Chemistry

Chemistry
Josh Hoppe
Block 3

Honors Chemistry
Class Policies and Grading
The students will receive a Unit Outline at the beginning of each Unit. It will have information
about the assignments that they will do, what it’s grade classification will be, what action
they will need to do to complete the assignment and when it is due.
The students will receive a Weekly Memo of the activities they will be responsible for that
week. It will serve to inform the students of the learning goal for the week. It will also give
the students any special information about that week.
The students will also receive daily lectures and assignments that are designed to teach and
re-enforce information related to the learning goal. This will be time in which new material
will be taught and reviewed and will give the students the opportunity to ask questions
regarding the concepts being taught.
The students will work with a Lab partner and also be in a Lab group, but it will be up to the
individual student to do his or her part of all assignments and the individual student will
ultimately be responsible for all information presented in the class.
The students will be required to follow all District and School Policies and to follow all Lab
Safety Procedures, which they will be given and will sign, while performing labs. Students
should come to class on time and with the supplies needed for that class.
The following grading policy will be used.
Units with Projects
Units without Projects
(2,3,4,5) (1,6)
Work 12% Work 18%
Quizzes 13% Quizzes 20%
Labs 25% Labs 31%
Test 25% Test 31%
Projects 25%
The students will be given a teacher generated Mid Term and a District Final.

Unit 1
Measurement Lab
Separation of Mixtures Lab with Lab Write Up
Unit 2
Flame Test Lab
Nuclear Decay Lab
Element Marketing Project
Unit 3
Golden Penny Lab with Lab Write Up
Molecular Geometry
Research Presentation on a Chemical
Mid Term
Unit 4
Double Displacement Lab
Stoichiometry Lab with Lab Write Up
Mole Educational Demonstration Project
Unit 5
Gas Laws Lab with Lab Write Up
States of Matter Lab
Teach a Gas Law Project
Unit 6
Dilutions Lab
Titration Lab
District Final

Unit 1 (14 days)
Chapter 1 Introduction to Chemistry
Honors Chemistry
2013/2014 Syllabus
1.1 The Scope of Chemistry 1.3 Thinking Like a Scientist
1.2 Chemistry and You 1.4 Problem Solving in Chemistry
Chapter 2 Matter and Change
2.1 Properties of Matter 2.3 Elements and Compounds
2.2 Mixtures 2.4 Chemical Reactions
Chapter 3 Scientific Measurement
3.1 Using and Expressing Measurements 3.3 Solving Conversion Problems
3.2 Units of Measurement
Unit 2 (15 days)
Chapter 4 Atomic Structure
4.1 Defining the Atom 4.3 Distinguishing Among Atoms
4.2 Structure of the Nuclear Atom
Chapter 5 Electrons in Atoms
5.1 Revising the Atomic Model 5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectrum and the Quantum Mechanical Model
Chapter 6 The Periodic Table
6.1 Organizing the Elements 6.3 Periodic Trends
6.2 Classifying Elements
Chapter 25 Nuclear Chemistry
25.1 Nuclear Radiation 25.3 Fission and Fusion
25.2 Nuclear Transformations 25.4 Radiation in Your Life
Unit 3 (11 days)
Chapter 7 Ionic and Metallic Bonding
7.1 Ions 7.3 Bonding in Metals
7.2 Ionic Bonds and Ionic Compounds
Chapter 8 Covalent Bonding
8.1 Molecular Compounds 8.3 Bonding Theories
8.2 The Nature of Covalent Bonding 8.4 Polar Bonds and Molecules
Chapter 9 Chemical Names and Formulas
9.1 Naming Ions 9.3 Naming and Writing Formulas for Molecular Compounds
9.2 Naming and Writing Formulas for Ionic Compounds 9.4 Names for Acids and Bases
3 days
5 days
6 days
3 days
5 days
3 days
3 days
4 days
4 days
3 days

Unit 4 (15 days)
Chapter 10 Chemical Quantities
10.1 The Mole: A Measurement of Matter 10.3 % Composition & Chem. Formulas
10.2 Mole-Mass and Mole-Volume Relationships
Chapter 11 Chemical Reactions
11.1 Describing Chemical Reactions 11.3 Reactions in Aqueous Solutions
11.2 Types of Chemical Reactions
Chapter 12 Stoichiometry
12.1 The Arithmetic of Equations 12.3 Limiting Reagent and % Yield
12.2 Chemical Calculations
Unit 5 (15 days)
Chapter 13 States of Matter
13.1 The Nature of Gases 13.3 The Nature of Solids
13.2 The Nature of Liquids 13.4 Changes in State
Chapter 14 The Behavior of Gases
14.1 Properties of Gases 14.3 Ideal Gases
14.2 The Gas Laws 14.4 Gases: Mixtures and Movement
Chapter 15 Water and Aqueous Systems
15.1 Water and its Properties 15.3 Heterogeneous Aqueous Systems
15.2 Homogeneous Aqueous Systems
Unit 6 (10 days)
Chapter 16 Solutions
16.1 Properties of Solutions 16.3 Colligative Properties of Solutions
16.2 Concentrations of Solutions 16.4 Calc. Involving Colligative Property
Chapter 17 Thermochemistry
17.1 The Flow of Energy 17.3 Heat in Changes of State
17.2 Measuring and Expressing Enthalpy Change 17.4 Calculating Heats in Reactions
Chapter 18 Reaction Rates and Equilibrium
18.1 Rates of Reactions 18.3 Reversible Reaction & Equilibrium
18.2 The Progress of Chemical Reactions 18.5 Free Energy and Entropy
Chapter 19 Acid and Bases
19.1 Acid-Base Theories 19.4 Neutralization Reactions
19.2 Hydrogen Ions and Acidity 19.5 Salts in Solutions
19.3 Strengths of Acids and Bases
5 days
5 days
5 days
4 days
5 days
3 days
4 days
2 days
2 days
2 days

Lorenzo Walker Technical High School
MUSTANG LABORATORIES
Chemistry Safety
Safety in the MUSTANG LABORATORIES - Chemistry Laboratory
Working in the chemistry laboratory is an interesting and rewarding experience. During your labs, you will be actively
involved from beginning to end—from setting some change in motion to drawing some conclusion. In the laboratory, you will
be working with equipment and materials that can cause injury if they are not handled properly.
However, the laboratory is a safe place to work if you are careful. Accidents do not just happen; they are caused—by
carelessness, haste, and disregard of safety rules and practices. Safety rules to be followed in the laboratory are listed
below. Before beginning any lab work, read these rules, learn them, and follow them carefully.
General
1. Be prepared to work when you arrive at the lab. Familiarize yourself with the lab procedures before beginning the lab.
2. Perform only those lab activities assigned by your teacher. Never do anything in the laboratory that is not called for in
the laboratory procedure or by your teacher. Never work alone in the lab. Do not engage in any horseplay.
3. Work areas should be kept clean and tidy at all times. Only lab manuals and notebooks should be brought to the work
area. Other books, purses, brief cases, etc. should be left at your desk or placed in a designated storage area.
4. Clothing should be appropriate for working in the lab. Jackets, ties, and other loose garments should be removed. Open
shoes should not be worn.
5. Long hair should be tied back or covered, especially in the vicinity of open flame.
6. Jewelry that might present a safety hazard, such as dangling necklaces, chains, medallions, or bracelets should not be
worn in the lab.
7. Follow all instructions, both written and oral, carefully.
8. Safety goggles and lab aprons should be worn at all times.
9. Set up apparatus as described in the lab manual or by your teacher. Never use makeshift arrangements.
10. Always use the prescribed instrument (tongs, test tube holder, forceps, etc.) for handling apparatus or equipment.
11. Keep all combustible materials away from open flames.
12. Never touch any substance in the lab unless specifically instructed to do so by your teacher.
13. Never put your face near the mouth of a container that is holding chemicals.
14. Never smell any chemicals unless instructed to do so by your teacher. When testing for odors, use a wafting motion to
direct the odors to your nose.
15. Any activity involving poisonous vapors should be conducted in the fume hood.
16. Dispose of waste materials as instructed by your teacher.
17. Clean up all spills immediately.
18. Clean and wipe dry all work surfaces at the end of class. Wash your hands thoroughly.
19. Know the location of emergency equipment (first aid kit, fire extinguisher, fire shower, fire blanket, etc.) and how to use them.
20. Report all accidents to the teacher immediately.
Handling Chemicals
21. Read and double check labels on reagent bottles before removing any reagent. Take only as much reagent as you need.
22. Do not return unused reagent to stock bottles.
23. When transferring chemical reagents from one container to another, hold the containers out away from your body.
24. When mixing an acid and water, always add the acid to the water.
25. Avoid touching chemicals with your hands. If chemicals do come in contact with your hands, wash them immediately.
26. Notify your teacher if you have any medical problems that might relate to lab work, such as allergies or asthma.
27. If you will be working with chemicals in the lab, avoid wearing contact lenses. Change to glasses, if possible, or notify
the teacher.
Handling Glassware
28. Glass tubing, especially long pieces, should be carried in a vertical position to minimize the likelihood of breakage and
to avoid stabbing anyone.

29. Never handle broken glass with your bare hands. Use a brush and dustpan to clean up broken glass. Dispose of the
glass as directed by your teacher.
30. Always lubricate glassware (tubing, thistle tubes, thermometers, etc.) with water or glycerin before attempting to insert it into a
rubber stopper.
31. Never apply force when inserting or removing glassware from a stopper. Use a twisting motion. If a piece of glassware
becomes "frozen" in a stopper, take it to your teacher.
32. Do not place hot glassware directly on the lab table. Always use an insulating pad of some sort.
33. Allow plenty of time for hot glass to cool before touching it. Hot glass can cause painful burns. (Hot glass looks cool.)
Heating Substances
34. Exercise extreme caution when using a gas burner. Keep your head and clothing away from the flame.
35. Always turn the burner off when it is not in use.
36. Do not bring any substance into contact with a flame unless instructed to do so.
37. Never heat anything without being instructed to do so.
38. Never look into a container that is being heated.
39. When heating a substance in a test tube, make sure that the mouth of the tube is not pointed at yourself or anyone else.
40. Never leave unattended anything that is being heated or is visibly reacting.
First Aid in the MUSTANG LABORATORIES - Chemistry Laboratory
Accidents do not often happen in well-equipped chemistry laboratories if students understand safe laboratory procedures
and are careful in following them. When an occasional accident does occur, it is likely to be a minor one.
The instructor will assist in treating injuries such as minor cuts and burns. However, for some types of injuries, you must take
action immediately. The following information will be helpful to you if an accident occurs.
1. Shock. People who are suffering from any severe injury (for example, a bad burn or major loss of blood) may be in a state
of shock. A person in shock is usually pale and faint. The person may be sweating, with cold, moist skin and a weak, rapid
pulse. Shock is a serious medical condition. Do not allow a person in shock to walk anywhere—even to the campus security
office. While emergency help is being summoned, place the victim face up in a horizontal position, with the feet raised about
30 centimeters. Loosen any tightly fitting clothing and keep him or her warm.
2. Chemicals in the Eyes. Getting any kind of a chemical into the eyes is undesirable, but certain chemicals are especially
harmful. They can destroy eyesight in a matter of seconds. Because you will be wearing safety goggles at all times in the lab,
the likelihood of this kind of accident is remote. However, if it does happen, flush your eyes with water immediately. Do NOT
attempt to go to the campus office before flushing your eyes. It is important that flushing with water be continued for a
prolonged time—about 15 minutes.
3. Clothing or Hair on Fire. A person whose clothing or hair catches on fire will often run around hysterically in an
unsuccessful effort to get away from the fire. This only provides the fire with more oxygen and makes it burn faster. For
clothing fires, throw yourself to the ground and roll around to extinguish the flames. For hair fires, use a fire blanket to
smother the flames. Notify campus security immediately.
4. Bleeding from a Cut. Most cuts that occur in the chemistry laboratory are minor. For minor cuts, apply pressure to the
wound with a sterile gauze. Notify campus security of all injuries in the lab. If the victim is bleeding badly, raise the bleeding
part, if possible, and apply pressure to the wound with a piece of sterile gauze. While first aid is being given, someone else
should notify the campus security officer.
5. Chemicals in the Mouth. Many chemicals are poisonous to varying degrees. Any chemical taken into the mouth should be
spat out and the mouth rinsed thoroughly with water. Note the name of the chemical and notify the campus office
immediately. If the victim swallows a chemical, note the name of the chemical and notify campus security immediately.
If necessary, the campus security officer or administrator will contact the Poison Control Center, a hospital emergency room,
or a physician for instructions.
6. Acid or Base Spilled on the Skin.
Flush the skin with water for about 15 minutes. Take the victim to the campus office to report the injury.
7. Breathing Smoke or Chemical Fumes.
All experiments that give off smoke or noxious gases should be conducted in a well-ventilated fume hood. This will make an
accident of this kind unlikely. If smoke or chemical fumes are present in the laboratory, all persons—even those who do not
feel ill—should leave the laboratory immediately. Make certain that all doors to the laboratory are closed after the last
person has left. Since smoke rises, stay low while evacuating a smoke-filled room. Notify campus security immediately.

