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<strong>Chemistry</strong><br />
Josh Hoppe<br />
Block 3
Honors <strong>Chemistry</strong><br />
Class Policies and Grading<br />
The students will receive a Unit Outline at the beginning of each Unit. It will have information<br />
about the assignments that they will do, what it’s grade classification will be, what action<br />
they will need to do to complete the assignment and when it is due.<br />
The students will receive a Weekly Memo of the activities they will be responsible for that<br />
week. It will serve to inform the students of the learning goal for the week. It will also give<br />
the students any special information about that week.<br />
The students will also receive daily lectures and assignments that are designed to teach and<br />
re-enforce information related to the learning goal. This will be time in which new material<br />
will be taught and reviewed and will give the students the opportunity to ask questions<br />
regarding the concepts being taught.<br />
The students will work with a Lab partner and also be in a Lab group, but it will be up to the<br />
individual student to do his or her part of all assignments and the individual student will<br />
ultimately be responsible for all information presented in the class.<br />
The students will be required to follow all District and School Policies and to follow all Lab<br />
Safety Procedures, which they will be given and will sign, while performing labs. Students<br />
should come to class on time and with the supplies needed for that class.<br />
The following grading policy will be used.<br />
Units with Projects<br />
Units without Projects<br />
(2,3,4,5) (1,6)<br />
Work 12% Work 18%<br />
Quizzes 13% Quizzes 20%<br />
Labs 25% Labs 31%<br />
Test 25% Test 31%<br />
Projects 25%<br />
The students will be given a teacher generated Mid Term and a District Final.
Unit 1<br />
Measurement Lab<br />
Separation of Mixtures Lab with Lab Write Up<br />
Unit 2<br />
Flame Test Lab<br />
Nuclear Decay Lab<br />
Element Marketing Project<br />
Unit 3<br />
Golden Penny Lab with Lab Write Up<br />
Molecular Geometry<br />
Research Presentation on a Chemical<br />
Mid Term<br />
Unit 4<br />
Double Displacement Lab<br />
Stoichiometry Lab with Lab Write Up<br />
Mole Educational Demonstration Project<br />
Unit 5<br />
Gas Laws Lab with Lab Write Up<br />
States of Matter Lab<br />
Teach a Gas Law Project<br />
Unit 6<br />
Dilutions Lab<br />
Titration Lab<br />
District Final
Unit 1 (14 days)<br />
Chapter 1 Introduction to <strong>Chemistry</strong><br />
Honors <strong>Chemistry</strong><br />
2013/2014 Syllabus<br />
1.1 The Scope of <strong>Chemistry</strong> 1.3 Thinking Like a Scientist<br />
1.2 <strong>Chemistry</strong> and You 1.4 Problem Solving in <strong>Chemistry</strong><br />
Chapter 2 Matter and Change<br />
2.1 Properties of Matter 2.3 Elements and Compounds<br />
2.2 Mixtures 2.4 Chemical Reactions<br />
Chapter 3 Scientific Measurement<br />
3.1 Using and Expressing Measurements 3.3 Solving Conversion Problems<br />
3.2 Units of Measurement<br />
Unit 2 (15 days)<br />
Chapter 4 Atomic Structure<br />
4.1 Defining the Atom 4.3 Distinguishing Among Atoms<br />
4.2 Structure of the Nuclear Atom<br />
Chapter 5 Electrons in Atoms<br />
5.1 Revising the Atomic Model 5.2 Electron Arrangement in Atoms<br />
5.3 Atomic Emission Spectrum and the Quantum Mechanical Model<br />
Chapter 6 The Periodic Table<br />
6.1 Organizing the Elements 6.3 Periodic Trends<br />
6.2 Classifying Elements<br />
Chapter 25 Nuclear <strong>Chemistry</strong><br />
25.1 Nuclear Radiation 25.3 Fission and Fusion<br />
25.2 Nuclear Transformations 25.4 Radiation in Your Life<br />
Unit 3 (11 days)<br />
Chapter 7 Ionic and Metallic Bonding<br />
7.1 Ions 7.3 Bonding in Metals<br />
7.2 Ionic Bonds and Ionic Compounds<br />
Chapter 8 Covalent Bonding<br />
8.1 Molecular Compounds 8.3 Bonding Theories<br />
8.2 The Nature of Covalent Bonding 8.4 Polar Bonds and Molecules<br />
Chapter 9 Chemical Names and Formulas<br />
9.1 Naming Ions 9.3 Naming and Writing Formulas for Molecular Compounds<br />
9.2 Naming and Writing Formulas for Ionic Compounds 9.4 Names for Acids and Bases<br />
3 days<br />
5 days<br />
6 days<br />
3 days<br />
5 days<br />
3 days<br />
3 days<br />
4 days<br />
4 days<br />
3 days
Unit 4 (15 days)<br />
Chapter 10 Chemical Quantities<br />
10.1 The Mole: A Measurement of Matter 10.3 % Composition & Chem. Formulas<br />
10.2 Mole-Mass and Mole-Volume Relationships<br />
Chapter 11 Chemical Reactions<br />
11.1 Describing Chemical Reactions 11.3 Reactions in Aqueous Solutions<br />
11.2 Types of Chemical Reactions<br />
Chapter 12 Stoichiometry<br />
12.1 The Arithmetic of Equations 12.3 Limiting Reagent and % Yield<br />
12.2 Chemical Calculations<br />
Unit 5 (15 days)<br />
Chapter 13 States of Matter<br />
13.1 The Nature of Gases 13.3 The Nature of Solids<br />
13.2 The Nature of Liquids 13.4 Changes in State<br />
Chapter 14 The Behavior of Gases<br />
14.1 Properties of Gases 14.3 Ideal Gases<br />
14.2 The Gas Laws 14.4 Gases: Mixtures and Movement<br />
Chapter 15 Water and Aqueous Systems<br />
15.1 Water and its Properties 15.3 Heterogeneous Aqueous Systems<br />
15.2 Homogeneous Aqueous Systems<br />
Unit 6 (10 days)<br />
Chapter 16 Solutions<br />
16.1 Properties of Solutions 16.3 Colligative Properties of Solutions<br />
16.2 Concentrations of Solutions 16.4 Calc. Involving Colligative Property<br />
Chapter 17 Thermochemistry<br />
17.1 The Flow of Energy 17.3 Heat in Changes of State<br />
17.2 Measuring and Expressing Enthalpy Change 17.4 Calculating Heats in Reactions<br />
Chapter 18 Reaction Rates and Equilibrium<br />
18.1 Rates of Reactions 18.3 Reversible Reaction & Equilibrium<br />
18.2 The Progress of Chemical Reactions 18.5 Free Energy and Entropy<br />
Chapter 19 Acid and Bases<br />
19.1 Acid-Base Theories 19.4 Neutralization Reactions<br />
19.2 Hydrogen Ions and Acidity 19.5 Salts in Solutions<br />
19.3 Strengths of Acids and Bases<br />
5 days<br />
5 days<br />
5 days<br />
4 days<br />
5 days<br />
3 days<br />
4 days<br />
2 days<br />
2 days<br />
2 days
Lorenzo Walker Technical High School<br />
MUSTANG LABORATORIES<br />
<strong>Chemistry</strong> Safety<br />
Safety in the MUSTANG LABORATORIES - <strong>Chemistry</strong> Laboratory<br />
Working in the chemistry laboratory is an interesting and rewarding experience. During your labs, you will be actively<br />
involved from beginning to end—from setting some change in motion to drawing some conclusion. In the laboratory, you will<br />
be working with equipment and materials that can cause injury if they are not handled properly.<br />
However, the laboratory is a safe place to work if you are careful. Accidents do not just happen; they are caused—by<br />
carelessness, haste, and disregard of safety rules and practices. Safety rules to be followed in the laboratory are listed<br />
below. Before beginning any lab work, read these rules, learn them, and follow them carefully.<br />
General<br />
1. Be prepared to work when you arrive at the lab. Familiarize yourself with the lab procedures before beginning the lab.<br />
2. Perform only those lab activities assigned by your teacher. Never do anything in the laboratory that is not called for in<br />
the laboratory procedure or by your teacher. Never work alone in the lab. Do not engage in any horseplay.<br />
3. Work areas should be kept clean and tidy at all times. Only lab manuals and notebooks should be brought to the work<br />
area. Other books, purses, brief cases, etc. should be left at your desk or placed in a designated storage area.<br />
