Dear AP Chemistry Student and Parents,
Welcome to AP Chemistry! I applaud you for wanting to challenge yourself and am very excited about embarking upon this
rewarding journey with you.
Important information to ensure your success in AP Chemistry:
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This is a College level Course so students can earn up to 8 hours college credit upon successful completion of this
course and a score of 3-5 (varies by college) on the AP Exam.
Students will be expected to work independently and utilize their resources. Problem-solving skills are emphasized in
this curriculum and will be developed throughout the year.
This is a challenging, fun and extremely rewarding class. The pace is fast and classroom attendance is imperative,
especially on lab days. Labs will be on the same day each week to facilitate scheduling of appointments.
To be successful on the AP Exam, students must be prepared to spend an average of 5-8 additional hours per week,
outside the classroom, working on AP Chemistry. This time will be spent on homework assignments, lab reports,
problem solving, etc. If students use their time efficiently, many tasks can be accomplished during class.
I will do my very best to provide a college course experience, which not only prepares you for the AP Exam but also
provides a solid foundation in Chemistry.
It may have been almost a year since some of you completed Chemistry. The fast pace of this class will not allow a lot
of time for review of basic concepts so attached is a Summer assignment packet. This will facilitate your review of
material you have already learned in Chemistry. So all will turn up on the first day of class totally fluent in the
language of Chemistry!
The SUMMER ASSIGNMENT is MANDATORY. You might get bored without elements and ions floating around your
heads!
Some parts of the Summer Assignment are designed to test your resourcefulness. Translation: Look it up if you don’t
know it! You have so many available resources. Students should not worry if their prior Chemistry course did not
cover some of the material – it is easily mastered.
The Summer Assignment will be taken up for a grade a WEEK AFTER SCHOOL starts.
Please read the following statement carefully: My signature below indicates that I am aware of the mandatory Summer
Assignment and the rigorous course expectations for AP Chemistry. I understand the high expectations regarding
coursework and attendance. I am aware that I will need to spend 1 – 1.5 hours per night doing homework and preparing
for the AP Course and exam
Best wishes for a safe and relaxing summer,
Mrs. Milam
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Student Name (Print)
Student Signature
Parent Name (Print)
Parent Signature
Student e-mail
Date
Parent e-mail
Date
Please return to Mrs. Milam, in Room 604, by Friday, May 20th
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AP CHEMISTRY SUMMER ASSIGNMENT
The Summer Assignments consists of two parts:
Part A: Material to be memorized by the second day of school
Part B: A packet to be turned in the second day of school, which involves review, and practice in:
 Nomenclature
 Solubility rules
 Balancing equations
 Problem-solving
 Oxidation numbers
PART A – MATERIAL TO BE MEMORIZED
There will be a test the second week of school on the concepts below.
Memorization is not encouraged in this course, as it is a problem-solving course. However, certain basic topics and rules
need to be memorized to facilitate instant recall in quickly solving problems on the AP exam in May.
Tips on Memorization:
 Make flashcards, bingo games, etc.
 Use quizlet.com - a flashcard website. It makes memorization fun by offering games and quizzes. AP
Chemistry review material can be accessed at: http://quizlet.com/subject/ap-chemistry
 Have your friends and family quiz you or form study groups with other students in the class.
 Be able to quickly recall the information below for success in this class.
Information to be Memorized by the 2nd week of School
Topic
Where do I find this information?
Element name & symbols:
The Periodic Table
Element symbols 1-40 along with Ag, Au, Cd, I, Xe, Cs, Ba,
W, Hg, Pb, Sb, Sn, Rn, Fr, U, Th, Pu, and Am
You should be able to locate elements quickly on the
periodic table. The table used on the AP exam does NOT
include element names
1. Monatomic ions
Listed at the end of this table
(& ones with multiple oxidation states)
2. Polyatomic ions and corresponding acids
Listed at the end of this table
(if you master the system for naming acids, you do not
have to memorize them)
3. Six Strong Acids
“CBSPIN”
(for practical purposes, all others are weak acids)
 Hydrochloric acid (HCl)
 Hydrobromic acid (HBr)
 Sulfuric Acid (H2SO4)
 Perchloric acid (HClO4)
 Hydroiodic acid (HI)
 Nitric acid (HNO3)
4. Strong bases
 Group 1 metal hydroxides (NaOH, KOH, etc.)