MUSTANG LABORATORIES
COMMITMENT TO SAFETY IN THE LABORATORY
As a student enrolled in Chemistry at Lorenzo Walker Technical High School, I agree to use
good laboratory safety practices at all times. I also agree that I will:
1. Conduct myself in a professional manner, respecting both my personal safety and the safety of others in the laboratory.
2. Wear proper and approved safety glasses or goggles in the laboratory at all times.
3. Wear sensible clothing and tie back long hair in the laboratory. Understand that open-toed shoes pose a hazard during
laboratory classes and that contact lenses are an added safety risk.
4. Keep my lab area free of clutter during an experiment.
5. Never bring food or drink into the laboratory, nor apply makeup within the laboratory.
6. Be aware of the location of safety equipment such as the fire extinguisher, eye wash station, fire blanket, first aid kit.
Know the location of the nearest telephone and exits.
7. Read the assigned lab prior to coming to the laboratory.
8. Carefully read all labels on all chemical containers before using their contents, remove a small amount of reagent
properly if needed, do not pour back the unused chemicals into the original container.
9. Dispose of chemicals as directed by the instructor only. At no time will I pour anything down the sink without prior
instruction.
10. Never inhale fumes emitted during an experiment. Use the fume hood when instructed to do so.
11. Report any accident immediately to the instructor, including chemical spills.
12. Dispose of broken glass and sharps only in the designated containers.
13. Clean my work area and all glassware before leaving the laboratory.
14. Wash my hands before leaving the laboratory.
NAME ____________________________
Josh Hoppe
PARENT NAME _______________________________________
Linda Hoppe
BLOCK ____________________________
3 PARENT NUMBER _____________________________________
(239) 537-2550
SIGNATURE _________________________________________________
8/20/13
DATE ____________________________________

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C
THE UNIVERSITY OF THE STATE OF NEW YORK• THE STATE EDUCATION DEPARTMENT• ALBANY, NY 12234
Reference Tables for Physical Setting/CHEMISTRY
Table A
Standard Temperature and Pressure
2011 Edition
Table D
Selected Units
Name Value Unit
Standard Pressure 101.3 kPa kilopascal
1 atm atmosphere
Standard Temperature 273 K kelvin
0°C degree Celsius
Table B
Physical Constants for Water
Heat of Fusion
Heat of Vaporization
Specific Heat Capacity of H 2
O()
Table C
Selected Prefixes
Factor Prefix Symbol
10 3 kilo- k
10 –1 deci- d
10 –2 centi- c
10 –3 milli- m
10 –6 micro- μ
10 –9 nano- n
10 –12 pico- p
334 J/g
2260 J/g
4.18 J/g•K
Symbol Name Quantity
m meter length
g gram mass
Pa pascal pressure
K kelvin temperature
mol
J
mole
joule
s second time
min minute time
h hour time
d day time
y year time
amount of
substance
L liter volume
energy, work,
quantity of heat
ppm parts per million concentration
M
molarity
solution
concentration
u atomic mass unit atomic mass
Reference Tables for Physical Setting/Chemistry – 2011 Edition 1

Table E
Selected Polyatomic Ions
Formula Name Formula Name
H 3
O +
hydronium
CrO 4
2–
chromate
Hg 2
2+
mercury(I)
Cr 2
O 7
2–
dichromate
NH 4
+
C 2
H 3
O
–
2 –}
CH 3
COO
CN –
CO 3
2–
HCO
–
3
C 2
O
2–
4
ClO –
ammonium
acetate
cyanide
carbonate
hydrogen
carbonate
oxalate
hypochlorite
MnO 4
–
NO
–
2
NO
–
3
O
2–
2
OH –
PO 4
3–
SCN –
SO 3
2–
permanganate
nitrite
nitrate
peroxide
hydroxide
phosphate
thiocyanate
sulfite
ClO 2
–
chlorite
SO 4
2–
sulfate
ClO 3
–
chlorate
HSO 4
–
hydrogen sulfate
ClO 4
–
perchlorate
S 2
O 3
2–
thiosulfate
Table F
Solubility Guidelines for Aqueous Solutions
Ions That Form
Soluble Compounds
Group 1 ions
(Li + , Na + , etc.)
ammonium (NH + 4
)
nitrate (NO – 3
)
acetate (C 2
H 3
O – 2
or
CH 3
COO – )
hydrogen carbonate
(HCO – 3
)
chlorate (ClO – 3
)
halides (Cl – , Br – , I – )
Exceptions
when combined with
Ag + , Pb 2+ , or Hg 2
2+
sulfates (SO 4 2– ) when combined with Ag + ,
Ca 2+ , Sr 2+ , Ba 2+ , or Pb 2+
Ions That Form
Insoluble Compounds* Exceptions
carbonate (CO 2– 3
) when combined with Group 1
ions or ammonium (NH + 4
)
chromate (CrO 2– 4
) when combined with Group 1
ions, Ca 2+ , Mg 2+ , or
ammonium (NH + 4
)
phosphate (PO 3– 4
) when combined with Group 1
ions or ammonium (NH + 4
)
sulfide (S 2– ) when combined with Group 1
ions or ammonium (NH + 4
)
hydroxide (OH – ) when combined with Group 1
ions, Ca 2+ , Ba 2+ , Sr 2+ , or
ammonium (NH + 4
)
*compounds having very low solubility in H 2 O
Reference Tables for Physical Setting/Chemistry – 2011 Edition 2

150.
140.
Table G
Solubility Curves at Standard Pressure
KI
NaNO 3
130.
120.
KNO 3
110.
100.
Solubility (g solute/100. g H 2
O)
90.
80.
70.
60.
HCl
NH 4
Cl
KCl
50.
40.
30.
NaCl
KClO 3
NH 3
20.
10.
SO 2
0
0 10. 20. 30. 40. 50. 60. 70. 80. 90. 100.
Temperature (°C)
Reference Tables for Physical Setting/Chemistry – 2011 Edition 3

Table H
Vapor Pressure of Four Liquids
200.
propanone
ethanol
150.
water
Vapor Pressure (kPa)
100.
101.3 kPa
ethanoic
acid
50.
0
0 25 50. 75 100. 125
Reference Tables for Physical Setting/Chemistry – 2011 Edition 4

Table I
Heats of Reaction at 101.3 kPa and 298 K
Reaction
ΔH (kJ)*
CH 4
(g) + 2O 2
(g) CO 2
(g) + 2H 2
O() –890.4
C 3
H 8
(g) + 5O 2
(g) 3CO 2
(g) + 4H 2
O() –2219.2
2C 8
H 18
() + 25O 2
(g) 16CO 2
(g) + 18H 2
O() –10943
2CH 3
OH() + 3O 2
(g) 2CO 2
(g) + 4H 2
O() –1452
C 2
H 5
OH() + 3O 2
(g) 2CO 2
(g) + 3H 2
O() –1367
C 6
H 12
O 6
(s) + 6O 2
(g) 6CO 2
(g) + 6H 2
O() –2804
2CO(g) + O 2
(g) 2CO 2
(g) –566.0
C(s) + O 2
(g) CO 2
(g) –393.5
4Al(s) + 3O 2
(g) 2Al 2
O 3
(s) –3351
N 2
(g) + O 2
(g) 2NO(g) +182.6
N 2
(g) + 2O 2
(g) 2NO 2
(g) +66.4
2H 2
(g) + O 2
(g) 2H 2
O(g) –483.6
2H 2
(g) + O 2
(g) 2H 2
O() –571.6
N 2
(g) + 3H 2
(g) 2NH 3
(g) –91.8
2C(s) + 3H 2
(g) C 2
H 6
(g) –84.0
2C(s) + 2H 2
(g) C 2
H 4
(g) +52.4
2C(s) + H 2
(g) C 2
H 2
(g) +227.4
H 2
(g) + I 2
(g) 2HI(g) +53.0
KNO 3
(s) H 2 O K + (aq) + NO 3 – (aq) +34.89
NaOH(s) H 2 O Na + (aq) + OH – (aq) –44.51
NH 4
Cl(s) H 2 O NH 4 + (aq) + Cl – (aq) +14.78
NH 4
NO 3
(s) H 2 O NH 4 + (aq) + NO 3 – (aq) +25.69
NaCl(s) H 2 O Na + (aq) + Cl – (aq) +3.88
LiBr(s) H 2 O Li + (aq) + Br – (aq) –48.83
H + (aq) + OH – (aq) H 2
O() –55.8
*The ΔH values are based on molar quantities represented in the equations.
A minus sign indicates an exothermic reaction.
Most
Active
Least
Active
Table J
Activity Series**
Metals Nonmetals Most
Active
Li
F 2
Rb Cl 2
K Br 2
Cs
I 2
Ba
Sr
Ca
Na
Mg
Al
Ti
Mn
Zn
Cr
Fe
Co
Ni
Sn
Pb
H 2
Cu
Ag
Au
**Activity Series is based on the hydrogen
standard. H 2 is not a metal.
Least
Active
Reference Tables for Physical Setting/Chemistry – 2011 Edition 5

Table K
Common Acids
Table N
Selected Radioisotopes
HCl(aq)
Formula
HNO 2
(aq)
HNO 3
(aq)
H 2
SO 3
(aq)
H 2
SO 4
(aq)
H 3
PO 4
(aq)
H 2
CO 3
(aq)
or
CO 2
(aq)
CH 3
COOH(aq)
or
HC 2
H 3
O 2
(aq)
Name
hydrochloric acid
nitrous acid
nitric acid
sulfurous acid
sulfuric acid
phosphoric acid
carbonic acid
ethanoic acid
(acetic acid)
Nuclide Half-Life Decay
Mode
Nuclide
Name
198 Au 2.695 d β – gold-198
14 C 5715 y β – carbon-14
37 Ca 182 ms β + calcium-37
60 Co 5.271 y β – cobalt-60
137 Cs 30.2 y β – cesium-137
53 Fe 8.51 min β + iron-53
220 Fr 27.4 s α francium-220
3 H 12.31 y β – hydrogen-3
131 I 8.021 d β – iodine-131
37 K 1.23 s β + potassium-37
42 K 12.36 h β – potassium-42
Table L
Common Bases
85 Kr 10.73 y β – krypton-85
16 N 7.13 s β – nitrogen-16
Formula
NaOH(aq)
KOH(aq)
Ca(OH) 2
(aq)
NH 3
(aq)
Name
sodium hydroxide
potassium hydroxide
calcium hydroxide
aqueous ammonia
19 Ne 17.22 s β + neon-19
32 P 14.28 d β – phosphorus-32
239 Pu 2.410 × 10 4 y α plutonium-239
226 Ra 1599 y α radium-226
222 Rn 3.823 d α radon-222
90 Sr 29.1 y β – strontium-90
Table M
Common Acid–Base Indicators
Approximate
Indicator pH Range Color
for Color Change
Change
methyl orange 3.1–4.4 red to yellow
bromthymol blue 6.0–7.6 yellow to blue
phenolphthalein 8–9 colorless to pink
litmus 4.5–8.3 red to blue
bromcresol green 3.8–5.4 yellow to blue
thymol blue 8.0–9.6 yellow to blue
99 Tc 2.13 × 10 5 y β – technetium-99
232 Th 1.40 × 10 10 y α thorium-232
233 U 1.592 × 10 5 y α uranium-233
235 U 7.04 × 10 8 y α uranium-235
238 U 4.47 × 10 9 y α uranium-238
Source: CRC Handbook of Chemistry and Physics, 91 st ed., 2010–2011,
CRC Press
Source: The Merck Index, 14 th ed., 2006, Merck Publishing Group
Reference Tables for Physical Setting/Chemistry – 2011 Edition 6

Table O
Symbols Used in Nuclear Chemistry
Name Notation Symbol
alpha particle
4
2
He or 4 2 α α
beta particle
0
–1
e or 0
–1 β β–
gamma radiation
0
0
γ γ
neutron
1
0
n n
proton
1
1
H or 1 1 p p
positron
0
+1
e or 0
+1 β β+
Name General Examples
Formula Name Structural Formula
alkanes C n
H 2n+2
ethane
alkenes C n
H 2n
ethene
alkynes C n
H 2n–2
ethyne
Table P
Organic Prefixes
Prefix
meth- 1
eth- 2
prop- 3
but- 4
pent- 5
hex- 6
hept- 7
oct- 8
Number of
Carbon Atoms
non- 9
dec- 10
Table Q
Homologous Series of Hydrocarbons
Note: n = number of carbon atoms
Reference Tables for Physical Setting/Chemistry – 2011 Edition 7
H
H
H
H
H
C
H
C
H
C
H
C
H
H
H
C C H