4. Clothing should be appropriate for working in the lab. Jackets, ties, and other loose garments should be removed. Open<br />
shoes should not be worn.<br />
5. Long hair should be tied back or covered, especially in the vicinity of open flame.<br />
6. Jewelry that might present a safety hazard, such as dangling necklaces, chains, medallions, or bracelets should not be<br />
worn in the lab.<br />
7. Follow all instructions, both written and oral, carefully.<br />
8. Safety goggles and lab aprons should be worn at all times.<br />
9. Set up apparatus as described in the lab manual or by your teacher. Never use makeshift arrangements.<br />
10. Always use the prescribed instrument (tongs, test tube holder, forceps, etc.) for handling apparatus or equipment.<br />
11. Keep all combustible materials away from open flames.<br />
12. Never touch any substance in the lab unless specifically instructed to do so by your teacher.<br />
13. Never put your face near the mouth of a container that is holding chemicals.<br />
14. Never smell any chemicals unless instructed to do so by your teacher. When testing for odors, use a wafting motion to<br />
direct the odors to your nose.<br />
15. Any activity involving poisonous vapors should be conducted in the fume hood.<br />
16. Dispose of waste materials as instructed by your teacher.<br />
17. Clean up all spills immediately.<br />
18. Clean and wipe dry all work surfaces at the end of class. Wash your hands thoroughly.<br />
19. Know the location of emergency equipment (first aid kit, fire extinguisher, fire shower, fire blanket, etc.) and how to use them.<br />
20. Report all accidents to the teacher immediately.<br />
Handling Chemicals<br />
21. Read and double check labels on reagent bottles before removing any reagent. Take only as much reagent as you need.<br />
22. Do not return unused reagent to stock bottles.<br />
23. When transferring chemical reagents from one container to another, hold the containers out away from your body.<br />
24. When mixing an acid and water, always add the acid to the water.<br />
25. Avoid touching chemicals with your hands. If chemicals do come in contact with your hands, wash them immediately.<br />
26. Notify your teacher if you have any medical problems that might relate to lab work, such as allergies or asthma.<br />
27. If you will be working with chemicals in the lab, avoid wearing contact lenses. Change to glasses, if possible, or notify<br />
the teacher.<br />
Handling Glassware<br />
28. Glass tubing, especially long pieces, should be carried in a vertical position to minimize the likelihood of breakage and<br />
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<br />
glass as directed by your teacher.<br />
30. Always lubricate glassware (tubing, thistle tubes, thermometers, etc.) with water or glycerin before attempting to insert it into a<br />
rubber stopper.<br />
31. Never apply force when inserting or removing glassware from a stopper. Use a twisting motion. If a piece of glassware<br />
becomes "frozen" in a stopper, take it to your teacher.<br />
32. Do not place hot glassware directly on the lab table. Always use an insulating pad of some sort.<br />
33. Allow plenty of time for hot glass to cool before touching it. Hot glass can cause painful burns. (Hot glass looks cool.)<br />
Heating Substances<br />
34. Exercise extreme caution when using a gas burner. Keep your head and clothing away from the flame.<br />
35. Always turn the burner off when it is not in use.<br />
36. Do not bring any substance into contact with a flame unless instructed to do so.<br />
37. Never heat anything without being instructed to do so.<br />
38. Never look into a container that is being heated.<br />
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.<br />
40. Never leave unattended anything that is being heated or is visibly reacting.<br />
First Aid in the MUSTANG LABORATORIES - <strong>Chemistry</strong> Laboratory<br />
Accidents do not often happen in well-equipped chemistry laboratories if students understand safe laboratory procedures<br />
and are careful in following them. When an occasional accident does occur, it is likely to be a minor one.<br />
The instructor will assist in treating injuries such as minor cuts and burns. However, for some types of injuries, you must take<br />
action immediately. The following information will be helpful to you if an accident occurs.<br />
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<br />
of shock. A person in shock is usually pale and faint. The person may be sweating, with cold, moist skin and a weak, rapid<br />
pulse. Shock is a serious medical condition. Do not allow a person in shock to walk anywhere—even to the campus security<br />
office. While emergency help is being summoned, place the victim face up in a horizontal position, with the feet raised about<br />
30 centimeters. Loosen any tightly fitting clothing and keep him or her warm.<br />
2. Chemicals in the Eyes. Getting any kind of a chemical into the eyes is undesirable, but certain chemicals are especially<br />
harmful. They can destroy eyesight in a matter of seconds. Because you will be wearing safety goggles at all times in the lab,<br />
the likelihood of this kind of accident is remote. However, if it does happen, flush your eyes with water immediately. Do NOT<br />
attempt to go to the campus office before flushing your eyes. It is important that flushing with water be continued for a<br />
prolonged time—about 15 minutes.<br />
3. Clothing or Hair on Fire. A person whose clothing or hair catches on fire will often run around hysterically in an<br />
unsuccessful effort to get away from the fire. This only provides the fire with more oxygen and makes it burn faster. For<br />
clothing fires, throw yourself to the ground and roll around to extinguish the flames. For hair fires, use a fire blanket to<br />
smother the flames. Notify campus security immediately.<br />
4. Bleeding from a Cut. Most cuts that occur in the chemistry laboratory are minor. For minor cuts, apply pressure to the<br />
wound with a sterile gauze. Notify campus security of all injuries in the lab. If the victim is bleeding badly, raise the bleeding<br />
part, if possible, and apply pressure to the wound with a piece of sterile gauze. While first aid is being given, someone else<br />
should notify the campus security officer.<br />
5. Chemicals in the Mouth. Many chemicals are poisonous to varying degrees. Any chemical taken into the mouth should be<br />
spat out and the mouth rinsed thoroughly with water. Note the name of the chemical and notify the campus office<br />
immediately. If the victim swallows a chemical, note the name of the chemical and notify campus security immediately.<br />
If necessary, the campus security officer or administrator will contact the Poison Control Center, a hospital emergency room,<br />
or a physician for instructions.<br />
6. Acid or Base Spilled on the Skin.<br />
Flush the skin with water for about 15 minutes. Take the victim to the campus office to report the injury.<br />
7. Breathing Smoke or Chemical Fumes.<br />
All experiments that give off smoke or noxious gases should be conducted in a well-ventilated fume hood. This will make an<br />
accident of this kind unlikely. If smoke or chemical fumes are present in the laboratory, all persons—even those who do not<br />
feel ill—should leave the laboratory immediately. Make certain that all doors to the laboratory are closed after the last<br />
person has left. Since smoke rises, stay low while evacuating a smoke-filled room. Notify campus security immediately.
