YORK MILLS COLLEGIATE INSTITUTE
SCIENCE DEPARTMENT
Jim Henson’s Muppets ©
Grade 11 Chemistry
University Level
SCH3U - ____
STUDENT WORKBOOK
For the school year:
2018 - 2019
Name: _________________________________________________
1
TABLE OF CONTENTS
Table of Contents.............................................................................................................................................................................. 2
Course Outline: Grade 11 Chemistry – University (SCH3U) ........................................................................................................... 4
Safety Rules in the Chemistry Lab.................................................................................................................................................... 6
Basic Rules for Multiple Choice Questions ....................................................................................................................................... 7
Lab Outline & Lab Format ................................................................................................................................................................. 8
Significant Digit Rules ..................................................................................................................................................................... 10
Problem Solving Format ................................................................................................................................................................. 12
Writing Study Notes ........................................................................................................................................................................ 13
Matter, Chemical Trends & Chemical Bonding ............................................................................................................................... 14
Atomic Notation............................................................................................................................................................................... 14
Isotope Problems ............................................................................................................................................................................ 14
Variation in Atomic Properties ......................................................................................................................................................... 15
Electronegativity & Types of Intramolecular Bonding ..................................................................................................................... 15
Ionic Bonding .................................................................................................................................................................................. 15
Covalent Bonding............................................................................................................................................................................ 16
VSEPR Structures & Polarity .......................................................................................................................................................... 16
Chemical Reactions Unit................................................................................................................................................................. 17
Chemical Nomenclature of Ionic Binary Compounds ..................................................................................................................... 17
Chemical Nomenclature of Multivalent Compounds ....................................................................................................................... 17
Chemical Nomenclature of Simple Polyatomic Compounds ........................................................................................................... 17
Chemical Nomenclature of More Polyatomic Compounds ............................................................................................................. 18
Chemical Nomenclature of Acidic & Basic Compounds ................................................................................................................. 18
Chemical Nomenclature of Covalent Binary Compounds ............................................................................................................... 18
Chemical Nomenclature Review ..................................................................................................................................................... 19
Balancing Skeleton Equations ........................................................................................................................................................ 19
Balancing Word Equations.............................................................................................................................................................. 20
Identifying Types of Reactions ........................................................................................................................................................ 20
Combustion, Synthesis & Decomposition Reactions ...................................................................................................................... 20
Single Displacement Reactions ...................................................................................................................................................... 21
Double Displacement Reactions ..................................................................................................................................................... 21
Net Ionic Equations ......................................................................................................................................................................... 21
Quantities in Chemical Reactions Unit............................................................................................................................................ 21
Molar Mass Problems ..................................................................................................................................................................... 21
Avogadro’s Number Problems ........................................................................................................................................................ 22
Mole Related Problems (# of Particles) .......................................................................................................................................... 22
Mole Related Problems (# of Moles)............................................................................................................................................... 22
Mole Related Problems (Advanced Calculations) ........................................................................................................................... 23
% Composition Problems (from Chemical Formulas) ..................................................................................................................... 23
% Composition Problems (from Experiment) .................................................................................................................................. 23
Simplest Formula Problems ............................................................................................................................................................ 24
Molecular Formula Problems .......................................................................................................................................................... 24
Hydrate Problems ........................................................................................................................................................................... 24
Simple Stoichiometry Problems (mol-mol) ...................................................................................................................................... 25
Excess Reagent/Reactant Mass-Mass Stoichiometry Problems .................................................................................................... 25
Limiting Reagent/Reactant Stoichiometry Problems ....................................................................................................................... 26
Stoichiometry Review Problems ..................................................................................................................................................... 26
% Yield Problems............................................................................................................................................................................ 27
% Purity Problems .......................................................................................................................................................................... 27
Solutions & Solubility Unit ............................................................................................................................................................... 28
Determining Solubility ..................................................................................................................................................................... 28
Concentration Unit Conversions ..................................................................................................................................................... 28
Concentration Formula Problems ................................................................................................................................................... 28
Molar Concentration Problems ....................................................................................................................................................... 29
Dilution Problems ............................................................................................................................................................................ 29
Solutions & Solubility Problems ...................................................................................................................................................... 30
2
Solubility Problems ......................................................................................................................................................................... 30
Solution Stoichiometry Problems .................................................................................................................................................... 31
Conjugate Acids & Bases ............................................................................................................................................................... 32
pH Calculations ............................................................................................................................................................................... 32
pH & pOH Calculations ................................................................................................................................................................... 32
pH & Strong Acid Problems ............................................................................................................................................................ 33
Strong Acid & Strong Base Titration Problems ............................................................................................................................... 33
Strong Acid & Strong Base Neutralization Problems ...................................................................................................................... 33
Gases & Atmospheric Chemistry Unit............................................................................................................................................. 34
Pressure & Temperature Conversions............................................................................................................................................ 34
Boyle’s Law Problems..................................................................................................................................................................... 34
Charles’ Law Problems ................................................................................................................................................................... 35
Gay-Lussac’s Law Problems .......................................................................................................................................................... 35
Combined Gas Law Problems ........................................................................................................................................................ 35
Ideal Gas Law Problems ................................................................................................................................................................. 36
Ideal Gas Law Molar Volume Problems.......................................................................................................................................... 36
Ideal Gas Law Molar Mass Problems ............................................................................................................................................. 37
Ideal Gas Law Density Problems .................................................................................................................................................... 37
Gas Stoichiometry Calculations ...................................................................................................................................................... 37
Dalton’s Law of Partial Pressures Problems ................................................................................................................................... 38
Water Vapour Pressure Problems .................................................................................................................................................. 38
Reference Material.......................................................................................................................................................................... 39
Conversion Factors for Calculations ............................................................................................................................................... 39
Constants ........................................................................................................................................................................................ 39
Useful Formulas .............................................................................................................................................................................. 39
Useful Polyatomic Ions ................................................................................................................................................................... 40
Metal & Halogen Activity Series ...................................................................................................................................................... 40
Solubility Table @ SATP................................................................................................................................................................. 40
Water Vapour Pressure Table ........................................................................................................................................................ 41
Periodic Table ................................................................................................................................................................................. 42
Last revised: August 20, 2018
3
COURSE OUTLINE: GRADE 11 CHEMISTRY – UNIVERSITY (SCH3U)
Chemistry Teachers:
E. Lindala, A. Salloum
Phone: 416-395-3340 x 20125
Text:
Jenkins, F., H. Van Kessel, L. Davies, O. Lantz, P. Thomas, D Tompkins, M. DiGiuseppe.
2002. Chemistry 11. Nelson Publishers, Toronto.
Prerequisite:
SNC2D
Course Description:
This course enables students to deepen their understanding of chemistry through the study of properties of
chemicals & chemical bonds; chemical reactions & quantitative relationships in those reactions; solutions &
solubility; & atmospheric chemistry & the behaviour of gases. Students will further develop their analytical skills &
investigate the qualitative & quantitative properties of matter, as well as the impact of some common chemical
reactions on society & the environment. Emphasis will also be placed on the importance of chemistry in other
branches of science.
Units of Study:
1. Matter, Chemical Trends & Chemical Bonding
2. Chemical Reactions
3. Quantities in Chemical Reactions
4. Solutions & Solubility
5. Gases & Atmospheric Chemistry
Evaluation Outline:
TERM WORK:
Knowledge/Understanding
Application
Communication
Thinking
70
30
100
Total – Term Work
SUMMATIVE EVALUATION:
COURSE TOTAL:
30 (± 5)
15 (± 5)
15 (± 5)
10 (± 5)
Final Exam
Policies & Expectations:
1. Attendance is critical for success in this course & the school attendance policy will be adhered to. If a test or
assignment due date is missed, inform your teacher in advance and provide documentation indicating that the
student is unable to attend school for the day & time the test or assignment is due, & acknowledge that the
test/assignment is being missed. The test/assignment will be accommodated when the student returns to
school, at the teacher’s convenience.
2. An assignment cannot be handed in for marking after the same marked assignment has been returned. All
assignments are due at the beginning of class. Late marks may be deducted.
3. Students are expected to come prepared to each class, with their course notes, student workbook, calculators,
periodic tables & other reference sheets as required.
4
4. Students are expected to do approximately 30 minutes of homework (on average) following every class. The
actual amount will depend on how much is accomplished during class time. Through much of the course,
understanding new concepts will depend on previous concepts & homework.
5. Plagiarism is cheating. Any copied work in which the source is not correctly cited is plagiarism, whether the
source is a book, the Internet, or another students’ work. All plagiarized work will receive a mark of zero, with no
opportunity to resubmit the assignment. In addition, the student’s name will be entered into the plagiarism
registry, in accordance with the school plagiarism policy.
6. Students in need of extra help are encouraged to ask for this help when needed. Please make an appointment
at a time that is mutually convenient for both the teacher & student. Extra help should not be expected on the
day of a formally announced evaluation. Students should be prepared to ask specific questions, e.g. question ##
on page ##.
7. On occasion, it will be required for students to email an electronic copy of work as part of their evaluation.
8. A complete version of the Course Outline for SCH3U can be found at: http://schools.tdsb.on.ca/yorkmillsci/
5
SAFETY RULES IN THE CHEMISTRY LAB
As a chemistry student, you MUST read & follow the instructions below:
1. Lab behaviour must be safe, calm & focussed at all times during a lab.
2. Eye protection must be worn over your eyes when directed by the teacher.
3. Be familiar with the location & use of safety equipment in the classroom & of the hazard information for all the
chemicals being used.
4. Report all accidents to the teacher at once, no matter how minor.
5. Report broken, damaged or defective equipment & unidentified chemicals to the teacher.
6. Place all cracked or broken glassware in the broken glass container at the front of the room.
7. Use separate scoopulas for separate chemicals.
8. Take only the amount of chemical needed & transport it in a safe container between the supply bench & your
work area.
9. Replace all container lids immediately after use.
10. Keep hands, scoopulas, pens, pencils, etc., away from your mouth
11. Leave all food out of the lab. During a lab, water bottles must also be left outside.
12. Maintain an uncluttered work area—the only things that should be at the lab bench are the lab outline, a pen &
your lab notebook/report.
13. Remain standing while performing any experiment.
14. Clean your lab area before leaving, including washing & putting away equipment & removal of all solids from
sinks & drains.
15. Beware of any liquids or solids on the lab benches—they may be corrosive.
16. Ensure that you understand the safe procedure to be followed before attempting any experiment.
17. Wash hands thoroughly after lab work.
In any lab exercise, safety is the first priority & is the responsibility of every individual
in the room. The purpose of these safety rules is to ensure a safe environment for
everyone & it is extremely important that they are followed during a lab.
Students who do not comply with these safety rules will be asked to leave & will
receive a zero on the assignment/lab.
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BASIC RULES FOR MULTIPLE CHOICE QUESTIONS
Here are some suggestions on improving your multiple choice testing skills.
1.
Read the question & try to answer it BEFORE looking at the answers
 By thinking of the answer first, you are less likely to be fooled by a wrong answer
 Read through all the answers before choosing…the best answer is not always the first correct answer you see
For example:
Toronto can be described as Ontario’s
a. Capital city
b. Largest city
c. Most multicultural city
d. All of the above
e. None of the above
2.
Do not spend too much time on any one question
 Sometimes the question will seem to have no right answer…your teacher may have made a mistake. There may not
be a right answer. It is important not to waste too much time trying to answer an impossible question. Choose an
answer at random, but circle the question number so that you can come back to it later if you have extra time.
 Similarly, if a question is too hard, or you just do not know the answer, choose an answer at random & come back
when you have completed all the questions you do know. Use whatever time is left over at the end of the test to tackle
these very difficult questions
For example:
The capital of Ontario is
a. Markham
b. Mississauga
c. Pickering
d. Richmond Hill
e. Windsor
3.
…the answer is (d)
…there is no correct answer
If the question asks you something you do not know, see if you can cross out any of the wrong answers before you guess
 If you eliminate answers you know are incorrect, you increase your chances of getting it right, just by guessing
For example:
The capital of British Columbia is
a. Edmonton
b. Ottawa
c. Vancouver
d. Victoria
e. Winnipeg
…answer is (d)
4.
Do not keep changing your answer.
 Research shows that your first choice was probably the right one. Most people who change their answers will change
from a correct one to a wrong one. Only change your answer if you are absolutely sure you made a mistake. (For
example, if another question on the test reminds you of the right answer)
5.
After you have finished the test, go back to those questions you circled as being too hard or as having no right answer
 Do not finish a test early unless you are sure you have answered everything to the best of your ability. Do not leave
any questions blank
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LAB OUTLINE & LAB FORMAT
All labs must be formatted as follows, with numbers 1 to 3 completed before the first lab:
1. All pages must be in (water insoluble) ink pen (preferably blue).
2. The date & title of the lab must head the first page of each lab.
3. All notations, including tables, observations, calculations & diagrams, must be made with pen only. All written sections
should be single-spaced.
4. Nothing should be whited out. If you want to make a change, draw a line neatly through the work to be changed &
rewrite it.
5. Only one side of the page should be used for the lab.
EVALUATION
The lab component of this course is 15% ±5%. You will be evaluated as follows:

for each lab, you will receive a performance mark based on preparedness, performance, understanding & safety
considerations

each lab will be submitted for marking for format & content—not every part of every lab will be marked

time permitting, there may be one formal lab write-up based on your notes
**A NOTE ABOUT PLAGIARISM**
As with any written work, it is considered plagiarism to copy work from another student, or to allow your work to be copied. While
it is acceptable to work with a partner on the lab, you must write every part of your lab report independently. You are each
expected to hand in your own work, with the report written in your own words. Copied lab reports will be given a mark of 0.
Lab Format
For every lab, a pre-lab must be completed before you arrive to the lab class; the pre-lab (purpose, introduction, chemical
hazards, method/procedure) will be checked at the beginning of the lab, & will count towards your lab mark. In general, lab
reports must be written in passive tense, with most sections are in past passive tense. Lab reports also have the same format
with the following headings & sections:
PURPOSE
• the purpose of the lab must be clearly stated, usually in point form
INTRODUCTION
• must be written as a paragraph, using complete sentences (about 5-6 sentences)
• begins with a description & explanation of the concept of the lab
• must include any special safety considerations at the end of the introduction
CHEMICAL HAZARDS
• a brief table (½ page maximum) containing the relevant health & first aid information for the lab’s major chemicals
• the information may be obtained from on-line MSDS (Material Safety Data Sheets)
METHOD/PROCEDURE
1. Consists of a flowchart showing how to do the lab
• a long vertical arrow (or series of arrows), drawn with a ruler, represents the order of events
• substances added are shown by a labeled arrow flowing into the main arrow from the side
• anything done to the experimental mixture is labeled beside the main arrow in a location appropriate to the time at
which it is done (see example below)
• if components of the mixture are separated &/or treated differently, branches must split off the main arrow in the
appropriate spot
Materials/Equipment
Actions Taken
100 mL beaker
Added 10.0 mL of water
Added 1.00 g salt
Stirring Rod
Stirred contents in beaker
8
2.