Table R
Organic Functional Groups
Class of
Compound
Functional
Group
General
Formula
Example
halide
(halocarbon)
F (fluoro-)
Cl (chloro-)
Br (bromo-)
I (iodo-)
R X
(X represents
any halogen)
CH 3
CHClCH 3
2-chloropropane
alcohol
OH
R
OH
CH 3
CH 2
CH 2
OH
1-propanol
ether
O
R O R′
CH 3
OCH 2
CH 3
methyl ethyl ether
aldehyde
O
C H
R
O
C H
O
CH 3
CH 2
C H
propanal
ketone
O
C
O
R C R′
O
CH 3
CCH 2
CH 2
CH 3
2-pentanone
organic acid
O
C OH
R
O
C OH
O
CH 3
CH 2
C OH
propanoic acid
ester
O
C O
O
R C O R′
O
CH 3
CH 2
COCH 3
methyl propanoate
amine
N
R
R′
N R′′
CH 3
CH 2
CH 2
NH 2
1-propanamine
amide
O
C NH
R
O R′
C NH
O
CH 3
CH 2
C NH 2
propanamide
Note: R represents a bonded atom or group of atoms.
Reference Tables for Physical Setting/Chemistry – 2011 Edition 8

0
6.941
+1
Li
3
2-1
Na
39.0983
K +1
19
2-8-8-1
Rb
Cs
(223)
Fr
87
-18-32-18-8-1
+1
Ra
88
-18-32-18-8-2
39
138.9055
La
57
2-8-18-18-9-2
+2 (227)
Ac
89
-18-32-18-9-2
47.867
Ti
22
2-8-10-2
91.224
Zr
40
2-8-18-10-2
+3 178.49
Hf
72
*18-32-10-2
+3 (261)
Rf
104
+2
+3
+4
+4
+4
50.9415
V
23
2-8-11-2
+2
+3
+4
+5
51.996
Cr
24
2-8-13-1
95.94
Mo
42
2-8-18-13-1
183.84
W
74
-18-32-12-2
+2
+3
+6
+6
+6
54.9380
Mn
25
2-8-13-2
+2
+3
+4
+7
55.845
Fe
26
2-8-14-2
+2
+3 58.9332
Co
27
2-8-15-2
+2
+3
58.693
Ni
28
2-8-16-2
+2
+3 63.546 Cu
2-8-18-1
107.868
Ag
47
2-8-18-18-1
79
+1
+2
+1
65.409
Zn
30
2-8-18-2
+3 12.011
B
5
2-3
26.98154
Al
13
2-8-3
+2 69.723
Ga
31
2-8-18-3
+3
+3
–4
+2
+4
C
6
2-4
28.0855
Si
14
2-8-4
72.64
Ge
32
2-8-18-4
Pb
–4
+2
+4
+2
+4
74.9216
As
33
2-8-18-5
Sb
–3
+3
15.9994
O
F
8
2-6
–2 18.9984
78.96
Se
34
2-8-18-6
127.60
Te
52
2-8-18-18-6
(209)
Po
84
-18-32-18-6
–2
+4
+6
–2
+4
+6
+2
+4
2-7
79.904
Br
35
2-8-18-7
126.904
l
53
2-8-18-18-7
(210)
At
85
-18-32-18-7
4.00260 0
He
2
2
–1 20.180
Ne
10
2-8
0
22.98977
11
2-8-1
1
1.00794 +1
–1
H
1
1
1
85.4678
37
2-8-18-8-1
30.97376
P
15
2-8-5
–3
+3
+5
32.065
S
16
2-8-6
–2
+4
+6
–1
+1
+5
–1
+1
+5
+7
39.948
Ar
18
2-8-8
83.798
Kr
36
2-8-18-8
131.29
Xe
54
2-8-18-18-8
(222)
Rn
86
-18-32-18-8
0
+2
0
+2
+4
+6
0
132.905
55
2-8-18-18-8-1
Symbol
Relative atomic masses are based
Group on 12 C = 12 (exact)
Group
2
13 14 15 16 17 18
Atomic Number
Note: Numbers in parentheses
10.81
are mass numbers of the most
Electron Configuration
stable or common isotope.
+1
+1
+1
9.01218 +2
Be
4
2-2
24.305
Mg
12
2-8-2
40.08
Ca
20
2-8-8-2
87.62
Sr
38
2-8-18-8-2
137.33
Ba
56
2-8-18-18-8-2
(226)
+2
+2
+2
+2
3
44.9559
Sc
21
2-8-9-2
88.9059
Y
2-8-18-9-2
+3
+3
4
KEY
+4
92.9064
Nb +3
+5
41
2-8-18-12-1
180.948
Ta
73
-18-32-11-2
(262)
105
5
Periodic Table of the Elements
Atomic Mass
Db
+5
6
(266)
Sg
106
12.011 2-4
–4
6
C
+2
+4
(98)
Tc
43
2-8-18-13-2
186.207
Re
75
-18-32-13-2
(272)
Bh
107
7
Group
+4
+6
+7
+4
+6
+7
8
101.07
Ru
44
2-8-18-15-1
190.23
Os
76
-18-32-14-2
(277)
Hs
108
+3
+3
+4
Selected Oxidation States
9
102.906
Rh
45
2-8-18-16-1
192.217
Ir
77
-18-32-15-2
(276)
Mt
109
+3
+3
+4
106.42
Pd
46
2-8-18-18
195.08
Pt
78
-18-32-17-1
+2
+4
+2
+4
196.967
Au
-18-32-18-1
(281)
Ds (280) Rg
110
10
29
111
11 12
+1
+3
112.41
Cd
48
2-8-18-18-2
200.59
Hg
80
-18-32-18-2
(285)
Cn
112
+2 114.818
In
+1
+2
49
2-8-18-18-3
204.383
Tl
81
-18-32-18-3
(284)
Uut
113**
+3
+1
+3
118.71
Sn
50
2-8-18-18-4
207.2
82
-18-32-18-4
(289)
Uuq
114
+2
+4
+2
+4
14.0067 –3
–2
N
–1
7
2-5
121.760
51
2-8-18-18-5
208.980
Bi
83
-18-32-18-5
(288)
Uup
115
+1
+2
+3
+4
+5
+5
–3
+3
+5
+3
+5
(292)
Uuh
116
35.453
Cl
17
2-8-7
( ? )
Uus
117
–1
+1
+5
+7
18
(294)
Uuo
118
140.116
Ce
58
232.038
Th
90
+3
+4
140.908
Pr +3
59
144.24
Nd
60
+4 231.036
Pa +4 238.029 +5
U +3
+4
+5
+6
91
92
+3
(145)
Pm
61
+3
150.36
Sm
62
+2
+3
151.964
Eu
63
+2
+3
157.25
Gd
64
+3
158.925
(237)Np (244) Pu (243) Am (247) Cm +3 (247) Bk +3
+3
+4
+5
+6
93 94
+3
+4
+5
+6
65
+3
+4
+5
+6
95 96 97
Tb
+3
+4
162.500
Dy
66
(251)
+3
164.930
Ho
67
+3
167.259
Er
68
Cf +3 (252) Es (257) Fm
100
98 99
+3
+3
+3
168.934
Tm +3
69
(258)
Md
101
+2
+3
173.04
Yb
70
(259)
No
102
+2
+3
+2
+3
174.9668
Lu
71
(262)
Lr
103
+3
+3
*denotes the presence of (2-8-) for elements 72 and above
**The systematic names and symbols for elements of atomic numbers 113 and above
will be used until the approval of trivial names by IUPAC.
Source: CRC Handbook of Chemistry and Physics, 91 st ed., 2010–2011, CRC Press
9
Period
1
2
3
4
5
6
7
Reference Tables for Physical Setting/Chemistry – 2011 Edition 9

Table S
Properties of Selected Elements
First
Atomic Symbol Name Ionization
Electro- Melting Boiling* Density** Atomic
Number Energy negativity Point Point (g/cm 3 ) Radius
(kJ/mol) (K) (K) (pm)
1 H hydrogen 1312 2.2 14 20. 0.000082 32
2 He helium 2372 — — 4 0.000164 37
3 Li lithium 520. 1.0 454 1615 0.534 130.
4 Be beryllium 900. 1.6 1560. 2744 1.85 99
5 B boron 801 2.0 2348 4273 2.34 84
6 C carbon 1086 2.6 — — .— 75
7 N nitrogen 1402 3.0 63 77 0.001145 71
8 O oxygen 1314 3.4 54 90. 0.001308 64
9 F fluorine 1681 4.0 53 85 0.001553 60.
10 Ne neon 2081 — 24 27 0.000825 62
11 Na sodium 496 0.9 371 1156 0.97 160.
12 Mg magnesium 738 1.3 923 1363 1.74 140.
13 Al aluminum 578 1.6 933 2792 2.70 124
14 Si silicon 787 1.9 1687 3538 2.3296 114
15 P phosphorus (white) 1012 2.2 317 554 1.823 109
16 S sulfur (monoclinic) 1000. 2.6 388 718 2.00 104
17 Cl chlorine 1251 3.2 172 239 0.002898 100.
18 Ar argon 1521 — 84 87 0.001633 101
19 K potassium 419 0.8 337 1032 0.89 200.
20 Ca calcium 590. 1.0 1115 1757 1.54 174
21 Sc scandium 633 1.4 1814 3109 2.99 159
22 Ti titanium 659 1.5 1941 3560. 4.506 148
23 V vanadium 651 1.6 2183 3680. 6.0 144
24 Cr chromium 653 1.7 2180. 2944 7.15 130.
25 Mn manganese 717 1.6 1519 2334 7.3 129
26 Fe iron 762 1.8 1811 3134 7.87 124
27 Co cobalt 760. 1.9 1768 3200. 8.86 118
28 Ni nickel 737 1.9 1728 3186 8.90 117
29 Cu copper 745 1.9 1358 2835 8.96 122
30 Zn zinc 906 1.7 693 1180. 7.134 120.
31 Ga gallium 579 1.8 303 2477 5.91 123
32 Ge germanium 762 2.0 1211 3106 5.3234 120.
33 As arsenic (gray) 944 2.2 1090. — 5.75 120.
34 Se selenium (gray) 941 2.6 494 958 4.809 118
35 Br bromine 1140. 3.0 266 332 3.1028 117
36 Kr krypton 1351 — 116 120. 0.003425 116
37 Rb rubidium 403 0.8 312 961 1.53 215
38 Sr strontium 549 1.0 1050. 1655 2.64 190.
39 Y yttrium 600. 1.2 1795 3618 4.47 176
40 Zr zirconium 640. 1.3 2128 4682 6.52 164
Reference Tables for Physical Setting/Chemistry – 2011 Edition 10

First
Atomic Symbol Name Ionization
Electro- Melting Boiling* Density** Atomic
Number Energy negativity Point Point (g/cm 3 ) Radius
(kJ/mol) (K) (K) (pm)
41 Nb niobium 652 1.6 2750. 5017 8.57 156
42 Mo molybdenum 684 2.2 2896 4912 10.2 146
43 Tc technetium 702 2.1 2430. 4538 11 138
44 Ru ruthenium 710. 2.2 2606 4423 12.1 136
45 Rh rhodium 720. 2.3 2237 3968 12.4 134
46 Pd palladium 804 2.2 1828 3236 12.0 130.
47 Ag silver 731 1.9 1235 2435 10.5 136
48 Cd cadmium 868 1.7 594 1040. 8.69 140.
49 In indium 558 1.8 430. 2345 7.31 142
50 Sn tin (white) 709 2.0 505 2875 7.287 140.
51 Sb antimony (gray) 831 2.1 904 1860. 6.68 140.
52 Te tellurium 869 2.1 723 1261 6.232 137
53 I iodine 1008 2.7 387 457 4.933 136
54 Xe xenon 1170. 2.6 161 165 0.005366 136
55 Cs cesium 376 0.8 302 944 1.873 238
56 Ba barium 503 0.9 1000. 2170. 3.62 206
57 La lanthanum 538 1.1 1193 3737 6.15 194
Elements 58–71 have been omitted.
72 Hf hafnium 659 1.3 2506 4876 13.3 164
73 Ta tantalum 728 1.5 3290. 5731 16.4 158
74 W tungsten 759 1.7 3695 5828 19.3 150.
75 Re rhenium 756 1.9 3458 5869 20.8 141
76 Os osmium 814 2.2 3306 5285 22.587 136
77 Ir iridium 865 2.2 2719 4701 22.562 132
78 Pt platinum 864 2.2 2041 4098 21.5 130.
79 Au gold 890. 2.4 1337 3129 19.3 130.
80 Hg mercury 1007 1.9 234 630. 13.5336 132
81 Tl thallium 589 1.8 577 1746 11.8 144
82 Pb lead 716 1.8 600. 2022 11.3 145
83 Bi bismuth 703 1.9 544 1837 9.79 150.
84 Po polonium 812 2.0 527 1235 9.20 142
85 At astatine — 2.2 575 — — 148
86 Rn radon 1037 — 202 211 0.009074 146
87 Fr francium 393 0.7 300. — — 242
88 Ra radium 509 0.9 969 — 5 211
89 Ac actinium 499 1.1 1323 3471 10. 201
Elements 90 and above have been omitted.
*boiling point at standard pressure
**density of solids and liquids at room temperature and density of gases at 298 K and 101.3 kPa
— no data available
Source: CRC Handbook for Chemistry and Physics, 91 st ed., 2010–2011, CRC Press
Reference Tables for Physical Setting/Chemistry – 2011 Edition 11

