Dickinson High School AP Chemistry
Summer Assignment & Class Info 2015-16
Dear AP Chemistry Student and Parent:
Welcome to AP Chemistry! I am so excited about this coming school year. For those of you that do
not know me, my name is Staci Wagner and I will be your instructor for AP Chemistry.
Overview of Course Expectations
2016 AP Chemistry Exam Date – Monday, May 2
Course Description – AP Chemistry is a college level course. Students can earn up to 8 hours college
credit for successful completion of this course and a good score on the AP exam. Check with the
University you plan on attending to see their requirements for AP Credit.
AP Chemistry is a time-consuming and challenging, yet extremely rewarding course. This course
moves at a fast pace and classroom attendance is a MUST. Students will be prepared to do college
level work of any type upon completion of this course due to the thought processes used and the
discipline/work habits required. This course meets daily, but due to the breadth of this course (we
cover two semesters of college chemistry), you will need to spend time on your own on homework,
assignments, pre-labs, lab reports, problem sets, etc. This class will be presented as a flipped class,
meaning that lecture will be presented in a video format for you to view at home. Class time will be
devoted to labs, practice problems and concept development. These statements are not meant to
discourage, but to point out and state the truth to avoid any misconceptions about the high
expectations for this course. I will do my very best to provide a college level course/experience
which not only prepares you for the AP exam, but provides a solid foundation in chemistry. I also
intend for it to be fun! We will do several labs throughout the year. Lab will begin during SMART
lunch and continue through class about once per week.
Summer Assignment:
Attached to this letter you will find a list of things that students need to
review, memorize, and/or practice prior to school starting August 24. The majority of the
materials required in this assignment are review material that students should have learned in
their first year chemistry class. Some parts of the summer assignment are meant to stretch
student thinking and resourcefulness. (This means look it up if you don’t know it off the top of
your head.  There are many resources available to you online.) Students should not panic if
they didn’t learn it in their first year – it is easily mastered. There will be a test the first
week of school over material covered in this packet.
Before summer begins, please come by my room (S-213) with a flash drive so
that I can give you an electronic copy of your textbook. I will also be
collecting email addresses from you.
Because this is a challenging problem-solving course, and for some of you, a year may have passed
since you have had a chemistry course, it is imperative that you come to class the first day with
Dickinson High School AP Chemistry
some of the jargon, etc. second nature for you because of the pace at which this course progresses.
Reviewing and committing to memory the topics in this summer assignment is not optional, and
completing the assignment in a thorough and focused manner will contribute to a student’s success
in this course and on the AP Exam. Do not procrastinate. You will be starting AP Chemistry behind
if you do not complete this assignment and you cannot afford for that to happen. I encourage you
to begin forming study groups now to help in learning this material and to continue to meet with
those study groups throughout the year. Again, you will be tested on the summer assignment the
first week of school. Failure to complete the summer assignment will result in removal form the
class.
I look forward to getting to know each of you! We will have fun and we will work hard. Students will
receive a detailed course syllabus when school resumes in the fall. Please feel free to email me with
any questions or comments. I will be checking my email frequently during the summer months.
My email address is : swagner1@dickinsonisd.org
A great website to use as reference:
http://apchemistrynmsi.wikispaces.com/AP+Chemistry+Class+Lecture+Notes+AND+instructional+vid
eos (or just google AP Chem NMSI to get here)
I hope you enjoy your summer and I look forward to seeing you in the fall.
Sincerely,
Mrs. Wagner
Supplies needed for this course:

A sturdy 3-ring binder (2 ½ inch)

Tab dividers- Objectives, Notes, Homework, Quizzes, Tests, Labs

Pencils, pens, highlighters, etc.

Student Lab Notebook with 100 Carbonless duplicate pages – You can order these from
Amazon

Scientific calculator for completing homework. A graphing calculator will also work, but the
scientific calculator is cheaper.

Flash Drive (not required, but will come in handy for lab reports or projects when
transferring/saving files and for your other classes too!)

A great attitude and a sense of humor. You will need it at times! Plan on forming study
groups.