Includes a diagram of any relevant apparatus or set-up
• all diagrams should be neatly drawn & labeled—use a ruler & avoid sketches
• diagrams are 2-dimensional figures
• each diagram must have a number & a descriptive title (i.e., Figure 1: Distillation Apparatus)
Burette
Stopcock
Retort Stand
Sample in Beaker
Figure 1: Setup of a Titration
OBSERVATIONS
• must include qualitative observations, in table form or in a series of point-form statements
• table format (see below) must be used if substances are being compared
• observations must follow the order of the procedure
• must include any quantitative observations, including all raw data (every measurement) organized into a table
• all tables must have a border, & lines separating columns & rows
• each table must have a number & a descriptive title (i.e., Table 1: Observations of Esters)
• rows & columns must be used logically & appropriately to organize the data
• for any calculated values included in a table, a sample calculation should be provided below the table
Table 1: Measurements of Length & Width
Object Length (m) Width (m)
RESULTS/ANALYSIS
• may include calculations
• correct problem-solving format & significant digits are required
• a heading must be used to indicate what is being determined
• if a series of calculations of the same type are being done, show one sample calculation & give the answers for
the rest
• may include balanced chemical equations &/or graphs
• correct graphing format must be used
• graphs may be completed on a computer & attached into the lab book
• every graph must have a number & descriptive title (i.e., Figure 1: Concentration of Iodine Solution)
• any questions in this section should be answered separately & in order, & the questions themselves should NOT be
copied from the lab sheet
DISCUSSION
• consists of answers to lab questions
• should be answered separately & in order
• written in complete, correct sentences which clearly indicate what the question was
• questions from the lab sheet should NOT be copied into the lab report
• must include reference to appropriate results where required
• state table or figure number & data or observation that supports the answer to the question
• usually requires thorough analysis of the data presented in the Results section
• avoid “human error” in error analysis & focus on the assumptions made in the lab
9
EXPERIMENTAL ERROR IS NOT THE SAME AS HUMAN ERROR!
• Experimental error is unavoidable! As you do your lab, consider the following:
• Assuming the technique is correct, what errors are associated with the lab or
concept behind the lab
• Are there any side reactions? (be specific!)
• Why do the experimental values do not match the “real” answer?
• Human error definition: “a source of error that can be avoided” (i.e., spillage, etc)
…in summary: if you can do something about it, it is a human error, if you can’t avoid it, it is
experimental error
CONCLUSION
• consists of one or two sentences with reference to the purpose of the lab
• should make specific reference to the significant result that addresses the purpose
SIGNIFICANT DIGIT RULES
Significant digits/figures (sig digs/figs) represent a degree of precision or accuracy in scientific measurements. The limitations
on precision must be reflected in calculations by using the appropriate number of significant digits.
For example, on a mass balance, mass can only be measured to 2 decimal places. We can measure 1.98 g
(3 sig digs) of Na2CO3, but not 1.975 g (4 sig digs) of Na2CO3.
Counting Significant Digits:
In order to use significant digits accurately, you need to know which digits are significant! If you are dealing with scientific
notation, count the digits in the first part.
1.
All non-zero digits are significant
 this means 68.3 mL has 3 sig digs & 165.35 cm has 5 sig digs
2.
Zeroes that have non-zero digits on either side are significant
 this means 307.5 L has 4 sig digs & 20.004 m have 5 sig digs
3.
Zeroes at the beginning of a number are not significant, because these zeroes are only placeholders
 this means both 0.0265 mm & 0.0000265 mm both have 3 sig digs
4.
Zeroes at the end (right side) of a number are significant if they appear after a decimal, otherwise they are insignificant
 this means 1.5000 km has 5 sig digs but 15000 km has only 2 sig digs
Rules for Using Significant Digits:
1. When multiplying or dividing measured values, the answer should be expressed with the same number of significant digits
as the data with the lowest number of significant digits.
2. When adding or subtracting, the answer should be expressed with the same number of decimal places as the data with the
least number of decimal places.
3. Some numbers are exact. For example, counted objects such as people, pens & calculators are considered to have an
infinite number of significant digits. This is because counting is not limited by a measuring device.
4. For all calculations, it is safer to carry at least ONE extra significant digit until you complete the final calculation. Rounding
too far too soon can lead to a significant impact on your answer.
5. Significant digits involving logarithms require a slightly different set of rules. See section below.
Example:
Suppose a group of students had to measure the dimensions of a box for the purpose of finding the total surface area. The
measurements were found to be: 1.1 m, 289 cm, 1500.5 mm. Show all your work.
10
l = 1.1 m
w = 289 cm = 2.89 m
h = 1500.5 mm = 1.5005 m
SA = 2lw + 2lh + 2hw
SA = 2(1.1)(2.89) + 2(1.1)(1.5005) + 2(2.89)(1.5005)
SA = 18.33199 m2  18 m2
 the surface area was 18 m2.
Notice the length has 2 sig digs whereas the width & height have 3 & 4 sig digs, respectively…the final answer
should be reported using the same number of significant digits as the data with the least number of significant
digits, in this case, 2 sig digs.
There is no value in saying it represents 18.33199 m2 (7 sig digs) because the precision of the initial
measurements were limited. This means we can only claim our calculated answer is as accurate as the least
accurate measurement, & we do this by reporting our final answer to the same number of significant digits as our
least precise measurement. The objective, when dealing with the quantitative aspects of science, is to be as
accurate as possible within the limitations of our data. Therefore, it is extremely important to use significant
digits correctly.
Significant Digits & Logarithms (pH Calculations):
When calculating the pH of a concentration (or vice versa), the number of significant digits in the concentration should equal
the mantissa (the number of significant digits to the right of the decimal place) in the pH value. Note the underlined values in the
following examples:
pH = -log (1.234 x 10-12) = 11.90868484
 11.9087
4 significant digits in the value 1.234 x 10-12 means 4
significant digits in the mantissa of the pH
pH = -log (0.0012) = 2.900818754
 2.90
2 significant digits in the value 0.0012 means 2 significant
digits in the mantissa of the pH
[H+] = 10-pH = 10-1.951 = 0.011194378 mol/L
 0.0112 mol/L
3 significant digits in the mantissa of the pH means 3
significant digits in the concentration
[H+] = 10-pH = 10-10.32 = 4.786300923 x 10-11 mol/L
 4.8 x 10-11 mol/L
2 significant digits in the mantissa of the pH means 2
significant digits in the concentration
The whole number in a pH value is not included in the significant digit count because it indicates the location of the decimal
point. Note the following example:
[H+] = 10-pH
= 10-9.87
= 10-0.87 x 10-9
= 0.134896288 x 10-9 mol/L
= 1.34896288 x 10-10 mol/L
 1.3 x 10-10 mol/L
11
PROBLEM SOLVING FORMAT
Chemistry is a cumulative subject, & much of the quantitative analysis in chemistry requires cumulative skills. Many
chemistry problems require a series of calculations & conversions in the solution.
Use of a logical, consistent problem solving strategy is important for many reasons:
 it helps with analysis of the problem & therefore directs the solution to the problem
 it ensures clarity in the communication of the solution
 it provides a tool by which errors can be isolated & corrected
Therefore, in solving quantitative problems encountered in this course, the process is more significant than the actual
numerical value in the solution. All problems, even simple ones, provide an opportunity to practice using a logical & consistent
strategy. With few exceptions, solutions to problems must include the following:
1.
Given Information
 this constitutes the initial analysis of the problem
 all quantitative information for the problem is listed neatly, with symbols & units (all numerical values must be
accompanied by a unit)
 any unknown quantities that are asked for or implied in the problem are also listed
2.
Formula or Clearly Defined Conversion Factors
 this shows the strategy chosen in order to solve the problem
3.
Substitution & Answer
 this is the actual process of manipulating the data to arrive at a final value
 communication is important here, so steps must be clearly indicated & units must be included
 the final answer must reflect the correct number of significant digits (usually 2 to 4 significant digits)
4.
Statement
 this answers the question posed in the problem if the answer is not the numerical value from the calculation
 every question must end with a concluding statement, even if it consists of a brief sentence
Example:
A ball is dropped from a height of 2.0 m. It rebounds one-half the distance on each bounce. What is the height the ball will
reach after the 2nd bounce?
Solution:
Initial height = yo = 2.0 m
b=½
n=2
y=?
y = yo(b)n
y = 2.0m(½)2
y = 0.50 m
 the ball would reach 0.50 m after the 2nd bounce.
12
WRITING STUDY NOTES
Writing study notes should not be difficult nor a big task. Study notes consist of important points that should be remembered.
This usually includes worked out examples/questions. Although everybody’s set of study notes will be different, there are some
points you should follow when creating study notes:






Keep it short (2 – 5 pages maximum…the better you can reword & condense, the more you remember)
Include worked out solutions to 2 or 3 important problem types (include call-outs to important parts where you can go
wrong)…again, the fewer the number of worked out problems, the more you realize that there isn’t much material to
learn/remember
Keep it organized (use headings, tables & strategic use of colour to help organize your information)
Space is important! A wall of text is difficult to study from!
Handwrite your study notes (studies have shown that the more times you write something, the better you assimilate the
information…also, handwriting allows you to write just about anywhere on a page without worrying about formatting
problems that typing gives)
Keep it relatively neat & legible (slight messiness is acceptable, but remember that the easier it is to read, the more
useful your study notes become at a future time, such as before an exam!) 
What NOT to do when writing study notes:
 Do NOT recopy your class notes…reword & use different examples!...design as if you were teaching someone else
How to Grade Study Notes:
Although everyone’s study notes will be different, all good study notes have similar qualities:
Level 1
Content
(x 1)
Organization, Colour
& Neatness
(x 1)
Length
(x 0.5)
Level 2
Work is incomplete
Work is recopied
from class notes
Sequence of
information is difficult
to follow; work is
illegible, no colour
Difficult to follow
work because of
organization,
indiscriminate use of
colour/no colour
Less than 2 pages or
more than 5 pages
Total
Level 3
Work is reworded,
using different
examples, without
detail (no callouts)
Work presents
information in a
relatively neat &
logical sequence,
used colour
Level 4
Work is reworded,
using different
examples, with full
detail (with callouts)
Information in logical
sequence which is
neat & can be easily
followed, strategic
use of colour
Points
Between 2 & 5
pages
/10
13
MATTER, CHEMICAL TRENDS & CHEMICAL BONDING
ATOMIC NOTATION
1.
Define the following terms:
a. atomic number
2.
For each of the 3 subatomic particles, answer the following:
a. Where is it located in the atom?
b. What is its mass?
c. What is the relative mass of particle (in comparison to the smallest subatomic particle)?
3.
Identify the elements which have the following numbers of protons in the nuclei of their atoms:
a. 7
b. 29
c. 47
b.
mass number
c.
isotope
4.
How many protons are in the atom of each of the following elements?
a. Boron (B)
d. Calcium (Ca)
b. Lead (Pb)
e. Silicon (Si)
c. Uranium (U)
f. Iron (Fe)
5.
Write the atomic notation for a neutral atom with the following:
a. 8 protons, 9 neutrons
c. 1 proton, 1 neutron
b. 6 protons, 8 neutrons
d. 45 neutrons, 35 electrons
6.
7.
d.
g.
h.
i.
Oxygen (O)
Helium (He)
Zinc (Zn)
e.
12 neutrons, 12 electrons
Determine the number of protons, electrons & neutrons for the following:
a. 25Mg
c. 80Br
e. 35P 3
37
d. 201Hg
f.
Cl 
b. 56Co
Name the neutral element & give the atomic notation for each when:
a. Z = 7 & A = 14
b. Z = 10 & A = 20
c. Z = 17 & A = 37
90
g.
h.
d.
3
H
55
Mn 7 
Z = 26 & A = 56
ISOTOPE PROBLEMS
1.
Natural potassium consists of 93.1% K-39 & 6.90% K-41. What is the average atomic mass of natural potassium? (39.1 u)
2.
Thallium consists of 29.50% Tl-203 & 70.50% Tl-205. Calculate the average atomic mass of thallium. (204.4 u)
3.
Neon consists of 90.92% Ne-20, 0.2600% Ne-21 & 8.820% Ne-22. What is the average atomic mass of neon? (20.18 u)
4.
Erbium consists of 33.41% Er-166, 22.94% Er-167, 27.07% Er-168 & 16.58% Er-170. What is the average atomic mass of
erbium?
(167.4 u)
5.
Mercury has 26 isotopes, but the 7 important ones with their corresponding percentage abundances are:
Hg-196
0.1400%
Hg-200
23.13%
Hg-204
Hg-198
10.02%
Hg-201
13.22%
Hg-199
16.84%
Hg-202
29.80%
What is the average atomic mass of naturally occurring mercury?
(200.6 u)
6.850%
6.
Gallium has two isotopes. If the two isotopes are: Ga-69 (68.95 u) & Ga-71 (70.95 u), find the isotopic abundances of each
isotope.
(61.50% & 38.50%)
7.
Natural argon has three isotopes: Ar-36 (35.967 u), Ar-38 (37.962 u) & Ar-40 (39.962 u). If the isotopic abundance of Ar-36
is 0.006000%, find the isotopic abundances of the other two isotopes.
(0.588% & 99.406%)
14
VARIATION IN ATOMIC PROPERTIES
1.
Define effective nuclear charge.
2.
Explain the trend for the effective nuclear charge across a given period.
3.
Discuss the effective nuclear charge for the elements carbon, silicon & germanium.
4.
For each pair of elements indicate & then explain which element has the greater atomic size:
a. Na or K
b. Na or Mg
c.
O or S
For each pair of elements, explain which element has the greater ionization potential energy:
a. Na or K
b. Na or Mg
c.
O or S
6.
For each pair of elements, explain which element has the greater electron affinity:
a. O or S
b. O or F
c.
F or Ne
7.
For each pair, indicate which element would be the more reactive. Explain in terms of either the ionization energies or
electron affinities of the elements:
a. Na or K
b. Na or Mg
c. O or F
8.
Explain what happens when you go down a group in the periodic table to:
a. The atomic size

The ionization energy
5.

The electron affinity
ELECTRONEGATIVITY & TYPES OF INTRAMOLECULAR BONDING
Determine whether the bond that would occur between the following pairs of elements is ionic or covalent. For each polar
covalent bond, label the atom that is slightly positive & slightly negative with the symbols + & -, respectively.
1. H & I
4. H & F
7. Mg & O
10. Mg & H
2. Li & Br
5. C & S
8. Ca & S
11. Si & B
3. P & H
9. K & F
6. Al & I
IONIC BONDING
1.
Each of the following pair of elements forms an ionic compound. For each pair, draw a Lewis diagram to show how each
are bonded to each other & write the ionic formula:
c. Ga & F
e. Na & O
a. Li & Cl
f. Ca & F
d. Al & F
b. K & S
2.
Ionic compounds always have relatively high melting points. What information does this give you concerning the strength of
the ionic bond? Explain.
3.
Oxygen & chlorine do not react by producing an ionic bond. Why not?
4.
Potassium metal (composed of atoms) is very reactive, but the potassium ion is stable. Briefly explain this difference in
properties.
5.
Element “X” reacts with sodium to form a compound with the formula Na2X. What group in the periodic table must “X”
belong to? Explain.
15
COVALENT BONDING
1.
Draw the structural diagram for each compound. (Note: the central atom is typically the first element listed in a compound,
otherwise it is underscored).
e. N2
i. F2
m. NF3
q. C2Br4
a. NCl3
j. SiBr4
n. CO
f. CCl4
b. CO2
k. C2H4F2
o. O2
c. C2H4
g. OF2
l. HBr
p. PCl3
d. HOBr
h. SH2
VSEPR STRUCTURES & POLARITY
1.
For each compound:
i. Calculate the ΔEN
ii. Draw the Structural diagram
iii. Draw the 3D VSEPR shape
iv. Is it polar or non-polar?
a. SiCl4
b. PCl3
c. NF3
d. AlCl3
e. CaF2
f. H2Se
g. SO3
h. CS2
i. * CH2Br2
j. * HCN
* simplified rule does not apply!)