MUSTANG LABORATORIES<br />
COMMITMENT TO SAFETY IN THE LABORATORY<br />
As a student enrolled in <strong>Chemistry</strong> at Lorenzo Walker Technical High School, I agree to use<br />
good laboratory safety practices at all times. I also agree that I will:<br />
1. Conduct myself in a professional manner, respecting both my personal safety and the safety of others in the laboratory.<br />
2. Wear proper and approved safety glasses or goggles in the laboratory at all times.<br />
3. Wear sensible clothing and tie back long hair in the laboratory. Understand that open-toed shoes pose a hazard during<br />
laboratory classes and that contact lenses are an added safety risk.<br />
4. Keep my lab area free of clutter during an experiment.<br />
5. Never bring food or drink into the laboratory, nor apply makeup within the laboratory.<br />
6. Be aware of the location of safety equipment such as the fire extinguisher, eye wash station, fire blanket, first aid kit.<br />
Know the location of the nearest telephone and exits.<br />
7. Read the assigned lab prior to coming to the laboratory.<br />
8. Carefully read all labels on all chemical containers before using their contents, remove a small amount of reagent<br />
properly if needed, do not pour back the unused chemicals into the original container.<br />
9. Dispose of chemicals as directed by the instructor only. At no time will I pour anything down the sink without prior<br />
instruction.<br />
10. Never inhale fumes emitted during an experiment. Use the fume hood when instructed to do so.<br />
11. Report any accident immediately to the instructor, including chemical spills.<br />
12. Dispose of broken glass and sharps only in the designated containers.<br />
13. Clean my work area and all glassware before leaving the laboratory.<br />
14. Wash my hands before leaving the laboratory.<br />
NAME ____________________________<br />
Josh Hoppe<br />
PARENT NAME _______________________________________<br />
Linda Hoppe<br />
BLOCK ____________________________<br />
3 PARENT NUMBER _____________________________________<br />
(239) 537-2550<br />
SIGNATURE _________________________________________________<br />
8/20/13<br />
DATE ____________________________________
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C<br />
THE UNIVERSITY OF THE STATE OF NEW YORK• THE STATE EDUCATION DEPARTMENT• ALBANY, NY 12234<br />
Reference Tables for Physical Setting/CHEMISTRY<br />
Table A<br />
Standard Temperature and Pressure<br />
2011 Edition<br />
Table D<br />
Selected Units<br />
Name Value Unit<br />
Standard Pressure 101.3 kPa kilopascal<br />
1 atm atmosphere<br />
Standard Temperature 273 K kelvin<br />
0°C degree Celsius<br />
Table B<br />
Physical Constants for Water<br />
Heat of Fusion<br />
Heat of Vaporization<br />
Specific Heat Capacity of H 2<br />
O()<br />
Table C<br />
Selected Prefixes<br />
Factor Prefix Symbol<br />
10 3 kilo- k<br />
10 –1 deci- d<br />
10 –2 centi- c<br />
10 –3 milli- m<br />
10 –6 micro- μ<br />
10 –9 nano- n<br />
10 –12 pico- p<br />
334 J/g<br />
2260 J/g<br />
4.18 J/g•K<br />
Symbol Name Quantity<br />
m meter length<br />
g gram mass<br />
Pa pascal pressure<br />
K kelvin temperature<br />
mol<br />
J<br />
mole<br />
joule<br />
s second time<br />
min minute time<br />
h hour time<br />
d day time<br />
y year time<br />
amount of<br />
substance<br />
L liter volume<br />
energy, work,<br />
quantity of heat<br />
ppm parts per million concentration<br />
M<br />
molarity<br />
solution<br />
concentration<br />
u atomic mass unit atomic mass<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 1
Table E<br />
Selected Polyatomic Ions<br />
Formula Name Formula Name<br />
H 3<br />
O +<br />
hydronium<br />
CrO 4<br />
2–<br />
chromate<br />
Hg 2<br />
2+<br />
mercury(I)<br />
Cr 2<br />
O 7<br />
2–<br />
dichromate<br />
NH 4<br />
+<br />
C 2<br />
H 3<br />
O<br />
–<br />
2 –}<br />
CH 3<br />
COO<br />
CN –<br />
CO 3<br />
2–<br />
HCO<br />
–<br />
3<br />
C 2<br />
O<br />
2–<br />
4<br />
ClO –<br />
ammonium<br />
acetate<br />
cyanide<br />
carbonate<br />
hydrogen<br />
carbonate<br />
oxalate<br />
hypochlorite<br />
MnO 4<br />
–<br />
NO<br />
–<br />
2<br />
NO<br />
–<br />
3<br />
O<br />
2–<br />
2<br />
OH –<br />
PO 4<br />
3–<br />
SCN –<br />
SO 3<br />
2–<br />
permanganate<br />
nitrite<br />
nitrate<br />
peroxide<br />
hydroxide<br />
phosphate<br />
thiocyanate<br />
sulfite<br />
ClO 2<br />
–<br />
chlorite<br />
SO 4<br />
2–<br />
sulfate<br />
ClO 3<br />
–<br />
chlorate<br />
HSO 4<br />
–<br />
hydrogen sulfate<br />
ClO 4<br />
–<br />
perchlorate<br />
S 2<br />
O 3<br />
2–<br />
thiosulfate<br />
Table F<br />
Solubility Guidelines for Aqueous Solutions<br />
Ions That Form<br />
Soluble Compounds<br />
Group 1 ions<br />
(Li + , Na + , etc.)<br />
ammonium (NH + 4<br />
)<br />
nitrate (NO – 3<br />
)<br />
acetate (C 2<br />
H 3<br />
O – 2<br />
or<br />
CH 3<br />
COO – )<br />
hydrogen carbonate<br />
(HCO – 3<br />
)<br />
chlorate (ClO – 3<br />
)<br />
halides (Cl – , Br – , I – )<br />
Exceptions<br />
when combined with<br />
Ag + , Pb 2+ , or Hg 2<br />
2+<br />
sulfates (SO 4 2– ) when combined with Ag + ,<br />
Ca 2+ , Sr 2+ , Ba 2+ , or Pb 2+<br />
Ions That Form<br />
Insoluble Compounds* Exceptions<br />
carbonate (CO 2– 3<br />
) when combined with Group 1<br />
ions or ammonium (NH + 4<br />
)<br />
chromate (CrO 2– 4<br />
) when combined with Group 1<br />
ions, Ca 2+ , Mg 2+ , or<br />
ammonium (NH + 4<br />
)<br />
phosphate (PO 3– 4<br />
) when combined with Group 1<br />
ions or ammonium (NH + 4<br />
)<br />
sulfide (S 2– ) when combined with Group 1<br />
ions or ammonium (NH + 4<br />
)<br />
hydroxide (OH – ) when combined with Group 1<br />
ions, Ca 2+ , Ba 2+ , Sr 2+ , or<br />
ammonium (NH + 4<br />
)<br />
*compounds having very low solubility in H 2 O<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 2
150.<br />
140.<br />
Table G<br />