A valid student email address that is checked.
Dickinson High School AP Chemistry
Reasons to take AP Chemistry:
1.
AP Chemistry will challenge you to the limits of your academic ability. It's like building a muscle: if you
never push it, it will never grow!
2.
AP Chemistry will teach you to think at higher levels. You will be forced to think and apply concepts to
new situations, and even derive your own theories from application. This is excellent preparation for
the higher levels of thinking required in college.
3.
Of course, one of the most obvious benefits to this course is that when you take and pass the national
AP Chemistry Exam given in May, you could receive college credit for the course when you enroll at
many colleges and universities in the United States. This will save you both time and money.
[Some students who have passed the AP Exam elect to take first year college chemistry anyway,
where they find the material an easy review, and achieve top grades while others around them are
frustrated and struggling in a class which is too large and/or the instructor is unavailable for help! I
especially recommend this approach for students considering majoring in any type of sciences.]
4.
AP Chemistry looks great on your transcript or on a letter of recommendation. More and more of the
best colleges and universities are looking for ways that students have distinguished themselves in high
school. Being a "straight A" student no longer carries the weight it once did, and many 4.0 grade
average students are finding themselves denied entry at the college of their choice. One of the first
things admissions officers ask counselors about a potential candidate for their university is ‘did this
student take the most challenging courses available?’ Taking AP Chemistry is a way of distinguishing
yourself in high school.
5.
AP Chemistry is an intense course of study where students and the teacher really get to know each
other. It is to the student’s advantage for the teacher to know them well when they need a letter of
recommendation.
6.
As difficult as AP Chemistry is, you will find that it will never be as easy to learn Freshman Chemistry
as it is now! There are several reasons for this:
1.
High school classes are generally much smaller than college classes.
2.
Most college professors don't regard teaching Freshman Chemistry as a priority; many
concentrate on their research, and consider teaching to be an interruption and distraction to
that end.
3.
At times Freshman Chemistry is used to "weed out" students. Most colleges prefer not to
have large class sizes in their upper division courses. Therefore the grades and difficulty
level of the freshman courses are adjusted so that only small numbers of very outstanding
students will be able to move on. This can result in a large portion of students in freshman
chemistry flunking the course!
Dickinson High School AP Chemistry
Chemistry - The Flipped Classroom!
This year in AP Chemistry we are going to be doing what is called a flipped classroom.
Traditionally classes have lectures during class time and then students have very little time to
complete assignments and have to finish them for homework. The problem with this
traditional set-up is that if a student gets stuck there is usually no one at home to get them
"unstuck". In a flipped class the lectures are recorded so that students can:
1.
2.
3.
4.
Listen to the lecture as many times as they want
Pause the lecture or repeat specific sections until they understand
Take notes at their own pace!
Have discussions with their peers about the information before they come to class
Then, most all of the homework and activities are done during class time. This way if a student
gets stuck during an activity there is someone around to help them along.
My hope is that this will remove a lot of the frustrations with learning a new and difficult subject
and allow us to get to all of the fun stuff in class! I want to move the lectures, which I feel are
essential for introducing information but very inefficient for advanced learning, out of class so
that we can focus more on student-centered activities that are very efficient in getting students to
higher levels of conceptual understanding and problem solving.
How the flipped class looks to the kids:
Step 1: Watch the online video at home and take notes to understand the basics. Do this actively!
There are questions to answer and if you do not understand something try watching it a few
times. You need to come to class ready to discuss the information. Take the online quiz.
Step 2: Come to class ready to start moving to higher levels of understanding and applications
Step 3: If you get stuck and don't understand something come to tutorials....
Dickinson High School AP Chemistry
Summer Assignment
The summer assignment consists of two parts –

Part A is the material you must have memorized by the first day of school.

Part B is practice with nomenclature, balancing equations, oxidation numbers, solubility
rules, and practicing problem solving.
Part A – Material to be memorized by the 1st day of school. Note: There will be at least one
exam during the first week of school covering the topics in this assignment!!

Memorization of material is not something that will be encouraged or emphasized in this
course because this is a problem-solving course and it is impossible to memorize everything
you will be asked to do. However, memorization of some topics/rules is necessary.

Master the memorization material listed in this assignment! DO whatever it takes to
commit this information to memory for instant recall. Make flashcards, bingo games,
etc. Have your friends and family quiz you, form study groups, etc. Again, this information
needs to be second nature to you to ensure success in this course. It will not be given
to you on your AP Exam.

I will be checking these assignments the first day of class to verify you have completed
them so that you get off to a good start in this class.

If you form a study group, do not split up the work and copy from each other. That will not
help any of you. Remember that you have to know this material!!!
Information to be Memorized by the 1st day of School!
Topic/Subject Matter
Where Do I Find It? (Reference)
Specified Element Names & Symbols
The Periodic Table
(Element symbols for elements 1-38 and Ag, Cd, I, Xe,
Cs, Ba, W, Hg, Pb, Sn, Rn, Fr, U, Th, Pu, and Am
written correctly. Students should be able to locate
these elements quickly on the periodic table provided
on the exam. This table will not have the element
names written on it.
Monatomic Ions
Listed at the end of this table
Polyatomic Ions & Corresponding Acids
Listed at the end of this table
6 Strong Acids
(for all practical purposes, all others are weak)
“CBSPIN”
(and ones with multiple oxidation states – know them!)
(if you master the system for naming acids, you do not
have to memorize them, only the ions!
Strong Bases
(all others are weak, such as NH3)
Solubility Rules








hydroChloric (HCl)
hydroBromic (HBr)
Sulfuric (H2SO4)
Perchloric (HClO4)
hydroIonic (HI)
Nitric (HNO3)
Group 1 metal hydroxides (NaOH, KOH, etc)
Group 2 metal hydroxides (Ba(OH)2, Sr(OH)2,
Ca(OH)2 – these are only slightly soluble, so
strong base – others (Mg, Be, and Ra) are
insoluble, so weak base)
Listed at the end of this table
Dickinson High School AP Chemistry
** Suggestion – Make flash cards or other study tools to aid in the memorization of
the material in the table above. Adding color can also help in memorization (write
the words in color).**
Examples of some Mnemonics for Chemistry
Electrochemistry

OIL RIG (Oxidation is Loss of electrons, Reduction is Gain of electrons)

RC cOlA (Reduction at the Cathode, Oxidation at the Anode)
Diatomic Elements

BrINClHOF (say “Brinkelhof)

Help Our Needy Class Finds Brains Immediately
Polyatomic Anions:

Nick the Camel ate a Clam Supper in Pheonix
o
Nick- N with 3 consonants and 1 vowel therefore NO3-1 (nitrate)
o
Camel – C with 3 consonants and 2 vowels therefore CO3-2 (carbonate)
o
Clam – Cl with 3 consonants and 1 vowel therefore ClO3-1 (chlorate)
o
Supper – S with 4 consonants and 2 vowels therefore SO4-2 (sulfate)
o
Phoenix - P with 4 consonants and 3 vowels therefore PO4-3 (phosphate)
SI Prefixes