2.
What two factors affect whether a molecule is going to be polar or non-polar overall?
16
CHEMICAL REACTIONS UNIT
CHEMICAL NOMENCLATURE OF IONIC BINARY COMPOUNDS
1.
Write the chemical formulas for the following:
a. calcium iodide
f. potassium bromide
b. sodium fluoride
g. zinc fluoride
c. aluminum bromide
h. barium bromide
d. calcium hydride
i. zinc hydride
e. calcium oxide
j. aluminum nitride
2.
Name the following:
a. AlF3
b. MgS
c. CaF2
d. Na2S
e.
f.
g.
h.
CaO
Li2O
KCl
K2O
i.
j.
k.
l.
k.
l.
m.
n.
o.
silver sulphide
barium oxide
zinc sulphide
magnesium chloride
magnesium carbide
m.
n.
o.
p.
MgCl2
BaI2
KI
ZnH2
p.
q.
r.
Ca3N2
AgBr
SrO
BeF2
lithium sulphide
silver phosphide
potassium oxide
q.
r.
AlN
Li4C
CHEMICAL NOMENCLATURE OF MULTIVALENT COMPOUNDS
1.
Write formulas for the following
a. iron (II) oxide
b. copper (I) sulphide
c. tin (II) fluoride
d.
e.
f.
copper (II) oxide
iron (II) nitride
lead (IV) fluoride
g.
h.
i.
tin (IV) iodide
iron (III) oxide
lead (II) chloride
2.
Write formulas for the following:
a. iron (III) sulphide
b. manganese (IV) oxide
c. platinum (II) nitride
d. mercury (I) sulphide
e. antimony (V) oxide
f. phosphorus (III) fluoride
g. platinum (IV) sulphide
h.
i.
j.
k.
l.
m.
n.
cobalt (III) phosphide
gold (III) oxide
copper (I) sulphide
cobalt (II) carbide
gold (I) chloride
lead (IV) sulphide
tin (II) bromide
o.
p.
q.
r.
s.
t.
iron (II) oxide
nickel (III) sulphide
phosphorus (III) oxide
chromium (II) hydride
chromium (II) phosphide
bismuth (V) oxide
3.
Name the following using IUPAC notation:
f. CoS
a. FeCl2
g. HgCl
b. SnO2
c. PI3
h. AuI
d. CuBr
i. Sb2S5
j. As2O3
e. FeCl3
k.
l.
m.
n.
o.
Hg2O
Pb3N4
Sb2O3
CuF
SnF2
p.
q.
r.
s.
t.
PbO2
Cu2O
AuN
SbCl3
PbCl2
u.
SnS2
CHEMICAL NOMENCLATURE OF SIMPLE POLYATOMIC COMPOUNDS
1.
Write formulas for the following polyatomic salts:
a. sodium chromate
h.
b. sodium acetate
i.
c. calcium sulphate
j.
d. potassium sulphate
k.
e. barium phosphate
l.
f. calcium dichromate
m.
g. lithium nitrate
n.
2.
Name the following polyatomic salts:
a. Ba(BrO3)2
f.
b. Ca(ClO3)2
g.
c. BaSO4
h.
d. Mg(BrO3)2
i.
e. Ca3(PO4)2
j.
copper (I) sulphate
tin (II) carbonate
mercury (II) hydroxide
magnesium bromate
ammonium carbonate
silver nitrate
lead (II) sulphate
k.
l.
m.
n.
o.
NaClO3
Al2(SO4)3
Na2SO4
PbSO4
Fe3(PO4)2
17
o.
p.
q.
r.
s.
t.
Zn(CH3COO)2
MgCrO4
Hg(ClO3)2
Ba3(PO4)2
Al(NO3)3
barium nitrate
magnesium acetate
silver carbonate
aluminum cyanide
iron (III) hydroxide
copper (II) manganate
p.
q.
r.
Cu(NO3)2
Al2(Cr2O7)3
Fe(MnO3)3
CHEMICAL NOMENCLATURE OF MORE POLYATOMIC COMPOUNDS
Write formulas for the following polyatomic compounds:
1. sodium phosphite
15. zinc carbonite
2. mercury (II) nitrite
16. magnesium percarbonate
3. iron (II) perchlorate
17. arsenic (V) nitrite
4. potassium perbromate
18. iron (II) hypochlorite
5. zinc sulphate
19. calcium hyposulphite
6. aluminum hypophosphite
20. potassium permanganate
7. copper (II) chlorite
21. antimony (III) chlorite
8. iron (III) hyponitrite
22. ammonium sulphite
9. calcium hypochlorite
23. silver bromite
10. tin (IV) chlorite
24. zinc chlorite
11. lead (II) perchlorate
25. tin (II) perchlorate
12. copper (II) hypobromite
26. gold (I) nitrite
13. antimony (V) sulphite
27. sodium thiosulphate
14. manganese (V) hypophosphite
28. copper (I) bicarbonate
Name the following polyatomic compounds:
41. Hg(NO)2
51. Fe3(PO3)2
52. MgCrO4
42. AlPO3
53. NaHSO4
43. Ba(ClO2)2
54. Zn(ClO)2
44. FeSO3
55. SbPO3
45. Sn(SO4)2
56. Cu(ClO)2
46. As(NO3)3
57. KBrO4
47. Hg(ClO)2
58. (NH4)2SO4
48. LiNO2
59. K3PO2
49. CuSO4
60. NaClO4
50. (NH4)2CO4
61.
62.
63.
64.
65.
66.
67.
68.
69.
70.
71.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
As(NO4)5
Sb2(SO3)3
Zn3(PO2)2
Ag2S2O3
Mg(BrO4)2
CaHPO4
MnSO3
Pb(IO2)2
AsPO5
Cu(BrO4)2
Fe2(SO3)3
mercury (II) perchlorate
manganese (II) sulphite
zinc hydrogen phosphate
potassium perbromate
barium hypomanganite
magnesium hypobromite
lead (IV) phosphite
mercury (II) perbromate
calcium sulphite
potassium perchlorate
ammonium thiosulphate
beryllium phosphate
72.
73.
74.
75.
76.
77.
78.
79.
80.
CuHPO4
KClO4
(NH4)3PO3
AgNO2
Zn3(PO3)2
Fe(MnO4)2
LiNO
SnS2O3
KHCO3
CHEMICAL NOMENCLATURE OF ACIDIC & BASIC COMPOUNDS
1.
2.
Write formulas for the following compounds:
a. perchloric acid
b. hyponitrous acid
c. iron (III) hydroxide
d. bromic acid
e. hydrophosphoric acid
f. iodic acid
g. sulphuric acid
h.
i.
j.
k.
l.
m.
n.
Name the following compounds:
e. HNO2 (aq)
a. HCl (aq)
f. H2S (aq)
b. HIO3 (aq)
g. HBrO (aq)
c. Cu(OH)2 (aq)
h. H2CO3 (aq)
d. H2SO3 (aq)
carbonous acid
hydroiodic acid
magnesium hydroxide
chromic acid
sodium hydroxide
chlorous acid
lithium hydroxide
i.
j.
k.
l.
Pb(OH)4(aq)
H3PO2 (aq)
HNO4 (aq)
HClO3 (aq)
o.
p.
q.
r.
s.
t.
m.
n.
o.
p.
pernitric acid
manganous acid
carbonic acid
phosphorous acid
hypobromous acid
chloric acid
HCH3COO (aq)
H2SO5 (aq)
HgOH (aq)
HI (aq)
q.
r.
HMnO3 (aq)
Sb(OH)5 (aq)
CHEMICAL NOMENCLATURE OF COVALENT BINARY COMPOUNDS
1.
Write the chemical formulas for the following:
a. hydrogen iodide
b. nitrogen trifluoride
c. carbon tetrabromide
d. phosphorus pentahydride
e.
f.
g.
h.
nitrogen dioxide
carbon dioxide
sulphur difluoride
diphosphorus pentoxide
18
i.
j.
k.
l.
dihydrogen sulphide
nitrogen trihydride
oxygen dichloride
carbon monoxide
m. dihydrogen monoxide
n. silicon dioxide
2.
Name the following:
a. SiF4
b. NCl3
c. H2S
d. SCl4
o.
p.
e.
f.
g.
h.
NO
CO2
N2O3
SCl2
dinitrogen trisulphide
tetraboron tricarbide
i.
j.
k.
l.
q.
r.
N2O5
NBr3
HBr
PCl3
m.
n.
o.
p.
hydrogen peroxide
carbon tetrahydride
HI
NH3
HBr
H3P
q.
r.
SF2
SiH4
CHEMICAL NOMENCLATURE REVIEW
Write formulas for the following compounds:
1. sodium iodide
2. iron (III) oxide
3. sulphur dichloride
4. potassium hypophosphite
5. hydrophosphoric acid
6. lithium persulphate
7. ammonium nitrate
8. mercury (I) sulphate
9. tin (II) carbonite
10. antimony (V) hydroxide
11. dinitrogen oxide
12. ammonium carbonate
13. cobalt (II) nitrate
14. strontium persulphate
15. sodium peroxide
16. nickel (II) acetate
17. copper (II) sulphate
pentahydrate
18. tin (II) hypocarbonite
19. hydrocyanic acid
20. iron (II) hydroxide
21. potassium permanganate
22. sodium acetate trihydrate
23. copper (II) sulphide
24. gold (I) fluoride
25. aluminum nitrite
26. lead (IV) acetate
27. sodium sulphate decahydrate
28. copper (I) hypophosphite
29. silver thiosulphate
30. carbonic acid
31. lithium perphosphate
32. ammonium hydroxide
Name the following compounds:
50. NaBrO3
58. Sb2(SO4)3
59. H2S (aq)
51. Sb(ClO)5
60. HCH3COO (aq)
52. H2SO4 (aq)
61. Cl2O
53. Mg(OH)2
54. SO3
62. MgCr2O7
55. H2O2
63. Ba(NO3)2
56. LiCl
64. AlN
57. H3PO2 (aq)
65. P2O3
66.
67.
68.
69.
70.
71.
72.
73.
GaH3
N2O5
H2S
HNO2 (aq)
HCl (aq)
Pb(CO2)2
Sb2O3
Zn(IO3)2
74.
75.
76.
77.
78.
79.
80.
81.
33.
34.
35.
36.
37.
38.
39.
40.
41.
42.
43.
44.
45.
46.
47.
48.
49.
Fe(NO4)2
CuCO3
ICl
LiClO
CaHPO4
SO2
NH4OH
BF3
phosphorus pentabromide
barium hypocarbonite
nitrous acid
sulphur difluoride
dinitrogen difluoride
bismuth (III) iodate
periodic acid
magnesium oxide
calcium bicarbonate
phosphoric acid
lithium peroxide
dinitrogen tetroxide
chromium (VI) chlorite
silicon dioxide
iron (III) iodide
platinum (II) carbide
beryllium fluoride
82.
83.
84.
85.
86.
87.
88.
OCl2
SiO2
K2O2
Zn2C
NaIO
MgSO4
Pb(MnO3)4
BALANCING SKELETON EQUATIONS
1.
Balancing chemical equations satisfies one of the most important chemical laws. What is the name of the law?
2.
Balance the following skeleton equations:
a. Mg + O2  MgO
b. N2 + H2  NH3
c. P4 + O2  P2O3
d. HgO  Hg + O2
e. Fe + H2O  Fe2O3 + H2
f. CH4 + O2  CO2 + H2O
g. Na + HCl  NaCl + H2
h. Al2O3  Al + O2
i. Mg + N2  Mg3N2
j. C3H8 + O2  CO2 + H2O
k. KClO3  KCl + O2
l.
m.
n.
o.
p.
q.
r.
s.
t.
u.
v.
19
MnO2 + HCl  MnCl2 + Cl2 + H2O
H3PO4 + NaOH  Na3PO4 + H2O
P4 + Cl2  PCl5
NO2 + H2O  HNO3 + NO
Na3PO4 + CaCl2  Ca3(PO4)2 + NaCl
C6H14 + O2  CO2 + H2O
Cu + HNO3  Cu(NO3)2 + NO2 + H2O
NaHSO3 + H2SO4  Na2SO4 + SO2 + H2O
C6H6 + O2  CO2 + H2O
FeCl3 + Na2S  Fe2S3 + NaCl
NO + H2  N2 + H2O
w. Fe2O3 + P4  Fe + P2O5
BALANCING WORD EQUATIONS
Balance the following word equations:
1. zinc + hydrochloric acid  zinc chloride + hydrogen
2. aluminum chloride + potassium  potassium chloride + aluminum
3. sodium hydroxide + phosphoric acid  sodium phosphate + water
4. iron (III) chloride + ammonium sulphide  iron (III) sulphide + ammonium chloride
5. zinc sulphide + oxygen  zinc oxide + sulphur dioxide
6. dicarbon dihydride (acetylene) + oxygen  carbon dioxide + water
7. lead (II) nitrate  lead (II) oxide + oxygen + nitrogen dioxide
8. nitric acid  water + nitrogen dioxide + oxygen
9. manganese dioxide + hydrochloric acid  manganese (II) chloride + water + chlorine
10. carbon + nitric acid  nitrogen dioxide + water + carbon dioxide
11. tin + sulphuric acid  tin (IV) sulphate + sulphur dioxide + water
12. sodium bromide + sulphuric acid  sodium sulphate + sulphur dioxide + bromine + water
13. calcium + water  calcium hydroxide + hydrogen
14. aluminum oxide + carbon + chlorine  aluminum chloride + carbon monoxide
15. aluminum hydroxide + sulphuric acid  aluminum sulphate + water
IDENTIFYING TYPES OF REACTIONS
For each of the following reactions, identify the type of reaction & balance the equation.
1. water  hydrogen + oxygen
2. lithium + water  lithium hydroxide + hydrogen
3. sulphuric acid + potassium hydroxide  water + potassium sulphate
4. copper + silver nitrate  silver + copper (II) nitrate
5. aluminum + oxygen  aluminum oxide
6. methane (CH4) + oxygen  carbon dioxide + water
7. hydrochloric acid + sodium hydroxide  water + sodium chloride
8. sodium carbonate + copper (II) sulphate  sodium sulphate + copper (II) carbonate
9. zinc + phosphoric acid  zinc phosphate + hydrogen
10. water + carbon dioxide  carbonic acid
11. calcium + oxygen  calcium oxide
12. butane (C4H10) + oxygen  carbon dioxide + water
COMBUSTION, SYNTHESIS & DECOMPOSITION REACTIONS
1.
Write out the balanced chemical equation for the following combustion reactions (include phase labels):
a. C + O2 
d. Zn + O2 
g. C2H6 + O2 
b. Mg + O2 
e. Li + O2 
h. C3H6 + O2 
c. S8 + O2 
f. CH4 + O2 
i. C6H6O6 + O2 
2.
Write the reaction when each of the products from question 1 is placed in water. Is the solution acidic, basic or neutral?
3.
Write out the balanced chemical equation for the following synthesis reactions (include phase labels):
a. C + H2 
c. N2 + H2 
e. Li + N2 
b. Mg + Br2 
d. Zn + S8 
f. H2 + P4 
4.
Write out the balanced chemical equation for the following decomposition reactions (include phase labels):
a. NH3 
b. N2H4 
c. CO2 
d.
20
H2O 
SINGLE DISPLACEMENT REACTIONS
Refer to the metal & halogen activity series to determine the products (if any) for the following reactants. If a reaction does not
occur, write NR. If a metal does react, assume the more stable ionic charge will form (refer to the periodic table). Take careful
note of the states of each compound.