Solubility Curves at Standard Pressure<br />
KI<br />
NaNO 3<br />
130.<br />
120.<br />
KNO 3<br />
110.<br />
100.<br />
Solubility (g solute/100. g H 2<br />
O)<br />
90.<br />
80.<br />
70.<br />
60.<br />
HCl<br />
NH 4<br />
Cl<br />
KCl<br />
50.<br />
40.<br />
30.<br />
NaCl<br />
KClO 3<br />
NH 3<br />
20.<br />
10.<br />
SO 2<br />
0<br />
0 10. 20. 30. 40. 50. 60. 70. 80. 90. 100.<br />
Temperature (°C)<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 3
Table H<br />
Vapor Pressure of Four Liquids<br />
200.<br />
propanone<br />
ethanol<br />
150.<br />
water<br />
Vapor Pressure (kPa)<br />
100.<br />
101.3 kPa<br />
ethanoic<br />
acid<br />
50.<br />
0<br />
0 25 50. 75 100. 125<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 4
Table I<br />
Heats of Reaction at 101.3 kPa and 298 K<br />
Reaction<br />
ΔH (kJ)*<br />
CH 4<br />
(g) + 2O 2<br />
(g) CO 2<br />
(g) + 2H 2<br />
O() –890.4<br />
C 3<br />
H 8<br />
(g) + 5O 2<br />
(g) 3CO 2<br />
(g) + 4H 2<br />
O() –2219.2<br />
2C 8<br />
H 18<br />
() + 25O 2<br />
(g) 16CO 2<br />
(g) + 18H 2<br />
O() –10943<br />
2CH 3<br />
OH() + 3O 2<br />
(g) 2CO 2<br />
(g) + 4H 2<br />
O() –1452<br />
C 2<br />
H 5<br />
OH() + 3O 2<br />
(g) 2CO 2<br />
(g) + 3H 2<br />
O() –1367<br />
C 6<br />
H 12<br />
O 6<br />
(s) + 6O 2<br />
(g) 6CO 2<br />
(g) + 6H 2<br />
O() –2804<br />
2CO(g) + O 2<br />
(g) 2CO 2<br />
(g) –566.0<br />
C(s) + O 2<br />
(g) CO 2<br />
(g) –393.5<br />
4Al(s) + 3O 2<br />
(g) 2Al 2<br />
O 3<br />
(s) –3351<br />
N 2<br />
(g) + O 2<br />
(g) 2NO(g) +182.6<br />
N 2<br />
(g) + 2O 2<br />
(g) 2NO 2<br />
(g) +66.4<br />
2H 2<br />
(g) + O 2<br />
(g) 2H 2<br />
O(g) –483.6<br />
2H 2<br />
(g) + O 2<br />
(g) 2H 2<br />
O() –571.6<br />
N 2<br />
(g) + 3H 2<br />
(g) 2NH 3<br />
(g) –91.8<br />
2C(s) + 3H 2<br />
(g) C 2<br />
H 6<br />
(g) –84.0<br />
2C(s) + 2H 2<br />
(g) C 2<br />
H 4<br />
(g) +52.4<br />
2C(s) + H 2<br />
(g) C 2<br />
H 2<br />
(g) +227.4<br />
H 2<br />
(g) + I 2<br />
(g) 2HI(g) +53.0<br />
KNO 3<br />
(s) H 2 O K + (aq) + NO 3 – (aq) +34.89<br />
NaOH(s) H 2 O Na + (aq) + OH – (aq) –44.51<br />
NH 4<br />
Cl(s) H 2 O NH 4 + (aq) + Cl – (aq) +14.78<br />
NH 4<br />
NO 3<br />
(s) H 2 O NH 4 + (aq) + NO 3 – (aq) +25.69<br />
NaCl(s) H 2 O Na + (aq) + Cl – (aq) +3.88<br />
LiBr(s) H 2 O Li + (aq) + Br – (aq) –48.83<br />
H + (aq) + OH – (aq) H 2<br />
O() –55.8<br />
*The ΔH values are based on molar quantities represented in the equations.<br />
A minus sign indicates an exothermic reaction.<br />
Most<br />
Active<br />
Least<br />
Active<br />
Table J<br />
Activity Series**<br />
Metals Nonmetals Most<br />
Active<br />
Li<br />
F 2<br />
Rb Cl 2<br />
K Br 2<br />
Cs<br />
I 2<br />
Ba<br />
Sr<br />
Ca<br />
Na<br />
Mg<br />
Al<br />
Ti<br />
Mn<br />
Zn<br />
Cr<br />
Fe<br />
Co<br />
Ni<br />
Sn<br />
Pb<br />
H 2<br />
Cu<br />
Ag<br />
Au<br />
**Activity Series is based on the hydrogen<br />
standard. H 2 is not a metal.<br />
Least<br />
Active<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 5
Table K<br />
Common Acids<br />
Table N<br />
Selected Radioisotopes<br />
HCl(aq)<br />
Formula<br />
HNO 2<br />
(aq)<br />
HNO 3<br />
(aq)<br />
H 2<br />
SO 3<br />
(aq)<br />
H 2<br />
SO 4<br />
(aq)<br />
H 3<br />
PO 4<br />
(aq)<br />
H 2<br />
CO 3<br />
(aq)<br />
or<br />
CO 2<br />
(aq)<br />
CH 3<br />
COOH(aq)<br />
or<br />
HC 2<br />
H 3<br />
O 2<br />
(aq)<br />
Name<br />
hydrochloric acid<br />
nitrous acid<br />
nitric acid<br />
sulfurous acid<br />
sulfuric acid<br />
phosphoric acid<br />
carbonic acid<br />
ethanoic acid<br />
(acetic acid)<br />
Nuclide Half-Life Decay<br />
Mode<br />
Nuclide<br />
Name<br />
198 Au 2.695 d β – gold-198<br />
14 C 5715 y β – carbon-14<br />
37 Ca 182 ms β + calcium-37<br />
60 Co 5.271 y β – cobalt-60<br />
137 Cs 30.2 y β – cesium-137<br />
53 Fe 8.51 min β + iron-53<br />
220 Fr 27.4 s α francium-220<br />
3 H 12.31 y β – hydrogen-3<br />
131 I 8.021 d β – iodine-131<br />
37 K 1.23 s β + potassium-37<br />
42 K 12.36 h β – potassium-42<br />
Table L<br />
Common Bases<br />
85 Kr 10.73 y β – krypton-85<br />
16 N 7.13 s β – nitrogen-16<br />
Formula<br />
NaOH(aq)<br />
KOH(aq)<br />
Ca(OH) 2<br />
(aq)<br />
NH 3<br />
(aq)<br />
Name<br />
sodium hydroxide<br />
potassium hydroxide<br />
calcium hydroxide<br />
aqueous ammonia<br />
19 Ne 17.22 s β + neon-19<br />
32 P 14.28 d β – phosphorus-32<br />
239 Pu 2.410 × 10 4 y α plutonium-239<br />
226 Ra 1599 y α radium-226<br />
222 Rn 3.823 d α radon-222<br />
90 Sr 29.1 y β – strontium-90<br />
Table M<br />
Common Acid–Base Indicators<br />
Approximate<br />
Indicator pH Range Color<br />
for Color Change<br />
Change<br />
methyl orange 3.1–4.4 red to yellow<br />
bromthymol blue 6.0–7.6 yellow to blue<br />
phenolphthalein 8–9 colorless to pink<br />
litmus 4.5–8.3 red to blue<br />
bromcresol green 3.8–5.4 yellow to blue<br />
thymol blue 8.0–9.6 yellow to blue<br />
99 Tc 2.13 × 10 5 y β – technetium-99<br />
232 Th 1.40 × 10 10 y α thorium-232<br />
233 U 1.592 × 10 5 y α uranium-233<br />
235 U 7.04 × 10 8 y α uranium-235<br />
238 U 4.47 × 10 9 y α uranium-238<br />
Source: CRC Handbook of <strong>Chemistry</strong> and Physics, 91 st ed., 2010–2011,<br />
CRC Press<br />
Source: The Merck Index, 14 th ed., 2006, Merck Publishing Group<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 6
Table O<br />
Symbols Used in Nuclear <strong>Chemistry</strong><br />
Name Notation Symbol<br />
alpha particle<br />
4<br />
2<br />
He or 4 2 α α<br />
beta particle<br />
0<br />
–1<br />
e or 0<br />
–1 β β–<br />
gamma radiation<br />
0<br />
0<br />
γ γ<br />
neutron<br />
1<br />
0<br />
n n<br />
proton<br />
1<br />
1<br />
H or 1 1 p p<br />
positron<br />
0<br />
+1<br />
e or 0<br />
+1 β β+<br />
Name General Examples<br />
Formula Name Structural Formula<br />
alkanes C n<br />
H 2n+2<br />
ethane<br />
alkenes C n<br />
H 2n<br />
ethene<br />
alkynes C n<br />