(all others are weak, such as NH3)
 Group 2 metal hydroxides (Ba(OH)2, Sr(OH)2)
5. Solubility rules
Listed at the end of this table
6. Intermolecular Forces
Listed at the end of this table
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1. & 2.Common Ions:
Positive ions (cations)
Negative ions (anions)
+1 Charge
-1 Charge
Ammonium (NH4+)
Copper (I) or cuprous (Cu+)
Hydrogen (H+) or “proton”
Hydronium ion (H3O+)
Silver (Ag+)
Group 1 (Li +, Na+, K+, Rb+, Cs+, Fr+)
Acetate (C2H3O2-)
Cyanide (CN-1)
dihydrogen phosphate (H2PO4-)
Hydrogen carbonate or bicarbonate (HCO3-)
Hydrogen sulfate or bisulfate (HSO4-)
Hydroxide (OH-)
Nitrate (NO3-)
Nitrite (NO2-)
Perchlorate (ClO4-)
Chlorate (ClO3-)
Chlorite (ClO2-)
Hypochlorite (ClO-)
Permanganate MnO4-)
Thiocyanate (SCN-)
Group 17 anions (F-, Cl-, Br-, I-,)
+2 Charge
Cadmium (Cd2+)
Chromium (II) or chromous (Cr 2+)
Cobalt (II) or cobaltous (Co2+)
Copper (II) or cupric (Cu2+)
Iron (II) or ferrous (Fe2+)
Lead (II) or plumbous (Pb2+)
Manganese (II) or manganous (Mn2+)
Mercury (I) or mercurous (Hg22+)
Mercury (II) or mercuric (Hg2+)
Nckel (Ni2+)
Tin (II) or stannous (Sn2+)
Zinc (Zn2+)
Group 2 (Be2+, Mg2+, Ca2+, Sr2+, Ba2+, Ra2+)
+3 Charge
Aluminum (Al3+)
Chromium (III) or chromic (Cr3+)
Iron (III) or ferric (Fe3+)
+4 Charge
lead (IV) or plumbic (Pb4+)
Tin (IV) or stannic (Sn4+)
-2 Charge
Carbonate (CO32-)
Chromate (CrO42-)
Dichromate (Cr2O72-)
Hydrogen phosphate (HPO42-)
Oxalate (C2O42-)
Oxide (O2-)
Peroxide (O22-)
Sulfate (SO42-)
Sulfite (SO32-)
Sulfide (S2-)
Thiosulfate (S2O32-)
-3 Charge
Arsenate (AsO43-)
Phosphate (PO43-)
Phosphite (PO33-)
Group 15 – nitride (N3-), phosphide (P3-)
Summary of metal ions with more than one charge:
Cu+, Cu2+; Hg22+, Hg2+; Co2+, Co3+; Cr2+, Cr3+; Fe2+, Fe3+; Mn2+, Mn3+; Pb2+, Pb4+; Sn2+, Sn4+
Manganese and several other metals can form several ions with different charge. You should know the ones listed.
Tips for Monoatomic ions: These can be organized into two groups:
1. Their position on the periodic table indicates the charge on the ion, since a neutral atom gains or loses a
predictable number of electrons (oxidation number) in order to obtain the noble gas configuration.
a. All group 1 elements (alkali metals) lose one electron to form a 1+ ion
b. All group 2 elements (alkaline earth metals) lose two electrons to form an ion with a 2+ charge
c. Group 13 metals lose three electrons to form an ion with a 3+ charge
d. All group 17 elements (halogens) gain one electron to form an ion with a 1- charge
e. All group 16 nonmetals gain two electrons to form an ion with a 2- charge
f. All group 15 nonmetals gain three electrons to form an ion with a 3- charge
Notice that cations keep their name (sodium ion, calcium ion) while the anions get an “-ide” ending
(chloride ion, oxide ion)
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2.
Metals that can form more than one ion will have their positive charge denoted by a roman numeral in
parenthesis immediately next to the name of the ion (iron (III) chloride)
Tips to Memorizing the Polyatomic Ions:
1. A quick way to memorize some of the polyatomic ions is to use:
“Nick the Camel ate a Clam for Supper in Phoenix on Crates”
Number of vowels = charge; Number of consonants = number of oxygen atoms
Nick – Nitrate; NO3Camel – Carbonate; CO32Clam – Chlorate; ClO3Supper – Sulfate; SO42Phoenix – Phosphate; PO43Crates – Chromate; CrO422.