1. Li (s) + H2O (l) 
8. Na (s) + H2O (l) 
16. Pb (s) + H2O (l) 
2. Al (s) + FeO (s) 
9. Zn (s) + PbI4 (s) 
17. Ni (s) + HCl (aq) 
3. F2(g) + HCl(aq) 
10. Pt (s) + H2SO4 (aq) 
18. Al (s) + H3PO4 (aq) 
11. Ni (s) + MgCO3 (s) 
4. Cu (s) + H2O (l) 
19. Ca (s) + H2O (l) 
12. Ba (s) + H2O (l) 
5. Mg (s) + CaCO3 (s) 
20. Zn (s) + Pb(NO3)2 (aq) 
13. Sn (s) + HgS (s) 
6. Ag (s) + HCl (aq) 
21. Ag (s) + CuSO4 (aq) 
14. Al (s) + HNO3 (aq) 
7. Fe (s) + HClO3 (aq) 
15. Ca (s) + FeCl2 (aq) 
DOUBLE DISPLACEMENT REACTIONS
Complete & balance the following equations.
1. NaCl (aq) + AgNO3 (aq) 
5.
6.
2. MgCl2 (aq) + KOH (aq) 
7.
3. BaCl2 (aq) + Na2SO4 (aq) 
8.
4. Ca(OH)2 (aq) + HCl (aq) 
Identify the precipitate using subscripts (refer to a solubility table).
NaOH (aq) + Fe(NO3)2 (aq) 
9. CuCl2 (aq) + Na2S (aq) 
Na2CO3 (aq) + MgSO4 (aq) 
10. Pb(NO3)2 (aq) + K2CO3 (aq) 
Pb(NO3)2 (aq) + H2S (aq) 
11. CaCl2 (aq) + H2S (aq) 
Fe2(SO4)3 (aq) + NaOH (aq) 
12. Al2(SO4)3 (aq) + NH4OH (aq) 
NET IONIC EQUATIONS
1.
2.
For each of the following,
 Complete & balance the chemical equation
 Write the ionic & net ionic equation
a. barium hydroxide & sodium sulphate
b. ammonium carbonate & lead (II) nitrate
c. silver nitrate & calcium chloride
d. potassium hydroxide & zinc sulphate
e. tin (II) acetate & barium hydroxide
f.
g.
h.
i.
j.
strontium nitrate & potassium sulphate
aluminum metal & silver nitrate
lead (II) nitrate & potassium chloride
barium hydroxide & sulphuric acid
chlorine gas & sodium iodide
What is the net ionic equation for any acid-base neutralization reaction?
QUANTITIES IN CHEMICAL REACTIONS UNIT
MOLAR MASS PROBLEMS
1.
What does it mean for carbon to have a molar mass of 12.01 g/mol?
2.
Find the molar masses of each of the following:
a. water
(18.02 g/mol)
b. carbon dioxide
(44.01 g/mol)
c. sodium bicarbonate
(84.01 g/mol)
d. lithium hydroxide
(23.95 g/mol)
e.
f.
g.
21
calcium perphosphate
silver nitrate
zinc chloride
(342.18 g/mol)
(169.88 g/mol)
(136.29 g/mol)
AVOGADRO’S NUMBER PROBLEMS
1.
How many objects are present in one mole?
2.
How many atoms can be found in:
a. 3.00 mol of lead
(1.81 x 1024)
b. 6.13 mol of silver
(3.69 x 1024)
c.
d.
0.00685 mol of argon
0.0320 mol of sodium
3.
How many molecules can be found in:
a. 0.0700 mol of carbon dioxide (4.21 x 1022)
b. 13.2 mol of ammonia
(7.95 x 1024)
c.
d.
4.56 mol of oxygen gas
(2.75 x 1024)
31.1 mol of sodium chloride (1.87 x 1025)
4.
How many moles are represented by:
a. 7.5 x 1023 atoms of gold
(1.2)
b. 8.24 x 1025 atoms of potassium (137)
c.
d.
1.45 x 1022 molecules of sucrose (0.0241)
1.73 x 1025 molecules of water (28.7)
5.
Sodium bromide, NaBr, dissociates into sodium ions & bromide ions. If 0.0125 mol of sodium bromide dissociated:
a. How many sodium ions were formed?
(7.53 x 1021)
b. How many bromide ions were formed?
(7.53 x 1021)
c. How many ions in total resulted from the dissociation?
(1.51 x 1022)
6.
Magnesium iodide, MgI2, dissociates into one magnesium ion for every two iodide ions. If 0.0350 mol of magnesium iodide
dissociated:
a. How many magnesium ions were formed?
(2.11 x 1022)
b. How many iodide ions were formed?
(4.21 x 1022)
c. How many ions in total resulted from the dissociation?
(6.32 x 1022)
7.
One molecule of methane consists of one carbon atom & four hydrogen atoms. How many atoms in total are in 0.00725
mol of methane?
(2.18 x 1022)
(4.12 x 1021)
(1.93 x 1022)
MOLE RELATED PROBLEMS (# OF PARTICLES)
1.
How many molecules are there in 1.25 mol of sulphuric acid?
(7.53 x 1023 molecules)
2.
How many helium atoms are there in 2.64 mol of helium gas?
(1.59 x 1024 atoms)
3.
How many moles of oxygen molecules are in 1.23 x 1024 diphosphorus pentoxide molecules? (5.11 moles)
4.
How many moles of magnesium hydroxide can be created using 2.23 x 1024 oxygen atoms? (1.85 moles)
5.
How many carbon atoms are in 1.25 mol of silver acetate?
(1.51 x 1024 atoms)
MOLE RELATED PROBLEMS (# OF MOLES)
1.
Which has a greater mass, one mole of lead atoms or ten moles of water molecules?
(Pb)
2.
What is the mass of 1.31 mol of silver nitrate?
(223 g)
3.
What is the mass of 2.44 mol of hydrogen peroxide?
(83.0 g)
4.
How many moles of molecules are in 53 g of sodium carbonate?
(0.50 mol)
5.
An unknown element is studied. If 8.65 mol of the element has a mass of 211 g, what is the element? (24.4 g/mol)
6.
A binary compound is studied. It is found that 1.75 mol of the compound has a mass of 29.8 g. If the compound contains
one nitrogen atom, what is the compound?
(17.0 g/mol)
7.
What mass of carbon are in 1.25 mol of silver acetate?
(30.0 g)
8.
What mass of chlorine gas is in 5.34 mol of iron (III) chloride?
(568 g)
22
MOLE RELATED PROBLEMS (ADVANCED CALCULATIONS)
1.
How many molecules are in 11.3 g of glucose, C6H12O6?
(3.78 x 1022 molecules)
2.
How many molecules are in 12.5 g of aluminum chloride?
(5.64 x 1022 molecules)
3.
How many atoms of carbon are in 17.1 g of sucrose, C12H22O11?
(3.61 x 1023 atoms)
4.
How many fluorine atoms can be made from 21.3 g of boron trifluoride?
(5.67 x 1023 atoms)
5.
How many molecules of oxygen can be made from 9.52 g of nickel (III) oxide?
(5.20 x 1022 molecules)
6.
What mass of carbon tetrahydride can be made from 2.66 x 1025 molecules of carbon tetrahydride? (709 g)
7.
What mass of nitric acid can be made from 1.23 x 1024 oxygen atoms?
(42.9 g)
8.
What mass of oxygen gas can be made from a total of 5.84 x 1024 oxygen atoms?
(155 g)
% COMPOSITION PROBLEMS (FROM CHEMICAL FORMULAS)
1.
Find the percentage composition of the indicated element in each of the following compounds:
a. Na in NaCl
(39.34%)
b. H in HBr
(1.248%)
c. O in SO2
(49.95%)
d. N in NH3
(82.22%)
e. O in Ca(NO3)2
(58.50%)
f. P in zinc phosphate
(16.04%)
g. Na in sodium oxide
(74.19%)
2.
Determine the % composition of oxygen & water present in Epsom salts, MgSO4 · 7 H2O.
3.
An ore of aluminum called bauxite has the formula Al2O3 · 2 H2O. What mass (in kg) of bauxite would be needed for the
extraction of one kilogram of aluminum?
(2.56 kg)
4.
The bones of an average adult person has a mass of approximately 11 kg & contains approximately 48% calcium
phosphate. Calculate the mass (in kg) of phosphorus present in the skeleton of an average adult. (1.1 kg)
(%O = 71.39%)
% COMPOSITION PROBLEMS (FROM EXPERIMENT)
1.
Students heated a sample of copper in excess oxygen & made the following observations:
Mass of crucible & lid
20.10 g
Mass of crucible, lid & sample
21.04 g
Mass of crucible, lid & product
21.20 g
Calculate the % composition (by weight) of copper oxide from this information.
(%Cu = 85.45%)
2.
10.0 g of Ca was placed in a large crucible & heated. The empty crucible weighed 186 g. After the Ca was heated, the
crucible & contents weighed 2.00 x 102 g. Calculate the % composition of CaO.
(%Ca = 71.4%)
3.
An oxide of mercury decomposes (into its constituent elements) when heated above 500°C. A 50.0 g sample of the oxide
was heated & 46.3 g of mercury was obtained. Calculate the % composition of the mercury oxide.
(%Hg = 92.6%)
4.
A student decomposes a sample of potassium bromide of mass 7.14 g. She finds that there are 2.35 g of potassium & the
rest is bromine. Use this information to calculate the % composition of potassium bromide.
(%K = 32.9%)
5.
A certain compound is known to be 32.8% lead. If 151 g of this compound are decomposed, what mass of lead will be
recovered?
(49.5 g)
6.
A compound is known to be 25.6% arsenic. If 23.8 g of arsenic is recovered, what was the mass of the original sample?
(93.0 g)
23
SIMPLEST FORMULA PROBLEMS
1.
Determine the empirical formula of a compound made up of 31.14% sulphur & 68.86% chlorine. (SCl2)
2.
What is the empirical formula of a compound that is 15.9% boron & 84.1% fluorine?
3.
An inorganic salt is composed of 17.6% sodium 39.7% chromium & the remainder oxygen. What is the empirical formula of
this salt?
(Na2Cr2O7)
4.
Compound X contains 69.90% carbon, 6.85% hydrogen & the remainder oxygen. Determine the empirical formula of
compound X.
(C12H14O3)
5.
A chemical compound is made up of 13.26 g chromium, 12.26 g sulphur & 24.48 g oxygen. What is the simplest formula of
the compound?
(Cr2(SO4)3)
6.
A 50.5 g sample of mercury oxide was heated & 48.6 g of mercury was obtained. Find the simplest formula of the mercury
oxide.
(Hg2O)
7.
Students heated a sample of iron in excess oxygen & made the following observations:
Mass of crucible & lid
19.76 g
Mass of crucible, lid & iron
22.05 g
Mass of crucible, lid & iron oxide
22.71 g
Find the simplest formula of iron oxide from this information.
(BF3)
(FeO)
MOLECULAR FORMULA PROBLEMS
1.
The simplest/empirical formula of butane, the fuel used in disposable lighters, is C2H5. In an experiment, the molar mass of
butane was determined to be 58 g/mol. What is the molecular formula of butane?
(C4H10)
2.
A compound has the simplest/empirical formula CHO. Its molecular mass is 116 g/mol. What is its molecular formula?
(C4H4O4)
3.
A compound having a molecular mass of 90.0 g/mol is composed of 39.95% carbon, 6.69% hydrogen & the remainder is
oxygen. Calculate the molecular formula of this compound.
(C3H6O3)
4.
A compound contains only carbon & hydrogen. It contains 80.0% carbon. It is also known that 1.45 mol of this gas has a
mass of 43.5 g. Calculate the molecular formula of this gas.
(C2H6)
5.
Ether contains 64.86% carbon, 13.62% hydrogen & 21.52% oxygen. In addition it is known that 1.25 x 10 -2 mol of this gas
has a mass of 0.926 g. Calculate the molecular formula of this gas.
(C4H10O)
6.
A compound has a molar mass of 170.2 g/mol. Its percentage composition is 49.38% carbon, 3.55% hydrogen, 9.40%
oxygen & the remainder is sulphur. Calculate the molecular formula.
(C7H6OS2)
HYDRATE PROBLEMS
1.
A 2.78 g sample of hydrated FeSO4 was heated to remove all the water of hydration. The mass of the anhydrous FeSO4
was 1.52 g. Calculate the number of water molecules associated with each formula unit of FeSO4. (7)
2.
A 19.76 g sample of hydrated sodium acetate was heated & 11.91 g of anhydrous salt remained. Calculate the number of
water molecules associated with each formula unit of sodium acetate.
(3)
3.
Copper (II) sulphate pentahydrate is a well-known hydrate. If a sample of the hydrate was heated & 7.24 g of anhydrous
copper (II) sulphate was recovered, what was the mass of the original hydrated sample?
(11.3 g)
4.
A 5.00 g sample of borax, Na2B4O7 · 10 H2O, was thoroughly heated to remove all the water of hydration. What mass of
anhydrous compound remained?
(2.64 grams)
5.
(Tricky) A sample of cobalt (II) nitrate hydrate was heated to remove all the water of hydration. The hydrate was found to
be 65.96% oxygen. Calculate the number of water molecules associated with each formula unit of cobalt (II) nitrate. (6)
24
6.
(Tricky) Epsom salts is MgSO4  x H2O. The hydrate was found to contain 71.4% oxygen. Calculate the number of water
molecules associated with each formula unit of magnesium sulphate hydrate.
(7)
7.
(Tricky) Zinc nitrate hydrate, Zn(NO3)2  x H2O, contains 21.98% zinc by mass. What is the value of x? (6)
SIMPLE STOICHIOMETRY PROBLEMS (MOL-MOL)
1.
Butane burns with the oxygen in air to give carbon dioxide & water. How many moles of carbon dioxide would be produced
from 0.150 mol of butane?
(0.600 mol)
_____ C4H10 (g) + _____ O2 (g)  _____ CO2 (g) + _____ H2O (g)
2.
Ethanol burns with the oxygen in air to give carbon dioxide & water. How many moles of water would be produced from
0.250 mol of ethanol?
(0.750 mol)
_____ C2H5OH (l) + _____ O2 (g)  _____ CO2 (g) + _____ H2O (l)
3.
Nickel (II) chloride reacts with sodium phosphate to precipitate nickel (II) phosphate. How many moles of nickel (II) chloride
are needed to produce 0.479 mol of nickel (II) phosphate?
(1.437 mol)
_____ NiCl2 (aq) + _____ Na3PO4 (aq)  _____ Ni3(PO4)2 (s) + _____ NaCl (aq)
4.
Nitric acid, HNO3, is manufactured by the Ostwald process, in which nitrogen dioxide, NO2, reacts with water. How many
moles of nitrogen dioxide are required in this reaction to produce 0.750 mol of HNO3?
(1.13 mol)
_____ NO2 (g) + _____ H2O (l)  _____ HNO3 (aq) + _____ NO (g)
5.
White phosphorus, P4, is prepared by fusing calcium phosphate, Ca3(PO4)2, with carbon, C, & sand, SiO2, in an electric
furnace. How many moles of Ca3(PO4)2 are required to give 1.35 mol of phosphorus?
(2.70 mol)
_____ Ca3(PO4)2 (s) + _____ SiO2 (s) + _____ C (s)  _____ P4 (g) + _____ CaSiO3 (l) + _____ CO (g)
EXCESS REAGENT/REACTANT MASS-MASS STOICHIOMETRY PROBLEMS
1.