The Great Monarch King Henry Died By Drinking Chocolate Milk Made Near Poland
o
The – Tera
o
Drinking – deci
o
Great – Giga
o
Chocolate – centi
o
Monarch – Mega
o
Milk – milli
o
King – Kilo
o
Made – micro - µ
o
Henry – Hecto
o
Near – nano
o
Died – Deca
o
Poland - pico
o
By – base unit – m, l, g, cal, J
Dickinson High School AP Chemistry
Common Ions
Ions Usually with One Oxidation State
Li+
lithium ion
N3Na+
sodium ion
O2K+
potassium ion
S22+
Mg
magnesium ion
Se2Ca2+
calcium ion
FSr2+
strontium ion
ClBa2+
barium ion
BrAg+
silver ion
IZn2+
zinc ion
2+
Cd
cadmium ion
Ni2+
nickel ion
Al3+
aluminum ion
3+
Ga
gallium ion
Cations with more than One Oxidation State
+1
1+
Cu
copper (I) or cuprous ion
Cu2+
Hg22+
mercury (I) or mercurous ion
Hg2+
+2
2+
Fe
iron (II) or ferrous ion
Fe3+
2+
Cr
chromium (II) or chromous ion
Cr3+
Mn2+
manganese (II) or manganous ion
Mn3+
Co2+
cobalt (II) or cobaltous ion
Co3+
+2
2+
Sn
tin (II) or stannous ion
Sn4+
Pb2+
lead (II) or plumbous ion
Pb4+
nitride
oxide
sulfide
selenide
fluoride
chloride
bromide
iodide
+2
copper (II) or cupric ion
mercury (II) or mercuric ion
+3
iron (III) or ferric ion
chromium (III) or chromic ion
manganese (III) or manganic ion
cobalt (III) or cobaltic ion
+4
tin (IV) or stannic ion
lead (IV) or plumbic ion
Dickinson High School AP Chemistry
Polyatomic Ions and Acids
Formula
Name
H2SO4
sulfuric acid
H2SO3
sulfurous acid
HNO3
nitric acid
HNO2
nitrous acid
H3PO4
phosphoric acid
H2CO3
carbonic acid
HMnO4
permanganic acid
HCN
hydrocyanic acid
HOCN
cyanic acid
HSCN
thiocyanic acid
HC2H3O2
acetic acid
H2C2O4
oxalic acid
H2CrO4
chromic acid
H2Cr2O7
dichromic acid
H2S2O3
thiosulfuric acid
H3AsO4
arsenic acid
H3AsO3
arsenous acid
HBrO3
bromic acid
Ion
SO42SO32NO31NO21PO43CO32MnO41CN1OCN1SCN1C2H3O21C2O42CrO42Cr2O72S2O32AsO43AsO33BrO31-
Ion Name
Sulfate ion
Sulfite ion
Nitrate ion
Nitrite ion
Phosphate ion
Carbonate ion
Permanganate ion
Cyanide ion
Cyanate ion
Thiocyanate ion
Acetate ion
Oxalate ion
Chromate ion
Dichromate ion
Thiosulfate ion
Arsenate ion
Arsenite ion
Bromate
Oxyhalogen Anids
Formula
Oxy name
Ion
Ion name
1HClO
Hypochlorous acid
ClO
Hypochlorite ion
1HClO2
Chlorous acid
ClO2
Chlorite ion
1HClO3
Chloric acid
ClO3
Chlorate ion
1HClO4
Perchloric acid
ClO4
Perchlorate ion
Br, I, can be substituted for Cl. F may form hypofluorous acid and the hypofluorite ion.
*Reminder – NH3 is ammonia
Dickinson High School AP Chemistry
Other Ions
Ion
Ion Name
2O2
Peroxide ion
1OH
Hydroxide ion
1HSO4
Bisulfate ion; hydrogen sulfate ion
1+
NH4
Ammonium ion
O21Superoxide ion
1HCO3
Bicarbonate ion; hydrogen carbonate ion
2HPO4
Hydrogen phosphate ion
1H2PO4
Dihydrogen phosphate ion
1N3
Azide ion
Colors of Common Ions in Aqueous Solution
Most common ions are colorless in solution. However, some have distinctive colors. These
colors have appeared in questions on the AP exam.
Fe 2+ and Fe3+
Various colors
2+
Cu
Blue to green
2+
Cr
Blue
3+
Cr
Green or violet
2+
Mn
Faint pink
2+
Ni
Green
Co2+
Pink
1MnO4
Dark purple
2CrO4
Yellow
2Cr2O7
Orange
Dickinson High School AP Chemistry
Solubility Rules:
Learn the following solubility rules:
Salts containing the following ions are normally soluble:
o All salts of Group IA (Li+, Na+, etc.) and the ammonium ion (NH4+) are soluble.
o All salts containing nitrate (NO31-), acetate (CH3COO1-) and percholates (ClO41-) are
soluble
o All chlorides (Cl1-), bromides (Br1-), and iodides (I1-) are soluble except those of Cu+,
Ag+, Pb2+, and Hg22+.
o All salts containing (SO42-, are soluble except those of Pb2+, Ca2+, Sr2+, and Ba2+).
Salts containing the following ions are normally insoluble:
o Most carbonates (CO32-) and phosphates (PO43-) are insoluble except those of
Group IA and the ammonium ion.
o Most sulfides (S2-) are insoluble except those of Group IA and IIA and the
ammonium ion.
o Most hydroxides (OH1-) are insoluble except those of Group IA, calcium, and
barium.
o Most oxides (O2-) are insoluble except for those of Group IA and IIA which react
with water to form the corresponding hydroxides
SI Base Units & Prefixes
Be sure to refamiliarize yourself with the SI pre-fixes for measurement pico (10-12)
through Tera (1012).
SI Unit for Length = meter; for Mass = kg; for volume = m3(1 liter)
Dickinson High School AP Chemistry
Summer Assignment – Part B: Practice with Nomenclature, Balancing Equations, Oxidation
Numbers, Solubility Rules and Problem Solving

Nomenclature: Simple Inorganic Formulas and Nomenclature – Complete Exercise 1 in
the Appendix of this packet. Review the naming rules and commit the naming prefixes
to memory!