“ate” anions have ONE more oxygen than the “ite” ion, but the same charge. If you memorize the “ate” ions,
then you should be able to derive the formula for the “ite’ and vice versa.
a. Sulfate is SO42-, so sulfite has the same charge but one less oxygen (SO32-)
b. Nitrate is NO3-, so nitrite has the same charge but one less oxygen (NO2-)
3.
If you know that a sulfate ion is SO42-, then to get the formula for the hydrogen sulfate ion, you add a
hydrogen ion to the front of the formula. Since hydrogen has a 1+ charge, the net charge on the new ion is
less one.
a.
PO43- 
HPO42
H2PO4phosphate
hydrogen phosphate
dihydrogen phosphate
4.
Learn the hypochlorite  Chlorite  chlorate  perchlorate series, and you also know the series
containing iodite/iodate as well as bromite/bromate.
a. The relationship between the “ite” and the “ate” is predictable, as always. Learn one and you know
the other.
b. The prefix hypo” means “under” or “too little” (think hypodermic, hypothermic or hypoglycemic)
i. hypochlorite is “under” chlorite, meaning it has one less oxygen
c. The prefix “hyper” means “above” or ‘too much” (think “hyperactive”)
i. The prefix “per” is derived from “hyper” so perchlorate (hyperchlorate) has one more
oxygen than chlorate.
d. Notice how this sequence increases in oxygen while retaining the same charge:
ClOhypochlorite
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ClO2chlorite
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ClO3chlorate

ClO4perchlorate
3. Acids
Rules for naming acids:
A. Binary acids – Contain hydrogen and one other atom (not oxygen)
 “Hydro(anion)ic acid”
o Ex. H2S is hydrosulfuric acid. Exception: HCN = hydrocyanic acid
B. Oxyacids – contain hydrogen and a polyatomic ion containing oxygen (oxyanion)
 (Anion name) suffix acid
o If the oxyanion ends with –ate, the suffix is –ic
 Ex. HNO3 is nitric acid (nitrate ion)
o If the oxyanion ends with –ite, the suffix is ous
 Ex. HNO2 is Nitrous acid (nitrite anion)
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Polyatomic Ions and Acids
Formula
Name
Ion
Ion Name
H2SO4
Sulfuric acid
SO42Sulfate ion
2H2SO3
Sulfurous acid
SO3
Sulfite ion
HNO3
Nitric acid
NO31Nitrate ion
HNO2
Nitrous acid
NO21Nitrite ion
H3PO4
Phosphoric acid
PO43Phosphate ion
H2CO3
Carbonic acid
CO32Carbonate ion
1HMnO4
Permanganic acid
MnO4
Permanganate ion
HCN
Hydrocyanic acid
CN1Cyanide ion
HOCN
Cyanic acid
OCN1Cyanate ion
HSCN
Thiocyanic acid
SCN1Thiocyanate ion
1HC2H3O2
Acetic acid
C2H3O2
Acetate ion
H2C2O4
Oxalic acid
C2O42Oxalate ion
H2CrO4
Chromic acid
CrO42Chromate ion
H2Cr2O7
Dichromic acid
Cr2O72Dichromate ion
H2S2O3
Thiosulfuric acid
S2O32Thiosulfate ion
3H3AsO4
Arsenic acid
AsO4
Arsenate ion
H3AsO3
Arsenous acid
AsO33Arsenite ion
Oxyhalogen Acids
Formula
Oxy name
Ion
Ion Name
HClO
Hypochlorous acid
ClO1Hypochlorite
HClO2
Chlorous acid
ClO21Chlorite
HClO3
Chloric acid
ClO31Chlorate ion
1HClO4
Perchloric acid
ClO4
Perchlorate ion
Br, I can be substituted for Cl. F may form hypofluorous acid and the hypofluorite ion.