What two things you must have in order to solve any stoichiometry problem?
2.
What is an excess reagent?
3.
Calcium reacts with water to produce calcium hydroxide & hydrogen. What mass of hydrogen is produced by the reaction of
12.4 g of calcium with excess water?
(0.625 g)
4.
Manganese (IV) oxide reacts with hydrochloric acid to produce manganese (II) chloride, water & chlorine. If 48.8 g of
manganese (IV) oxide was used, what mass of water would be produced?
(20.2 g)
5.
Iron (II) oxide reacts with sulphur, S8, to produce iron metal & sulphur dioxide. If 193 g of iron (II) oxide was used, what
mass of sulphur dioxide will be produced?
(86.1 g)
6.
Sodium peroxide reacts with water to produce sodium hydroxide & oxygen. What mass of oxygen can be prepared by
reacting 18.0 g of sodium peroxide with water?
(3.69 g)
7.
What mass of mercury (II) oxide must be decomposed to produce 2.56 x 1023 molecules of oxygen? (184 g)
8.
What mass of methane, CH4, must be burned to produce 12.3 g of water?
9.
Sodium oxide & water produce sodium hydroxide in a synthesis reaction. A student had 31.2 g of sodium oxide. What
mass of water should be reacted with this so that no water & no sodium oxide will be left over? (9.07 g)
(5.48 g)
10. Titanium metal is made by reacting titanium (IV) chloride with magnesium metal. Magnesium chloride is also produced.
How many kilograms of magnesium are required to produce 1.00 kg of titanium?
(1.02 kg)
25
LIMITING REAGENT/REACTANT STOICHIOMETRY PROBLEMS
1.
Hydrazine, N2H4, is used as a rocket fuel. It reacts with oxygen to form nitrogen gas & water. In a particular rocket engine,
2.29 kg of hydrazine & 3.14 kg of oxygen are available for reaction. Calculate the mass of water produced. (2.57 kg)
2.
Sodium & chlorine react to form sodium chloride. If 12.5 g of sodium & 25.5 g of chlorine gas are available for reaction,
determine the mass of NaCl produced.
(31.8 g)
3.
One type of stomach antacid consists of magnesium hydroxide. This reacts with stomach acid, which is mainly hydrochloric
acid, to form magnesium chloride & water. If 2.86 g of magnesium hydroxide & 5.15 g of hydrochloric acid are available for
reaction, what mass of magnesium chloride is produced?
(4.67 g)
4.
Hydrogen gas reacts with nitrogen gas to form ammonia. What mass of ammonia could be produced if 25.0 g of hydrogen
gas is mixed with 95.0 g of nitrogen?
(116 g)
5.
Aluminum reacts with iron (III) oxide in a single displacement reaction. If 46.6 g of aluminum are mixed with 145 g of iron
(III) oxide, what mass of iron would result?
(96.5 g)
6.
Ammonia reacts with oxygen to form nitrogen monoxide, NO, & water. What mass of nitric oxide could be produced if 68.0
g of ammonia is mixed with 35.0 g of oxygen?
(26.3 g)
7.
Calcium hydroxide & carbon dioxide react to form calcium carbonate & water. If 11.20 g of calcium hydroxide is reacted
with 16.65 g of carbon dioxide, what mass of calcium carbonate will form?
(15.20 g)
8.
Bismuth (III) oxide reacts with carbon to form bismuth & carbon dioxide. A student mixed 23.3 g of bismuth (III) oxide with
0.900 g of carbon. What is the mass of carbon dioxide produced?
(3.30 g)
9.
Ethane, C2H6, burns in oxygen to form carbon dioxide & water. If 0.0170 g of ethane is mixed with 0.0672 g of oxygen, what
is the combined mass of the two products?
(0.0803 g)
10. AsCl3 reacts with H2S to form As2S3 & HCl. If 4.58 g of AsCl3 are mixed with 1.36 g of H2S, how much of which reactant
will be left over after the reaction?
(0.0680 g)
STOICHIOMETRY REVIEW PROBLEMS
1.
Tungsten metal, W, is used to make incandescent bulb filaments. The metal is produced by a reaction between WO3 & H2.
Water is the other product. What mass of tungsten metal can be obtained from 4.81 kg of WO3? (3.81 kg)
2.
CCl4 is a strong organic solvent. It is produced when carbon disulphide reacts with chlorine gas to produce CCl4 & S2Cl2.
What mass of carbon tetrachloride would be produced by the reaction of 62.7 g of Cl2?
(45.3 g)
3.
Dinitrogen pentoxide decomposes to form nitrogen dioxide & oxygen gas. If 1.62 g of oxygen gas are produced, how many
grams of nitrogen dioxide would also be produced?
(9.32 g)
4.
Ammonia & hydrogen chloride react to form ammonium chloride. If 1.00 g of ammonia & 1.00 g of hydrogen chloride react,
identify the limiting reagent & calculate the mass of ammonium chloride that will be formed. (1.47 g)
5.
Commercial bleach is sodium hypochlorite. It is prepared when chlorine gas reacts with sodium hydroxide to produce
sodium chloride, sodium hypochlorite & water. If 4.5 g of chlorine gas is mixed with 5.6 g of sodium hydroxide, what mass
of sodium hypochlorite will be produced?
(4.7 g)
6.
MnO2 reacts with HCl to produce MnCl2, water & Cl2. If 24.0 g of MnO2 react with 38.2 g of HCl, calculate the mass of
excess reagent remaining after the reaction.
(1.23 g)
7.
Ammonia & oxygen gas react to form nitrogen monoxide & water. A chemist places 1.22 kg of ammonia & 1.96 kg of
oxygen gas into a reacting vessel. Identify the limiting reagent & calculate the mass of water produced. (1.32 kg)
8.
A student reacts 14.0 g of chlorine gas with 8.30 g of water to produce hydrogen chloride & oxygen gas. Identify the limiting
reagent & calculate the total mass of products produced.
(17.6 g)
26
% YIELD PROBLEMS
1.
The following reaction proceeds with a 70.0% yield: C6H6 + HNO3  C6H5NO2 + H2O
Calculate the mass of C6H5NO2 expected if 12.8 g of C6H6 reacts with excess HNO3.
(14.1 g)
2.
Sulphur trioxide & water react to form sulphuric acid. If 65.7 g of sulphur trioxide react, what is the theoretical yield of the
sulphuric acid? If the % yield is 85.0%, what was the actual yield of sulphuric acid?
(68.4 g)
3.
Nitric acid is one of the components of acid rain. It is produced when nitrogen dioxide reacts with water to form nitric acid &
nitrogen monoxide. If 80.0 g of nitrogen dioxide react, what is the theoretical yield of nitric acid? If 65.0 g are produced,
what is the % yield?
(89.0%)
4.
Yeasts can act on a sugar, such as glucose, in the following reaction: C6H12O6  2 C2H5OH + 2 CO2
If 223 g of C2H5OH are recovered after 1.63 kg of glucose react, what is the percentage yield of the reaction? (26.7%)
5.
Mercury, in its elemental form or in a chemical compound is highly toxic. Water-soluble mercury compounds, such as
Hg(NO3)2, can be removed from industrial wastewater by adding Na2S to the water, which forms a precipitate of HgS, which
can then be filtered out. If 3.45 x 1023 molecules of Hg(NO3)2 are reacted with excess Na2S, what mass of HgS can be
expected if this process occurs with 97.0% yield?
(129 g)
Hg(NO3)2 (aq) + Na2S (aq)  HgS (s) + 2 NaNO3 (aq)
% PURITY PROBLEMS
1.
Sodium carbonate is synthesized as follows: CaCO3 (s) + 2 NaCl (aq)  Na2CO3 (s) + CaCl2 (aq)
What mass of Na2CO3 can be produced from 3.40 kg of 87.1% pure CaCO3 & an unlimited amount of NaCl? (3.14 kg)
2.
AgNO3 reacts with copper metal in a single displacement reaction to produce silver metal & copper (II) nitrate. If 45.0 g of
81.0% pure copper react, what mass of silver should be produced?
(124 g)
3.
Pure iron can be produced by the following reaction: Fe3O4 (s) + 4 H2 (g)  3 Fe (s) + 4 H2O
What mass of iron can be produced if you have 1.70 kg of 35.0% pure Fe3O4 & unlimited hydrogen? (431 g)
4.
Given the following reaction: 4 FeS2 + 11 O2  2 Fe2O3 + 8 SO2
What mass of 73.5% pure oxygen gas is needed to produce 16.0 g of iron (III) oxide?
5.
Chlorine gas can be produced by the following reaction: MnO2 (s) + 4 HCl (aq)  MnCl2 (aq) + Cl2 (g) + 2 H2O (l)
What mass of 63.0% pure HCl would you need to prepare 8.00 kg of chlorine gas?
(26.1 kg)
6.
Copper metal is produced by the following reaction: 2 Cu2O (s) + Cu2S (s)  6 Cu (s) + SO2 (g)
If 2.50 kg of 35.0% pure copper (I) oxide is heated with 1.40 kg of 20.0% pure copper (I) sulphide, find the mass of copper
metal produced.
(671 g)
27
(24.0 g)
SOLUTIONS & SOLUBILITY UNIT
DETERMINING SOLUBILITY
1.
What does it mean for a substance to be soluble?
2.
Describe what happens during:
a. ionic dissociation b. covalent ionization
3.
For each of the following compounds, determine if the compound will dissolve (if at all) when added to water? State the
reason why.
a. K2O
g. NaBr
j. HF
d. MgCl2
h. BaSO4
k. BF3
b. NCl3
e. HNO3
i. Na2SO4
c. CH3OH
f. CH4
4.
Why is dissolving normally considered to be a physical change?
5.
Given any chemical compound, outline the steps necessary to determine if a compound is soluble or insoluble in water.
c. simple dissolving
CONCENTRATION UNIT CONVERSIONS
1.
What are the six concentration units? Where might each concentration unit be useful? (In other words, under which context
would the use of each concentration unit make more sense?)
2.
Perform the indicated conversions (assume the solvent is water where d(water) = 1.0 g/mL):
a. 12.5 ppm to ppb
b. 5.0 m/v% NaOH to mol/L
c. 2.30 mol/L NaCH3COO to m/v%
d. 0.00333 m/v% to ppm
e. 1.51 m/v% HCl to v/v%; d(HCl) = 1.19 g/mL
f. 0.750 m/m% CH3OH to mol/L; d(solution) = 0.800 g/mL
g. 536 ppm AgNO3 to mol/L
(1.25 x 104 ppb)
(1.25 mol/L)
(18.9 m/v%)
(33.3 ppm)
(1.27 v/v%)
(0.187 mol/L)
(0.00316 mol/L)
3.
The concentration of chlorine, Cl2, in a swimming pool is generally kept in the range of 1.4 to 4.0 mg/L. The water in a
certain pool has 3.0 mg/L of chlorine. Express this value as ppm & ppb.
(3.0 ppm , 3.0 x 103 ppb)
4.
Some municipalities add sodium fluoride to drinking water to help protect the teeth of children. The concentration of sodium
fluoride, NaF, is maintained at 1.2 x 10-3 g/L. Express this concentration in m/v %.
(1.2 x 10-4)
CONCENTRATION FORMULA PROBLEMS
1.
A solution of hydrochloric acid, HCl, was formed by dissolving 1.52 g of hydrogen chloride gas in enough water to make
24.1 mL of solution. What is the m/v% concentration of the solution?
(6.31%)
2.
Milk fat is present in milk. Whole milk usually contains about 5.000% milk fat (v/v). If you drink a glass of milk with a volume
of 250.0 mL, what volume of milk fat have you consumed?
(12.50 mL)
3.
If 55.00 g of KOH, is dissolved in 100.0 g of water, what is the m/m% concentration of the solution? (35.48%)
4.
Steel is an alloy of iron & carbon. The % composition of carbon in steel is about 1.7% (m/m). It also contains small
amounts of other materials, such as manganese & phosphorus. What mass of carbon is needed to make a 5.0 kg sample
of steel?
(85 g)
5.
At 25.0°C, a saturated solution of carbon dioxide gas has a concentration of 0.145% (m/v). What mass of carbon dioxide is
present in 250.0 mL of the solution?
(0.363 g)
6.
Calculate the m/m % of each of the following solutions:
a. 8.60 g of sodium chloride, NaCl, dissolved in 95.0 g of water.
b. 375 mg of calcium chloride, CaCl2, in 50.0 g of solution.
28
(8.30%)
(0.750%)
c.
225 mg of sulphuric acid, H2SO4, in 20.0 g of solution.
(1.13%)
7.
A 25.0 g sample of 14 karat gold contains 10.4 g of copper metal. What is the m/m % of copper in 14 karat gold? (41.6%)
8.
A solution of potassium iodide, KI, has a concentration of 2.50% (m/m). If this solution contains 258 mg of potassium
iodide, what is the mass of the solution?
(10.3 g)
9.
Coffee beans contain 1.45% caffeine (m/m). What mass of caffeine is present in 125 g of coffee beans? (1.81 g)
10. Water is added to 21.0 mL of ethanol, CH3CH2OH, until the total volume of the mixture is 100.0 mL. Calculate the v/v % of
the alcohol present in the mixture.
(21.0%)
MOLAR CONCENTRATION PROBLEMS
1.
A student mixed 4.25 mol of magnesium sulphide, MgS, in water to a final volume of 6.00 L. What is the molar
concentration of the solution?
(0.708 mol/L)
2.
A solution was made by mixing 76.3 g of calcium nitrate, Ca(NO3)2, in water to a total volume of 4.25 L. What was the molar
concentration?
(0.109 mol/L)
3.
A student needs to make 2.50 L of a solution with a molar concentration of 0.275 mol/L. How many moles of solute should
be used?
(0.688 mol)
4.
A student was asked to mix 350.0 mL of a 0.150 mol/L solution of sodium carbonate, Na2CO3. What mass of sodium
carbonate is required?
(5.56 g)
5.
Suppose 2.50 L of 0.125 mol/L sodium hydroxide, NaOH, solution is needed for a lab. However, only 0.500 mol/L sodium
hydroxide is available. What volume of concentrated sodium hydroxide is needed?
(625 mL)
6.
What concentration of solution is obtained by diluting 50.0 mL of 0.720 mol/L aqueous NaNO3, to 0.400 L? (0.0900 mol/L)
7.
A solution is prepared by adding 600 mL of distilled water to 100.0 mL of 0.150 mol/L ammonium nitrate, NH4NO3.
Calculate the molar concentration of the solution.
(0.0214 mol/L)
8.
Sulphuric acid, H2SO4, can be sold as a 18.0 mol/L solution. What volume of water is needed to dilute 10.0 mL of the
sulphuric acid to obtain a concentration of 1.50 mol/L?
(1.10  102 mL)
9.
Water is used to dilute 8.00 mol/L potassium nitrate, KNO3, solution to produce 700.0 mL of a solution with a concentration
of 6.00 mol/L. What volumes of water & potassium nitrate solution are used?
(525 mL, 175 mL)
10. Concentrated hydrochloric acid, HCl, has a density of 1.19 g/mL & a percentage purity of 39.1%. Suppose a lab required
7.25 L of 0.500 mol/L solution. What volume of concentrated hydrochloric acid (in mL) is needed for the dilution? (284 mL)
DILUTION PROBLEMS
1.
A student mixed 24.5 mL of a stock solution of potassium permanganate, KMnO4, with a concentration of 5.25 mol/L with
enough water to make 750 mL of a new solution.
a. What was the final concentration?
(0.172 mol/L)
b. How much additional water is necessary to create a solution of concentration 0.125 mol/L? (279 mL)
2.