More Nomenclature: Ternary Nomenclature: Acids and Salts – Complete Exercise 2
in the Appendix of this packet.

Balancing Equations: Balancing Molecular Equations – Complete Exercise 3 in the
Appendix of this packet.

Solubility Rules – Using Solubility Rules Table – You must Memorize these! Complete
Exercise 4 in the Appendix of this packet using the solubility rules listed in Part A of
the Summer Assignment.

Reaction Types and Reaction Prediction – Predict the products, write the equation
and then balance, etc.-Complete Exercise 5 in the Appendix of this packet.

Review Problem Set – Complete the problems on the attached Review Problem Set –
show your work and clearly mark your final answers!! Use correct significant figures for
math answers and units where needed!
Arrive at Dickinson High School with all of these items completed and you will be well on
your way to a terrific year in AP Chemistry!
Again, do not hesitate to contact me if you have questions!! 
Dickinson High School AP Chemistry
Nomenclature Review
Forming binary ionic compounds
A. In a binary ionic compound the total positive charges must equal the total negative
charges. The best way to write correct formula units for ionic compounds is to use the
“Criss-cross method”.
B. Sample Problem: What ionic compound would form when calcium ions combine with
bromide ion?
Steps to the Criss-cross Method:
1. Write the ions with their charges, cations are always first.
Ca2+ Br1-
2. Cross over the charges by using the absolute value of each ion’s charge as the
subscript for the other ion.
Ca1Br2
3. Check to make sure the subscripts are in the lowest whole number ratio possible.
Then write the formula.
CaBr2
Naming binary ionic compounds
A. Combine the names of the cation and the anion, and change the second half of the anion
to -ide.
B. Example: BaBr2 is named barium bromide
Naming binary ionic compounds that contain polyatomic ions.
A. The polyatomic ions on your common ion list should be memorized.
B. The most common oxyanions (polyatomic ions that contain oxygen) end in –ate.
Oxyanions with one less oxygen end in –ite. For example:
NO31- is nitrate
SO42- is sulfate
NO21- is nitrite
SO32- is sulfite
C. Anions with one less oxygen than the –ite ion are given the prefix hypo-.
D. Anions with one more oxygen than the –ate ion are given the prefix per-.
ClO1- is hypochlorite
ClO31- is chlorate
ClO21- is chlorite
ClO41- is perchlorate
E. Naming compounds with polyatomics is the same as naming other compounds, just name
the cation and then the anion. If there is a transition metal involved, be sure to check
the charges to identify which ion (+1, +2, +3, +4…) it may be so that you can put the
correct Roman numeral in the name.
Dickinson High School AP Chemistry
Polyatomic ions ending in –ate
BO33-
2-
NO31PO43AsO43-
CO3
SiO44-
O
SO42SeO42TeO42-
F
ClO31BrO31IO31-
Notes on Observations:

The individual locations of the elements in the table correspond to their relative
locations on the periodic table

The “legs” - grey shaded areas all end in “O3”

The “interior” – lighter grey shaded areas all end in “O4”

The charges of the ions become more positive as you go across a “period”