H 2n–2<br />
ethyne<br />
Table P<br />
Organic Prefixes<br />
Prefix<br />
meth- 1<br />
eth- 2<br />
prop- 3<br />
but- 4<br />
pent- 5<br />
hex- 6<br />
hept- 7<br />
oct- 8<br />
Number of<br />
Carbon Atoms<br />
non- 9<br />
dec- 10<br />
Table Q<br />
Homologous Series of Hydrocarbons<br />
Note: n = number of carbon atoms<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 7<br />
H<br />
H<br />
H<br />
H<br />
H<br />
C<br />
H<br />
C<br />
H<br />
C<br />
H<br />
C<br />
H<br />
H<br />
H<br />
C C H
Table R<br />
Organic Functional Groups<br />
Class of<br />
Compound<br />
Functional<br />
Group<br />
General<br />
Formula<br />
Example<br />
halide<br />
(halocarbon)<br />
F (fluoro-)<br />
Cl (chloro-)<br />
Br (bromo-)<br />
I (iodo-)<br />
R X<br />
(X represents<br />
any halogen)<br />
CH 3<br />
CHClCH 3<br />
2-chloropropane<br />
alcohol<br />
OH<br />
R<br />
OH<br />
CH 3<br />
CH 2<br />
CH 2<br />
OH<br />
1-propanol<br />
ether<br />
O<br />
R O R′<br />
CH 3<br />
OCH 2<br />
CH 3<br />
methyl ethyl ether<br />
aldehyde<br />
O<br />
C H<br />
R<br />
O<br />
C H<br />
O<br />
CH 3<br />
CH 2<br />
C H<br />
propanal<br />
ketone<br />
O<br />
C<br />
O<br />
R C R′<br />
O<br />
CH 3<br />
CCH 2<br />
CH 2<br />
CH 3<br />
2-pentanone<br />
organic acid<br />
O<br />
C OH<br />
R<br />
O<br />
C OH<br />
O<br />
CH 3<br />
CH 2<br />
C OH<br />
propanoic acid<br />
ester<br />
O<br />
C O<br />
O<br />
R C O R′<br />
O<br />
CH 3<br />
CH 2<br />
COCH 3<br />
methyl propanoate<br />
amine<br />
N<br />
R<br />
R′<br />
N R′′<br />
CH 3<br />
CH 2<br />
CH 2<br />
NH 2<br />
1-propanamine<br />
amide<br />
O<br />
C NH<br />
R<br />
O R′<br />
C NH<br />
O<br />
CH 3<br />
CH 2<br />
C NH 2<br />
propanamide<br />
Note: R represents a bonded atom or group of atoms.<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 8
0<br />
6.941<br />
+1<br />
Li<br />
3<br />
2-1<br />
Na<br />
39.0983<br />
K +1<br />
19<br />
2-8-8-1<br />
Rb<br />
Cs<br />
(223)<br />
Fr<br />
87<br />
-18-32-18-8-1<br />
+1<br />
Ra<br />
88<br />
-18-32-18-8-2<br />
39<br />
138.9055<br />
La<br />
57<br />
2-8-18-18-9-2<br />
+2 (227)<br />
Ac<br />
89<br />
-18-32-18-9-2<br />
47.867<br />
Ti<br />
22<br />
2-8-10-2<br />
91.224<br />
Zr<br />
40<br />
2-8-18-10-2<br />
+3 178.49<br />
Hf<br />
72<br />
*18-32-10-2<br />
+3 (261)<br />
Rf<br />
104<br />
+2<br />
+3<br />
+4<br />
+4<br />
+4<br />
50.9415<br />
V<br />
23<br />
2-8-11-2<br />
+2<br />
+3<br />
+4<br />
+5<br />
51.996<br />
Cr<br />
24<br />
2-8-13-1<br />
95.94<br />
Mo<br />
42<br />
2-8-18-13-1<br />
183.84<br />
W<br />
74<br />
-18-32-12-2<br />
+2<br />
+3<br />
+6<br />
+6<br />
+6<br />
54.9380<br />
Mn<br />
25<br />
2-8-13-2<br />
+2<br />
+3<br />
+4<br />
+7<br />
55.845<br />
Fe<br />
26<br />
2-8-14-2<br />
+2<br />
+3 58.9332<br />
Co<br />
27<br />
2-8-15-2<br />
+2<br />
+3<br />
58.693<br />
Ni<br />
28<br />
2-8-16-2<br />
+2<br />
+3 63.546 Cu<br />
2-8-18-1<br />
107.868<br />
Ag<br />
47<br />
2-8-18-18-1<br />
79<br />
+1<br />
+2<br />
+1<br />
65.409<br />
Zn<br />
30<br />
2-8-18-2<br />
+3 12.011<br />
B<br />
5<br />
2-3<br />
26.98154<br />
Al<br />
13<br />
2-8-3<br />
+2 69.723<br />
Ga<br />
31<br />
2-8-18-3<br />
+3<br />
+3<br />
–4<br />
+2<br />
+4<br />
C<br />
6<br />
2-4<br />
28.0855<br />
Si<br />
14<br />
2-8-4<br />
72.64<br />
Ge<br />
32<br />
2-8-18-4<br />
Pb<br />
–4<br />
+2<br />
+4<br />
+2<br />
+4<br />
74.9216<br />
As<br />
33<br />
2-8-18-5<br />
Sb<br />
–3<br />
+3<br />
15.9994<br />
O<br />
F<br />
8<br />
2-6<br />
–2 18.9984<br />
78.96<br />
Se<br />
34<br />
2-8-18-6<br />
127.60<br />
Te<br />
52<br />
2-8-18-18-6<br />
(209)<br />
Po<br />
84<br />
-18-32-18-6<br />
–2<br />
+4<br />
+6<br />
–2<br />
+4<br />
+6<br />
+2<br />
+4<br />
2-7<br />
79.904<br />
Br<br />
35<br />
2-8-18-7<br />
126.904<br />
l<br />
53<br />
2-8-18-18-7<br />
(210)<br />
At<br />
85<br />
-18-32-18-7<br />
4.00260 0<br />
He<br />
2<br />
2<br />
–1 20.180<br />
Ne<br />
10<br />
2-8<br />
0<br />
22.98977<br />
11<br />
2-8-1<br />
1<br />
1.00794 +1<br />
–1<br />
H<br />
1<br />
1<br />
1<br />
85.4678<br />
37<br />
2-8-18-8-1<br />
30.97376<br />
P<br />
15<br />
2-8-5<br />
–3<br />
+3<br />
+5<br />
32.065<br />
S<br />
16<br />
2-8-6<br />
–2<br />
+4<br />
+6<br />
–1<br />
+1<br />
+5<br />
–1<br />
+1<br />
+5<br />
+7<br />
39.948<br />
Ar<br />
18<br />
2-8-8<br />
83.798<br />
Kr<br />
36<br />
2-8-18-8<br />
131.29<br />
Xe<br />
54<br />
2-8-18-18-8<br />
(222)<br />
Rn<br />
86<br />
-18-32-18-8<br />
0<br />
+2<br />
0<br />
+2<br />
+4<br />
+6<br />
0<br />
132.905<br />
55<br />
2-8-18-18-8-1<br />
Symbol<br />
Relative atomic masses are based<br />
Group on 12 C = 12 (exact)<br />
Group<br />
2<br />
13 14 15 16 17 18<br />
Atomic Number<br />
Note: Numbers in parentheses<br />
10.81<br />
are mass numbers of the most<br />
Electron Configuration<br />
stable or common isotope.<br />
+1<br />
+1<br />
+1<br />
9.01218 +2<br />
Be<br />
4<br />
2-2<br />
24.305<br />
Mg<br />
12<br />
2-8-2<br />
40.08<br />
Ca<br />
20<br />
2-8-8-2<br />
87.62<br />
Sr<br />
38<br />
2-8-18-8-2<br />
137.33<br />
Ba<br />
56<br />
2-8-18-18-8-2<br />
(226)<br />
+2<br />
+2<br />
+2<br />
+2<br />
3<br />
44.9559<br />
Sc<br />
21<br />
2-8-9-2<br />
88.9059<br />
Y<br />
2-8-18-9-2<br />
+3<br />
+3<br />
4<br />
KEY<br />
+4<br />
92.9064<br />
Nb +3<br />
+5<br />
41<br />
2-8-18-12-1<br />
180.948<br />
Ta<br />
73<br />
-18-32-11-2<br />
(262)<br />
105<br />
5<br />
Periodic Table of the Elements<br />
Atomic Mass<br />
Db<br />
+5<br />
6<br />
(266)<br />
Sg<br />
106<br />
12.011 2-4<br />
–4<br />
6<br />
C<br />
+2<br />
+4<br />
(98)<br />
Tc<br />
43<br />
2-8-18-13-2<br />
186.207<br />
Re<br />
75<br />
-18-32-13-2<br />
(272)<br />
Bh<br />
107<br />
7<br />
Group<br />
+4<br />
+6<br />
+7<br />
+4<br />
+6<br />
+7<br />
8<br />
101.07<br />
Ru<br />
44<br />
2-8-18-15-1<br />
