. The concentration of ozone in the upper atmosphere is maintained by the following
reactions.
12. Explain why ozone can be decomposed by light with a longer wavelength than that
required to decompose oxygen.
43
Unit 5: Entropy & Spontaneity Homework
SECTION 5.1: Entropy
1. Identify the process expected to have a value of ΔS closest to zero?
a. C2H4 (g) + H2(g) C2H6 (g)
b. H2(g) + Cl2(g) 2HCl (g)
c. CaCO3 (s) CaO (s) + CO2(g)
d. H2O (l) H2O (g)
2. Identify the process which have an associated increase in entropy.
I. Br2(g) Br2(l)
II. Br2(g) 2Br (g)
III. KBr (s) K+ (aq) + Br- (aq)
a. I and II c. II and III
b. I and III d. I, II, and III
3. Which is the best description of entropy and enthalpy changes accompanying the
sublimation of iodine: I2 (s) I2 (g)
a. ΔS+, ΔH+, reaction is endothermic
b. ΔS+, ΔH, reaction is exothermic
c. ΔS-, ΔH+, reaction is endothermic
d. ΔS-, ΔH-, reaction is exothermic
4. Identify the reaction which has the largest increase in entropy?
a. AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)
b. H2(g) + Cl2(g) 2HCl (g)
c. C2H4 (g) + H2(g) C2H6(g)
d. Mg(s) + H2SO4 (aq) MgSO4(aq) + H2 (g)
5. Predict the entropy change ΔS for the following reactions.
a. N2 (g) + 3H2 (g) 2NH3 (g)
b. Ba(OH)2 · 8H2O (s) + 2NH4SCN (s) Ba(SCN)2 (aq) + 2NH3 (aq) + 10H2O (l)
6. Sketch a graph to show how the entropy of a solid changes as the temperature increases.
44
7. Calculate the standard entropy change associated with the formation of methane from its
elements.
SECTION 5.2: Spontaneity
1. Ammonium chloride dissolves in water spontaneously in an endothermic process.
Identify the best explanation for these observations
a. Endothermic processes are energetically favorable
b. The bonds in solid NH4Cl are very weak
c. The entropy change of the system drives the process
d. The entropy change of the surroundings drives the process.
2. (a) Use the data from Table 12 in your Data Booklet and additional data (ΔH(H2O(s)) =
-292 kJ mol-1) given to calculate the enthalpy change that occurs when ice melts
.
(b) The entropy change when ice melts is 22.0 J K-1 mol-1. Deduce the value for the
melting point of ice.
3. Identify the combination which leads a reaction that is NOT spontaneous at low
temperatures but becomes spontaneous at higher temperatures.
a. ΔS – and ΔH –
b. ΔS + and ΔH –
c. ΔS - and ΔH +
d. ΔS + and ΔH +
4. The enthalpy and entropy changes for the reaction:
A(s) + B(aq) C(aq) + D(g)
Are:
ΔH = 100 kJ mol-1 and ΔS = 100 kJ mol-1
a. The reaction is not spontaneous at any temperature
b. The reaction is spontaneous at all temperatures
c. The reaction is spontaneous at all temperatures below 1000°C
d. The reaction is spontaneous at all temperatures above 1000°C
45
5. The decomposition of limestone can be represented by the equation:
CaCO3 (s) CaO (s) + CO2(g)
a. Predict the sign for the enthalpy of the reaction
b. Predict the sign for entropy change of the reaction
c. Deduce how stability of limestone changes with temperature.
6. Magnesium carbonate MgCO3 is a white solid that occurs in nature as the mineral
magnesite. Magnesite decomposes to the oxide at temperatures above 540°C.
MgCO3(s) MgO(s) + CO2 (g)
Identify the correct description of this reaction at 800°C.
ΔG ΔH ΔS
A + + +
B + - -
C - + +
D - + -
7. Calculate ΔGreaction for the thermal decomposition of calcium carbonate
CaCO3(s) CaCO(s) + CO2(g) From the following data, and comment on the significance of the value obtained.
Compound ΔGf (kJ mol-1)
CaCO3 (s) -1129
CaO (s) -604
CO2(g) -394
46
8. Calculate ΔGreaction at 2000K for the thermal decomposition of calcium carbonate from the data given below
Compound ΔHf (kJ mol-1) S (J K-1 mol-1)
CaCO3 (s) -1207 92.9
CaO (s) -635 39.7
CO2(g) -392 214
9. Which property of an element has a value of zero in its standard state? I. ΔHf II. Sf III. ΔGf
a. I and II c. II and III b. II and III d. I, II, and III
10. The standard enthalpy change for the formation of ethanol C2H5OH (l) and its molar entropy are in the IB Data Tables.
a. Write an equation for the formation of ethanol.
b. Calculate the entropy change for this process. The entropies of its constituent elements are C(graphite) 5.7 J K-1 mol-1, O2(g) 102.5 J K-1 mol-1.
c. Calculate the standard free energy change of formation of ethanol at 500K.
d. Predict the effect, if any, of an increase in temperature on the spontaneity of this reaction.
47
11. What signs of ΔHᶱreaction and ΔSᶱ
reaction for a reaction result in a complete reaction at all
temperatures?