For ions with the same root containing oxygen, the suffixes and prefixes are:
(Using chlorate as an example)
o
Ions starting with “per” will have one more oxygen. Ex. ClO41-= perchlorate
o
Ions ending with “ite” will have one less oxygen. Ex. ClO21- = chlorite
o
Ions starting with “hypo” and ending in “ite” will have two less oxygens.
Ex. ClO1-. = hypochlorite
Naming binary molecular formulas
A. With molecules, the prefix system is used.
Number
Prefix
Number
Prefix
1
mono7
hepta2
di8
octa3
tri9
nona4
tetra10
deca
5
penta 11
undeca6
hexa 12
dodecaB. The less-electronegative element is always written first. It only gets a prefix if it has
more than one atom in the molecule.
C. The second element gets the prefix and the ending –ide.
D. The o or a at the end of the prefix is dropped when the word following the prefix
begins with another vowel, for example monoxide or pentoxide.
Dickinson High School AP Chemistry
Exercise 1 – Nomenclature: Simple Inorganic Formulas and Nomenclature
I. In the first column, classify each of the following as molecular (covalent) (M) or ionic (I). In
the second column, name each compound:
M or
I
Name
M or
I
1) CaF2
10)SrI2
2)P4O10
11)CO
3)K2S
12)Cs2Po
4)NaH
13)ZnAt2
5)Al2Se3
14)P2S3
6)N2O
15)AgCl
7)O2F
16)Na3N
8)SBr6
17)Mg3P2
9)Li2Te
18)XeF6
Name
II. In the first column, write the chemical formula (formula unit) for the compound formed
between the two given elements. In the second column, write the name for the compound:
1
Elements
Magnesium and iodine
2
Potassium and sulfur
3
Chlorine and sulfur
4
Zinc and bromine
5
Strontium and oxygen
6
Calcium and nitrogen
7
Calcium and oxygen
8
Copper (I) and oxygen
9
Copper (II) and chlorine
10
Mercury (II) and oxygen
11
Nitrogen and aluminum
12
Sulfur and cesium
Formula Unit
Name
Dickinson High School AP Chemistry
Exercise 2 – More Nomenclature – Including some ternary nomenclature: Acids and Salts
I. Name the following substances:
Formula
1. FeSO3
Name
Formula
16. Fe2O3
2. Cu(NO3)2
17. (NH4)2SO3
3. Hg2Cl2
18. Ca(MnO4)2
4. AgBr
19. PF5
5. KClO3
20. LiH
6. MgCO3
21. HIO3
7. BaO2
22. NaBrO2
8. KO2
23. Ca3(PO4)2
9. SnO2
24. HIO4
10. Ni3(PO4)2
25. Fe(IO2)3
11. Pb(OH)2
26. HAt(aq)
12. CuCH3COO
27. C6H5COOH
13. N2O4
28. Hg2(IO)2
14. Rb3P
29. H3PO4
15. S8
30. NH4BrO3
Name
II. Write formulas for the following substances:
Name
1. vanadium (V) oxide
Formula
Name
16. francium dichromate
2. dihydrogen monoxide
17. calcium carbide
3. ammonium oxalate
18. Mercury (I) nitrate
4. polonium (VI) thiocyanate
19. cerium (IV) benzoate
5. tetraphosphorous decoxide
20. potassium hydrogen phthlate
6. zinc hydroxide
21. carbonic acid
7. potassium cyanide
22. calcium hypochlorite
8. cesium thiosulfate
23. hydrotelluric acid
9. oxygen molecule
24. copper (II) nitrite
10. mercury (II) acetate
25. nitrous acid
11. silver chromate
26. hypoiodous acid
12. tin (II) carbonate
27. cyanic acid
13. manganese (VII) oxide
28. phthalic acid
14. sodium hydrogen carbonate
29. tin (IV) chromate
15. Copper (II) dihydrogen
phosphate
30. hydrocyanic acid
Formula
Dickinson High School AP Chemistry
III. Practice with acids! Remember:
-IC from –ATE
Complete the following table:
Name of Acid
Hydrochloric acid
Sulfuric acid
-OUS from –ITE
Formula of Acid
HCl
H2SO4
HI
HYDRO-,-IC from -IDE
Name of Anion
Chloride
Sulfate
Sulfite
Chlorous acid
Nitrate
HC2H3O2 or CH3COOH
Hydrobromic acid
Sulfide
HNO2
Chromic acid
Phosphate
Exercise 3 – Balancing Equations
I. Balance the following equations by adding coefficients as needed. Some equations may already be balanced.
1. ___C6H6+___O2 →___ H2O+ ___CO2
2.___NaI+___Pb(SO4)2→___PbI4+___Na2SO4
3. ___NH3+___O2 → ___NO+ ___H2O
4. ___HNO3+___Mg(OH)2 → ___H2O +___Mg(NO3)2
5. ___H3PO4+___NaBr
→___HBr + Na3PO4
6. ___CaO+___MnI4→ ___MnO2+___CaI2
7. ___C2H2+___H2
→___C2H6
8. ___VF5+ ___HI → ___V2I10+___HF
9. ___OsO4+___PtCl4
→ ___PtO2 + ___OsCl8
10. ___Hg2I2+___O2
→___Hg2O + ___I2
Dickinson High School AP Chemistry
Exercise 4 – Solubility Rules
For the compounds in the table, write the formula for each compound in the first column and
then use the solubility rules to determine if each compound is soluble or insoluble in water. In
the second column, write and (S) for those that are soluble and an (I) for those that are
insoluble.
Name
Formula
Silver nitrate
cobalt (II) sulfate
Zinc hydroxide
Iron (III) iodide
Nickel (II) chloride
Lead (II) iodide
Sodium carbonate
Barium sulfate
Lead (II) sulfide
Silver phosphate
Lithium phosphate
Nickel (II) carbonate
Copper (II) hydroxide
Tin (IV) sulfate
Lead (II) nitrate
Exercise 5 – Reaction Prediction Practice
I. Predict the products, write the equation and then balance.
Combustion
1. C4H9OH + oxygen →
2. C7H14+ oxygen →
Synthesis
1. Sodium + oxygen→
2. calcium + nitrogen→
3. potassium + bromine→
(S) or (I)
Dickinson High School AP Chemistry
Decomposition
1. Strontium carbonate→
2. Mercury (II) oxide→
3. Aluminum chlorate→
Double Replacement
1. Iron (III) sulfate + calcium hydroxide→
2. Sodium hydroxide + sulfuric acid→
3. sodium sulfide + manganese (VI) acetate→
4. chromium (III) bromide + sodium sulfite→
5. barium hydroxide + chlorous acid
→
Single Replacement
Use the activity series in this packet (or online) to complete and balance these equations. If no
reaction occurs, write NR.
1. Nickel + steam
2. chlorine + aluminum iodide→
3. potassium + water→
4. lead + copper (II) chloride→
5. zinc + hydrochloric acid
→
Dickinson High School AP Chemistry
Chapter 1
1. For each of the following glassware, discuss the number of significant figures and
uncertainty for each.
a. Beaker
c. Buret
b. Graduated cylinder
d. Volumetric flask
2. A student performed an analysis of a sample for its calcium content and got the following
results: 14.92%, 14.91%, 14.88%, and 14.91%
The actual amount of calcium in the sample is 15.70%. What conclusion can you draw about the
accuracy and precision of these results?
3. How many significant figures are in each of the following?
a. 12
f. 0.0000101
b.1098
g. 1000.
c. 2001
h. 22.04030
d. 2.001 x 103
i. 1.00 x103
e. 100