190.23<br />
Os<br />
76<br />
-18-32-14-2<br />
(277)<br />
Hs<br />
108<br />
+3<br />
+3<br />
+4<br />
Selected Oxidation States<br />
9<br />
102.906<br />
Rh<br />
45<br />
2-8-18-16-1<br />
192.217<br />
Ir<br />
77<br />
-18-32-15-2<br />
(276)<br />
Mt<br />
109<br />
+3<br />
+3<br />
+4<br />
106.42<br />
Pd<br />
46<br />
2-8-18-18<br />
195.08<br />
Pt<br />
78<br />
-18-32-17-1<br />
+2<br />
+4<br />
+2<br />
+4<br />
196.967<br />
Au<br />
-18-32-18-1<br />
(281)<br />
Ds (280) Rg<br />
110<br />
10<br />
29<br />
111<br />
11 12<br />
+1<br />
+3<br />
112.41<br />
Cd<br />
48<br />
2-8-18-18-2<br />
200.59<br />
Hg<br />
80<br />
-18-32-18-2<br />
(285)<br />
Cn<br />
112<br />
+2 114.818<br />
In<br />
+1<br />
+2<br />
49<br />
2-8-18-18-3<br />
204.383<br />
Tl<br />
81<br />
-18-32-18-3<br />
(284)<br />
Uut<br />
113**<br />
+3<br />
+1<br />
+3<br />
118.71<br />
Sn<br />
50<br />
2-8-18-18-4<br />
207.2<br />
82<br />
-18-32-18-4<br />
(289)<br />
Uuq<br />
114<br />
+2<br />
+4<br />
+2<br />
+4<br />
14.0067 –3<br />
–2<br />
N<br />
–1<br />
7<br />
2-5<br />
121.760<br />
51<br />
2-8-18-18-5<br />
208.980<br />
Bi<br />
83<br />
-18-32-18-5<br />
(288)<br />
Uup<br />
115<br />
+1<br />
+2<br />
+3<br />
+4<br />
+5<br />
+5<br />
–3<br />
+3<br />
+5<br />
+3<br />
+5<br />
(292)<br />
Uuh<br />
116<br />
35.453<br />
Cl<br />
17<br />
2-8-7<br />
( ? )<br />
Uus<br />
117<br />
–1<br />
+1<br />
+5<br />
+7<br />
18<br />
(294)<br />
Uuo<br />
118<br />
140.116<br />
Ce<br />
58<br />
232.038<br />
Th<br />
90<br />
+3<br />
+4<br />
140.908<br />
Pr +3<br />
59<br />
144.24<br />
Nd<br />
60<br />
+4 231.036<br />
Pa +4 238.029 +5<br />
U +3<br />
+4<br />
+5<br />
+6<br />
91<br />
92<br />
+3<br />
(145)<br />
Pm<br />
61<br />
+3<br />
150.36<br />
Sm<br />
62<br />
+2<br />
+3<br />
151.964<br />
Eu<br />
63<br />
+2<br />
+3<br />
157.25<br />
Gd<br />
64<br />
+3<br />
158.925<br />
(237)Np (244) Pu (243) Am (247) Cm +3 (247) Bk +3<br />
+3<br />
+4<br />
+5<br />
+6<br />
93 94<br />
+3<br />
+4<br />
+5<br />
+6<br />
65<br />
+3<br />
+4<br />
+5<br />
+6<br />
95 96 97<br />
Tb<br />
+3<br />
+4<br />
162.500<br />
Dy<br />
66<br />
(251)<br />
+3<br />
164.930<br />
Ho<br />
67<br />
+3<br />
167.259<br />
Er<br />
68<br />
Cf +3 (252) Es (257) Fm<br />
100<br />
98 99<br />
+3<br />
+3<br />
+3<br />
168.934<br />
Tm +3<br />
69<br />
(258)<br />
Md<br />
101<br />
+2<br />
+3<br />
173.04<br />
Yb<br />
70<br />
(259)<br />
No<br />
102<br />
+2<br />
+3<br />
+2<br />
+3<br />
174.9668<br />
Lu<br />
71<br />
(262)<br />
Lr<br />
103<br />
+3<br />
+3<br />
*denotes the presence of (2-8-) for elements 72 and above<br />
**The systematic names and symbols for elements of atomic numbers 113 and above<br />
will be used until the approval of trivial names by IUPAC.<br />
Source: CRC Handbook of <strong>Chemistry</strong> and Physics, 91 st ed., 2010–2011, CRC Press<br />
9<br />
Period<br />
1<br />
2<br />
3<br />
4<br />
5<br />
6<br />
7<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 9
Table S<br />
Properties of Selected Elements<br />
First<br />
Atomic Symbol Name Ionization<br />
Electro- Melting Boiling* Density** Atomic<br />
Number Energy negativity Point Point (g/cm 3 ) Radius<br />
(kJ/mol) (K) (K) (pm)<br />
1 H hydrogen 1312 2.2 14 20. 0.000082 32<br />
2 He helium 2372 — — 4 0.000164 37<br />
3 Li lithium 520. 1.0 454 1615 0.534 130.<br />
4 Be beryllium 900. 1.6 1560. 2744 1.85 99<br />
5 B boron 801 2.0 2348 4273 2.34 84<br />
6 C carbon 1086 2.6 — — .— 75<br />
7 N nitrogen 1402 3.0 63 77 0.001145 71<br />
8 O oxygen 1314 3.4 54 90. 0.001308 64<br />
9 F fluorine 1681 4.0 53 85 0.001553 60.<br />
10 Ne neon 2081 — 24 27 0.000825 62<br />
11 Na sodium 496 0.9 371 1156 0.97 160.<br />
12 Mg magnesium 738 1.3 923 1363 1.74 140.<br />
13 Al aluminum 578 1.6 933 2792 2.70 124<br />
14 Si silicon 787 1.9 1687 3538 2.3296 114<br />
15 P phosphorus (white) 1012 2.2 317 554 1.823 109<br />
16 S sulfur (monoclinic) 1000. 2.6 388 718 2.00 104<br />
17 Cl chlorine 1251 3.2 172 239 0.002898 100.<br />
18 Ar argon 1521 — 84 87 0.001633 101<br />
19 K potassium 419 0.8 337 1032 0.89 200.<br />
20 Ca calcium 590. 1.0 1115 1757 1.54 174<br />
21 Sc scandium 633 1.4 1814 3109 2.99 159<br />
22 Ti titanium 659 1.5 1941 3560. 4.506 148<br />
23 V vanadium 651 1.6 2183 3680. 6.0 144<br />
24 Cr chromium 653 1.7 2180. 2944 7.15 130.<br />
25 Mn manganese 717 1.6 1519 2334 7.3 129<br />
26 Fe iron 762 1.8 1811 3134 7.87 124<br />
27 Co cobalt 760. 1.9 1768 3200. 8.86 118<br />
28 Ni nickel 737 1.9 1728 3186 8.90 117<br />
29 Cu copper 745 1.9 1358 2835 8.96 122<br />
30 Zn zinc 906 1.7 693 1180. 7.134 120.<br />
31 Ga gallium 579 1.8 303 2477 5.91 123<br />
32 Ge germanium 762 2.0 1211 3106 5.3234 120.<br />
33 As arsenic (gray) 944 2.2 1090. — 5.75 120.<br />
34 Se selenium (gray) 941 2.6 494 958 4.809 118<br />
35 Br bromine 1140. 3.0 266 332 3.1028 117<br />
36 Kr krypton 1351 — 116 120. 0.003425 116<br />
37 Rb rubidium 403 0.8 312 961 1.53 215<br />
38 Sr strontium 549 1.0 1050. 1655 2.64 190.<br />
39 Y yttrium 600. 1.2 1795 3618 4.47 176<br />
40 Zr zirconium 640. 1.3 2128 4682 6.52 164<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 10
First<br />
Atomic Symbol Name Ionization<br />
Electro- Melting Boiling* Density** Atomic<br />
Number Energy negativity Point Point (g/cm 3 ) Radius<br />
(kJ/mol) (K) (K) (pm)<br />
41 Nb niobium 652 1.6 2750. 5017 8.57 156<br />
42 Mo molybdenum 684 2.2 2896 4912 10.2 146<br />
43 Tc technetium 702 2.1 2430. 4538 11 138<br />
44 Ru ruthenium 710. 2.2 2606 4423 12.1 136<br />
45 Rh rhodium 720. 2.3 2237 3968 12.4 134<br />
46 Pd palladium 804 2.2 1828 3236 12.0 130.<br />
47 Ag silver 731 1.9 1235 2435 10.5 136<br />
48 Cd cadmium 868 1.7 594 1040. 8.69 140.<br />
49 In indium 558 1.8 430. 2345 7.31 142<br />
50 Sn tin (white) 709 2.0 505 2875 7.287 140.<br />
51 Sb antimony (gray) 831 2.1 904 1860. 6.68 140.<br />