5. Solubility Rules
1. All compounds of Group 1 and ammonium ions are soluble
2. All nitrates, acetate, and chlorates are soluble
3. All binary compounds of the halogens (except F) are soluble except those of Ag, Hg(I), and Pb
4. All sulfates are soluble except those of barium, strontium, calcium and lead
5. Sulfides and hydroxides are INsoluble except for Ca, Ba, Sr, ammonium and the alkali metals
6. Except for rule 1, carbonates, oxides, silicates, and phosphates are INsoluble
Note on Solubility Rules:
Solubility rules allow you to predict whether the product of a double replacement reaction will be a
precipitate or not:
 If the compound is soluble  it will dissociate into free ions into solution
 If the compound is INSOLUBLE  it will form a precipitate
 Ex. AgNO3(aq) + NaCl (aq)  AgCl(s) + NaNO3(aq)
 According to rule 3 – AgCl is insoluble and will form a precipitate
 According to rules 1 & 2, NaNO3 is soluble and will remain in solution
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6. Intermolecular Forces
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Attractive forces between solid or liquid molecules.
Network Covalent
Ionic (electrostatic attraction)
Metallic
Hydrogen Bonding
Dipole-dipole
London Dispersion Force
Direction covalent bond
C (graphite, diamond)
Si, SiO2 (sand)
Forces between adjacent ions
(Na+ --- Cl-)
Forces between metal nuclei
(Cu, Ag) sea of electrons
Forces between adjacent
molecules with H & F, O, N or Cl.
(H2O, NH3)
Forces between adjacent polar
molecules
Forces between adjacent
nonpolar molecules (CO2, Cl2)
*These can be strong in atoms &
molecules w/large numbers of
electrons
Strongest (generally)
Weakest
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AP CHEMISTRY SUMMER ASSIGNMENT
PART B – REVIEW PACKET
I.
Due - 2nd week of school
NOMENCLATURE REVIEW
IONIC COMPOUNDS
a. Formula for Binary Ionic Compounds
 The total positive charge must equal the total negative charge. The best way to write the
formula is to use the “Criss Cross Method“(also called “Swap & Drop”)
 Example: What ionic compound is formed when calcium combines with fluorine?
 Steps to the Criss Cross Method:
a. Write the ions with their charges; cations always are first  Ca2+ F1b. Cross over absolute value of charges and drop as subscript  Ca1 F2
c. Check to ensure the subscripts are the lowest whole number ratio
possible. Then write the formula:

CaF2
b. Naming Binary Ionic Compounds
 Combine the name of the cation (1st) and anion (2nd). The cation name stays the same and the
ending of the anion changes to –ide.
 Example: BaCl2 is named barium chloride
 If the cation has more than one oxidation state, this is indicated by using roman numerals
 Example : FeCl2 is iron (II) chloride
c.
Naming Ionic Compounds with Polyatomic Ions
 The rules are the same as above except the ending of the polyatomic is not changed.
 Example: Na2CO3 is sodium carbonate
II. MOLECULAR COMPOUNDS
a. Naming Binary Molecular Compounds – use prefixes
Number
1
2
3
4
5
6
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Prefix
monoditritetrapentahexa
Number
7
8
9
10
11
12
Prefix
heptaocta
nonadecaundecadodeca
The less electronegative atom (farthest from F) is written first. It only gets a prefix if it has more
than one atom.
The 2nd element gets a prefix and the ending –ide.
The o or a at the end of the prefix is dropped when the following word begins with a vowel, for
example monoxide or pentoxide.
Exercise 1 – Nomenclature: Simple, inorganic formulas & nomenclature
1. In the first column, classify each of the following as molecular(=covalent) (C) or ionic (I). In the 2nd
column, name each compound:
M or I
1)CaF2
2) P4O10
3) K2S
4) NaH
5) Al2Se3
6) N2O
7) O2F
8) SBr6
9) Li2Te
Name
C or I
Name
10) SrI2
11) CO
12) Cs2Po
13) ZnAt2
14) P2S3
15) AgCl
16) Na3N
17) Mg3P2
18)XeF6
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2. In the first column, write the chemical formula (formula unit) for the compound formed between the
two given elements. In the second column, name the compound:
Elements
Formula Unit
Name
1 magnesium & iodine
2 potassium & sulfur
3 chlorine and aluminum
4 zinc & bromine
5 strontium & oxygen
6 calcium & nitrogen
7 calcium & oxygen
8 copper (I) & oxygen
9 copper (II) & chlorine
10 mercury (II) & oxygen
11 nitrogen & aluminum
12 sulfur & cesium
Exercise 2 – Nomenclature: Oxidation numbers: anions & cations
Summary of Rules for Oxidation Numbers (you did not learn this in 1st year chemistry):
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Rule 1: Atoms in a pure element have an oxidation number of zero
Rule 2: The most electronegative element in a binary compound is assigned the number equal to the
negative charge it would have as an anion. The less electronegative atom is assigned the number equal to
the positive charge it would have as a cation.
o Example: group 1 is +1; Group 2 is +2; Group 13 is +3; Group 17 is -1; Group 16 is -2; Group 15 is 3, etc
Rule 3: Fluorine has an oxidation # of -1 in all its compounds because it is the most electronegative
element.