4. Round each of the following numbers to two significant figures, and write the answers in
scientific notation.
a. 0.00031254
c. 35900
b. 31254000
d. 0.00000399
5. Use scientific notation to express the number 480 to
a. One significant figure
b. Two significant figures
c. Three significant figures
6. Calculate the percent error for the following measurements.
a. The density of an aluminum block determined in an experiment was 2.64 g/cm3.
(Accepted value = 2.70 g/cm3)
b. The experimental determination of iron in ore was 16.48%. (Accepted value was 16.12%)
7. Perform the following mathematical operations, and express each result to the correct
number of significant figures.
a. 97.381 + 4.2502 + 0.99195
c. 0.102 * 0.0821 * 273.5
b. 171.5 + 72.915 – 8.23
1.2
d. (9.04 – 8.23+ 21.954 + 81.0)/3.1416
Dickinson High School AP Chemistry
8. Precious metals and gems are measured in troy weights in the English system:
24 grains = 1 pennyweight
20 pennyweights = 1 troy ounce
12 troy ounces = 1 troy pound
1 grain = 0.0648 gram
1 carat = 0.200 gram
a. Diamonds are measured in carats. If a lucky girl receives a 5 carat diamond how many
pennyweights is it?
b. What is the mass of 2.3 troy ounces of gold in grams?
c. The density of gold is 19.3 g/cm3. What is the volume of a troy pound of gold?
9. Apothecaries (druggists) use the following set of measures:
20 grains ap = 1 scruple
3 scruples = 1 dram ap
8 dram ap = 1 oz. ap
1 dram ap= 3.888 g
a. An aspirin tablet contains 5.00 x 102 mg of active ingredients. How many grains ap of
active ingredients does it contain?
b. From (a), how many scruples?
c. What is the mass of 1.00 scruples in grams?
10. The world record for the hundred meter dash is 9.79s. What is the corresponding speed in
units of m/s, km/hr, ft/s, and mi/hr?
a. At this speed how long would it take to run a mile (5,820 ft)?
11. You’re planning to buy a new car. One model that you’re considering gets 32 miles per gallon
of gasoline in highway travel. The one that your spouse likes gets 14 kilometers to the liter.
Which car has the better gas mileage? (1 gal = 4 qt., 1.057 qt = 1L)
12. You pass a road sign saying “New York – 112km.” If you drive at a constant speed of 65
mi/hr., how long should it take you to reach New York?
a. If your car gets 28 miles to gallon, how many liters of gasoline are necessary to travel
112 km?
13. You have a 1.0 cm3 sample of lead and a 1.0 cm3 sample of glass. You drop each in a separate
beakers of water. How do the volumes of water displaced by each sample compare? Explain.
Density of lead = 11.35g/ cm3
Denisty of glass = 3.00 g/ cm3
14. A person has a temperature of 102.5 F. What is the temperature on the Celsius scale?
Kelvin scale?
Dickinson High School AP Chemistry
15. Convert the following Celsius temperatures to Kelvin
a. The boiling point of ethyl alcohol, 78.1 C
b. A cold winter day, -25 C
c. The lowest possible temperature, -273 C
d. The melting point of sodium chloride, 801 C
16. The density of diamond is 3.51 g/ cm3. What is the volume of a 4.5 carat diamond? 1 carat =
0.200g
17. The volume of a diamond is found to be 2.8 mL. What is the mass of the diamond in carats?
18. A sample containing 33.42 g of metal pellets is poured into a graduated cylinder initially
containing 12.7mL of water, causing the water level in the cylinder to rise to 21.6 mL. Calculate
the density of the metal.
19. Two spherical objects have the same mass. One floats on water; the other sinks. Which
object has the greater diameter? Explain you answer.
20. What are some of the differences between a solid, a liquid, and a gas?
21. What is the difference between homogeneous or heterogeneous matter.
22. Classify each of the following as homogeneous or heterogeneous.
a. soil
d. gasoline
b. the atmosphere
e. gold
c. a carbonated soft drink
f. a solution of ethanol and water
23. Classify each of the following as a mixture or a pure substance. Of the pure substances,
which are elements and which are compounds?
a. water
f. uranium
b. blood
g. wine
c. the oceans
h. leather
d. iron
i. table salt (NaCl)
e. brass
24. Distinguish between physical and chemical changes.
25. List four indications that a chemical change (reaction) has occurred.
Dickinson High School AP Chemistry
26. If you place a glass rod over a burning candle, the glass appears to turn black. What is
happening to each of the following (physical change, chemical change, both, or neither) as the
candle burns? Explain each answer.
a. the wax
b. the wick
c. the glass rod
27. The properties of a mixture are typically averages of the properties of its components.
The properties of a compound may differ dramatically from the properties of the elements
that combine to produce the compound. For each process described below, state whether the
material being discussed is most likely a mixture or a compound, and state whether the process
is a chemical change or a physical change.
a. An orange liquid is distilled, resulting in the collection of a yellow liquid and a red solid.
b. A colorless, crystalline solid is decomposed, yielding a pale yellow-green gas and a soft,
shiny metal.
c. A cup of tea becomes sweeter as sugar is added to it.
Chapter 2
1. Describe Dalton’s atomic theory.
2. What discoveries were made by J. J. Thomson, Henri Becquerel, and Lord Rutherford? How
did Dalton’s model of the atom have to be modified to account for these discoveries?
3. What is the distinction between atomic number and mass number?
4. What is the difference between atomic mass and average atomic mass?
5. What is an isotope?
6. How many protons and neutrons are contained in the nucleus of each of the following atoms?
a. 22Ti42
d. 36Kr86
b. 30Zn64
e. 33As75
c. 32Ge76
f. 19K41
7. Write the isotopic symbol for each of the isotopes below.
a. Atomic number = 8, number of neutrons = 9
b. The isotope of chlorine in which the mass = 37
c. Atomic number = 27, mass = 60
d. Number of protons = 26, number of neutrons = 31
e. The isotope of I with a mass number of 131
f. Atomic number = 3, number of neutrons = 4
Dickinson High School AP Chemistry