52 Te tellurium 869 2.1 723 1261 6.232 137<br />
53 I iodine 1008 2.7 387 457 4.933 136<br />
54 Xe xenon 1170. 2.6 161 165 0.005366 136<br />
55 Cs cesium 376 0.8 302 944 1.873 238<br />
56 Ba barium 503 0.9 1000. 2170. 3.62 206<br />
57 La lanthanum 538 1.1 1193 3737 6.15 194<br />
Elements 58–71 have been omitted.<br />
72 Hf hafnium 659 1.3 2506 4876 13.3 164<br />
73 Ta tantalum 728 1.5 3290. 5731 16.4 158<br />
74 W tungsten 759 1.7 3695 5828 19.3 150.<br />
75 Re rhenium 756 1.9 3458 5869 20.8 141<br />
76 Os osmium 814 2.2 3306 5285 22.587 136<br />
77 Ir iridium 865 2.2 2719 4701 22.562 132<br />
78 Pt platinum 864 2.2 2041 4098 21.5 130.<br />
79 Au gold 890. 2.4 1337 3129 19.3 130.<br />
80 Hg mercury 1007 1.9 234 630. 13.5336 132<br />
81 Tl thallium 589 1.8 577 1746 11.8 144<br />
82 Pb lead 716 1.8 600. 2022 11.3 145<br />
83 Bi bismuth 703 1.9 544 1837 9.79 150.<br />
84 Po polonium 812 2.0 527 1235 9.20 142<br />
85 At astatine — 2.2 575 — — 148<br />
86 Rn radon 1037 — 202 211 0.009074 146<br />
87 Fr francium 393 0.7 300. — — 242<br />
88 Ra radium 509 0.9 969 — 5 211<br />
89 Ac actinium 499 1.1 1323 3471 10. 201<br />
Elements 90 and above have been omitted.<br />
*boiling point at standard pressure<br />
**density of solids and liquids at room temperature and density of gases at 298 K and 101.3 kPa<br />
— no data available<br />
Source: CRC Handbook for <strong>Chemistry</strong> and Physics, 91 st ed., 2010–2011, CRC Press<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 11
Table T<br />
Important Formulas and Equations<br />
d = density<br />
m<br />
Density d = m = mass<br />
V<br />
V = volume<br />
Mole Calculations number of moles =<br />
given mass<br />
gram-formula mass<br />
measured value – accepted value<br />
Percent Error % error = × 100<br />
accepted value<br />
mass of part<br />
Percent Composition % composition by mass = × 100<br />
mass of whole<br />
mass of solute<br />
parts per million = × 1000000<br />
mass of solution<br />
Concentration<br />
molarity =<br />
moles of solute<br />
liter of solution<br />
Combined Gas Law<br />
P<br />
P = pressure<br />
1<br />
V 1<br />
P<br />
= 2<br />
V 2<br />
V = volume<br />
T 1<br />
T 2 T = temperature<br />
M A<br />
= molarity of H + M B<br />
= molarity of OH –<br />
Titration M A<br />
V A<br />
= M B<br />
V B<br />
V A<br />
= volume of acid V B<br />
= volume of base<br />
q = mCΔT q = heat H f<br />
= heat of fusion<br />
Heat q = mH f<br />
m = mass H v<br />
= heat of vaporization<br />
q = mH v<br />
C=specific heat capacity<br />
ΔT = change in temperature<br />
Temperature<br />
K = °C + 273<br />
K = kelvin<br />
°C = degree Celsius<br />
DET 609 ADU<br />
Reference Tables for Physical Setting/<strong>Chemistry</strong> – 2011 Edition 12
Common Lab Equipment Uses
SCIENTIFIC NOTATION RULES<br />
How to Write Numbers in Scientific Notation<br />
Scientific notation is a standard way of writing very large and very small numbers so that they're<br />
easier to both compare and use in computations. To write in scientific notation, follow the form<br />
N X 10 ᴬ<br />
where N is a number between 1 and 10, but not 10 itself, and A is an integer (positive or negative<br />
number).<br />
RULE #1: Standard Scientific Notation is a number from 1 to 9 followed by a decimal and the<br />
remaining significant figures and an exponent of 10 to hold place value.<br />
Example:<br />
5.43 x 10 2 = 5.43 x 100 = 543<br />
8.65 x 10 – 3 = 8.65 x .001 = 0.00865<br />
****54.3 x 10 1 is not Standard Scientific Notation!!!<br />
RULE #2: When the decimal is moved to the Left the exponent gets Larger, but the value of the<br />
number stays the same. Each place the decimal moves Changes the exponent by one (1). If you<br />
move the decimal to the Right it makes the exponent smaller by one (1) for each place it is moved.<br />
Example:<br />
6000. x 10 0 = 600.0 x 10 1 = 60.00 x 10 2 = 6.000 x 10 3 = 6000<br />
(Note: 10 0 = 1)<br />
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.<br />
Example:<br />
(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.<br />
(3.0 x 10 2 )<br />
+ (64. x 10 2 )<br />
67.0 x 10 2 = 6.70 x 10 3 = 6.7 x 10 3<br />
67.0 x 10 2 is mathematically correct, but a number in standard scientific notation can only<br />
have one number to the left of the decimal, so the decimal is moved to the left one place and<br />
one is added to the exponent.<br />
Following the rules for significant figures, the answer becomes 6.7 x 10 3 .<br />
RULE #4: To multiply, find the product of the numbers, then add the exponents.<br />
Example:<br />
(2.4 x 10 2 ) (5.5 x 10 –4 ) = ? [2.4 x 5.5 = 13.2]; [2 + -4 = -2], so<br />
(2.4 x 10 2 ) (5.5 x 10 –4 ) = 13.2 x 10 –2 = 1.3 x 10 – 1<br />
RULE #5: To divide, find the quotient of the number and subtract the exponents.<br />
Example:<br />
(3.3 x 10 – 6 ) / (9.1 x 10 – 8 ) = ? [3.3 / 9.1 = .36]; [-6 – (-8) = 2], so<br />
(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<br />
2. 884.2 7. 3.622<br />
3. 0.00002887 8. 0.00001735<br />
4. 0.05893 9. 0.9736<br />
5. 0.006162 10. 0.08558<br />
11. 6.633 X 10−⁴ 16. 1.937 X 10⁴<br />
12. 4.445 X 10−⁴ 17. 3.457 X 10⁴<br />
13. 2.182 X 10−³ 18. 3.948 X 10−⁵<br />
14. 4.695 X 10² 19. 8.945 X 10⁵<br />
15. 7.274 X 10⁵ 20. 6.783 X 10²
Convert each number from Scientific Notation to real numbers:<br />
1. 7.485 X 10³ 6. 1.683 X 10⁰<br />
2. 8.842 X 10² 7. 3.622 10⁰<br />
3. 2.887 X 10−⁵ 8. 1.735 X 10−⁵<br />
4. 5.893 X 10−² 9. 9.736 X 10−¹<br />
5. 6.162 X 10−³ 10. 8.558 X 10−²<br />
Convert each number from a real number to Scientific Notation:<br />
11. 0.0006633 16. 1,937,000<br />
12. 0.0004445 17. 34,570<br />
13. 0.002182 18. 0.00003948<br />
14. 469.5 19. 894,500<br />
15. 727,400 20. 678.3
Significant Figures Rules<br />
There are three rules on determining how many significant figures are in a<br />
number:<br />
1. Non-zero digits are always significant.<br />
2. Any zeros between two significant digits are significant.<br />