Rule 4: Oxygen has an oxidation # of -2 in almost all compounds
o Exceptions:
 Peroxides such as H2O2 in which its oxidation # is -1
 When oxygen is with halogens such as OF2, its oxidation # is +2
Rule 5: Hydrogen has an oxidation # of +1 in all compounds that are more electronegative that it is; it has
an oxidation # of -1 in all compounds with metals
Rule 6: The algebraic sum of the oxidation numbers of all atoms in a neutral compound is zero
Rule 7: The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge
on the ion
Rule 8: Rules 1-7 apply to covalently bonded atoms; however, oxidation numbers can also be assigned to
atoms in ionic compounds.
Determine the Oxidation Number of each of the underlined elements in the table below:
1) K2S
6) S8
11) C60
2) NaClO4
7) Mg
12) ZrO2
3) BrCl
8) K2W4O13
13) K2Cr2O7
4) Li2CO3
9) Mg(BF4)2
14) Al2(CrO4)3
5) OF2
10) Au2O3
15) Cs2TeF8
Exercise 3: More Nomenclature – Ternary Nomenclature: Acids & Salts
1. Name the following substances:
Formula
Name
Formula
Name
1) FeSO4
16) Fe2O3
2) Cu(NO3)2
17) (NH4)2SO3
3) Hg2Cl2
18) Ca(MnO4)2
8
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
12) CuCH3COO
27) C6H5COOH
13) N2O4
28) Hg2(IO)2
14) Rb3P
29) H3PO3
15) S8
30) NH4BrO3
2. Write the formulas for the following substances:
Name
1) vanadium (V) oxide
Formula
Name
16) francium dichromate
Formula
2) dihydrogen monoxide
17) calcium carbide
3) ammonium oxalate
18) mercury (I) nitrate
4) polonium (VI) thiocyanate
19) cerium (IV) benzoate
5) tetraphosphorus decoxide
20) potassium hydrogen phthalate
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) sodium hydrogen carbonate
28) phthalic acid
14) manganese (VII) oxide
29) tin (IV) chromate
15) copper (II) dihydrogen phosphate
30) hydrocyanic acid
3. Practice with Acids. Recall:
-IC from –ATE
-OUS from –ITE
HYDRO-, -IC from –IDE
Compete the Following Table:
Name of Acid
hydrochloric
Formula of Acid
HCl
Name of Anion
chloride
sulfuric acid
H2SO4
sulfate
HI
9
sulfite
chlorous acid
nitrate
HC2H3O2 or CH3COOH
hydrobromic acid
sulfide
HNO2
chromic acid
phosphate
Exercise 4 – Balancing Equations
1. 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) ___ VF6 + ___ HI  V2I12 + ___ HF
9) ___ OsO4 + ___ PtCl4  ___ PtO2 + ___ OsCl8
10) ___ Hg2I2 + ___ O2  ___ Hg2O + ___ I2
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 
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3.
Potassium + bromine 
DECOMPOSITION
1.
Strontium carbonate 
2.
Mercury (II) oxide 
3.
Aluminum chlorate 
DOUBLE REPLACEMENT (all are aqueous solutions)
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 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 
Exercise 6 – Solubility Rules: using the Solubility Rules Table on page 5-6 of this handout
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 an (S) for those
that are soluble and an (I) for those that are insoluble in water.
Name
Formula
(S) or (I)
Silver nitrate
cobalt (II) sulfate
Zinc hydroxide
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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 7 : Calculations Review
Complete all the problems below – show your work and circle your final answers. Record your answers with the
correct number of significant figures and remember units!