8. The element copper has naturally occurring isotopes with mass number of 63 and 65. The
relative abundance of the isotopes are 69.2% for mass = 62.93 amu, and 30.8% for mass =
64.93 amu. Calculate the average atomic mass of copper.
9. An element consists of 1.40% of an isotope with mass 203.973 amu, 24.10% of an isotope
with the mass 205.9745, 22.10% of an isotope with the mass 206.9759 amu, and 52.40% of an
isotope with the mass 207.9766 amu. Calculate the average atomic mass and identify the
element.
10. Distinguish between the terms family and period in connection to the period table. For
which of these terms is the term group also used?
11. In the periodic table, what is the name of the following groups?
a. Group (2)
b. Group (18)
12. An ion contains 50 protons, 68 neutrons, and 48 electrons. What is its symbol and charge?
13. Which of the following sets of elements are all in the same group on the periodic table?
a. N, P, O
c. Rb, Sn
b. C, Si, Ge
d. Mg, Ca
14. Identify each of the following elements:
a. A member of the same family as Oxygen whose most stable ion contains 54 electrons
b. A member of the alkali metal family whose most stable ion contains 36 electrons
c. A noble gas with 18 protons in the nucleus
d. A halogen with 85 protons and 85 electrons
15. Would you expect each of the following atoms to gain or lose electrons when forming ions?
What ion is the most likely in each case?
a. Na
d. Ba
g. Al
b. Sr
e. I
h. S
c. P
f. O
16. For each of the following ions, indicate the total number of protons and electrons in the
ion. For the positive ions, predict the formula of the simplest compound formed between itself
and oxide. For the negative ions predict the simplest compound formed between itself and
Aluminum.
a. Fe2+
d. Cs1+
g. Br13+
2b. Fe
e. S
h. N3c. Ba2+
f. P3-
Dickinson High School AP Chemistry
17. An element’s most stable ion forms an ionic compound with bromine, having the formula
XBr2. If the ion of element X has a mass number of 230 and 86 electrons, what is the identity
of the element, and how many neutrons does it have?
Chapter 3
1. The molecular formula of aspartame, the artificial sweetener marketed as NutraSweet, is
C14H18N2O5.
a. What is the molar mass of aspartame?
b. How many moles of aspartame are present in 3769.4 grams of aspartame?
c. How many molecules of aspartame are present in 345.9 grams of aspartame?
d. How many oxygen atoms are present in 23.6 grams of aspartame?
2. How many moles of ammonium ions are in 0.557 g of ammonium carbonate?
3. What is the mass, in grams, of 0.0438 moles of iron (III) phosphate?
4. What is the mass, in grams, of 2.69 x 1023 molecules of aspirin, C9H8O4?
5. What is the molar mass of diazepam (Valium) if 0.05570 mol has a mass of 15.86 g?
6. Determine the empirical formulas of the following compounds.
a. 10.4% C, 27.8% S, and 61.7% Cl.
b. Monosodium glutamate (MSG), a flavor enhancer on certain foods, 35.51 g C, 4.77 g H,
37.85 g O, 8.29 g N, and 13.60 g Na.
7. Find the molecular formulas of the following compounds.
a. 73.8% carbon, 8.7% hydrogen, 17.5% nitrogen, molar mass = 166.0 g/mol
b. 80.0 % carbon, 20.0 % hydrogen, molar mass = 30.0 g/mol.
8. 4 FeCr2O7 + 8 K2CO3 + O2 → 2 Fe2O3 + 8 K2CrO4 + 8 CO2
a. How many grams of FeCr2O7 are required to produce 44.0 g of CO2?
b. How many grams of O2 are required to produce 100.0 g of Fe2O3?
c. If 300.0 g of FeCr2O7 react, how many grams of O2 will be consumed?
d. How many grams of Fe2O3 will be produced from 300.0 g of FeCr2O7?
e. How many grams of K2CrO4 are formed per gram of K2CO3used?
9. Given the reaction S + O2→SO2
a. How many grams of sulfur must be burned to give 100.0g of SO2?
b. How many grams of oxygen must be required for the reaction in part (a)?
Dickinson High School AP Chemistry
10. 6 NaOH + 2Al → 2 Na3AlO3 + 3 H2
a. How much aluminum is required to produce 17.5 g of hydrogen?
b. How much Na3AlO3 can be formed from 165.0 g of sodium hydroxide?
c. How many moles of NaOH are required to produce 3 g of hydrogen?
d. How many moles of hydrogen can be prepared from 1g of aluminum?
11. The following unblalanced reaction takes place at high temperatures.
Cr2O3(s) +
Al(l) →
Cr(l) +
Al2O3(l)
If 42.7 g Cr2O3 and 9.8 g Al are mixed and reacted until one of the reactants is used up.
a. Which reactant will be left over?
b. How much will be left?
c. How many grams of chromium will be formed?
12. Calculate the mass of water produced when 42.0 g of propane, C3H8, is burned with 115g of
oxygen.
Congratulations!!! You have made it! Be proud of yourself, and get ready for a fun-filled and
challenging year which will push you to your limits, but make you a better student, get you very
prepared for college, and prove to yourself how very brilliant you are!
Remember, I am on your side, and just want to help! I am trying to give you the tools to
succeed. If you need anything, please do not hesitate to email me!
See you in the fall!
Mrs. Wagner
Dickinson High School AP Chemistry
Some preliminary notes from Chapters 1-3
Chapter 1 – Introduction: Matter and Measurement
A. Classification of Matter
1.
States of Matter
a.
b.
Gas(vapor)
i.
Has no fixed volume or shape
ii.
Takes the shape of its container
iii.
Can be compressed or expanded
iv.
Molecules are far apart and moving at high speeds
Liquid
i.
Definite volume, cannot be compressed
ii.
Takes the shape of its container
iii.
Molecules are much closer than in a gas but still move rapidly (they can slide past each
other)
c.
2.
Solid
i.
Definite shape and volume, cannot be compressed
ii.
Molecules are held tightly together, typically in definite arrangements
Pure substances and mixtures
a.
Pure substances – matter that has a fixed composition and distinct properties
i.
Two types
1.
Elements – substances that cannot be decomposed into simpler substances
2.
Compounds – composed of two or more elements chemically bonded together
a.
Law of Constant Composition – (Joseph Proust) the makeup of compounds
is always the same
b.
Mixtures – combination of two or more substances in which each substance retains its own chemical
identity and properties
i.