Table T
Important Formulas and Equations
d = density
m
Density d = m = mass
V
V = volume
Mole Calculations number of moles =
given mass
gram-formula mass
measured value – accepted value
Percent Error % error = × 100
accepted value
mass of part
Percent Composition % composition by mass = × 100
mass of whole
mass of solute
parts per million = × 1000000
mass of solution
Concentration
molarity =
moles of solute
liter of solution
Combined Gas Law
P
P = pressure
1
V 1
P
= 2
V 2
V = volume
T 1
T 2 T = temperature
M A
= molarity of H + M B
= molarity of OH –
Titration M A
V A
= M B
V B
V A
= volume of acid V B
= volume of base
q = mCΔT q = heat H f
= heat of fusion
Heat q = mH f
m = mass H v
= heat of vaporization
q = mH v
C=specific heat capacity
ΔT = change in temperature
Temperature
K = °C + 273
K = kelvin
°C = degree Celsius
DET 609 ADU
Reference Tables for Physical Setting/Chemistry – 2011 Edition 12

Common Lab Equipment Uses

SCIENTIFIC NOTATION RULES
How to Write Numbers in Scientific Notation
Scientific notation is a standard way of writing very large and very small numbers so that they're
easier to both compare and use in computations. To write in scientific notation, follow the form
N X 10 ᴬ
where N is a number between 1 and 10, but not 10 itself, and A is an integer (positive or negative
number).
RULE #1: Standard Scientific Notation is a number from 1 to 9 followed by a decimal and the
remaining significant figures and an exponent of 10 to hold place value.
Example:
5.43 x 10 2 = 5.43 x 100 = 543
8.65 x 10 – 3 = 8.65 x .001 = 0.00865
****54.3 x 10 1 is not Standard Scientific Notation!!!
RULE #2: When the decimal is moved to the Left the exponent gets Larger, but the value of the
number stays the same. Each place the decimal moves Changes the exponent by one (1). If you
move the decimal to the Right it makes the exponent smaller by one (1) for each place it is moved.
Example:
6000. x 10 0 = 600.0 x 10 1 = 60.00 x 10 2 = 6.000 x 10 3 = 6000
(Note: 10 0 = 1)
All the previous numbers are equal, but only 6.000 x 10 3 is in proper Scientific Notation.

RULE #3: To add/subtract in scientific notation, the exponents must first be the same.
Example:
(3.0 x 10 2 ) + (6.4 x 10 3 ); since 6.4 x 10 3 is equal to 64. x 10 2 . Now add.
(3.0 x 10 2 )
+ (64. x 10 2 )
67.0 x 10 2 = 6.70 x 10 3 = 6.7 x 10 3
67.0 x 10 2 is mathematically correct, but a number in standard scientific notation can only
have one number to the left of the decimal, so the decimal is moved to the left one place and
one is added to the exponent.
Following the rules for significant figures, the answer becomes 6.7 x 10 3 .
RULE #4: To multiply, find the product of the numbers, then add the exponents.
Example:
(2.4 x 10 2 ) (5.5 x 10 –4 ) = ? [2.4 x 5.5 = 13.2]; [2 + -4 = -2], so
(2.4 x 10 2 ) (5.5 x 10 –4 ) = 13.2 x 10 –2 = 1.3 x 10 – 1
RULE #5: To divide, find the quotient of the number and subtract the exponents.
Example:
(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = ? [3.3 / 9.1 = .36]; [-6 – (-8) = 2], so
(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = .36 x 10 2 = 3.6 x 10 1

1. 7,485 6. 1.683
2. 884.2 7. 3.622
3. 0.00002887 8. 0.00001735
4. 0.05893 9. 0.9736
5. 0.006162 10. 0.08558
11. 6.633 X 10−⁴ 16. 1.937 X 10⁴
12. 4.445 X 10−⁴ 17. 3.457 X 10⁴
13. 2.182 X 10−³ 18. 3.948 X 10−⁵
14. 4.695 X 10² 19. 8.945 X 10⁵
15. 7.274 X 10⁵ 20. 6.783 X 10²

Convert each number from Scientific Notation to real numbers:
1. 7.485 X 10³ 6. 1.683 X 10⁰
2. 8.842 X 10² 7. 3.622 10⁰
3. 2.887 X 10−⁵ 8. 1.735 X 10−⁵
4. 5.893 X 10−² 9. 9.736 X 10−¹
5. 6.162 X 10−³ 10. 8.558 X 10−²
Convert each number from a real number to Scientific Notation:
11. 0.0006633 16. 1,937,000
12. 0.0004445 17. 34,570
13. 0.002182 18. 0.00003948
14. 469.5 19. 894,500
15. 727,400 20. 678.3

Significant Figures Rules
There are three rules on determining how many significant figures are in a
number:
1. Non-zero digits are always significant.
2. Any zeros between two significant digits are significant.
3. A final zero or trailing zeros in the DECIMAL PORTION ONLY are
significant.
Please remember that, in science, all numbers are based upon measurements (except for a very few
that are defined). Since all measurements are uncertain, we must only use those numbers that are
meaningful.
Not all of the digits have meaning (significance) and, therefore, should not be written down. In
science, only the numbers that have significance (derived from measurement) are written.
Rule 1: Non-zero digits are always significant.
If you measure something and the device you use (ruler, thermometer, triple-beam, balance, etc.)
returns a number to you, then you have made a measurement decision and that ACT of measuring
gives significance to that particular numeral (or digit) in the overall value you obtain.
Hence a number like 46.78 would have four significant figures and 3.94 would have three.
Rule 2: Any zeros between two significant digits are significant.
Suppose you had a number like 409. By the first rule, the 4 and the 9 are significant. However, to
make a measurement decision on the 4 (in the hundred's place) and the 9 (in the one's place), you
HAD to have made a decision on the ten's place. The measurement scale for this number would have
hundreds, tens, and ones marked.
Like the following example:
These are sometimes called "captured zeros."
If a number has a decimal at the end (after the one’s place) then all digits (numbers) are significant
and will be counted.
In the following example the zeros are significant digits and highlighted in blue.
960.
70050.

Rule 3: A final zero or trailing zeros in the decimal portion ONLY are
significant.
This rule causes the most confusion among students.
In the following example the zeros are significant digits and highlighted in blue.
0.07030
0.00800
Here are two more examples where the significant zeros are highlighted in blue.
When Zeros are Not Significant Digits
4.7 0 x 10−³
6.5 0 0 x 10⁴
Zero Type # 1 : Space holding zeros in numbers less than one.
In the following example the zeros are NOT significant digits and highlighted in red.
0.09060
0.00400
These zeros serve only as space holders. They are there to put the decimal point in its correct
location.
They DO NOT involve measurement decisions.
Zero Type # 2 : Trailing zeros in a whole number.
In the following example the zeros are NOT significant digits and highlighted in red.
200
25000
For addition and subtraction, look at the decimal portion (i.e., to the right of the decimal point)
of the numbers ONLY. Here is what to do:
1) Count the number of significant figures in the decimal portion of each number in the problem. (The
digits to the left of the decimal place are not used to determine the number of decimal places in the
final answer.)
2) Add or subtract in the normal fashion.
3) Round the answer to the LEAST number of places in the decimal portion of any number in the
problem
The following rule applies for multiplication and division:
The LEAST number of significant figures in any number of the problem determines the number of
significant figures in the answer.
This means you MUST know how to recognize significant figures in order to use this rule.

How Many Significant Digits for Each Number?
1) 2359 = ______
2) 2.445 x 10−⁵= ______
3) 2.93 x 10⁴= ______
4) 1.30 x 10−⁷= ______
5) 2604 = ______
6) 9160 = ______
7) 0.0800 = ______
8) 0.84 = ______
9) 0.0080 = ______
10) 0.00040 = ______
11) 0.0520 = ______
12) 0.060 = ______
13) 6.90 x 10−¹= ______
14) 7.200 x 10⁵= ______
15) 5.566 x 10−²= ______
16) 3.88 x 10⁸= ______
17) 3004 = ______
18) 0.021 = ______
19) 240 = ______
20) 500 = ______

For addition and subtraction, look at the decimal portion (i.e., to the right of the decimal point) of the
numbers ONLY. Here is what to do:
1) Count the number of significant figures in the decimal portion of each number in the problem. (The
digits to the left of the decimal place are not used to determine the number of decimal places in the
final answer.)
2) Add or subtract in the normal fashion.
3) Round the answer to the LEAST number of places in the decimal portion of any number in the
problem.
Solve the Problems and Round Accordingly...
1) 43.287 + 5.79 + 6.284 = _______
2) 87.54 - 3.3 = _______
3) 99.1498 + 6.5397 + 9.7 = _______
4) 5.868 - 5.1 = _______
5) 59.9233 + 86.21 + 99.396 = _______
6) 7.7 + 26.756 = _______
7) 66.8 + 2.3 + 4.8516 = _______
8) 9.7419 + 43.545 = _______
9) 4.8976 + 48.4644 + 1.514 = _______
10) 4.335 + 35.85 = _______
11) 9.448 - 1.7 = _______
12) 75.826 - 8.6555 = _______
13) 57.2 + 23.814 = _______
14) 77.684 - 4.394 = _______
15) 26.4496 + 3.339 = _______
16) 9.6848 + 29.85 = _______
17) 63.11 + 2.5412 + 4.93 = _______
18) 11.2471 + 75.4 = _______
19) 73.745 - 8.755 = _______
20) 6.5238 + 1.7 + 27.79 = _______

The following rule applies for multiplication and division:
The LEAST number of significant figures in any number of the problem determines the number of
significant figures in the answer.
This means you MUST know how to recognize significant figures in order to use this rule.
Solve the Problems and Round Accordingly...
1) 0.6 x 65.0 x 602 = __________
2) 720 ÷ 7.7 = __________
3) 929 x 0.3 = __________
4) 300 ÷ 44.31 = __________
5) 608 ÷ 9.8 = __________
6) 0.06 x 0.079 = __________
7) 0.008 x 72.91 x 7000 = __________
8) 73.94 x 67 x 780 = __________
9) 0.62 x 0.097 x 40 = __________
10) 600 x 10 x 5030 = __________
11) 5200 ÷ 4.46 = __________
12) 0.0052 x 0.4 x 107 = __________
13) 0.099 x 8.8 = __________
14) 0.0095 x 5.2 = __________
15) 8000 ÷ 4.62 = __________
16) 0.6 x 0.8 = __________
17) 2.84 x 0.66 = __________
18) 0.5 x 0.09 = __________
19) 8100 ÷ 34.84 = __________
20) 8.24 x 6.9 x 8100 = __________

Question Sig Figs Question Add & Subtract Question Multiple & Divide
1 4 1 55.36 1 20,000
2 4 2 84.2 2 94
3 3 3 115.4 3 300
4 3 4 0.8 4 7
5 4 5 245.53 5 62
6 3 6 34.5 6 0.005
7 3 7 74.0 7 4,000
8 2 8 53.287 8 3,900,000
9 2 9 54.876 9 2
10 2 10 40.19 10 30,000,000
11 3 11 7.7 11 1,200
12 2 12 67.170 12 0.2
13 3 13 81.0 13 0.87
14 4 14 73.290 14 0.049
15 4 15 29.789 15 2,000
16 3 16 39.53 16 0.5
17 4 17 70.58 17 1.9
18 2 18 86.6 18 0.05
19 2 19 64.990 19 230
20 1 20 36.0 20 460,000