A solution was made by mixing 275 mL of a stock solution of sodium nitrate, NaNO3, with a concentration of 6.25 mol/L, into
water. The final concentration was 1.20 mol/L. What was the final volume of the solution? (1.43 L)
3.
A lab required 250.0 mL of LiOH solution with a concentration of 0.125 mol/L. If the stock solution was 5.50 mol/L, how
much of the stock solution was required to make the dilute solution?
(5.68 mL)
4.
A solution was made by mixing 325 mL of stock solution, with a concentration of 2.50 mol/L, with 525 mL of water.
a. What was the final concentration?
(0.956 mol/L)
b. How much additional water is necessary to create a solution of concentration 0.305 mol/L? (1.81 L)
5.
A solution was made by mixing 175 mL of stock solution with water. The new solution had a volume of 1.50 L & a
concentration of 0.250 mol/L. What was the concentration of the stock solution?
(2.14 mol/L)
29
6.
A stock solution was made by mixing 37.5 g of calcium nitrate, Ca(NO3)2, in water to a final volume of 250.0 mL.
Afterwards, 35.0 mL of the stock solution were mixed in water to a final volume of 500 mL.
a. What was the concentration of stock solution?
(0.914 mol/L)
b. What was the concentration of the new solution?
(0.0640 mol/L)
7.
A stock solution was made by mixing 78.4 g of (NH4)2CO3, in water to a final volume of 425 mL. Then some of the stock
solution was used to make another solution with a volume of 75.0 mL & a concentration of 1.25 mol/L.
a. What was the concentration of the stock solution?
(1.92 mol/L)
b. What volume of the stock solution was used to make the final solution?
(48.8 mL)
8.
A student made 500.0 mL of a stock solution of potassium phosphate, K3PO4. Then she mixed 35.25 mL of the stock
solution with water to make a new solution with a volume of 125.0 mL & a concentration of 0.1250 mol/L. What mass of
potassium phosphate did the student use to make the stock solution?
(47.06 g)
SOLUTIONS & SOLUBILITY PROBLEMS
For each of the following compounds, use the solubility table to predict if it will form a solution. If yes, calculate the molar
concentration.
1. 50.0 g of sodium phosphate, Na3PO4, in 500 mL of solution.
2.
75.0 g of magnesium carbonate, MgCO3, 300 mL of solution.
3.
85.5 g of aluminum chloride, AlCl3, in 450 mL of solution.
4.
102.5 g of barium hydroxide, Ba(OH)2, in 600.0 mL of solution.
5.
86.5 g of calcium hydroxide, Ca(OH)2, in 350 mL of solution.
6.
95.2 g of aluminum phosphate, AlPO4, in 425 mL of solution.
7.
35.5 g of magnesium bromide, MgBr2, in 275 mL of solution.
8.
57.3 g of lead (II) chloride, PbCl2, in 175 mL of solution.
9.
48.0 g of barium chloride, BaCl2, in 215 mL of solution.
10. 27.5 g of ammonium carbonate, (NH4)2CO3, in 55.0 mL of solution.
Answers:
1. 0.610 mol/L
3. 1.43 mol/L
4. 0.9970 mol/L
7. 0.701 mol/L
9. 1.07 mol/L
10. 5.20 mol/L
Take the compounds that will form solutions & pair them up in order. For each pair:
 write & balance an equation to show the reaction between the two.
 identify which product is the precipitate, using phase labels.
 write an ionic equation.
 write a net ionic equation.
SOLUBILITY PROBLEMS
1.
In general, as temperature increases, what happens to the solubility of each substance in water? (assuming it is soluble)
a. solid
b. liquid
c. gas
2.
Calculate the solubility, in grams per 100 mL of water, of potassium sulphate if 1.20 g of the solid dissolves in 10.5 mL of
water at 20°C.
(11.4 g/100 mL)
3.
The solubility of aluminum fluoride is 0.559 g/100 mL water at 25°C. Is it possible to dissolve 3.0 g of solid in 500 mL of
water at 25°C? Show your calculations.
(2.80 g)
4.
Molar solubility refers to the solubility of a substance in mol/L. The solubility of potassium bromide at 50°C is 80 g/100 mL.
Convert this value to molar solubility.
(6.72 mol/L)
30
5.
A glass of cold water left sitting on a counter at room temperature usually develops many small bubbles on the inside of the
glass. Describe what is likely happening.
6.
The solubility of AlF3 is listed as 2.0 g/100 mL at a certain temperature. List two ways to dissolve 3.0 g of AlF3.
7.
Potassium alum, KAl(SO4)2 · 12 H2O, is used to stop bleeding from small cuts. The solubility of potassium alum, at various
temperatures, is given in the following table:
Solubility (g/100 g of water)
Temperature (°C)
4
0
10
10
15
20
23
30
31
40
49
50
67
60
101
70
135
80
a. Plot a graph of solubility against temperature.
b. From your graph, interpolate the solubility of potassium alum at 67°C.
c. By extrapolation, estimate the solubility of potassium alum at 82°C.
d. Look at your graph. At what temperature will 120 g of potassium form a saturated solution in 100 g of water?
e. Look at your graph. How much of 35 g of potassium alum will not dissolve at 35°C?
SOLUTION STOICHIOMETRY PROBLEMS
1.
A student mixed 35.0 mL of a 0.525 mol/L solution of potassium iodide with excess lead (II) nitrate in a double displacement
reaction. What mass of lead (II) iodide was produced in the reaction?
(4.24 g)
2.
In a single displacement reaction, a copper wire was placed in 125 mL of a solution of silver nitrate with a concentration of
0.100 mol/L, & the reaction proceeded until all the silver nitrate reacted. Suppose 1.05 g of silver metal was recovered in an
experiment. What was the % yield of the reaction?
(77.9%)
3.
Excess CuSO4 reacted with 32.5 mL of sodium carbonate solution in a double displacement reaction. If 0.750 g of copper
(II) carbonate were produced, what was the concentration of the sodium carbonate solution? (0.187 mol/L)
4.
Copper metal reacts with nitric acid. If 275 mL of nitric acid with a concentration of 6.00 mol/L react, what mass of copper
would be required to react with it? The unbalanced chemical equation is below.
(39.3 g)
Cu (s) + HNO3 (aq)  Cu(NO3)2 (aq) + NO (g) + H2O (l)
5.
Nickel (II) chloride reacts with sodium phosphate in a double displacement reaction. If 1.37 g of nickel (II) phosphate is
produced, what volume of 0.125 mol/L nickel (II) chloride is needed for the reaction?
(89.8 mL)
6.
The concentration of magnesium ions (assume magnesium chloride, MgCl2) in seawater was analyzed & found to be
0.0500 mol/L. What minimum volume of 0.200 mol/L sodium hydroxide, NaOH, would be needed in an industrial process to
precipitate all of the magnesium ions, Mg2+, from 1000 L of sea water?
(5.00 x 102 L)
7.
Aqueous solutions that contain silver ions are usually treated with chloride ions to recover silver chloride, AgCl. What is the
minimum volume of 0.25 mol/L magnesium chloride, MgCl2, needed to precipitate all the silver ions in 60.0 mL of 0.30 mol/L
silver nitrate, AgNO3?
(36 mL)
8.
Ammonium phosphate, (NH4)3PO4, can be used as a fertilizer. 6.00 g of (NH4)3PO4 is dissolved in sufficient water to
produce 300 mL of solution. What are the concentrations of NH4+ & PO43- present?
(0.402 mol/L & 0.134 mol/L)
9.
8.76 g of sodium sulphide, Na2S, is added to 350 mL of 0.250 mol/L lead (II) nitrate solution, Pb(NO 3)2. Calculate the
maximum mass of precipitate that can form.
(20.9 g)
10. Silver chromate, Ag2CrO4, is insoluble. It forms a brick-red precipitate. Calculate the mass of silver chromate that forms
when 50.0 mL of 0.100 mol/L silver nitrate, AgNO3, reacts with 25.0 mL of 0.150 mol/L sodium chromate. (0.829 g)
11. Suppose that you want to remove the barium ions from 120 mL of 0.0500 mol/L barium nitrate solution, Ba(NO3)2. What is
the minimum mass of sodium carbonate, Na2CO3, that you should add?
(0.636 g)
31
CONJUGATE ACIDS & BASES
1.
What is the Arrhenius definition of acids & bases?
2.
What is the Bronsted-Lowry definition of acids & bases?
3.
For each of the following acid-base reactions:
 label each reactant as an acid or base
 label each product as a conjugate acid or conjugate base
 link the conjugate acid-base pairs
a. HSO4- + NH3  SO42- + NH4+
d.
b. HPO42- + NH4+  H2PO4- + NH3
e.
c. SO32- + NH4+  HSO3- + NH3
f.
H2PO4- + HCO3-  HPO42- + H2CO3
F- + HSO4-  HF + SO42HSO4- + H2O  SO42- + H3O+
4.
What is meant by the conjugate acid of a base?
5.
H2PO3- is an amphiprotic substance. Write two equations, one in which it acts like an acid, & one in which it acts like a
base. (Hint: review your notes to determine what it could react with in each case).
PH CALCULATIONS
1.
What does pH stand for?
2.
For each of the following, determine whether they are acidic or basic, & calculate the indicated value to the correct number
of significant digits:
a. Calculate the pH of a solution with [H3O+] = 0.00270 mol/L
(2.569)
b. [H3O+] in a cola drink is about 5.000 x 10-3 mol/L. Calculate the pH of the drink.
(2.3010)
c. A glass of orange juice has [H3O+] of 2.900 x 10-4 mol/L. Calculate the pH of the juice. (3.5376)
d. [H3O+] of a solution of sodium hydroxide is 6.59 x 10-10 mol/L. Calculate the pH of the solution. (9.181)
e. The pH of a solution is 3.34. What is the [H3O+]?
(4.6 x 10-4 mol/L)
f. Which of the above five questions is the strongest acid? Which is the weakest acid?
PH & POH CALCULATIONS
1.
What does pOH stand for? Given pH, how would you find pOH?
2.
Calculate the pOH of a solution with [OH-] = 0.0125 mol/L.
(1.903)
3.
Calculate the [H+] of a solution with [OH-] = 2.5 x 10-6 mol/L.
(4.0 x 10-9 mol/L)
4.
Calculate the [OH-] of a solution with [H+] = 3.27 x 10-10 mol/L.
(3.06 x 10-5 mol/L)
5.
The pH of a solution is 1.20. What is the pOH?
(12.80)
6.
The pH of a solution is 4.00. What is the [OH-]?
(1.0 x 10-10 mol/L)
7.
The pOH of a solution is 7.50. What is the [H3O+]?
(3.2 x 10-7 mol/L)
8.
The pH of a solution is 3.15. What is the [OH-]?
(1.4 x 10-11 mol/L)
9.
The pH of a hydrochloric acid solution is 5.620. What is [HCl]?
(2.40 x 10-6 mol/L)
10. The pOH of a magnesium hydroxide solution is 5.81. What is [Mg(OH)2]?
32
(7.7 x 10-7 mol/L)
PH & STRONG ACID PROBLEMS
1.
Define the following terms:
a. Strong acid
b.
Weak acid
c.
Strong base
d.
Weak base
2.
A student prepared 1250 mL containing 1.34 x 10-4 mol of nitric acid. What was the pH of the solution? (3.970)
3.
A solution was prepared by bubbling 0.145 g of hydrochloric acid into enough water to make 1.75 L of solution. What was
the pH of the acid?
(2.643)
4.
If 12.5 mL of hydrobromic acid with concentration 0.0500 mol/L was combined with enough water to make 2.00 L of
solution, what was the pH?
(3.505)
5.
A student was asked to create 2.00 L of perchloric acid with a pH of 2.500. If the original stock solution had a concentration
of 3.00 mol/L, what volume of stock would be required to prepare the solution?
(2.11 mL)
6.
A solution of hydroiodic acid had a pH of 4.500. It was prepared using 40.0 mL of stock solution with a concentration of
2.00 mol/L. What was the volume of the acid solution?
(2.53 x 103 L)
STRONG ACID & STRONG BASE TITRATION PROBLEMS
1.
Describe or draw the apparatus needed for an acid-base titration experiment. Label all equipment & identify the purpose of
each part of the titration set-up.
2.
What happens to the chemicals during a titration?
3.
What is the concentration of potassium hydroxide solution, KOH, if 12.8 mL of this solution is required to react with 25.0 mL
of 0.110 mol/L sulphuric acid, H2SO4?
(0.430 mol/L)
4.
What volume of 0.125 mol/L calcium hydroxide, Ca(OH)2, is required to react completely with 15.0 mL of 0.100 mol/L
sulphuric acid?
(12.0 mL)
5.
In a chemical analysis, a 10.0 mL sample of phosphoric acid, H3PO4, was reacted with 18.2 mL of 0.259 mol/L sodium
hydroxide. Calculate the concentration of the phosphoric acid.
(0.157 mol/L)
6.
A student used a sodium hydroxide solution with concentration 0.100 mol/L to titrate 25.0 mL of hydrochloric acid with an
unknown concentration. In three tests, the average volume of NaOH added was 14.65 mL. What was the concentration of
the acid?
(0.0586 mol/L)
7.
A solution of KOH with a concentration of 0.0225 mol/L was used to titrate a 12.5 mL of HNO3 with an unknown
concentration. The average volume of titrant was 13.27 mL. What was the concentration of the acid? (0.0239 mol/L)
8.
A solution of nitric acid was prepared by mixing 15.0 mL of acid with 25.0 mL of water. The resulting solution was then
titrated with sodium hydroxide with a concentration of 0.125 mol/L. The average volume of titrant added was 7.64 mL.
What was the original concentration of the acid?
(0.0637 mol/L)
9.
A student mixed 12.5 mL of barium hydroxide solution with water to a final volume of 50.0 mL. The solution was titrated
with 0.0250 mol/L hydrochloric acid. In three trials, the volumes of titrant added were 15.4 mL, 15.5 mL & 15.5 mL. What
was the original concentration of the barium hydroxide solution?
(0.0155 mol/L)
10. A solution of aluminum hydroxide was prepared by combining 10.5 mL of stock solution with 49.5 mL of water. The solution
was titrated with 0.100 mol/L hydrochloric acid. The average volume of titrant added was 15.47 mL. What was the original
concentration of aluminum hydroxide?
(0.0491 mol/L)
STRONG ACID & STRONG BASE NEUTRALIZATION PROBLEMS
Note: the equivalence point may or may not have been reached…
1.
30.0 mL of 0.150 mol/L hydrochloric acid was added to 20.0 mL of 0.200 mol/L sodium hydroxide. Find the [OH -] of the
resultant solution.
(1.00 x 10-12 mol/L)
33
2.
In a titration experiment, 25.0 mL of 0.100 mol/L HCl solution was used as the sample solution. The endpoint was overshot
when 27.0 mL of 0.100 mol/L NaOH solution was added. What is the pH of the resultant solution? (11.585)
3.
30.0 mL of 0.150 mol/L HCl was added to 40.0 mL of 0.200 mol/L NaOH. Find the pH of the resultant solution. (12.699)
4.
45.0 mL of 0.100 mol/L hydrochloric acid was added to 10.0 mL of 0.150 mol/L calcium hydroxide. Find the [H +] of the
resultant solution.
(0.0273 mol/L)
5.
45.0 mL of 0.100 mol/L hydrochloric acid was added to 10.0 mL of 0.225 mol/L calcium hydroxide. Find the pH of the
resultant solution.
(7.000)
6.