ΔHᶱreaction ΔSᶱ
reaction
A - -
B + -
C - +
D + +
12. Which conditions correspond to a system of equilibrium?
I. the entropy of the system is at a maximum
II. the free energy of a system is at a maximum
III. ΔGᶱreaction = 0
a. I and II only c. II and III only
b. I and III only d. I, II, and III
13. Which values correspond to a reaction that can be reversed by changing the temperature?
ΔHᶱreaction ΔSᶱ
reaction
I - -
II + -
III + +
a. I and II only c. II and III only
b. I and III only d. I, II, and III
14. Propene reacts with hydrogen in the presense of nickel catalyst to form propane.
C3H6(g) + H2(g) C3H8(g)
ΔHᶱreaction = -123 kJ mol-1, ΔSᶱ
reaction = -128 J K-1 mol-1
Estimate the temperature range in which a mixture of all three gases will be present.
15. The Haber process is an important process in which ammonia is formed from nitrogen
and hydrogen.
N2(g) + H2(g) NH3(g)
ΔHᶱreaction = -198 kJ mol-1, ΔSᶱ
reaction = -93 J K-1 mol-1
Estimate the temperature range in which a mixture of all three gases will be present.
48
Unit 6: Kinetics Homework
SECTION 6.1: Collision Theory
1. Which statement is correct for a collision between reactant particles leading to a reaction?
a. Colliding particles must have different energy
b. All reactant particles must have the same energy
c. Colliding particles must have a kinetic energy higher than the activation
energy
d. Colliding particles must have the same velocity
2. Which of the following is (are) important in determing whether a reaction occurs?
I. Orientation of the molecules
II. Energy of the molecules
III. The volume of the container
a. I and II
b. I and III
c. II only
d. I, II, and III
3. If we compare two reactions, one which requires the simultaneous collision of three
molecules and the other which requires a collision between two molecules, which
reaction would you expect to be faster. Explain why.
4. Which of the following statements is correct?
a. A catalyst increases the rate of the forward reaction only
b. A catalyst increase the rate of the forward and backward reactions
c. A catalyst increases the yield of product formed.
d. A catalyst increases the activation energy
5. The rate of a reaction between two gases increases when the temperature is increased and
a catalyst is added. Which statements are both correct for the effect of these changes on
the reaction?
Increasing the temperature Adding a catalyst
A Collision frequency increases Activation energy increases
B Activation energy increases Activation energy does not change
C Activation energy does not change Activation energy decreases
D Activation energy increases Collision theory increases
49
6. The reaction between marble (calcium carbonate) and hydrochloric acid, which set of
conditions, would give the highest rate of reaction?
CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)
a. Marble chips and 10 mol dm-3 HCl
b. Marble powder and 1.0 mol dm-3 HCl
c. Marble chips and 0.1 mol dm-3 HCl
d. Marble powder and 0.1 mol dm-3 HCl
7. A sugar cube cannot be ignited with a match, but a sugar cube coated in ashes will ignite.
Suggest a reason for this observation.
8. Catalytic converters are now used in most cars to convert some components of exhaust
gases into less environmentally damaging molecules. One of these reactions convers
carbon monoxide and nitrogen monoxide into carbon dioxide and nitrogen. The catalyst
usually consists of metals such as platinum or rhodium.
a. Write an equation for this reaction.
b. Suggest why it is important to reduce the concentrations of carbon monoxide
and nitrogen monoxide released into the atmosphere
c. Why do you think the converter sometimes consist of small ceramic beads
coated with the catalyst?
d. Suggest why the converter usually does not work effectively until the car
engine has warmed up.
e. Discuss whether the use of catalytic converters in cars solves the problem of
pollution from cars.
50
SECTION 6.2: Rates of reaction
9. Consider the following reaction
2MnO4- (aq) + 5C2O4
2- (aq) + 16H+ (aq) 2Mn2+ (aq) + 10CO2 (aq) + 8H2O (l)
Describe three ways in which you could measure the rate of this reaction.
10. Based on the definition for rate of reaction, which units are used for a rate?
a. mol dm-3 c. dm3 time-1
b. mol dm-1 d. mol dm-3 time-1
11. The reaction between calcium carbonate and hydrochloric acid, carried out in an open
flask, can be represented by the following equation:
CaCO3 (s) + 2HCl (aq) CaCl2 (aq) + H2O(l) + CO2(g)
a. Suggest three different types of data that could be collected to measure the rate of
this reaction.
b. Explain how you would expect the rate of the reaction to change with time and
why.
12. The following data were collected for the reaction
2H2O2 (aq) 2H2O(l) + O2(g) at 390°C
H2O2 (mol dm-3) Time (s)
0.200 0
0.153 20
0.124 40
0.104 60
0.090 80
0.079 100
0.070 120
0.063 140
0.058 160
0.053 180
0.049 200
Draw a graph of concentration against time and determine the reaction rate after 60 s and
after 120 s.
51
SECTION 6.3: Rate expression
1. Write the rate expression for each of the reactions in the table below.
Equation for the reaction Order with
respect to
1st reactant
Order with
respect to
2nd reactant
Overall
order of
reaction
Rate expression
H2(g) + I2(g) 2HI (g) 1 1 2
2H2O2 (aq) 2H2O(l) + O2 (g) 1 - 1
S2O82-(aq) + 2- (aq) 2SO4
2-
(aq) + I2(aq)
1 1 2
2N2O5 (g) 4NO2 (g) + O2 (g) 1 - 2
2. The rate expression for the reaction NO(g) + O3 (g) NO2 (g) + O2(g) is:
rate = k[NO][O3]
What is the order with respect to each reactant and what is the overall order?
3. The reaction CH3Cl (aq) + OH- (aq) CH3OH (aq) + Cl- (aq) is found to be second
order overall. Give three possible rate expressions consistent with this finding.