To use the Stair-Step method, find the prefix the original measurement starts with. (ex. milligram)
If there is no prefix, then you are starting with a base unit.
Find the step which you wish to make the conversion to. (ex. decigram)
Count the number of steps you moved, and determine in which direction you moved (left or right).
The decimal in your original measurement moves the same number of places as steps you moved and in the
same direction. (ex. milligram to decigram is 2 steps to the left, so 40 milligrams = .40 decigrams)
If the number of steps you move is larger than the number you have, you will have to add zeros to hold the
places. (ex. kilometers to meters is three steps to the right, so 10 kilometers would be equal to 10,000 m)
That’s all there is to it! You need to be able to count to 6, and know your left from your right!
1) Write the equivalent
a) 5 dm =_______m b) 4 mL = ______L c) 8 g = _______mg
d) 9 mg =_______g e) 2 mL = ______L f) 6 kg = _____g
g) 4 cm =_______m h) 12 mg = ______ g i) 6.5 cm 3 = _______L
j) 7.02 mL =_____cm 3 k) .03 hg = _______ dg l) 6035 mm _____cm
m) .32 m = _______cm n) 38.2 g = _____kg

2. One cereal bar has a mass of 37 g. What is the mass of 6 cereal bars? Is that more than or less
than 1 kg? Explain your answer.
3. Wanda needs to move 110 kg of rocks. She can carry l0 hg each trip. How many trips must she
make? Explain your answer.
4. Dr. O is playing in her garden again She needs 1 kg of potting soil for her plants. She has 750 g.
How much more does she need? Explain your answer.
5. Weather satellites orbit Earth at an altitude of 1,400,000 meters. What is this altitude in kilometers?
6. Which unit would you use to measure the capacity? Write milliliter or liter.
a) a bucket __________
b) a thimble __________
c) a water storage tank__________
d) a carton of juice__________
7. Circle the more reasonable measure:
a) length of an ant 5mm or 5cm
b) length of an automobile 5 m or 50 m
c) distance from NY to LA 450 km or 4,500 km
d) height of a dining table 75 mm or 75 cm
8. Will a tablecloth that is 155 cm long cover a table that is 1.6 m long? Explain your answer.
9. A dollar bill is 15.6 cm long. If 200 dollar bills were laid end to end, how many meters long would
the line be?
10. The ceiling in Jan’s living room is 2.5 m high. She has a hanging lamp that hangs down 41 cm.
Her husband is exactly 2 m tall. Will he hit his head on the hanging lamp? Why or why not?

Using SI Units
Match the terms in Column II with the descriptions in Column I. Write the letters of the correct term in
the blank on the left.
Column I Column II
_____ 1. distance between two points
a. time
_____ 2. SI unit of length
_____ 3. tool used to measure length
_____ 4. units obtained by combining other units
_____ 5. amount of space occupied by an object
_____ 6. unit used to express volume
_____ 7. SI unit of mass
_____ 8. amount of matter in an object
_____ 9. mass per unit of volume
_____ 10. temperature scale of most laboratory thermometers
_____ 11. instrument used to measure mass
_____ 12. interval between two events
_____ 13. SI unit of temperature
_____ 14. SI unit of time
_____ 15. instrument used to measure temperature
b. volume
c. mass
d. density
e. meter
f. kilogram
g. derived
h. liter
i. second
j. Kelvin
k. length
1. balance
m. meterstick
n. thermometer
o. Celsius
Circle the two terms in each group that are related. Explain how the terms are related.
16. Celsius degree, mass, Kelvin _____________________________________________________
________________________________________________________________________________
17. balance, second, mass __________________________________________________________
________________________________________________________________________________
18. kilogram, liter, cubic centimeter __________________________________________________
________________________________________________________________________________
19. time, second, distance __________________________________________________________
________________________________________________________________________________
20. decimeter, kilometer, Kelvin _____________________________________________________
________________________________________________________________________________

1. How many meters are in one kilometer? __________
2. What part of a liter is one milliliter? __________
3. How many grams are in two dekagrams? __________
4. If one gram of water has a volume of one milliliter, what would the mass of one liter of water be in
kilograms?__________
5. What part of a meter is a decimeter? __________
In the blank at the left, write the term that correctly completes each statement. Choose from the terms
listed below.
Metric SI standard ten
prefixes ten tenth
6. An exact quantity that people agree to use for comparison is a ______________ .
7. The system of measurement used worldwide in science is _______________ .
8. SI is based on units of _______________ .
9. The system of measurement that was based on units of ten was the _______________ system.
10. In SI, _______________ are used with the names of the base unit to indicate the multiple of ten
that is being used with the base unit.
11. The prefix deci- means _______________ .

Standards of Measurement
Fill in the missing information in the table below.
S.I prefixes and their meanings
Prefix
Meaning
0.001
0.01
deci- 0.1
10
hecto- 100
1000
Circle the larger unit in each pair of units.
1. millimeter, kilometer 4. centimeter, millimeter
2. decimeter, dekameter 5. hectogram, kilogram
3. hectogram, decigram
6. In SI, the base unit of length is the meter. Use this information to arrange the following units of
measurement in the correct order from smallest to largest.
Write the number 1 (smallest) through 7 - (largest) in the spaces provided.
_____ a. kilometer
_____ b. centimeter
_____ c. meter
_____ e. hectometer
_____ f. millimeter
_____ g. decimeter
_____ d. dekameter
Use your knowledge of the prefixes used in SI to answer the following questions in the spaces
provided.
7. One part of the Olympic games involves an activity called the decathlon. How many events do you
think make up the decathlon?_____________________________________________________
8. How many years make up a decade? _______________________________________________
9. How many years make up a century? ______________________________________________
10. What part of a second do you think a millisecond is? __________________________________

Dimensional Analysis
This is a way to convert from one unit of a given substance to
another unit using ratios or conversion units. What this video
www.youtube.com/watch?v=aZ3J60GYo6U
Let’ look at a couple of examples:
1. Convert 2.6 qt to mL.
First we need a ratio or conversion unit so that we can go from quarts to milliliters. 1.00 qt = 946 mL
Next write down what you are starting with
2.6 qt
Then make you conversion tree
2.6 qt
Then fill in the units in your ratio so that you can cancel out the original unit and will be left with the
unit you need for the answer. Cross out units, one at a time that are paired, and one on top one on
the bottom.
2.6 qt mL
qt
Now fill in the values from the ratio.
2.6 qt 946 mL
1.00 qt
Now multiply all numbers on the top and multiply all numbers on the bottom and write them as a
fraction.
2.6 qt 946 mL = 2,459.6 mL
1.00 qt 1.00
Now divide the top number by the bottom number and write that number with the unit that was not
crossed out.

1qt=32 oz 1gal = 4qts 1.00 qt = 946 mL 1L = 1000mL
2. Convert 8135.6 mL to quarts
=
3. Convert 115.2 oz to mL
=
4. Convert 2.3 g to Liters
=
5. Convert 8.42 L to oz
=
Go to http://science.widener.edu/svb/tutorial/ chose #7 “Converting Volume” and do 5 more in the
space provided.
1. Convert _________ to _________
=
2. Convert _________ to _________
=
3. Convert _________ to _________
=
4. Convert _________ to _________
=
5. Convert _________ to _________
=

Converting Real Things
Table 1
Using the scale, come up with a conversion ratio just by looking at the scale and prove that it works
but converting 10 grams to ounces. (Hint: 7grams and 15 grams)
Table 2
Using a 50 mL graduated cylinder, fill a 600mL beaker 66.7% full.
Table 3
Using the measuring cup determine how many mL are in 4 ounces of water. Make a conversion ratio
you could use to do other conversions.
Table 4
Using a meter stick, measure the back table in inches, feet and yards and convert them into
centimeters and meters.
Table 5
Using the scale, find the mass of a text book on your table in ounces (round to the nearest ounce)
and cover it to grams.
Table 6
Using the ruler, measure the length of a piece of paper in inches and then convert that into meters.
Make a conversion ratio you could use to do other conversions.

Atoms Are Building Blocks
Atoms are the basis of chemistry. They are the basis for everything in the Universe. You
should start by remembering that matter is composed of atoms. Atoms and the study of
atoms are a world unto themselves. We're going to cover basics like atomic structure
and bonding between atoms.
Smaller Than Atoms?
Are there pieces of matter that are smaller than atoms?
Sure there are. You'll soon be learning that atoms are
composed of pieces like electrons, protons, and neutrons.
But guess what? There are even smaller particles moving
around in atoms. These super-small particles can be found
inside the protons and neutrons. Scientists have many
names for those pieces, but you may have heard of
nucleons and quarks. Nuclear chemists and physicists
work together at particle accelerators to discover the
presence of these tiny, tiny, tiny pieces of matter.
Even though super-tiny atomic particles exist, you only
need to remember the three basic parts of an atom: electrons, protons, and neutrons.
What are electrons, protons, and neutrons? A picture works best to show off the idea.
You have a basic atom. There are three types of pieces in that atom: electrons, protons,
and neutrons. That's all you have to remember. Three things! As you know, there are
almost 120 known elements in the periodic table. Chemists and physicists haven't
stopped there. They are trying to make new ones in labs every day. The thing that
makes each of those elements different is the number of electrons, protons, and
neutrons. The protons and neutrons are always in the center of the atom. Scientists call
the center region of the atom the nucleus. The nucleus in
a cell is a thing. The nucleus in an atom is a place where
you find protons and neutrons. The electrons are always
found whizzing around the center in areas called shells or
orbitals.
You can also see that each piece has either a "+", "-", or a
"0." That symbol refers to the charge of the particle. Have
you ever heard about getting a shock from a socket, static
electricity, or lightning? Those are all different types of
electric charges. Those charges are also found in tiny particles of matter. The electron
always has a "-", or negative, charge. The proton always has a "+", or positive, charge. If
the charge of an entire atom is "0", or neutral, there are equal numbers of positive and
negative pieces. Neutral means there are equal numbers of electrons and protons. The
third particle is the neutron. It has a neutral charge, also known as a charge of zero. All
atoms have equal numbers of protons and electrons so that they are neutral. If there are
more positive protons or negative electrons in an atom, you have a special atom called
an ion.

http://www.learner.org/interactives/periodic/basics_interactive.html

Looking at Ions
We haven’t talked about ions before, so let’s get down to basics. The
atomic number of an element, also called a proton number, tells you the
number of protons or positive particles in an atom. A normal atom has a
neutral charge with equal numbers of positive and negative particles.
That means an atom with a neutral charge is one where the number of
electrons is equal to the atomic number. Ions are atoms with extra
electrons or missing electrons. When you are missing an electron or
two, you have a positive charge. When you have an extra electron
or two, you have a negative charge.
What do you do if you are a sodium (Na) atom? You have eleven
electrons — one too many to have an entire shell filled. You need to
find another element that will take that electron away from you. When you lose that
electron, you will you’ll have full shells. Whenever an atom has full shells, we say it is
"happy." Let's look at chlorine (Cl). Chlorine has seventeen electrons and only needs
one more to fill its third shell and be "happy." Chlorine will take your extra sodium
electron and leave you with 10 electrons inside of two filled shells. You are now a happy
atom too. You are also an ion and missing one electron. That missing electron gives you
a positive charge. You are still the element sodium, but you are now a sodium ion (Na + ).
You have one less electron than your atomic number.
Ion Characteristics
So now you've become a sodium ion. You have ten electrons.
That's the same number of electrons as neon (Ne). But you
aren't neon. Since you're missing an electron, you aren't really
a complete sodium atom either. As an ion you are now
something completely new. Your whole goal as an atom was
to become a "happy atom" with completely filled electron
shells. Now you have those filled shells. You have a lower
energy. You lost an electron and you are "happy." So what
makes you interesting to other atoms? Now that you have
given up the electron, you are quite electrically attractive.
Other electrically charged atoms (ions) of the opposite charge
(negative) are now looking at you and seeing a good partner to
bond with. That's where the chlorine comes in. It's not only chlorine. Almost any ion with
a negative charge will be interested in bonding with you.

Electrovalence
Don't get worried about the big word. Electrovalence is just another word for something
that has given up or taken electrons and become an ion. If you look at the periodic table,
you might notice that elements on the left side usually become positively charged ions
(cations) and elements on the right side get a negative charge (anions). That trend
means that the left side has a positive valence and the right side has a negative
valence. Valence is a measure of how much an atom wants to bond with other atoms. It
is also a measure of how many electrons are excited about bonding with other atoms.
There are two main types of bonding, covalent and electrovalent. You may have heard
of the term "ionic bonds." Ionic bonds are electrovalent bonds. They are just groups of
charged ions held together by electric forces. When in the presence of other ions, the
electrovalent bonds are weaker because of outside electrical forces and attractions.
Sodium and chlorine ions alone have a very strong bond, but as soon as you put those
ions in a solution with H + (Hydrogen ion), OH - (Hydroxide), F - (Fluorine ion) or Mg ++
(Magnesium ion), there are charged distractions that break the Na-Cl bond.
Look at sodium chloride (NaCl) one more time. Salt is a very strong bond when it is
sitting on your table. It would be nearly impossible to break those ionic/electrovalent
bonds. However, if you put that salt into some water (H 2 O), the bonds break very
quickly. It happens easily because of the electrical attraction of the water. Now you have
sodium (Na + ) and chlorine (Cl - ) ions floating around the solution. You should remember
that ionic bonds are normally strong, but they are very weak in water.