50.0 mL of 0.150 mol/L hydrochloric acid was added to 20.0 mL of 0.200 mol/L aluminum hydroxide. Find the pH of the
resultant solution.
(12.808)
7.
30.0 mL of 0.150 mol/L phosphoric acid was added to 20.0 mL of 0.200 mol/L sodium hydroxide. Not all of the acid was
neutralized. Find the pH of the resultant solution.
(0.721)
8.
45.0 mL of 0.100 mol/L sulphuric acid was added to 10.0 mL of 0.150 mol/L aluminum hydroxide. Find the pH of the
resultant solution.
(1.09)
GASES & ATMOSPHERIC CHEMISTRY UNIT
PRESSURE & TEMPERATURE CONVERSIONS
1.
Define pressure.
2.
What does STP & SATP stand for?
3.
Convert the following pressure values to kilopascals:
a. 1.40 atm
b. 0.987 atm
c.
754 torr
d.
792 mmHg
4.
Convert the following pressure values to atmospheres:
a. 105 kPa
b. 9.57 x 103 Pa
c.
745 torr
d.
815 mmHg
5.
Convert the following pressure values to mmHg or torr:
a. 98.5 kPa
b. 104 kPa
c.
1.10 atm
d.
0.975 atm
Convert the following temperature values to degrees Celsius:
a. 298 K
b.
245 K
Convert the following temperature values to Kelvin:
a. 20.0°C
b.
-45.5°C
6.
7.
BOYLE’S LAW PROBLEMS
1.
A 50.00 cm3 sample of nitrogen gas is collected at 101.3 kPa. If the volume is reduced to 5.000 cm 3, & the temperature
remains constant, what will the final pressure of the nitrogen be?
(1013 kPa)
2.
A weather balloon has a volume of 1.000  103 L at a pressure of 740.0 torr. The balloon rises to a height of 1.000 km
where the atmospheric pressure is measured as 450.0 torr. Assuming there is no change in temperature, what is the final
volume of the weather balloon?
(1644 L)
3.
A 45.0 cm3 sample of nitrogen gas is collected at 1.0 atm. The nitrogen is compressed to a pressure of 10.0 atm. What is
the final volume of the nitrogen if the temperature remains constant?
(4.5 cm3)
4.
A 1.00 L helium balloon is floating in the air on a day when the atmospheric pressure is 102.5 kPa & the temperature is
20.0°C. Suddenly, clouds appear & the pressure rapidly drops to 98.6 kPa at a temperature of 20.0°C. By how much did
the volume of the balloon increase?
(3.96%)
34
5.
0.750 L of oxygen gas is trapped at 101.3 kPa in a cylinder with a moveable piston. The piston is moved & the gas is
compressed to a volume of 0.500 L. What was the final pressure, in atm, that was applied to the oxygen gas if the
temperature remains unchanged?
(1.50 atm)
CHARLES’ LAW PROBLEMS
1.
Methane gas can be condensed by cooling & increasing the pressure. A 6.00  102 L sample of methane gas at 25.0°C &
1.00  102 kPa is cooled to -20.0°C. What will be the final volume?
(509 L)
2.
A sample of gas at 15.9°C is placed into a syringe. By what amount will the gas volume change (relative to the initial
volume) after the gas is subjected to a temperature of 65.5°C?
(1.17)
3.
A sample of nitrogen gas has a volume of 0.400 L at 100°C. At what temperature will it have a volume of 0.200 L if the
pressure does not change?
(-86.6°C)
4.
A 14.5 cm3 sample of oxygen gas at 24.3°C is drawn into a syringe with a maximum volume of 60.0 cm3. What is the
maximum change in temperature that the oxygen can be subjected to before the plunger pops out of the syringe? (933°C)
5.
A balloon is filled with 2.50 L of dry helium at 23.5°C. The balloon is placed in a freezer overnight. The next morning, the
balloon is removed & the volume is found to be 2.15 L. What was the temperature (in °C) inside the freezer if the pressure
remained constant?
(-18.0°C)
GAY-LUSSAC’S LAW PROBLEMS
1.
A glass bottle has an internal volume of 0.50 L & contains air at STP. It will burst if the pressure of the trapped gas exceeds
5.5 atm. To what temperature can the bottle be heated before it bursts?
(1.2 x 103°C)
2.
A metal cylinder, which has a safety valve that opens at a pressure of 1.00  102 atm, it to be filled with nitrogen gas &
heated to 300.0°C. What is the maximum pressure to which it can be filled at 25.0°C?
(52.0 atm)
3.
At 18.0°C, a sample of helium gas is stored in a metal cylinder exerts a pressure of 17.5 atm. What will the pressure
become if the tank is placed in a closed room where the temperature increases to 40.0°C? (18.8 atm)
4.
A cylinder of chlorine gas is stored in a concrete-lined room for safety. The cylinder is designed to withstand 50.0 atm of
pressure. The pressure gauge reads 35.0 atm at 23.2°C. An accidental fire in the room next door causes the temperature
in the storage room to increase to 87.5°C. What will the pressure gauge read at this temperature & will the cylinder
explode?
(42.6 atm)
5.
A truck leaves Yellowknife in January when the temperature is -30.0°C. The tires of the truck are inflated to 210.0 kPa.
Four days later, the truck arrives in California where the temperature is 30.0°C. What is the air pressure in the tires when
the truck arrives?
(262 kPa)
6.
Before leaving on a trip to Florida, the pressure inside the tires of a car at a gas station was 206.5 kPa. The temperature
was -7.50°C. The day after arriving in Florida, the tire pressure was 34.3 psi. Most pressure gauges in the United States
are calibrated in psi. What is the approximate temperature in Florida? (1 psi = 6.894 kPa)
(31.0°C)
COMBINED GAS LAW PROBLEMS
1.
A gas originally had a volume of 6.35 L at STP. If the conditions are altered to 400 mmHg & 25.0°C, what volume does the
gas have now?
(13.2 L)
2.
Sandra had a birthday party on a winter day. She tied 4.2 L balloons to the front of the house. The temperature was -2.0°C
& the pressure was 100.8 kPa. Unfortunately, the weather changed & a higher-pressure (103 kPa) cold front (-25°C)
rushed into town. What will happen to the balloons after the weather change?
(3.8 L)
3.
A sample of gas has a volume of 150 mL at 260 K & 92.3 kPa. What will the new volume be at 376 K & 123 kPa? (0.16 L)
4.
A cylinder at 48 atm pressure & 290 K releases 35 mL of carbon dioxide gas into a 4.0 L container at 297 K. What is the
pressure inside the container?
(0.43 atm)
35
5.
An automated instrument has been developed to help drug-research chemists determine the amount of nitrogen in a
compound. Any compound containing carbon, nitrogen & hydrogen is reacted with copper (II) oxide to produce carbon
dioxide, water & nitrogen gases. The gases are collected separately & analyzed. In an analysis of 39.8 mg of caffeine
using this instrument, 10.1 mL of nitrogen gas is produced at 23.0°C & 746 torr. What must the new temperature of the
nitrogen gas be, in °C, if the volume is increased to 12.0 mL & the pressure is increased to 780 torr? (94.7°C)
6.
A container of gas has a volume of 2.50 L. Its pressure is 450.0 mmHg & its temperature is unknown. The gas is now
changed to a volume of 5.50 L at 650.0 mmHg & 30.0°C. Calculate what the initial temperature must have been. (-178°C)
7.
A 3.84 mL sample of nitrogen gas at 23.0°C & 785 mmHg is cooled to 0.00°C. The pressure changes to 761 mmHg. What
is the new volume?
(3.65 mL)
8.
A cylinder of gas has a volume of 135 L, a pressure of 15.0 atm & a temperature of 22.0°C. The gas now escapes into the
room, which has a volume of 2.00 x 105 L. The temperature is lowered to 18.0°C. What is the pressure of the gas in this
room?
(9.99  10-3 atm)
9.
A weather balloon with a volume of 55.0 L is filled with hydrogen gas at a pressure of 98.5 kPa & a temperature of 13.0°C.
When the balloon is released, it rises to the stratosphere where the temperature is -48.0°C & the pressure is 19.7 kPa.
What is the volume of the balloon under these conditions?
(216 L)
10. A 10.1 mL sample of nitrogen gas is prepared at 23.0°C & 746 mmHg. If the volume changes to 9.14 mL & the pressure
increases to 762 mmHg, what is the temperature of the gas?
(0.599°C)
IDEAL GAS LAW PROBLEMS
1.
A 14.0 L balloon holds 2.04 moles of carbon dioxide gas at 259 kPa & 30.2°C. If 3.56 moles of carbon dioxide is placed at
800 torr & 45.0°C into a different balloon, what would be volume become?
(62.2 L)
2.
What mass of sulphur dioxide, SO2, is in 36.2 L of gas at 1.22 atm & 9.20°C?
(122 g)
3.
How many moles of gas are present in 11.2 L at SATP? How many molecules?
(2.72 x 1023 molecules)
4.
What is the volume of 3.45 mol of argon gas at SATP?
(85.5 L)
5.
A quantity of carbon dioxide gas has a volume of 19.5 L at a temperature of 400°C & a pressure of 23.8 atm. How many
moles of carbon dioxide are present? What is the mass of the gas?
(0.370 kg)
6.
By how much will a balloon change (relative to the initial volume) when 39.2 g of oxygen gas at 492 kPa & 15.5°C is
changed to contain 15.4 g at STP?
(1.81)
7.
An experiment calls for 3.50 mol of chlorine gas. What volume would this be if the gas is at 34.0°C & 2.45 atm? (36.0 L)
8.
The maximum safe pressure that a certain 4.00 L container can hold is 355 kPa. If the container is filled with 0.410 mol of
gas, what is the maximum temperature to which the container can be subjected?
(143°C)
9.
A 2.50 L flask was used to collect a 5.65 g sample of propane gas, C3H8. After the sample was collected, the gas pressure
was found to be 956 mmHg. What was the temperature of the propane in the flask?
(26.0°C)
10. What mass of oxygen gas is in a 50.0 L tank at 21.0°C when the pressure is 15.7 atm?
(1.04 kg)
IDEAL GAS LAW MOLAR VOLUME PROBLEMS
Date: __________________________________________
1.
Manipulate the Ideal Gas Law to derive an equation to calculate the molar volume of a gas.
2.
What is the molar volume of any gas at STP?
(22.4 L/mol)
3.
What is the molar volume of any gas at SATP?
(24.8 L/mol)
36
4.
What is the molar volume of oxygen gas at 19.0°C & 100.0 kPa?
(24.3 L/mol)
IDEAL GAS LAW MOLAR MASS PROBLEMS
1.
Using the relationships n = m/M & the Ideal Gas Law, derive an equation to calculate the molar mass of a gas.
2.
A 1.28 g sample of liquid was vapourized in a 250.0 mL flask at 121°C & 786 torr. What is the molar mass of the
substance?
(1.60 x 102 g/mol)
3.
What is the molar mass of a gas, 0.842 g of which occupies 0.450 L at a pressure of 100.0 kPa & a temperature of
100.0°C?
(58.1 g/mol)
4.
A scientist isolates 2.366g of a gas. The sample occupies a volume of 800 mL at 78.0°C and 103 kPa. Which gas is it likely
to be?
(Kr)
5.
As geologists study the area where an ancient marsh was located, they discover an unknown gas seeping from the ground.
They collect a sample of the gas & take it to a lab for analysis. Lab technicians find that the gas is made up of 80.0%
carbon & 20.0% hydrogen. They also find that a 4.60 g sample occupies a volume of 2.50 L at 1.50 atm & 25.0°C. What is
the molecular formula of the gas?
(C2H6)
6.
A gas that consists of only nitrogen & oxygen atoms is found to contain 30.0% nitrogen. A 9.23 g sample of the gas
occupies 2.20 L at STP. What is the gas?
(N2O4)
1.
Using the relationships n = m/M, d = m/V & the Ideal Gas Law, derive an equation to calculate the density of a gas.
2.
The atmosphere consists largely of nitrogen gas.
a. Calculate the densities of N2, Cl2 & He at 20.0°C & 101.3 kPa.
(1.16 g/L, 2.95 g/L, 0.166 g/L)
b. Explain why chlorine gas would appear as a yellow cloud creeping along the ground when it was used in World War I.
c. Explain why helium balloons float.
3.
Calculate the density of hydrogen sulphide gas, H2S, at 56.0°C & 967 mmHg.
4.
An unknown compound has a density of 1.585 g/L at 90.0°C & 753 mmHg. What is the molar mass of the compound?
(47.7 g/mol)
5.
Oxygen gas makes up about 20.00% of the Earth’s atmosphere. Find the density of pure oxygen gas, in g/L, at 12.50°C &
126.63 kPa.
(1.706 g/L)
6.
On Mars, the atmosphere has a pressure of about 15.0 torr. The daytime temperature is about -40.0°C. How many moles
are there in 1.00 L of this atmosphere? How many molecules is this per litre?
(6.21 x 1020 molecules/L)
IDEAL GAS LAW DENSITY PROBLEMS
(1.61 g/L)
GAS STOICHIOMETRY CALCULATIONS
1.
Ammonia, NH3, is produced by a reaction of nitrogen gas, N2, & hydrogen gas, H2. Suppose that 12.0 L of nitrogen gas
reacts with hydrogen gas & produces ammonia, all at the same temperature & pressure. What volume of ammonia is
produced? What volume of hydrogen is consumed?
(24.0 L, 36.0 L)
2.
Water vapour is produced when hydrogen gas, H2, is reacted with oxygen gas, O2. What volume of water vapour can be
produced if 40 L of hydrogen gas is mixed with 15 L of oxygen gas to produce water vapour, all at the same pressure &
temperature? What volume of the excess gas will remain?
(30 L, 10 L)
3.
Calcium carbide, CaC2, reacts with water to produce acetylene (C2H2) & calcium hydroxide, Ca(OH)2. Calculate the volume
of acetylene produced at 26.0°C & 104 kPa from 10.2 g of calcium carbide & excess water. (3.81 L)
4.
Lithium hydroxide, LiOH, is used in spacecraft to absorb carbon dioxide, CO2, exhaled by astronauts, which then produces
lithium carbonate, Li2CO3, & water. What volume of carbon dioxide at 21.1°C & 1.15 atm could be absorbed by 348 g of
lithium hydroxide?
(153 L)
37
5.
A 24.9 mL sample of hydrochloric acid, HCl, reacts with excess sodium carbonate, Na2CO3, to produce carbon dioxide,
CO2, water & sodium chloride, NaCl. The volume of carbon dioxide formed is 141 mL at 27.0°C & 742 torr. What is the
molar concentration of hydrochloric acid?
(0.449 mol/L)
6.
Methanol, CH3OH, has potential to be used as an alternative fuel. It burns in the presence of oxygen, O2, to produce carbon
dioxide, CO2, & water. 10.0 L of oxygen gas is completely consumed at 775 torr & 35.1°C. What volume of carbon dioxide
at 750 torr & 23.4°C is produced?
(6.63 L)
7.
When ammonia, NH3, & oxygen, O2, react, they produce nitrogen monoxide, NO, & water. What volume of oxygen gas at
35.0°C & 2.15 atm is needed to produce 46.1 g of nitrogen monoxide?
(22.6 L)
8.
Magnesium nitride reacts with water to produce ammonia gas & magnesium hydroxide solid. What volume of ammonia gas
at 24.0°C & 743 mmHg will be produced from 4.56 g of magnesium nitride?
(2.25 L)
DALTON’S LAW OF PARTIAL PRESSURES PROBLEMS
1.