Rules for Significant Figures
THREE rules:
 RULE 1: Non-zero digits are ALWAYS significant
 RULE 2: Zeroes between two significant digits are significant
 RULE 3: If a decimal is PRESENT, only the zeroes AFTER the non-zero numbers count
Number
35,456
415,002
0.001010
0.300000
100.0000
12303000
Rule
Follow Rule 1
Follow Rule 2
Decimal point present - Follow Rule 3
Decimal point present: only count zeroes after 3
Decimal point present : Follow rule 3
Follow rule 1 and 2
Significant Figures
5
6
4
6
7
5
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Rules for Using Significant Figures in Calculations:
1. When adding or subtracting, the answer should have the same number of figures to the right of the
decimal as the value with the fewest decimal places. For example, 3.4 + 5.023 = 8.423. Round to 8.4,
because 3.4 has only one digit to the right of the decimal.
2. When multiplying or dividing, the answer should have the same number of significant figures as the
value with the fewest significant figures. For example, 1.220 x 3.4870 = 4.25414. Round this answer to
4.254, because 1.220 has only four significant figures.
Significant Figure learning Aid – Atlantic – Pacific
If decimal is
PRESENT, start
counting from LEFT
(Pacific side) at first
NON-ZERO digit &
keep counting
If decimal is ABSENT,
start counting from
RIGHT (Atlantic side)
at first NON-ZERO
digit & keep counting
Significant Figures
1.
perform the following calculations with correct significant figures:
a. 22.3411 x 42,68
8.36
b.
(9.24 + 43.1120 + 12.0) / 5.7821
Density
2. Calculate the mass of a sample of copper that occupies 5.3 x 10-2 cm3 if the density of copper is 8.94 g.cm3.
3.
A 9.46 g sample of a solid is placed in a 25.00 mL flask. The remaining volume of the flask is filled with
benzene in which the solid is insoluble. The solid and the benzene together weigh 26.83 g. The density of the
benzene is 0.879 g/ mL. What is the density of the solid?
Dimensional Analysis
4.
2.54cm = 1 inch; 1Kg = 2.205 Lb; 1 meter = 1.094 yards; 1 mile = 1760 yards
A pencil is 7.00 inches long. What is its length in centimeters?
5.
A student has entered a 10.0 km run. How long is the run in miles?
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Exercise 8: Atomic Theory, Electron Configuration & Periodicity
1.
Fill in the following table:
Element/ion
Fe
# of Protons
# of Neutrons
# of Electrons
Na+
27
25
S2Cr3+
2.
Write the electron configuration for:
a. Ca2+
b.
3.
4.
For Se write:
a. the complete electron configuration
b.
the noble gas configuration
c.
the orbital diagram
d.
the dot diagram
Draw the Lewis dot structures for the following elements:
a.
5.
Br-1
Pb
b.
N
c.
F
d.
Ca
e.
He
Place the elements S, Se, I, Ca and Be in:
a. order of increasing atomic radius
b.
decreasing ionization energy
Exercise 9: The Mole & Stoichiometry
1. Convert 3.47 x 1020 molecules of SO2 to moles. What is the mass of this quantity?
2.
A solution containing 4.5g of sodium phosphate is mixed with a solution containing 3.75 g of barium
nitrate. How many grams of barium phosphate can be produced?
3.
The thermite reaction (mixing solid iron (III) oxide with aluminum metal) ihas been used over the years to
weld railroad rails, in incendiary bombs, and to ignite solid-fuel rocket motors:
a. Write a balanced equation for the reaction
b.
What masses of iron (III) oxide and aluminum must be used to produce 15.0g of iron?
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Exercise 10: Molarity (M = mol solute/L solution)
1.
How many grams of solute are present in 100 mL of 1.50 M MgSO4?
2.
What is the molarity of 35 g of iron (II) acetate dissolved in enough water to make 250 mL of solution?
Exercise 11: Thermochemistry
Q = mCT
Q – Heat energy in joules or calories; C = Specific heat capacity of substance; T = Final – initial temperature
Note: The heat capacity of water should be known and is 4.184J/gC or is 1 cal/g C
1.
The specific heat capacity of graphite is 0.71J/gC. Calculate the energy (in calories) required to raise the
temperature of 1.8 kg of graphite by 100 C.
Resources to help you review:

http://www.khanacademy.org/science/chemistry

http://www.brightstorm.com/science/chemistry

http://quizlet.com/subject/ap-chemistry

http://www.bozemanscience.com/ap-chemistry/
Highly Recommended AP Chemistry Review Books:
Barron’s*, Princeton Review, Adrian Dingle’s Crash Course Chemistry
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