Neutron Madness
We have already learned that ions are atoms that are
either missing or have extra electrons. Let's say an atom
is missing a neutron or has an extra neutron. That type of
atom is called an isotope. An atom is still the same
element if it is missing an electron. The same goes for
isotopes. They are still the same element. They are just a
little different from every other atom of the same element.
For example, there are a lot of carbon (C) atoms in the
Universe. The normal ones are carbon-12. Those atoms have 6 neutrons. There are a
few straggler atoms that don't have 6. Those odd ones may have 7 or even 8 neutrons.
As you learn more about chemistry, you will probably hear about carbon-14. Carbon-14
actually has 8 neutrons (2 extra). C-14 is considered an isotope of the element carbon.
Messing with the Mass
If you have looked at a periodic table, you may have noticed that the atomic mass of
an element is rarely an even number. That happens because of the isotopes. If you are
an atom with an extra electron, it's no big deal. Electrons don't have much of a mass
when compared to a neutron or proton.
Atomic masses are calculated by figuring out the
amounts of each type of atom and isotope there are in
the Universe. For carbon, there are a lot of C-12, a
couple of C-13, and a few C-14 atoms. When you
average out all of the masses, you get a number that is a
little bit higher than 12 (the weight of a C-12 atom). The
average atomic mass for the element is actually 12.011.
Since you never really know which carbon atom you are
using in calculations, you should use the average mass
of an atom.
Bromine (Br), at atomic number 35, has a greater variety of isotopes. The atomic mass
of bromine (Br) is 79.90. There are two main isotopes at 79 and 81, which average out
to the 79.90amu value. The 79 has 44 neutrons and the 81 has 46 neutrons. While it
won't change the average atomic mass, scientists have made bromine isotopes with
masses from 68 to 97. It's all about the number of neutrons. As you move to higher
atomic numbers in the periodic table, you will probably find even more isotopes for
each element.

I did not use publisher for this part because my computer does
not have it to be used. I used a PDF file and used as much of it
as I could.

I did not use publisher for this part because my computer does
not have it to be used. I used a PDF file and used as much of it
as I could.

I did not use publisher for this part because my computer does
not have it to be used. I used a PDF file and used as much of it
as I could.

I did not use publisher for this part because my computer does
not have it to be used. I used a PDF file and used as much of it
as I could.

P
N
P
N

P
N
P
N

Electron Configuration
Color the sublevel:
s = Red
d = Green
p = Blue
f = Orange
Write in sublevels
Write period, sublevel and super scripts.
Ctrl Shift =
gives you super scripts

www.youtube.com/watch?v=jtYzEzykFdg
www.youtube.com/watch?
annotation_id=annotation_2076&feature=iv&src_vid=jtYzEzykFdg&v=cOlac8ruD_0
www.youtube.com/watch?
annotation_id=annotation_570977&feature=iv&src_vid=cOlac8ruD_0&v=lR2vqHZWb5A

Electron Configuration
In order to write the electron configuration for an atom you must know the 3 rules of
electron configurations.
1. Aufbau
Notation
nO e
where
n is the energy level
O is the orbital type (s, p, d, or f)
e is the number of electrons in that orbital shell
Principle
electrons will first occupy orbitals of the lowest energy level
2. Hund rule
when electrons occupy orbitals of equal energy, one electron enters each orbital until
all the orbitals contain one electron with the same spin.
3. Pauli exclusion principle
an orbital contains a maximum of 2 electrons and
paired electrons will have opposite spin

In the space below, write the unabbreviated electron configurations of the following elements:
1) sodium ________________________________________________
2) iron ________________________________________________
3) bromine ________________________________________________
4) barium ________________________________________________
5) neptunium ________________________________________________
In the space below, write the abbreviated electron configurations of the following elements:
6) cobalt ________________________________________________
7) silver ________________________________________________
8) tellurium ________________________________________________
9) radium ________________________________________________
10) lawrencium ________________________________________________
Determine what elements are denoted by the following electron configurations:
11) 1s²s²2p⁶3s²3p⁴ ____________________
12) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s¹ ____________________
13) [Kr] 5s²4d¹⁰5p³ ____________________
14) [Xe] 6s²4f¹⁴5d⁶ ____________________
15) [Rn] 7s²5f¹¹ ____________________
Identify the element or determine that it is not a valid electron configuration:
16) 1s²2s²2p⁶3s²3p⁶4s²4d¹⁰4p⁵ ____________________
17) 1s²2s²2p⁶3s³3d⁵ ____________________
18) [Ra] 7s²5f⁸ ____________________
19) [Kr] 5s²4d¹⁰5p⁵ ____________________
20) [Xe] ____________________
1)sodium 1s 2 2s 2 2p 6 3s 1 2)iron 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6
3)bromine 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 4)barium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2
5)neptunium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 5 6)cobalt [Ar] 4s 2 3d 7
7)silver [Kr] 5s 2 4d 9 8)tellurium[Kr] 5s 2 4d 10 5p 4
9)radium [Rn] 7s 2 10)lawrencium[Rn] 7s 2 5f 14 6d 1
1s 2 2s 2 2p 6 3s 2 3p 4 sulfur 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 rubidium
[Kr] 5s 2 4d 10 5p 3 antimony [Xe] 6s 2 4f 14 5d 6 osmium
[Rn] 7s 2 5f 11 einsteinium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4d 10 4p 5 not valid (take a look at “4d”)
1s 2 2s 2 2p 6 3s 3 3d 5 not valid (3p comes after 3s) [Ra] 7s 2 5f 8 not valid (radium isn’t a noble gas)
[Kr] 5s 2 4d 10 5p 5 valid iodine
20)[Xe] not valid (an element can’t be its own electron configuration)

Using Wikipedia define the 8 categories of elements on pages 168 and 169 on the left
page.
Color your periodic table similar to the one on Pages 168—169.

Create groups for these Scientist and explain your groupings
(use the information you got from your research)

Research these Scientist and summarize their contributions to chemistry
Antoine Henri Becquerel
Niels Bohr
Louis de Barogilie
Glenn Seaborg
Hantaro Nagaoka
Democritus
Marie and Pierre Curie
Eugene Goldstein
Dmitri Mendeleev
J.J. Thomson
James Chadwick
Erwin Shrodinger
John Dalton
Lothar Meyer
Robert Millikan
J.W. Dobereiner
Ernest Rutherford

Chapter 25 Notes
Vocab:
-25.1 Nuclear Radiation-
Radioactivity – reference to the spontaneous emission of rays or particles from certain
elements
Nuclear radiation – the name of the rays and particles emitted from a radioactive source
Radiosiotopes – unstable isotopes
Alpha particle – the helium nuclei emitted by some radioactive sources, has 2 protons
and 2 neutrons
Beta particles – an electron resulting from the breaking apart of an atom, has 1 electron
and 1 proton, but no neutron
Gamma Ray – a high-energy photon emitted by a radioisotope, electromagnetic
radiation with high penetration through the human body
Notes:
Antoine Henri Bequerel and Marie and Pierre Curie accidentally discovered that uranium
salt emitted radioactive rays because of a fog left on a photographic plate, and used
radioactivity to explain why this happened.
Atoms become more stable in nuclear or chemical reactions.
Nuclear reactions begin with unstable radioisotopes, which later become much more
stable after the reaction.
Unlike chemical reactions, nuclear reactions are not affected by changes in
temperature, pressure, or the presence of catalysts. Nuclear reactions of a given
radioisotope can’t be slowed down, sped up, or stopped.
Radioactive decay doesn’t require input energy. Radioactivity is continued until unstable
isotopes are changed to become stable isotopes of a different element.
Type Particles Symbol Charge Mass Source Penetration
Alpha Radiation Helium nuclei α 2+ 4 Radium Low (0.05
mm body
tissue)
Beta Radiation Electrons β 1- 1/1837 Carbon Moderate (4
mm body
tissue)

Gamma
Radiation
High energy
electromagnetic
radiation
γ 0 0 Cobalt Very High
(penetrates
body easily)
-Chapter 25.2 Nuclear Transformations-
Vocab:
Nuclear force – an attractive force that acts between all nuclear particles that are
extremely close together
Positron – a particle with the mass of an electron but a positive charge
Half-life – the time required for one-half of the nuclei in a radioisotope sample to decay
to products
Transmutation – the conversion of one atom of one element into another atom of
another element
Transuranium elements – elements with atomic numbers above 92 (uranium)
Notes:
The neutron-to-proton ratio in a radioscope determines the type of decay that occurs.
All nuclei that have an atomic number over 83 are radioactive, these contain too many
neutrons and protons that make the atom unstable.
During each half-life, half of the remaining radioactivity atoms decay into atoms of a
new element.
Half-Life Amount Remaining
0 A^0 X (1/2)^0 = A^O
1 A^0 X (1/2)^1 = A^0
2 A^0 X (1/2)^2 = A^0
Radiocarbon dating measures the amount of the isotope carbon-14 an object contains.
Can only be used the measure the time line of a carbon-based material or life form.
Transmutation can occur by radioactive decay or when particles attack the nucleus of
an atom.

Learning Goal: The students will learn how molecular bonding is different than ionic bonding and
electrons affect the shape of a molecule and its properties.
Metallic or Ionic Bonds
CCSTD HS Chemistry 2.a..
A chemical bond is the “glue” that holds atoms together. Chemical bonds
can be weak or strong, and everything in between!
One thing that all chemical bonds share is that they are just an atom’s
attempt to complete its outer (valence) shell of electrons. The closer an atom’s
valence shell comes to being completely full, then the closer that atom comes to
resemble an inert, or noble, gas.
It turns out that all atoms strive to have the electronic configuration of an
inert gas atom! Why is that such a desirable state? Inert gases are extremely stable.
In fact, inert gases are so stable that they don’t bond, ever, to anything. They are
perfectly content to remain just as they are.
Ionic Bonds
Another type of chemical bond, the ionic bond, involves a somewhat
different arrangement of electrons within a molecule. Ionic bonds form when a
“give and take” situation exists between atoms. One atom within the molecule

gives an electron (to leave behind a positively charged atom), and another atom
takes an electron (to result in a negatively charged atom).
You can just imagine what will happen if you put a positively charged atom
(cation) next to a negatively charged atom (anion)! They are attracted to each other
and we call this type of attraction an ionic bond.
Think of an ionic bond as one valence shell cloud with a positive charge
being attracted to another valence shell cloud with a negative charge. When those
two clouds are next to each other, they “stick” together and the resultant molecule
has no net charge.
Table salt, or Sodium Chloride (NaCl), is a compound that has an ionic bond
(see Figure 3.1.2). The Sodium within the molecule donates its electron and the
Chloride accepts the electron. This condition gives the Sodium atom the electronic
configuration of a Neon atom (an inert gas), and the Chloride atom the electronic
configuration of an Argon atom (another inert gas!).
` An ionic bond can involve the loss of more than one electron, with the other
atom in the ionic compound gaining the same number of electrons. Magnesium
Chloride (MgCl2) is an example of a doubly ionic compound. Once again,
Magnesium Chloride has no net charge, and in this case, one Magnesium atom has
lost two electrons (achieving the electronic configuration of Neon), with each of
two Chloride atoms accepting one electron.
Ionic bonds have distinctive properties of:
When in solid form, all ionic compounds are crystalline solids at room temperature
(the alternating positive and negative ions are arranged in a crystal lattice
structure),

Strong ionic bonding forces making the solid structures hard with high melting and
boiling points,
Multiple charges within the ionic bond lead to stronger bonding forces (i.e. the
bonding of MgCl2 is stronger than the bonding of NaCl), and
Most ionic compounds are soluble in water (once they are fused (melted), or
dissolved in water, ionic compounds conduct electricity because the anions are
then free to move freely). Most often, ionic compounds are combinations of
cations of atoms from the furthest left of the periodic (groups IA and IIA), with
anions of atoms from the right of the periodic table (group VIIA), immediately
preceding the inert gases.
Metallic Bonds
The metallic bond is actually a special case of an ionic bond. Think of metallic
bonding as a more or less static arrangement of positive ions within a moving array
of mobile electrons.
This unique arrangement of cations and electrons give metals their characteristic
properties. Metallic properties include.
Heat conductivity: mobile electrons can carry the kinetic energy of heat.
Shiny appearance: the rapidly moving electrons emit energy in the form of light.
Electricity conductors: electricity is the flow of electrons.
Malleability: ability to be easily shaped into flat sheets or drawn into wires.