To speed up a reaction in a vessel pressurized at 98.0 kPa, a chemist added 202.65 kPa of hydrogen gas. What was the
resultant pressure?
(301 kPa)
2.
A gas mixture contains 12% neon, 23% helium & 65% radon. If the total pressure is 116 kPa, what is the partial pressure of
each gas?
(14 kPa, 27 kPa, 75 kPa)
3.
A mixture of nitrogen & carbon dioxide gas is at a pressure of 1.00 atm & a temperature of 278 K. If 30.0% of the mixture is
nitrogen, what is the partial pressure of the carbon dioxide?
(0.700 atm)
4.
The partial pressure of argon gas, making up 40.0% of a mixture, is 325 torr. What is the total pressure of the mixture in
kPa?
(108 kPa)
5.
A mixture of non-reactive gases in a 5.00 L flask at 15.0°C includes 0.835 g of xenon gas. What is the partial pressure of
xenon?
(3.05 kPa)
6.
A 200.0 mL flask contains 1.03 mg of oxygen gas & 0.410 mg of helium gas at 15.0°C. Calculate the total pressure in the
flask.
(1.61 kPa)
7.
A 1.00 L sample of a gas mixture contains nitrogen gas & oxygen gas at 25.0°C. The mass of nitrogen gas is 1.05 g, & the
total pressure in the flask is 115 kPa. What is the partial pressure of oxygen in the flask?
(22.1 kPa)
8.
The atmosphere in a sealed diving bell with a volume of 1.00 L contains oxygen gas & helium gas at 20.0°C. If the gas
mixture has 0.200 atm of helium & a total pressure of 3.00 atm, calculate the mass of oxygen. (3.72 g)
WATER VAPOUR PRESSURE PROBLEMS
1.
In an experiment, 0.750 L of hydrogen gas is collected over water at 25.0°C & 101.6 kPa. What volume will the dry
hydrogen occupy at 103.3 kPa & 25.0°C?
(0.715 L)
2.
A chemist collects 8.15 L of oxygen gas, collected over water, at 22.0°C & 105.0 kPa. What mass of oxygen gas was
collected?
(10.9 g)
3.
Formic acid, HCOOH, can decompose to form carbon monoxide gas, CO, & liquid water. If 3.85 L of CO was collected by
downward displacement of water at 25.0°C & 689 mmHg, how many grams of formic acid were reacted? (6.34 g)
4.
Ammonium nitrite, NH4NO2, decomposes to produce nitrogen gas & liquid water. What mass of ammonium nitrate needs to
react if 4.10 L of nitrogen gas were collected over water at 19.0°C & 97.8 kPa?
(10.3 g)
5.
Magnesium metal reacts with excess dilute hydrochloric acid, HCl, to produce hydrogen gas & aqueous magnesium
chloride, MgCl2. Suppose 0.150 g of magnesium metal was used & the hydrogen gas is collected over water at 28.0°C &
101.8 kPa. What volume will the hydrogen occupy?
(0.158 L)
38
REFERENCE MATERIAL
CONVERSION FACTORS FOR CALCULATIONS
1 metre (m) = 100 centimetres = 1000 millimetres = 39.370 inches
1 centimetre (cm) = 10 millimetres (mm) = 0.39370 inches
1 Angstrom (A) = 10-8 cm
1 micron (µ) = 10-3 millimetre
1 kilogram (kg) = 1000 grams = 2.2046 pounds
1 gram (g) = 1000 milligrams (mg)
1 atomic mass unit (amu) = 1.6604 x 10-24 g
1 litre (L) = 1000 millilitres (mL) = 1.0567 quarts
1 gallon (gal) = 3.7854 L
1 cubic inches (in3) = 16.387 mL
1 cubic foot (ft3) = 28317 mL
1 electron volt (eV) = 1.6021 x 10-12 erg = 23.061 kcal/mol
1 calorie (cal) = 4.1840 x 107 erg
1 joule = 1 x 107 erg
°C = 5/9(F – 32)
Kelvin = °C + 273.15
1 atm = 760 mmHg = 13.6 x 760 mmH2O
CONSTANTS
(assume all constants have infinite number of significant digits)
R = 8.3145 L.kPa/mol.K
R = 0.08206 L.atm/mol.K
NA = 6.023  1023 particles/mol
KW = 1  10-14 at 25.0°C
USEFUL FORMULAS
(not all formulas may be used)
m = nM
N = nNA
massofcomponent
%composition 
 100%
totalmolarmass
exp erimentalyield
 100%
theoreticalyield
pureyield
% purity 
 100%
impuresample
% yield 
m %  massofsolute( g )  100%
m
massofsolution( g )
v %  volumeofsolute(mL) 100%
v
volumeofsolution (mL)
C1V1 = C2V2
Kw = [H+][OH-]
pH = -log[H+]
pOH = -log[OH-]
V1 V2
P1 P2


T1 T2
T1 T2
PV = nRT
n = CV
massofsolu
te( g )
m %
100%
V
volumeofsolution (mL)
massofsolute( g )
10 6
massofsolution( g )
massofsolute( g )
ppb 
10 9
massofsolution( g )
massofsolute( g )
so lub ility 
100mLofsolution
ppm 
pH + pOH = pKw
KaKb = Kw
 [ A ] 

pH  pK a  log 
 [ HA] 
P1V1 = P2V2
P1V1 P2V2

T1
T2
Ptotal = P1 + P2 + …
Mass # = # protons + # neutrons
Average Atomic Mass =
(%abundance)(isotopicmass)
formal charge = valence electrons – ½(bonding pair electrons)
39
P1 P2

T1 T2
USEFUL POLYATOMIC IONS
Formula
NH4+
CH3COO- or C2H3O2BrO3CNOH-
Name
ammonium
acetate
bromate
cyanide
hydroxide
Formula
MnO3CrO42Cr2O72AsO43-
Name
manganate
chromate
dichromate
arsenate
METAL & HALOGEN ACTIVITY SERIES
Metal Activity Series:
Li
K
Ba
Ca
Na
Mg
Al
Zn
Cr
Fe
Cd
Co
Ni
Sn
Pb
H
Cu
Hg
Ag
Pt
Au
A possible phrase which may help memorize this list goes as follows:
“Linda, please bring Carmen Sandiego more awesome zebra crackers in ceramic containers,
not the large heavy crate marked striped perishable goods!”
Halogen Activity Series:
F
Cl
Br
I
Each element in an activity series (metal or halogen) will displace any element listed below it in the same series.
In the metal activity series, hydrogen is included, even though it is not a metal, because it forms positively charged ions like
metals do. Many metals will therefore displace hydrogen from compounds such as water or acids. Metals listed below hydrogen
do not react with water or acids.
SOLUBILITY TABLE @ SATP
(other versions are possible)
High
Solubility
(>0.1 mol/L at
SATP)
Cl-, Br-, I-
S2-
OH-
SO42-
CO32-,
PO43-,
SO32-
CH3COO-
NO3-
ClO3-
O2-
most
Group 1,
NH4+,
Group 2
Group 1,
NH4+,
Sr2+,
Ba2+, Tl+
most
Group 1,
NH4+
most
all
most
Group 1,
NH4+,
Ba2+
Ag+, Pb2+,
Ca2+,
Ag+, Pb2+,
Ba2+,
Tl+, Hg2+,
most
most
most
Ag+
none
Ca2+
Sr2+,
+
Cu
Hg2+,
Ra2+
All Group 1 compounds, including acids & all ammonium compounds are assumed to have high solubility in water
Low Solubility
(<0.1 mol/L at
SATP)
Cations
Anions
40
most
WATER VAPOUR PRESSURE TABLE
Temperature
(°C)
0.0
1.0
2.0
3.0
4.0
5.0
6.0
7.0
8.0
9.0
10.0
11.0
12.0
13.0
Vapour Pressure
(kPa)
0.61
0.66
0.71
0.76
0.81
0.87
0.93
1.00
1.07
1.15
1.23
1.31
1.40
1.50
Temperature
(°C)
14.0
15.0
16.0
17.0
18.0
19.0
20.0
21.0
22.0
23.0
24.0
25.0
26.0
27.0
Vapour Pressure
(kPa)
1.60
1.70
1.82
1.94
2.06
2.20
2.34
2.49
2.64
2.81
2.98
3.17
3.36
3.57
41
Temperature
(°C)
28.0
29.0
30.0
31.0
32.0
33.0
34.0
35.0
36.0
37.0
38.0
39.0
40.0
Vapour Pressure
(kPa)
3.78
4.01
4.24
4.49
4.75
5.03
5.32
5.62
5.94
6.27
6.62
6.99
7.37
PERIODIC TABLE
(Electronegativities, Common Valences, Atomic Masses)
1 (IA)
1
18 (VIIIA)
2.20
H
2
1+, 1-
Hydrogen
1.01
3
0.98
Li
1+
Lithium
6.94
11
0.93
Na
1+
Sodium
22.99
19
0.82
K
1+
Potassium
39.10
37
0.82
Rb
1+
Rubidium
85.47
55
0.79
Cs
1+
Cesium
132.91
87
0.70
Fr
1+
Francium
223.02
0
He
2 (IIA)
4
Legend
Atomic #
1.57
Be
2+
Symbol
2+
Calcium
40.08
38
0.95
Sr
2+
Strontium
87.62
56
0.89
Ba
2+
Barium
137.33
88
0.90
Ra
Radium
226.03
5
 Most stable valence is bolded
2+
14 (IVA)
2.04
3+
B
Atomic Name
Atomic Mass
Boron
10.81
13
1.61
2+
Magnesium
24.31
20
1.00
Ca
EN
Common Valences
Beryllium
9.01
12
1.31
Mg
13 (IIIA)
3+
3 (IIIB)
21
4 (IVB)
1.36
Sc
3+
Scandium
44.96
39
1.22
3+
Y
Yttrium
88.91
71
1.27
Lu
Lutetium
174.97
103
Lr
2+
3+
22
5 (VB)
1.54
Ti
3+, 4+
Titanium
47.88
40
1.33
4+
Zr
Zirconium
91.22
72
1.30
4+
Hf
Hafnium
178.49
104
23
6 (VIB)
1.63
2+, 3+, 4+, 5+
24
7 (VIIB)
1.66
2+, 3+, 6+
V
Cr
Vanadium
50.94
41
1.60
Chromium
52.00
42
2.16
3+, 5+
Nb
Niobium
92.91
73
1.50
5+
Ta
Tantalum
180.95
105
2+, 3+, 4+, 5+, 6+
25
8 (VIIIB)
1.55
2+, 3,+, 4+, 6+, 7+
27
1.88
2+, 3+
28
11 (IB)
1.91
2+, 3+
Mn
Fe
Co
Ni
Iron
55.85
44
2.20
Cobalt
58.93
45
2.28
Nickel
58.69
46
2.20
Mo
Tc
Technetium
(98.91)
75
1.90
1-, 2+, 4+, 6+, 7+
W
Tungsten
183.85
106
1.83
2+, 3+
10 (VIIIB)
Manganese
54.94
43
1.90
Molybdenum
95.94
74
2.36
2+, 3+, 4+, 5+, 6+
26
9 (VIIIB)
Re
Rhenium
186.21
107
2+, 3+, 4+, 6+, 8+
2+, 3+, 4+
2+, 4+
29
12 (IIB)
1.90
1+, 2+
Zn
Zinc
65.39
48
1.69
1+
Rh
Pd
Ag
Rhodium
102.91
77
2.20
Palladium
106.42
78
2.28
Silver
107.87
79
2.54
Os
Osmium
190.23
108
Ir
Iridium
192.22
109
Pt
2+, 4+
Platinum
195.88
110
2+
Cu
Ru
2+, 3+, 4+, 6+
1.65
Copper
63.55
47
1.93
Ruthenium
101.07
76
2.20
2+, 3+, 4+, 6+, 8+
30
1+, 3+
Cd
Cadmium
112.41
80
2.00
1+, 2+
Au
Gold
196.97
111
2+
2.55
4+
C
Carbon
12.01
14
1.90
4+
Al
Si
Aluminum
26.98
31
1.81
Silicon
28.09
7
16 (VIA)
3.04
3-
N
Nitrogen
14.01
15
2.19
3-
P
2.01
Phosphorus
30.97
33
2.18
Gallium
69.72
49
1.78
Germanium
72.61
50
1.96
Arsenic
74.92
51
2.05
In
Sn
Ga
3+
3+
Indium
114.82
81
2.04
1+, 3+
Hg
Mercury
200.59
112
6
15 (VA)
32
4+
Ge
2+, 4+
Tin
118.71
82
2.33
Tl
Pb
Thallium
204.38
Lead
207.2
2+, 4+
3-
As
Sb
3+, 5+
Antimony
121.75
83
2.02
Bi
3+, 5+
Bismuth
208.98
114
8
3.44
2-
O
Oxygen
16.00
16
2.58
2-
S
Sulphur
32.07
34
2.55
2-
Se
Selenium
78.96
52
2.10
2-
Te
Tellurium
127.60
84
2.00
Po
2+, 4+
Polonium
(208.98)
116
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
Cn
Fl
Lv
Lawrencium
(260.11)
Rutherfordium
(261.11)
Dubium
(262.12)
Seaborgium
(263.12)
Bohrium
(262.12)
Hassium
(264)
Meitnerium
(266.14)
Darmstadtium
(269)
Roentgenium
(272)
Copernicium
(277)
Flerovium
(285)
Livermorium
(289)
57
58
63
64
65
66
1.10
La
3+
1.12
59
1.13
60
1.14
61
1.13
62
1.17
2+, 3+
1.20
2+, 3+
1.20
3+
1.20
3+, 4+
1.22
3+
67
1.23
68
1.24
69
1.25
70
1.11
Eu
Gd
Tb
Dy
Europium
151.96
95
1.30
Gadolinium
157.24
96
1.30
Terbium
158.93
97
1.30
Dysprosium
162.50
98
1.30
Holmium
164.93
99
1.30
Erbium
167.26
100
1.30
Thulium
168.93
101
1.30
Ytterbium
173.04
102
1.30
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Americum
243.06
Curium
247.07
Berkelium
247.07
Californium
251.08
Actinium
227.03
Thorium
232.04
Protactinium
231.04
Uranium
238.03
Neptunium
237.05
Plutonium
224.06
42
3+
3+, 4+
3+
3+
Einsteinium
252.08
Fermium
257.10
3+
2+, 3+
Mendelevium
258.10
Yb
2+, 3+
Sm
3+, 4+, 5+, 6+
Tm
2+, 3+
Samarium
150.36
94
1.28
3+, 4+, 5+, 6+
Er
3+
Promethium
144.91
93
1.36
3+, 4+, 5+, 6+
Ho
3+
Neodymium
144.24
92
1.38
3+, 4+, 5+, 6+
Pm
3+
Praseodynium
140.91
91
1.50
4+, 5+
Nd
3+
Cerium
140.12
90
1.30
4+
Pr
3+, 4+
Lanthanum
138.91
89
1.10
3+
Ce
3+, 4-
17 (VIIA)
2+, 3+
Nobelium
259.10
9
Helium
4.00
3.98
F
1-
Fluorine
19.00
17
3.16
1-
10
0
Ne
Neon
20.18
18
0
Cl
Ar
Chlorine
35.45
35
2.96
Argon
39.95
Br
1-
Bromine
79.90
53
2.66
I
1-
Iodine
126.90
85
2.20
At
Astatine
(209.99)
1-
36
0
Kr
Krypton
83.80
54
0
Xe
Xenon
131.29
86
0
Rn
Radon
(221.02)