4. Give the units of k in each of the rate expression below:
a. rate = k[NO3]2 d. rate = k[NO]2[Br]
b. rate = k[CH3CH2Br] e. rate = k[H2][I2]
c. rate = k[NH3]0
5. The reaction:
2N2O5(g) 4NO2 (g) + O2(g)
Has a value of k = 6.9 x 10-4 s-1 at a certain temperature.
Deduce the rate expression for this reaction.
52
6. A reaction involving A and B is found to be zero order with respect to A and second
order with respect to B. When the initial concentrations of A and B are 1.0 x 10-3 mol dm-
3 and 2.0 x 10-3 mol dm-3 respectively, the initial rate of the reaction is 4.5 x 10-4 mol dm-3
min-1. Calculate the value of the rate constant for the reaction.
7. The reaction between NO2 and F2 gives the following rate data at a certain temperature.
What’s the order of reaction with respect to NO2 and F2.
NO3 (mol dm-3) F2 (mol dm-3) Rate (mol dm-3 min-1)
0.1 0.2 0.1
0.2 0.2 0.4
0.1 0.4 0.2
NO2 order F2 order
a. First First
b. First Second
c. Second First
d. Second Second
8. The following data were obtained for the reaction of NO(g) with O2(g) to form NO2(g) at
25°C.
Experiment NO (mol dm-3) O2 (mol dm-3) Initial rate (mol dm-3 s-1)
1 0.30 0.20 2.0 x 10-3
2 0.30 0.40 4.0 x 10-3
3 0.60 0.80 3.2 x 10-2
Calculate the order with respect to the two reacts and write the rate expression for the
reaction.
53
9. If the reaction:
A + 2B products
Has the rate expression:
rate = k[A]2
Deduce the rates in experiments 2 and 3 in the table below.
Experiment A (mol dm-3) B (mol dm-3) Initial rate (mol dm-3 s-1)
1 0.01 0.01 3.8 x 10-3
2 0.02 0.01
3 0.02 0.02
10. If the reaction NO2(g) + CO(g) CO2(g) + NO(g) occurs by a one-step collision
process, what is the expected rate expression for the reaction?
11. 2NO2(g) 2NO(g) + O2(g)
is shown experimentally to be second order with respect to NO2. Is his consistent with
the mechanism shown below?
NO2 + NO2 NO3 + NO slow
NO3 + NO 2NO + O2 fast
12. Which statement about the following reaction at 450°C is correct?
I. the reaction must involve a collision between one O2 and two SO2 molecules’
II. every collision between SO2 and O2 will produce SO3
III. the rate-determing step is the slowest step of the reaction
a. I and III
b. II only
c. III only
d. None of the statements is correct
54
13. If the mechanism of a reaction is:
AB2 + AB2 A2B4 slow
A2B4 A2 + 2B2 fast
a. What is the overall equation for the reaction?
b. What is the rate expression for this reaction?
c. What units will the rate constant have in this expression?
SECTION 6.4: Activation energy
1. Consider the following statements
I. The rate constant of a reaction increases with increase in temperature
II. Increase in temperature decreases the activation energy of the reaction
III. The term A in the Arrhenius equation (k = Ae-Ea/RT) relates to the energy
requirements of the collisions.
Which statement(s) is/are correct?
a. I only c. I and III only
b. II only d. II and III only
2. To what does A refer in the Arrhenius equation, k = Ae-Ea/RT?
a. Activation energy c. Gas constant
b. Rate constant d. Collision geometry
3. The rate of a chemical reaction increases with increasing temperature. This increase in
rate is due to:
I. an increase in the collision rate
II. an increase in the activation energy
III. an increase in the rate constant
a. I and II only c. II and III only
b. I and III only d. I, II, and III
4. Rate constraints for the reaction
NO2(g) + CO(g) NO(g) + CO2(g) are given below.
At 700 K, k = 1.3 mol dm-3 s-1
At 800 K, k = 2.3 mol dm-3 s-1
Calculate the value of activation energy in kJ mol-1.
55
Specialized Tools Lab
Background:
When it comes to measuring volumes of liquids, general chemistry has relied on
graduated cylinders. Graduated cylinders come in a variety of sciences ranging from 10.0 mL to
500.0 mL. Each contains their own accuracy and precision. However, analytical chemists use a
wider range of tools when it comes to measure volumes of liquids. The first are volumetric
pipette or bulb pipettes, which allows extremely accurate (of four significant figures)
measurement of a volume of solution. These pipettes have a large bulb with a long narrow
portion above with a single graduation mark as it is calibrated for a single volume (like a
volumetric flask). Typical volumes are 10, 25, and 50 mL. Volumetric pipettes are commonly
used in analytical chemistry to make laboratory solutions from a base stock as well as prepare
solutions for titration. On how to use one can best be seen in the video at the following link:
https://www.youtube.com/watch?v=2WZ6sREzCsc
Similarly, a burette (also buret) is a device used in analytical chemistry for the dispensing
of variable, measured amounts of a chemical solution. A volumetric burette delivers measured
volumes of liquid. Since it uses gravity to deliver its measured volume, the scale of a buret is
listed inverse to the graduated cylinder. The link below is a video on how best to use a buret:
https://www.youtube.com/watch?v=ZmCTUB91HB4
Given the wide range of measurement tools for just liquids, they can categorized by their
accuracy and precision. Accuracy is determined by how close one can get to the target. Precision
is determined by how reproducible one can be to a designated spot. The two can be seen in the
picture below:
You need to be both accurate and precise to accomplish a lab. Likewise, uncertainty can
be a large factor in quantitative labs. Uncertainty is determined by the variation in readings of a
scale. For example, when reading a graduated cylinder you record what you can see and then
estimate the next lowest digit. Thus, this is where your uncertainty lies. The general rule of
thumb is that your uncertainty is a range consisting of +1 above the estimated lowest digit and -1
below the estimated lowest digit. The numerical value of a "plus or minus" (±) uncertainty value
tells you the range of the result. For example a result reported as 1.23 ± 0.05 means that the
experimenter has some degree of confidence that the true value falls in between 1.18 and 1.28.