More notes…
Ionic Bonds
Ionic bonds are the fusing together of metals and nonmetals
Cross multiply the charges, and the charges become subscripts
Names: Metal Nonmetal (NaCl), the nonmetal adds the suffix ‘ide’ for the full
name (Sodium Chloride)
The metal atoms give away their electrons to become a positively charged ion, and
the nonmetal takes electrons to become a negatively charged ion.
H 1 0
H 2 0
H 3 0
H 4 0
1-No pronunciation (Hydrogen Oxide)
2-Di (Hydrogen Dioxide)
3- Tri (Hydrogen Trioxide)
4-Tetra (Hydrogen Tetraoxide)

Learning Goal: The students will learn how molecular bonding is different than ionic bonding and
electrons affect the shape of a molecule and it's properties.
Covalent Bond Notes
When 2 nonmetals bond together, they share electrons
Charges are not needed, because the mixtures are always stabilized.
Sometimes they are called strong bonds because of the high difficulty of breaking these atoms
apart aside from using special enzymes.
Hydrogen and ionic bonds are easily broken by anything as small as temperature change.
Covalent bonds form definite and easily predictable shapes.
Another form of covalent bonding is polar bonding, where an atom shares atoms with
hydrogen atoms. Water is an example of polar bonding.
The Octet rule includes that all atoms excluding hydrogen and helium have to have 8 total
electrons in their valence shell to become stable.
This process can be achieved through many forms of bonding, with covalent being the most
powerful.

A single bond is when two pair of electrons are shared between two atoms.
A Double pair is when an atom shares two pairs of electrons to another atom
A Triple pair is when an atom shares 3 pairs of electrons (total of 6) with another atom
A polar covalent bond is a bond made, but is not equally shared because of the higher
electronegativity of one of the atoms is higher than the other, which causes higher attraction
with the electron
A nonpolar covalent bond is bond made and is equally shared because of the equal
electronegativity and therefore equally magnetic pull of the electron between the atoms

Name Formula Charge
Dichromate Cr₂O₇ 2-
Sulfate SO₄ 2-
Hydrogen Carbonate HCO₃ 1-
Hypochlorite ClO 1-
Phosphate PO₄ 3-
Nitrite NO₂ 1-
Chlorite ClO₂ 1-
Dihydrogen phosphate H₂PO₄ 1-
Chromate CrO₄ 2-
Carbonate CO₃ 2-
Hydroxide OH 1-
Hydrogen phosphate HPO₄ 2-
Ammonium NH₄ 1+
Acetate C₂H₃O₂ 1-
Perchlorate ClO₄ 1-
Permanganate MnO₄ 1-
Chlorate ClO₃ 1-
Hydrogen Sulfate HSO₄ 1-
Phosphite PO₃ 3-
Sulfite SO₃ 2-
Silicate SiO₃ 2-
Nitrate NO₃ 1-
Hydrogen Sulfite HSO₃ 1-
Oxalate C₂O₄ 2-
Cyanide CN 1-
Hydronium H₃O 1+
Thiosulfate S₂O₃ 2-

Orbital Equation Lone Pairs Angle Name

Linear AX 2
sp
Bent AX 2 E
Bent AX 2 E 2
One pair of electrons Two pairs of electrons.
sp 2 sp 3
Tetrahedral AX 4
Octahedral

Trig.plannar
AX 3 sp 3 Trig. Pyramidal AX 3 E
sp 3 d
T-shaped AX 3 E 2
sp 3 d
Trig.Bipyramidal
AX 5
sp 3 d
Square plannar
AX 4 E 2 sp 3 d 2

www.youtube.com/watch?v=AsqEkF7hcII
www.youtube.com/watch?v=tEn0N4R2dqA
www.youtube.com/watch?v=Pft2CASl0M0
www.youtube.com/watch?v=rwhJklbK8R0
The Mole

www.youtube.com/watch?v=BTRm8PwcZ3U
www.youtube.com/watch?v=F9NkYSKJifs
www.youtube.com/watch?v=xPdqEX_WMjo
Molar Mass

Mole Conversions

Steps for Mole Conversions

Answer the following questions:
1) How many moles are in 25 grams of water?
Mole Calculation Practice
2) How many grams are in 4.5 moles of Li 2 O?
3) How many molecules are in 23 moles of oxygen?
4) How many moles are in 3.4 x 10 23 molecules of H 2 SO 4 ?
5) How many molecules are in 25 grams of NH 3 ?
6) How many grams are in 8.2 x 10 22 molecules of N 2 I 6 ?

7) How many grams does 0.500 moles of CuBr weigh?
8) How many molecules are there in 0.655 moles of C6H14?
9) How many moles are there in 2.35 x 1024 molecules of water?
10) How many grams does 5.60 x 1022 molecules of SiO2 weigh?
11) How many molecules are there in 21.6 grams of CH4?
1) 1.39 moles
2) 134.1 grams
3) 1.38 x 10 25 molecules
4) 0.56 moles
5) 8.85 x 10 23 molecules
6) 106.7 grams
7) How many grams does 0.500 moles of CuBr weigh? 31.8 grams
8) How many molecules are there in 0.655 moles of C 6H 14? 3.94 x 10 23 molecules
9) How many moles are there in 2.35 x 10 24 molecules of water? 3.90 moles
10) How many grams does 5.60 x 10 22 molecules of SiO 2 weigh? 5.59 grams
11) How many molecules are there in 21.6 grams of CH 4? 8.13 x 10 23

How to Balance Chemical Equations
A chemical equation is a theoretical or written representation of what happens during a chemical
reaction. The law of conservation of mass states that no atoms can be created or destroyed in a
chemical reaction, so the number of atoms that are present in the reactants has to balance the
number of atoms that are present in the products. Follow this guide to learn how to balance chemical
equations.
Step 1
Write down your given equation. For this example, we will use:
C 3 H 8 + O 2 --> H 2 O + CO 2
Step 2
Write down the number of atoms that you have on each side of the equation. Look at the subscripts
next to each atom to find the number of atoms in the equation.
Left side: 3 carbon, 8 hydrogen and 2 oxygen
Right side: 1 carbon, 2 hydrogen and 3 oxygen

Step 3
Always leave hydrogen and oxygen for last. This means that you will need to balance the carbon
atoms first.
Step 4
Add a coefficient to the single carbon atom on the right of the equation to balance it with the 3 carbon
atoms on the left of the equation.
C 3 H 8 + O 2 --> H 2 O + 3CO 2
The coefficient 3 in front of carbon on the right side indicates 3 carbon atoms just as the subscript 3
on the left side indicates 3 carbon atoms.
In a chemical equation, you can change coefficients, but you should never alter the subscripts.

Step 5
Balance the hydrogen atoms next. You have 8 on the left side, so you'll need 8 on the right side.
C 3 H 8 + O 2 --> 4H 2 O + 3CO 2
On the right side, we added a 4 as the coefficient because the subscript showed that we already
had 2 hydrogen atoms.
When you multiply the coefficient 4 times the subscript 2, you end up with 8.
Step 6
Finish by balancing the oxygen atoms.
Because we've added coefficients to the molecules on the right side of the equation, the number of
oxygen atoms has changed. We now have 4 oxygen atoms in the water molecule and 6 oxygen
atoms in the carbon dioxide molecule. That makes a total of 10 oxygen atoms.
Add a coefficient of 5 to the oxygen molecule on the left side of the equation. You now have 10
oxygen molecules on each side.
C 3 H 8 + 5O 2 --> 4H 2 O + 3CO 2.
The carbon, hydrogen and oxygen atoms are balanced. Your equation is complete.

1) ___ NaNO 3 + ___ PbO ___ Pb(NO 3 ) 2 + ___ Na 2 O
2) ___ AgI + ___ Fe 2 (CO 3 ) 3 ___ FeI 3 + ___ Ag 2 CO 3
3) ___ C 2 H 4 O 2 + ___ O 2 ___ CO 2 + ___ H 2 O
4) ___ ZnSO 4 + ___ Li 2 CO 3 ___ ZnCO 3 + ___ Li 2 SO 4
5) ___ V 2 O 5 + ___ CaS ___ CaO + ___ V 2 S 5

6) ___ Mn(NO 2 ) 2 + ___ BeCl 2 ___ Be(NO 2 ) 2 + ___ MnCl 2
7) ___ AgBr + ___ GaPO 4 ___ Ag 3 PO 4 + ___ GaBr 3
8) ___ H 2 SO 4 + ___ B(OH) 3 __ B 2 (SO 4 ) 3 + ___ H 2 O
9) ___ S 8 + ___ O 2 ___ SO 2
10) ___ Fe + ___ AgNO 3 ___ Fe(NO 3 ) 2 + ___ Ag

1) 2 NaNO 3 + PbO Pb(NO 3 ) 2 + Na 2 O
2) 6 AgI + Fe 2 (CO 3 ) 3 2 FeI 3 + 3 Ag 2 CO 3
3) C 2 H 4 O 2 + 2 O 2 2 CO 2 + 2 H 2 O
4) ZnSO 4 + Li 2 CO 3 ZnCO 3 + Li 2 SO 4
5) V 2 O 5 + 5 CaS 5 CaO + V 2 S 5
6) Mn(NO 2 ) 2 + BeCl 2 Be(NO 2 ) 2 + MnCl 2
7) 3 AgBr + GaPO 4 Ag 3 PO 4 + GaBr 3
8) 3 H 2 SO 4 + 2 B(OH) 3 B 2 (SO 4 ) 3 + 6 H 2 O
9) S 8 + 8 O 2 8 SO 2
10) Fe + 2 AgNO 3 Fe(NO 3 ) 2 + 2 Ag
Additional Notes:

Categories of Reactions
All chemical reactions can be placed into one of six categories. Here they are, in no
particular order:
1) Synthesis: A synthesis reaction is when two or more simple compounds combine to form a
more complicated one. These reactions come in the general form of:
A + B ---> AB
One example of a synthesis reaction is the combination of iron and sulfur to form iron (II) sulfide:
8 Fe + S 8 ---> 8 FeS
If two elements or very simple molecules combine with each other, it’s probably a synthesis reaction.
The products will probably be predictable using the octet rule to find charges.
2) Decomposition: A decomposition reaction is the opposite of a synthesis reaction - a
complex molecule breaks down to make simpler ones. These reactions come in the general form:
AB ---> A + B
One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen
gas:
2 H 2O ---> 2 H 2 + O 2
If one compound has an arrow coming off of it, it’s probably a decomposition reaction. The products
will either be a couple of very simple molecules, or some elements, or both.
3) Single displacement: This is when one element trades places with another element in a
compound. These reactions come in the general form of:
A + BC ---> AC + B
One example of a single displacement reaction is when magnesium replaces hydrogen in water to
make magnesium hydroxide and hydrogen gas:
Mg + 2 H 2O ---> Mg(OH) 2 + H 2
If a pure element reacts with another compound (usually, but not always, ionic), it’s probably a single
displacement reaction. The products will be the compounds formed when the pure element switches
places with another element in the other compound.
Important note: these reactions will only occur if the pure element on the reactant side of the equation
is higher on the activity series than the element it replaces.

4) Double displacement: This is when the anions and cations of two different molecules
switch places, forming two entirely different compounds. These reactions are in the general form:
AB + CD ---> AD + CB
One example of a double displacement reaction is the reaction of lead (II) nitrate with potassium
iodide to form lead (II) iodide and potassium nitrate:
Pb(NO 3) 2 + 2 KI ---> PbI 2 + 2 KNO 3
If two ionic compounds combine, it’s probably a double displacement reaction. Switch the cations
and balance out the charges to figure out what will be made.
Important note: These reactions will only occur if both reactants are soluble in water and only one
product is soluble in water.
5) Acid-base: This is a special kind of double displacement reaction that takes place when an
acid and base react with each other. The H + ion in the acid reacts with the OH - ion in the base,
causing the formation of water. Generally, the product of this reaction is some ionic salt and water:
HA + BOH ---> H 2O + BA
One example of an acid-base reaction is the reaction of hydrobromic acid (HBr) with sodium
hydroxide:
HBr + NaOH ---> NaBr + H 2O
If an acid and a base combine, it’s an acid-base reaction. The products will be an ionic compound
and water.
6) Combustion: A combustion reaction is when oxygen combines with another compound to
form water and carbon dioxide. These reactions are exothermic, meaning they produce heat. An
example of this kind of reaction is the burning of napthalene:
C 10H 8 + 12 O 2 ---> 10 CO 2 + 4 H 2O
If something that has carbon and hydrogen reacts with oxygen, it’s probably a combustion reaction.
The products will be CO 2 and H 2 O.
Follow this series of questions. When you can answer "yes" to a question, then
stop!
1) Does your reaction have two (or more) chemicals combining to form one chemical? If yes, then it's
a synthesis reaction
2) Does your reaction have one large molecule falling apart to make several small ones? If yes, then
it's a decomposition reaction
3) Does your reaction have any molecules that contain only one element? If yes, then it's a single
displacement reaction
4) Does your reaction have water as one of the products? If yes, then it's an acid-base reaction
5) Does your reaction have oxygen as one of it's reactants and carbon dioxide and water as
products? If yes, then it's a combustion reaction
6) If you haven't answered "yes" to any of the questions above, then you've got a double
displacement reaction.