Properties can vary
1.
Example – adding sugar to coffee is a mixture, you can make it very sweet, add a
little, or none at all.
ii.
Two types
1.
Heterogeneous – different composition throughout
a.
2.
Homogeneous (aka solutions) – uniform composition throughout
a.
c.
Rocks, sand, wood, chocolate chip cookies
Air(gaseous solution), gasoline(liquid solution), brass(solid solution)
Separation of Mixtures
i.
Filtration – separating a solid component from a liquid component using a funnel, filter
paper, and gravity
ii.
Distillation – separating liquid components utilizing different boiling points
iii.
Chromatography – separating substances by how they adhere to surfaces ( used frequently
for ink)
3.
Properties of Matter
a.
Physical properties – description of what something looks like
i.
b.
Color, odor, density, melting point, boiling point, hardness
Chemical properties – how a chemical reacts with other chemicals
Dickinson High School AP Chemistry
i.
4.
Flammability, reactivity with other chemicals
Changes in Matter
a.
Physical changes – physical appearance is changed
i.
b.
Ripping paper, melting wax, ALL CHANGES OF STATE (BOILING, EVAPORATING)
Chemical changes (reactions) – chemically transformed into a new substance\
i.
Sodium metal reacts with chlorine gas to form salt
B. Units of Measurement
1. Metric System / Significant Figures / Dimensional Analysis
a. you should ALREADY know this
Chapter 2 – Atoms, Molecules, Ions
A. The Atomic Theory of Matter
1. History of the Atom
a.
Democritus – first person to speculate that matter was mass of atoms. Greek philosopher
i.
Plato and Aristotle refuted this idea, atomic theory faded for many centuries
b. John Dalton – came up with first atomic theory, English school teacher
i.
Each element is composed of extremely small particles called atoms
ii.
All atoms of a given element are identical; the atoms of different elements are different
and have different properties (including different masses)
iii.
Atoms of an element are not changes into different types of atoms by chemical reactions;
atoms are neither created nor destroyed in chemical reactions
iv.
Compounds are formed when atoms of more than one element combine; a given compound
always has the same relative number and kind of atom.
Dalton thought that atoms could not be broken down any further, this was expressed in the atomic
model – Billiard Ball Model
Laws from the time period
a. Law of Constant Composition
b. Law of Conservation of Mass (LeChatelier) – matter and energy cannot be created or
destroyed
c. Law of Multiple Proportions – if elements combine to form more than one compound they
must be different by whole numbers.
i. Carbon monoxide, CO, carbon dioxide, CO2
c. Cathode Rays – a high voltage electricity passed through partially evacuated tubes produced radiation
and mass glass fluoresce, called cathode rays because they originated from the cathode
i. Rays were deflected by electric and magnetic fields, suggesting that the rays were charged
ii. J.J. Thomson – observed that the rays were the same no matter what type of material was used,
concluded that the rays were actually particles with mass, these particles were called electro
1.
Able to calculate the charge to mass ration of an electron, 1.76 x 108 Coulombs/gram
2.
Came up with second atomic model – Plum Pudding Model
d. Robert Millikan – performed the oil drop experiment and determined the charge of an electron (1.60 x
10-19) and then determined the mass of an electron (9.11 x 10-28g)
e. Henri Becquerel – studied an ore of Uranium called pitchblende and discovered the spontaneous emission
of radiation called radioactivity
i. Marie Curie and her husband, Pierre also studied this
Dickinson High School AP Chemistry
f. Ernest Rutherford – studied radiation and discovered three types of radiation: alpha, beta, and gamma
i. Utilizing alpha particles, Rutherford performed the Gold Foil Experiment and determined that
the atom had a nucleus
ii. Also discovered protons
g. James Chadwick – discovered neutrons
2. Modern View of Atomic Structure
a. Atoms are made of protons, neutrons, and electrons
b. Electronic charge is measured in Coulombs (C)
i. Electrons have a charge of -1.60 x 10-19 C
ii. Protons have a charge of +1.60 x 10-19 C
iii. For simplicity we change this to +1 and -1, but you should still know what the value is
iv. Neutrons have no charge
3.
c.
Atoms are typically neutral, which means they have the same number of protons and electrons
d.
Protons and neutrons are in the nucleus, electrons circle around
e.
Vast majority of an atom’s volume is the space where the electrons are found
f.
Isotopes – atoms of a given element that differ in the number of neutrons
g.
Protons – all atoms of an element have the same number of protons in the nucleus, aka atomic number
h.
Mass number – number of protons + number of neutrons
Periodic Table
a.
You should know the general layout of the periodic table (groups, rows, where the metals, nonmetals,
and metalloids are)
4.
Writing chemical formulas (reviewed earlier in packet)
Chapter 3 – Stoichiometry: Calculations with Chemical Formulas and Equations
1.
All chemical equations need to be written correctly and balanced appropriately (kind of redundant I know)
2.
We will go over all of the types of chemical reactivity but below are some for review
3.
a.
Most common involve oxygen as a reactant
b.
Often involve hydrocarbons (compounds that contain hydrogen and carbon)
Atomic and Molecular Weight
a.
Atomic Mass Scale – is based off of Carbon – 12, mass of carbon – 12 = 12 amu
b.
Amu = atomic mass unit, 1g = 6.022 x 1023 amu
4. Average Atomic Masses
a. the masses listed on the periodic table are weighted averages based on the abundance in nature
b. see example problems in book
5. Percent Composition from Formulas
a. part/whole x 100%
b. used to determine how much of a compound is a particular kind of element
6. The Mole
a. used to convert between the microscopic and the macroscopic
b. Avogadro’s number = 6.02 x 1023
7. Problems – work through chapter problems if you need extra help