See Ms. Manakul for clarification.
Furthermore, doing calculations with uncertainty requires its own set of calculations.
Generally, these would follow the same set of rules for significant figures and any general types
of calculation. If values with uncertainty are averaged. The uncertainty should also be averaged
separately. Then the reported number will included the average of the data and the average of the
uncertainty. For example:
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Trials 1 2 3 Average
Ruler Measurement
(cm)
2.67 ± 0.01 4.45 ± 0.01 3.32 ± 0.01 3.48 ± 0.01
Objective:
Students will gain best practice knowledge of specialized laboratory tools.
Students will be able to calculate uncertainty.
Methods:
In this lab you will be conducting two experiments. In the first experiment you will
establish how to use a volumetric pipette and an electronic balance beam. To do this, you will
need to determine as accurately as possible the weight of 10 mL of tap water. The best technique
to do this, is to obtain a 250 mL beaker and fill it with 200 mL of tap water. Next calibrate your
electronic balance beam by placing a clean and empty 250 mL beaker on the scale and zeroing it
out. We only want to know the mass of the water, and not the water and the beaker. Next place
the bulb at the end of the pipette, and place bottom end of the pipette into your tap water. Gently
measure out 10 mL of water into the pipette MAKING SURE NO WATER ENTERS THE
BULB AS THAT CAN DESTROY THE BULB. Once you have 10 mL of water in your pipette
you may place this water into the 250 mL beaker located on your scale. Once all your water is on
the scale record the weight in grams. Then you are repeat this step four more times for a total of
five trials. You do the same measurements for a 10 mL graduated cylinder and a burette. Each
will need a total of five trials.
In the second lab you will be determining the molarity of hydrochloric acid. To do this,
you will need to conduct a titration. Hydrochloric acid is a strong acid and will produce you a
salt and water when mixed with equal portions of a strong base. Furthermore, it will cause a pH
change, which can be seen using an indicator. Thus, you will need to reset your buret with 1.0 M
sodium hydroxide. Remember to rinse your buret with sodium hydroxide, so there are no
impurities in the buret before the titration. In a 250 mL flask, place 10.0 mL of the unknown
molarity of hydrochloric acid, and add 1 – 2 drops of a wide range indicator. Since we don’t
know exactly at what pH the hydrochloric acid will be neutralized, we will use an indicator
which will go through a range of colors depending on the pH of the solution.
Once everything has been set up, you may begin adding 1 drop of 1 M sodium hydroxide
to your 10 mL of unknown molarity of hydrochloric acid. Always, before recording your amount
of sodium hydroxide added, make sure to touch the tip of the burette to the side of the beaker.
This ensures that the volume you record was actually placed in the beaker. You can stop once
your solution turns a bluish/purple. This is where the solution has become neutralized and is
safer to handle. If it is dark purple, you have gone too far. Add a few drops of hydrochloric acid
to ensure proper neutralization. Be very careful when doing this, as it will affect your results.
When done, neutralized solutions may go down the drain with lots of water, all materials are
cleaned with soap and water, and placed in their designated areas.
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Safety:
Full lab gear is needed for this lab (i.e. goggles, closed toed shoes and lab coats). Sodium
hydroxide and hydrochloric acid are hazardous to eyes, skin, and lungs. Both solutions can easily
permeate or absorbed into the skin, cause ingestion, cause eye irritation, and can make it hard to
breath. Inflammation of the eye is characterized by redness, watering, and itching. Skin
inflammation is characterized by itching, scaling, reddening, or, occasionally, blistering. Should
your skin or eyes comes into contact with sodium hydroxide immediately rinse with water then
contact Ms. Manakul for further health precautions. If you spill occurs, contact Ms. Manakul
immediately, who will clean up the material or instruct you on proper clean up procedures.
Environmental Safety:
Ensuring proper neutralization of hydrochloric acid and sodium hydroxide will reduce the
change in pH of the environment. Furthermore, the particulates from the graduated cylinder,
buret, and pipette will be minimized by the dilution of water. Thus causing little to no harm to
the environment.
Data Collection:
1. Create a table to record the weights from the three trials of the water using the three tools
(i.e. pipette, buret, and graduated cylinder)
Data Processing:
1. Write the balanced equation for the neutralization of NaOH with HCl.
2. Calculate the amount of sodium hydroxide needed to neutralize 10.0 mL HCl.
3. Calculate the uncertainty within your trials.
Challenge Questions:
1. Which tool is more accurate in distilling accurate amounts of sodium hydroxide?
Conclusion & Evaluation:
1. Write your own conclusion and evaluation of this lab.
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Empirical & Molecular Formula Lab
Background:
An empirical formula gives the simplest whole number ratio of the different atoms in a
compound. The empirical formula does not necessarily indicate the exact number of atoms in a
single molecule. This information is given by the molecular formula, which is always a simple
multiple of the empirical formula.