1) double displacement
2) combustion
3) single displacement
4) double displacement
5) acid-base
6) synthesis
7) decomposition
List what type the following reactions are: (answers)
Predicting Reaction Products – Answers
Predict the products of each of the following chemical reactions. If a reaction will not occur, explain why not:
1) ____ Ag 2 SO 4 + ____ NaNO 3 → no reaction!
Examining this reaction, it appears that a double displacement reaction will occur. This would lead to the conclusion that
the products would be AgNO3 and Na2SO4. However, for this reaction to occur, both reactants and only one of the
products must be soluble in water. If you look up the solubilities on a chart, you’ll find that Ag2SO3 is partly soluble in
water, and all of the other compounds are totally soluble in water. This tells us that this reaction will not occur.
2) ____ NaI + ____ CaSO 4 → no reaction!
Another double displacement reaction, this time with Na2SO4 and CaI2 as products. Because both products are soluble
in water and CaSO4 is only partially soluble in water, the conditions for a successful double displacement reaction are not
met.
3) 2 HNO 3 + 1 Ca(OH) 2 → 1 Ca(NO 3 ) 2 + 2 H 2 O
It’s an acid-base reaction, and acid-base reactions occur readily whether or not the reactants are both soluble in water.
4) 1 CaCO 3 → 1 CaO + 1 CO 2
It’s a decomposition reaction. If you didn’t guess that these were the products, you should have at least known that it was
a decomposition reaction and predicted that this would have broken into its constituent elements, Ca, C, and O2.
5) 1 AlCl 3 (aq) + 1 (NH 4 ) 3 PO4(aq) → AlPO 4 (s) + 3 NH 4 Cl(aq)
This is a double displacement reaction, except in this case both of the reactants and only one product are soluble in
water. Because the conditions for a successful reaction are met, the reaction does occur!
6) ____ Pb + ____ Fe(NO 3 ) 3 →
Though this is a single displacement reaction, lead is lower on the activity series than the iron it would replace. As a
result, this reaction does not occur.
7) 2 C 3 H 6 + 9 O 2 → 6 CO 2 + 6 H 2 O
The reactants suggest that this is a combustion reaction, meaning that the products must be carbon dioxide and water.
Once you figure this out, the only thing left to do is balance it, as shown.
8) 2 Na + 1 CaSO 4 → 1 Na 2 SO 4 + 1 Ca
This should clearly be a single displacement reaction. Because sodium is higher on the activity series than calcium, this
reaction does occur.

Balancing Equations Practice Problems
From Widener Chemistry Practice
1 page

Stoichiometry Practice Problems
From Widener Chemistry Practice
2 pages

Percent Yield Practice Problems
from Widener Chemistry Practice
1 pages

Cody Nichols
Between the particles in gas, it is empty; this allows for more
movement and kinetic energy that is caused by the particles hitting
each other
Because of no attractive and repulsive force between them, it
allows for a balance state of gas. Due to this, the particles stay in a
random walk .
The collisions that occur during a random walk are perfectly
elastic and the kinetic energy is transferred without any lose.
Gases are very small particles, but humans
have learned to harness the thousands
of types of gases that exists,
which aid them in everyday life, like
propane, nitrogen, co 2, and other gases.

The particles in liquid have kinetic
energy , this makes the particles
flow past each other. The flow of theses particles is
referred to as a fluid.
Unlike gas though, liquid particles are atattracted
to each other, depending on weather the
attraction is disrupted determines the physical
Liquid evaporates and then turns into a vapor
or a gas, then it condensates and turns back
into a liquid. The vapor pressure or the force exerted
by a gas above a liquid is caused by the collision
of particles in a container that contains evaporated
liquid.
Water is the most abundant
source of liquid on
earth. H 2 O can take the
form of all three states of
matter( solid, liquid, and
gases). Examples of this
are Ice as a solid, regular
water as a liquid , and a
rain cloud.

The particles in a solid are packed tightly together, they are dense
and not easy to compress.
When heating a solid, the particles start to break down and at
freezing point, the solid is stuck in a fixed position.
There are crystal, allotropes, and non-crystalline solids, which either
have simple cubic unit cells, body centered unit cells or face
centered unit cells.

Changes from a solid to a gas without going through liquid is called
sublimation. Solid goes through sublimation to a vapor, then the vapor goes
to deposition, back to a solid. A phase diagram shows the three states and
what points that they will change at, the only point where they are equal is
the triple point.
Learning goal
The students will learn what are the factors that determine and characteristics that distinguish
gases liquids and solids and how substances change from one state to another.

Boyles law: V1 + P2 = P 2 + V2
rearrange the equation to isolate V2:
substitute the known values for P1 V1 and P2 into the equation and solve
The finished equation is
V2 = P1 X V1 / P2 = 104kPa X 2.50 L/ 12.0 L = 6.48 L

http://www.youtube.com/watch?v=N5xft2fIqQU

The volume of a fixed mass of
directly proportional to its Kelvin
temperature if the pressure is
kept constant.
As temperature increases, so does
the volume of the substance being
exposed to the heat.

V1
T1 = V2
T2
Cross multiply the
numbers
11.
6.80 L
325 C = X
25 C = 6.80
L* 25 C/ 325 C = 170
L/325=0.52 L

http://www.youtube.com/watch?v=7JKVtbe-hV8

Pressure of a gas
corresponds to the kelvin
temperature if the volume
remains constant.

you put the numbers
where they need to be
108/41 = / 22Cross
multiple 2376/41 then
divide=57.9512195,
then you use the least
significant figures, so
the least is 2, 58

http://www.youtube.com/watch?v=gFDS3bwAU-s

Indirect: 2*3=6 3*2=6 P1 x T1 = P2 x T2
Direct: 6/1=6 12/2=6 P1 / T1 = P2 / T2
Combined Gas Law
P1 x V1/ T1 = P2 x V2/ T2
V1 = 3110 mL T1 = x V2 = 651 mL T2 = 80.60 K Indirect equation
3110 / x = 651 / 80.6 3110 * 80.6 / 651 = x 250, 666 / 651 x=385 K
V1 = 2600 mL P1 = x V2 = 6570 mL P2 = 1500 torr Direct equation
2600x = 6570 * 1500 2600x = 985500 x = 3790 torr 3790 torr = 4.965 atm x = 5.0 atm
P1 = 1320 torr V1 = x T1 = 615 C P2 = 1.28 atm V1 = 7.11 L T2 = 905 K CGL equation
Conversions: T1615 C = 888 K P11320 torr = 1.7292 atm
1.7292*x/888 = 1.28*7.11/9051.7292x/888=0.010051.7292x=8.9244x=5.16 L
P=Pressure V=Volume T=Tempurature n=# in mol r=8.31kpq OR 0.082 atm
Ideal Gas Law PxV=nrT P=nrT/V
P = 1470 torr V = 7.87 L T = 712 K n = x r = 0.082 atm IGL equation
Conversions: P1470 torr = 195.51
195.51 = 58.384x/7.87 1538.6637=58.384x 26.354=x x = 26.4
P = 1.36 atm V = 8.53 L T = 499 K n = x r = 0.082 atm
1.36*8.53 = 0.082*499*x 11.6008=40.918x 0.2835 = x x = 0.284 mol

P = Pressure must be in same units
V = Volume must be in same units
T = Temperature must be in Kelvin
Direct relationship = as the temperature increases, the volume increases
V1/T1 = V2/T2
P1/T1 = P2/T2
Indirect relationship = as volume increases, pressure increases
V1 x P1 = V2 x P2
Combined Gas Law
P = Pressure in atm or kPa
V = Volume must be in Liters
N = number of moles
R = Constant either 0.082 atm OR 8.31 kPa, depending on pressure
T = Tempurature in Kelvin
PV = NRT
Ideal Gas Law
P1 = 2.27 atm V1 = 7.94 L P2 =? V2 = 10.8 L
2.27 x 7.94/ 10.8 = x x = 1.67 atm 1.67 atm = 1270 torr P2 = 1270 torr
V1 = 12.6 L T1 = 383 K V2 = ? T2 = 45.3 K
12.6/383=x/45.3 1.49 = x V2 = 1.5 L
V1 = 2.83L T1 = 832K P1 = ? V2 = 3.85L T2 = 411K P2 = 0.396 atm
2.83*x / 832 = 3.85 * 0.396 / 411 x = 3.85 * 0.396 * 832 / 411 * 2.83 x = 1.09056 P1 = 1.09 atm
P = 1.6 atm N = 0.433 mol T = 346 K R = 0.082 atm V = ?
1.6*x = 0.433*346*0.082 x= 7.678 7.678 L = 7678 mL V = 7680 mL
P1 = 0.494 atm V1 = 2230 mL T1 = ? P2 = 0.150 atm V2 = 3590 mL T2 = 279 K
0.494*2230/ x = 0.15*3590 / 279 x=0.00175 T1=0.002 K

Unit 5 & 6 Test Review
Heat of Fusion
q = m ∙ΔH f
Specific Heat
c = ____q____
m ∙ ΔT ᵒC
Heat of Fusion of Water
334 J/g = 80 cal/g
Specific Heat of Water
4.18 J/g ∙ᵒC
Heat of fusion is the amount of heat energy required to change the state of a substance from solid to
liquid. This example problem demonstrates how to calculate the amount of energy required to melt a
sample of water ice.
With Graham's Law, you can find the effusion rates for two gases or the molecular mass of a gas.
This ratio of effusion rates follows the pattern that the gas with the lesser molecular mass has a
greater rate of effusion.
Ideal Gas Law
Combined Gas Law
p∙v=n∙R∙T V 1 ∙P 1 = V 2 ∙P 2
T 1 T 2
R = .082 atm or 8.31kpa
Boyle’s Law
P 1 ∙V 1 = P 2 ∙V 2
Charles’ Law
Gay – Lussac’s Law
V 1 = V 2 P 1 = P 2
T 1 T 2 T 1 T 2
Molarity
M = moles
L
Gas Solubility
S 1 = S 2
P 1 P 2

Unit 5
Chapter 13 States of Matter
Chapter 14 The Behavior of Gases
Chapter 15 Water and Aqueous Systems
Unit 6
Chapter 16 Solutions
Chapter 17 Thermochemistry
Chapter 18 Reaction Rates and Equilibrium
Chapter 19 Acid and Bases
Learning Goals
The students will learn what are the factors that determine and characteristics that distinguish gases
liquids and solids and how substances change from one state to another.
The students will learn how gases respond to changes in pressure, volume, and temperature and why
the ideal gas law is useful even though ideal gases do not exist.
The students will learn how the interactions between water molecules account for the unique
properties of water and how aqueous solutions form.
The student will learn how energy is converted in a chemical or physical process and how to
determine the amount of energy is absorbed or released in that process.

#1
What is the heat in Joules required to melt 25 grams of ice? What is the heat in calories?
#2
How much energy is released when 20 g of water is frozen at 0ᵒ C?

#3
A sample of water of unknown mass undergoes a temperature change from 58.85ᵒC to 16.06 ᵒC with
a concomitant change in heat content of -123200 Joules. What is the mass of the water?
#4
A sample of water with a mass of 597.9g and an unknown temperature loses 18890 Joules. If the
final temperature is found to be 35.76ᵒC .What was the initial temperature?

#5
A gas system has volume, moles and temperature of 9.76L, 0.264moles and 308.0C, respectively.
What is the pressure in atm?
#6
A gas system has pressure, volume and temperature of 1.43atm, 6.86L and 394K, respectively. How
many moles of gas are present?

#7
A gas system has pressure, volume and moles of 1.44atm, 8.16L and 0.427moles, respectively. What
is the temperature in K?
#8
What is the molarity of a solution of sodium bromide which contains 64.1 grams of solute in a total
volume of 0.436L?

#9
What volume in mL of a 0.942M solution of sodium bromide would be required if you wanted
0.692moles of solute?
#10
What is the percent by mass of water in CuSO 4 ∙ 5H 2 0?