In this experiment, you will determine the empirical formula of a magnesium-oxygen
product, a compound that is formed when magnesium metal reacts with oxygen gas. According
to the law of conservation of mass, the total mass of the products must equal the total mass of the
reactants in a chemical reaction. Therefore,
mass Mg + mass O2 = mass MgxOy
Since you will measure the mass of magnesium and the magnesium-oxygen product, you
will be able to calculate the mass of oxygen consumed during the reaction. Then, the ratio
between the moles of magnesium and the moles of oxygen consumed can be calculated. Finally,
the empirical formula can be written on the basis of this ratio.
Objective:
You will be able to determine the empirical and molecular formula of a compound using
experimental data.
Method:
1. Clean and dry a crucible and lid and place them on a clay triangle as shown in Figure 1. To
dry them, heat strongly for 1 to 2 minutes over the burner. Then let them cool to room
temperature. CAUTION: Hot porcelain. Use tongs to handle the crucible and lid
throughout the experiment.
2. Record the combined mass of the crucible and lid in your data table.
3. Polish a strip of magnesium ribbon with steel wool until it is shiny. Cut the strip into small
pieces and place them in the crucible. Record the combined mass of the crucible, lid, and
magnesium.
4. Heat the covered crucible gently over the burner. Lift the lid about every 20 seconds to allow
air in. CAUTION: Do not look directly at the burning magnesium metal. Avoid inhaling any
fumes.
5. When the magnesium appears to fully reacted (no longer producing an intense light),
partially remove the crucible lid as shown in Figure 2. Continue heating for 1 minute.
6. Turn off the burner. After the crucible has cooled to room temperature, carefully add a few
drops of water to decompose any nitrides that may have formed. CAUTION: Use care when
adding water. Too much water can cause the crucible to crack.
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7. Cover the crucible completely. Resume heating with the burner for 1 minute.
8. Turn off the burner. Cool the crucible, lid, and contents to room temperature. Record the
combined mass of the crucible, cover, and product.
9. Clean up your lab station. Discard the solid product in the designated waste beaker. Wash
and towel-dry the crucible and lid.
Safety:
You will be heating glassware. Hence, you will need to wear full lab gear (goggles, lab coat, and
closed toed shoes) throughout the entire experiment. Remember to always make sure to roll up
your sleeves instead of pushing up sleeves, since pushed up sleeves are more likely to fall down.
All hair or loose clothing should be tied back for the duration of the lab. Magnesium can give off
an intense light, causing damage to retinas. Thus, during the heating process, never look directly
at the magnesium.
Environmental Safety:
Magnesium oxide is a common salt, which when diluted with water causes little to no harm to
the environment. Excess magnesium will be collected and eliminated according to state and
national standards.
Data Collection:
Construct a data table to display the measurements taken during the procedure. Make
sure that each value is clearly labeled, including units.
Data processing:
Show all calculations and label all values.
1. Use your data to calculate the mass of magnesium and the mass of oxygen in the product.
2. Determine the empirical formula of the magnesium-oxygen product. When calculating the
mole ratio, round to the nearest whole number. In book problems, you should multiply by 2,
3, etc. to get a whole number ratio. In this case, you need to round in order to compensate for
experimental error.
3. Use the masses from #1 to calculate the percent composition (%Mg and %O) of the product.
4. The literature value for the %Mg in this magnesium-oxygen compound is 60.3%. Use this
value to calculate the percent error of your experimental %Mg.
5. Determine the limiting reactant, either showing calculations or by justifying your answer.
6. Using your data calculate the theoretical yield for your Magnesium-oxygen product. (must
use limiting reactant)
7. Using your data and the answer to number 6, calculate the actual yield and % yield for your
product.
8. Calculate the uncertainty within this lab.
Challenge Questions:
1. Explain how a clean and dry crucible was necessary for this lab.
2. Explain why the use of steel wool on the magnesium was essential for this lab.
Conclusion/Evaluation:
1. Write your own conclusion and evaluation for this lab.
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The Ideal Gas Law Lab
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
idea 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) and T is the temperature of the gas. You will
need to know the atmospheric pressure when you perform this lab. Go to
htt://www.wunderground.com and place the school’s 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.
Objective:
You will be able to calculate the molar mass of an ideal gas from experimental data.
You will be able to explain why most gases are not ideal.
Methods:
In this lab you will be testing to see if we change the amount of volume released from a
can of compressed gas will it affect the number of moles being used. Thus, you will need to
create a data table to record the mass and volume of the gas for this experiment. Be sure to
include the three trials you must accomplish for this lab. The mass and volume will be your
independent variables. The dependent variable will be the resulting measurements. Your
controlled variables will be the temperature of the room, pressure of the room, and the ideal gas
constant. To determine the mass of the gas you will need to record the change in mass of the
compressed can of gas by recording the before and after of the mass of the can. Thus before
setting up your apparatus you will need to weigh the compressed can of gas on a triple balance
beam before starting each trial. To determine the volume of your gas you will need to make the
following apparatus. Your apparatus is shown in the picture below:
In the picture you will be given a compressed can of gas, a hose, and clay. You will need
to attach the hose to the can with clay. Make sure that there are not gaps in the clay, because any
gaps will lead to gas escaping unintentionally. Next you will need to fill a tub with 3 inches of
tap water. Once done you will need to fill a 50 mL graduated cylinder to the very top so that you
create a concave curve on the very top with water. Place a watch glass curve side down on top of
the graduated cylinder. Carefully turn the graduated cylinder and watch glass open side down
into the tub of water, make sure not to spill any water. This will be how you can accurately
measure the volume of your gas. Gently remove the watch glass from the tub while keeping all
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the water in the graduated cylinder. After that, gently insert the open end of the hose into the
graduated cylinder, all the while making sure no water escapes from the graduated cylinder.
Now you can begin testing. In trial 1 you will need to make the volume of your gas 10 cm3 by
letting out the gas in the can until there is an air pocket of 10 cm3 in your graduated cylinder. In
trial 2 you will need to make the volume of your gas 15 cm3. In trail three you will need to make
the volume of your gas 20 cm3. When you are done separate your apparatus into its individual
pieces and put items in their designated spots.
Safety:
Full lab gear will be worn at all times. Never shake or tilt the can of compressed gas before or
during usage. Never use the gas or canister around a possible ignition source as it flammable.
Avoid contact of gas with your skin. DO NOT WASTE ANY OF THE COMPRESSED GAS!
Violations of any of these safety precautions is grounds for expulsion from the lab.
Environmental Safety:
In the lab, the gas obtained is a common gas, and will cause little to no environmental harm
when diluted with the atmosphere. The gas canister will be disposed of and recycled following
proper state and national standards to ensure little to no environmental harm.
Data Processing:
1. The atmospheric pressure you recorded earlier was in inches of mercury. To do the next
calculations will you will need to calculate your pressure into kPa. Hint: 1 in = 25.4 mm
and 760 mm Hg = 101.3 kPa.
2. Show the calculations of the number of moles for each trial. Make sure to label which
calculation goes with which trial.
3. Calculate the uncertainty in this lab.
Challenge Questions:
1. How does doing calculations in standard units affects the validity of our results? (Be very
clear in your explanation)
2. In this lab, the gas you obtained was not completely ideal. Explain how our data would be
affected if we used an ideal gas.
Conclusion/Evaluation:
1. Write your own conclusion and evaluation for this lab.
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Calorimetry Lab
Introduction:
Specific heat capacity is the amount of heat it takes to increase 1 gram of a substance
1°C. Every compound has different specific heat capacities. In this lab you will be determining
what type of metal the washer is. To do this, you will be using a calorimeter. Remember, a
calorimeter is a great vehicle for measuring temperature changes. By insulating the exchange in
temperature we can assume that the heat of the system is equivalent to the heat of surroundings.
You will be using a simplified version of a calorimeter with two Styrofoam cups. The foam cups
are good insulators, and will be able to simulate proper heat transfer with minimal error. Below
is a table of metals, which your washer could be with each specific heat capacity values.
Substance Specific Heat
Capacity (in J/g°C)
Water 4.18
Aluminum 0.897
Iron 0.449
Zinc 0.385
Tin 0.222
Lead 0.129
Stainless steel 0.927
Objective:
You will be able to do heat calculations.
Methods:
To create a calorimeter you will need two Styrofoam cups. At the bottom of one cup you
will need to poke a hole (some have been done for you). This will allow you to take the
temperature inside the cups without losing heat to the surrounding environment. You will stack
the Styrofoam cups on top of each other (the cup with the hole goes on top). They should look
like the picture below.
Place 50.0 mL of distilled water into the Styrofoam cups, then take the temperature of the
Styrofoam cups. This will be the initial temperature for the water. In a 250 mL beaker heat 100.0
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mL of distilled water to boiling. The hot plate should never be set higher than 7. Once the water
has begun to boil set the hot plate to 5. This will ensure the temperature of the water is hot
without losing too much water in the process. Weigh your washer on an electronic balance beam.
Place your metal washer into the hot water for 3 minutes. Take the temperature of your boiling
water. HINT: To get accurate results, hold the thermometer in the water. DO NOT let it sit on
the bottom of the beaker or leave the thermometer in the water for too long. It is made to
withstand high heat, but can break if subjugated to high heat for long periods of time.
Furthermore, you don’t want to take the temperature of the hot plate. Hence, hold the
thermometer with the bottom just above the bottom of the beaker. This will ensure you are
measuring the water’s temperature and not the hot plate. This temperature of hot water will be
the initial temperature of your metal washer, because the heat transferred to the boiling water is
the same amount of heat taken in by the metal washer.
Remove the washer from the hot water with tongs and IMMEDIATELY place the washer
into your calorimeter (Styrofoam cups). Wait 1 minute to allow proper heat movement. Record
this temperature. This is the final temperature of both the metal washer and your water. To repeat
this process again: take your washer out of the calorimeter, dry it, then place it back into the hot
water for another 3 minutes, then repeat the process again. You will need to do 3 trials.
Safety: You will be heating glassware. Hence, you will need to wear full lab gear (goggles, lab coat, and
closed toed shoes) throughout the entire experiment. Remember to always make sure to roll up
your sleeves instead of pushing up sleeves, since pushed up sleeves are more likely to fall down.
All hair or loose clothing should be tied back for the duration of the lab.
Environmental Safety:
There is little to no environmental safety for this lab, since all components are common
materials, and particulates from any metal transfer will be minimal in this lab.
Data collection:
Create a table to collect all the data from your three trials. Do not forget uncertainty.
Data Processing:
1. Calculate the specific heat capacity of the metal washer for EACH trial.
2. Calculate the average specific heat capacity for your trials.
3. Calculate the uncertainty for the average.
4. Calculate your percent error for this lab. Ms. Manakul will tell you the real composition
of the metal at the end of the lab.
Challenge Questions:
1. Explain how using distilled water yields more accurate results instead of tap water.
2. How would the use of a proper calorimeter affect the results of the lab?
3. If the washer was composed of a mixture of two metals, propose a method for
determining the percentage of metals within the washer.
Conclusion/Evaluation:
1. Write your own conclusion and evaluation for this lab.
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Entropy Lab
Background:
Reactions in nature are driven by two forces, which in combination determine whether or
not the reaction will be spontaneous. Firstly, reactions that are exothermic (give off heat) are
generally favored by nature. However, some endothermic reactions, such as the melting of ice,
are spontaneous and thus favored by nature. The second driving force that also determines
whether or not a reaction will occur is entropy. Entropy can be defined as a measure of the
degree of randomness of the particles, such as molecules, in a system. As might be expected
from the chaotic world in which we live, nature favors an increase in entropy. In other words,
reactions that increase the disorder of the system tend to be spontaneous. The amount of entropy
of a system is best understood by considering the three principle states of matter. In a solid, the
particles vibrate in place and are not free to switch places with each other. As such, solids are
considered to have very low entropy because very little randomness exists in them. Liquids, in
general, are more disorderly than solids, and thus have higher entropy. Gases, the most
disorderly of the three states possesses the highest entropy. These are general guideline as some
liquids (mercury, for instance) have lower entropy than that certain solids. In general, the
dissolution process increases the entropy of a system. Entropy of substances can be determined
quantitatively in the lab and are measured in molar values with units kJ/(mol·K). In this class, we
will be most concerned with the change in entropy of a system, denoted ΔS. In this lab you will
be conducting three tests to determine the enthalpy and entropy changes.
Objective:
You will be able to determine the entropy of a reaction given your knowledge of how
disorder affects entropy.
Methods:
You will first need to read through the following tests, and make a prediction for each what the
sign of enthalpy & entropy would be, then follow the steps for each test. When done clean up
your station by putting all materials away and spraying your station down with soap and water,
and drying it with a paper towel.
Test #1
1. Measure out 10 ml of 1.0 M sodium bicarbonate into a test tube. Carefully add 10 ml of 3
M acetic acid to the test tube.
2. Record your observations of the reaction.
3. Light a wood splint using matches. Blow out the flame and place the glowing splint in the
mouth of the test tube. (Be careful not to drop the splint into the test tube).
4. Record your observations
5. Dispose of your solution as directed by your teacher and clean your glassware before
proceeding.
Test #2
1. Add 10 ml of ethanol to a clean small plastic bottle.
2. Swirl the contents of the bottle for approximately 2 minutes, and drain the ethanol down
the drain.
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3. Let the bottle stand for 30 seconds. Light a wood splint with a match and carefully pass it
over the mouth of the bottle.
4. Record your observations.
Test #3
1. Measure out 100 ml of distilled water in a 250 ml beaker and record its temperature.
2. Measure out 2.50 grams of ammonium chloride. Add the ammonium chloride to the
distilled water.
3. Record the temperature of the solution after the dissolution process is complete.
Safety:
You will lighting items on fire, thus, you will need full lab gear for this lab. Meaning you will
need to wear a lab coat, goggles, and closed toed shoes at all time. All dangly jewelry should be
tucked away, and long hair tied back. Ethanol is highly flammable, so keep it far away from any
fire source. Ammonium chloride is a strong base and can cause skin, eye, and respiratory
irritation. Do not break the ammonium directly. If your skin comes in contact with ammonium
chloride or ethanol rise with water and then notify Ms. Manakul for any further medical
consultation. If you spill ethanol or ammonium chloride notify Ms. Manakul immediately, and
she will either clean it up or direct you on proper disposal.
Environmental Safety:
Ammonium is a highly environmental toxin and will be disposed of individually, and following
the strict national and state mandated standards. This will ensure the toxic chemical will not
affect the environment.
Data Collection:
Create a data table to collect your observations for each test.
Data Processing:
There is no data processing for this lab. (HINT: Write this entire statement out in your lab
notebook)
Challenge Questions:
Test #1
1. What gas was produced in this test? How do you know? (Hint: write out the reaction)
2. Write out two complete balanced equations for the reactions. (The first reaction should be
a double displacement, and the second reaction should be a decomposition of carbonic
acid).
3. Were the reactions exothermic or endothermic?
a. How might you change the experiment to determine this?
4. Based on your chemical equation, did the system’s entropy increase or decrease in each
reaction. Explain your reasoning?
Test #2
5. Write a complete balanced equation for the combustion of ethanol. (The products of this
combustion are CO2 (g) and H2O (l)).
6. Is this product exothermic or endothermic? How can you tell?
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7. Based on the number of gas molecules on each side of the equation, did the entropy of the
system increase or decrease?
Test #3
8. Write out a complete balanced equation for this dissociation.
9. Was this reaction exothermic or endothermic? Explain your reasoning.
10. Did entropy of the system increase or decrease in this reaction? Explain your reasoning.
Conclusion/Evaluation:
1. Write your own conclusion and evaluation for this lab.
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Antacid Rate of Reaction Lab
Background:
Antacids are substances which neutralizes stomach acids. They are generally found in
tablet form, and cause a reaction to produce gas when mixed with water. The most common
antacid is Alka-Seltzer. Alka-Seltzer is a combination of sodium bicarbonate, aspirin, and
anhydrous citric acid. The bicarbonate in sodium bicarbonate is the main component for
neutralizing stomach acids. The reaction is as follows:
NaHCO3 (s) + H2O (l) NaOH (aq) + CO2 (g)
The reaction rate (rate of reaction) or speed of reaction for a reactant or product in a
particular reaction is intuitively defined as how fast or slow a reaction takes place. For example,
the oxidative rusting of iron under Earth's atmosphere is a slow reaction that can take many
years, but the combustion of cellulose in a fire is a reaction that takes place in fractions of a
second. There are four things which greatly affect rate of reaction: temperature, particle size,
concentration, and use of a catalyst. In this lab you will be graphing, and comparing how particle
size affects the rate of reaction.
Objective:
You will be able to determine the rate of a reaction of an antacid tablet.
You will be able to see how rate of reaction is affected by particle size and concentration.
Methods:
You will be determining the rate of reaction for an Alka-Seltzer tablet (whole and
crushed) placed in water. You will need to first write a hypothesis on which will have a faster
rate of reaction, a whole tablet, full crushed tablet, ½ tablet, or ¼ tablet. To conduct the
experiment, you will need to obtain three Alka-Seltzer tablets. You will need to keep one full
tablet, cut one tablet into ½ and ¼ sized piece, and fully crush one full tablets. You will also need
two clean 250 mL flasks. Place 100 mL of distilled water in each flask. Before you put anything
else in the flask, be prepared to place a mouth of a balloon over the mouth of the flask, so that no
air can be gained or lost. You will need to do this quickly since the gas evolution will begin as
soon as the tablet/pieces encounter the water. REMEMBER to check the balloon for holes,
otherwise, you will have to redo the trial.
In beaker 1 place the whole antacid tablet, then record the size of the diameter of the
balloon after every 30 seconds for five minutes or until 90% of the table has been dissolved.
In beaker 2 place the crushed antacid tablet, then record the size of the diameter of the balloon
after every 30 seconds for five minutes or until 90% of the table has been dissolved.
After beaker 1 is complete, pour contents down the drain with lots of water, and reset the beaker
for another experiment. In this experiment, you will add half of an antacid tablet. Then record the
size of the diameter of the balloon after every 30 seconds for five minutes or until 90% of the
table has been dissolved.
After beaker 2 is complete, pour contents down the drain with lots of water, and reset the
beaker for another experiment. In this experiment, you will add a quarter of an antacid tablet.
Then record the size of the diameter of the balloon after every 30 seconds for five minutes or
until 90% of the table has been dissolved.
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Safety:
You will be working with an effervescent compound (gives off gas). Even though Alka-Seltzer is
a commonly used compound and gives off a harmless gas, large quantities of carbon dioxide can
be harmful. Make sure the reaction occurs in a well ventilated area, and never place your face
directly at the mouth of the reaction. Furthermore, the chemical reaction can cause skin and eye
irritation, thus full lab gear should be worn at all times. Full lab gear includes goggles, lab coat,
closed toed shoes, long hair tied back, loose jewelry put away, and long sleeves rolled.
Environmental safety:
All chemicals are common and non-toxic chemicals when diluted with water. Thus, the disposal
with lots of water will ensure little to no harm to the environment.
Data Collection:
1. Create a data table to record the size of the balloon for every 30 seconds for your
independent variables (whole tablet, full crushed tablet, ½ tablet, and ¼ tablet).
Remember to include uncertainty.
Data Processing:
1. Determine the rate expression for this reaction.
2. Make a line graph for the rate of reaction for the whole antacid tablet, full crushed tablet,
½ tablet, and ¼ tablet. The x-axis is the time, and the y-axis should be amount of gas
obtained. Remember to create a key.
3. Calculate the average uncertainty for this lab.
Challenge Questions:
1. Explain how the rate was affected by the size of the antacid tablet. Use your graphs from
the data processing section to support your claim.
Conclusion/Evaluation:
1. Write your own conclusion and evaluation for this lab.
Appendix
A
Stoichiometry Mole Map
B
VESPR Theory
Species Type Geometry Shape Bond Angle Example
A2 Linear Linear 180° H2
AX2 Linear Linear 180° CO2
AX3
Planar
Triangular
Planar
Triangular 120° BF3
AX2E1
Planar
Triangular V-Shape 104.5° SO2
AX4 Tetrahedral Tetrahedral 109.5° CH4
AX3E1 Tetrahedral Pyramidal 109.5° NH3
AX2E2 Tetrahedral V-Shape 104.5 H2O
AX5 Triangular
Bipyramidal
Triangular
Bipyramidal
90°, 120°,
180° PCl5
AX4E1
Triangular
Bipyramidal See Saw
90°,
120°¸180°
SF4
AX3E2
Triangular
Bipyramidal T-Shape 90°, 180° ClF3
A3E3 or
AX2E3
Triangular
Bypiramidal Linear 180° I3
-
AX6 Octahedral Octahedral 90°, 180° SF6
AX5E1 Octahedral Square
Pyramidal 90°, 180° BrF5
AX4E2 Octahedral Square Planar 90°, 180° XeF4
C
IUPAC Tables
Number of carbon atoms in longest chain Stem in IUPAC name
1 Meth-
2 Eth-
3 Prop-
4 But-
5 Pent-
6 Hex-
Homologous Series Functional Group Suffix in IUPAC Name
Carboxylic acid R-COOH -oic acid
Ester* R-COOR -oate
Amide* R-CONH2 -amide
Nitrile R-C=N -itrile
Aldehyde R-CHO -al
Ketone R-C=O -one
Alcohol R-OH -ol
Amine* R-NH2 -amine
Alkane R-C-C-R -ane
Alkene R-C=C-R -ene
Side chain/ substituent group Prefix in IUPAC name Example of Compound
-CH3 Methyl CH3CH(CH3)CH3 2-methylpropane
-C2H5 Ethyl CH(C2H5)3 3-ethylpentane
-C3H7 Propyl CH(C3H7)3 4-propylheptane
-F, -Cl, -Br, I- Fluoro, Chloro, Bromo, Iodo
CCl4 tetrachloromethane
-NO2 Nitro CH2(NO2)COOH 2-nitroethanoic acid
D
International Bachelorette Data Tables