3. A final zero or trailing zeros in the DECIMAL PORTION ONLY are<br />
significant.<br />
Please remember that, in science, all numbers are based upon measurements (except for a very few<br />
that are defined). Since all measurements are uncertain, we must only use those numbers that are<br />
meaningful.<br />
Not all of the digits have meaning (significance) and, therefore, should not be written down. In<br />
science, only the numbers that have significance (derived from measurement) are written.<br />
Rule 1: Non-zero digits are always significant.<br />
If you measure something and the device you use (ruler, thermometer, triple-beam, balance, etc.)<br />
returns a number to you, then you have made a measurement decision and that ACT of measuring<br />
gives significance to that particular numeral (or digit) in the overall value you obtain.<br />
Hence a number like 46.78 would have four significant figures and 3.94 would have three.<br />
Rule 2: Any zeros between two significant digits are significant.<br />
Suppose you had a number like 409. By the first rule, the 4 and the 9 are significant. However, to<br />
make a measurement decision on the 4 (in the hundred's place) and the 9 (in the one's place), you<br />
HAD to have made a decision on the ten's place. The measurement scale for this number would have<br />
hundreds, tens, and ones marked.<br />
Like the following example:<br />
These are sometimes called "captured zeros."<br />
If a number has a decimal at the end (after the one’s place) then all digits (numbers) are significant<br />
and will be counted.<br />
In the following example the zeros are significant digits and highlighted in blue.<br />
960.<br />
70050.
Rule 3: A final zero or trailing zeros in the decimal portion ONLY are<br />
significant.<br />
This rule causes the most confusion among students.<br />
In the following example the zeros are significant digits and highlighted in blue.<br />
0.07030<br />
0.00800<br />
Here are two more examples where the significant zeros are highlighted in blue.<br />
When Zeros are Not Significant Digits<br />
4.7 0 x 10−³<br />
6.5 0 0 x 10⁴<br />
Zero Type # 1 : Space holding zeros in numbers less than one.<br />
In the following example the zeros are NOT significant digits and highlighted in red.<br />
0.09060<br />
0.00400<br />
These zeros serve only as space holders. They are there to put the decimal point in its correct<br />
location.<br />
They DO NOT involve measurement decisions.<br />
Zero Type # 2 : Trailing zeros in a whole number.<br />
In the following example the zeros are NOT significant digits and highlighted in red.<br />
200<br />
25000<br />
For addition and subtraction, look at the decimal portion (i.e., to the right of the decimal point)<br />
of the numbers ONLY. Here is what to do:<br />
1) Count the number of significant figures in the decimal portion of each number in the problem. (The<br />
digits to the left of the decimal place are not used to determine the number of decimal places in the<br />
final answer.)<br />
2) Add or subtract in the normal fashion.<br />
3) Round the answer to the LEAST number of places in the decimal portion of any number in the<br />
problem<br />
The following rule applies for multiplication and division:<br />
The LEAST number of significant figures in any number of the problem determines the number of<br />
significant figures in the answer.<br />
This means you MUST know how to recognize significant figures in order to use this rule.
How Many Significant Digits for Each Number?<br />
1) 2359 = ______<br />
2) 2.445 x 10−⁵= ______<br />
3) 2.93 x 10⁴= ______<br />
4) 1.30 x 10−⁷= ______<br />
5) 2604 = ______<br />
6) 9160 = ______<br />
7) 0.0800 = ______<br />
8) 0.84 = ______<br />
9) 0.0080 = ______<br />
10) 0.00040 = ______<br />
11) 0.0520 = ______<br />
12) 0.060 = ______<br />
13) 6.90 x 10−¹= ______<br />
14) 7.200 x 10⁵= ______<br />
15) 5.566 x 10−²= ______<br />
16) 3.88 x 10⁸= ______<br />
17) 3004 = ______<br />
18) 0.021 = ______<br />
19) 240 = ______<br />
20) 500 = ______
For addition and subtraction, look at the decimal portion (i.e., to the right of the decimal point) of the<br />
numbers ONLY. Here is what to do:<br />
1) Count the number of significant figures in the decimal portion of each number in the problem. (The<br />
digits to the left of the decimal place are not used to determine the number of decimal places in the<br />
final answer.)<br />
2) Add or subtract in the normal fashion.<br />
3) Round the answer to the LEAST number of places in the decimal portion of any number in the<br />
problem.<br />
Solve the Problems and Round Accordingly...<br />
1) 43.287 + 5.79 + 6.284 = _______<br />
2) 87.54 - 3.3 = _______<br />
3) 99.1498 + 6.5397 + 9.7 = _______<br />
4) 5.868 - 5.1 = _______<br />
5) 59.9233 + 86.21 + 99.396 = _______<br />
6) 7.7 + 26.756 = _______<br />
7) 66.8 + 2.3 + 4.8516 = _______<br />
8) 9.7419 + 43.545 = _______<br />
9) 4.8976 + 48.4644 + 1.514 = _______<br />
10) 4.335 + 35.85 = _______<br />
11) 9.448 - 1.7 = _______<br />
12) 75.826 - 8.6555 = _______<br />
13) 57.2 + 23.814 = _______<br />
14) 77.684 - 4.394 = _______<br />
15) 26.4496 + 3.339 = _______<br />
16) 9.6848 + 29.85 = _______<br />
17) 63.11 + 2.5412 + 4.93 = _______<br />
18) 11.2471 + 75.4 = _______<br />
19) 73.745 - 8.755 = _______<br />
20) 6.5238 + 1.7 + 27.79 = _______
The following rule applies for multiplication and division:<br />
The LEAST number of significant figures in any number of the problem determines the number of<br />
significant figures in the answer.<br />
This means you MUST know how to recognize significant figures in order to use this rule.<br />
Solve the Problems and Round Accordingly...<br />
1) 0.6 x 65.0 x 602 = __________<br />
2) 720 ÷ 7.7 = __________<br />
3) 929 x 0.3 = __________<br />
4) 300 ÷ 44.31 = __________<br />
5) 608 ÷ 9.8 = __________<br />
6) 0.06 x 0.079 = __________<br />
7) 0.008 x 72.91 x 7000 = __________<br />
8) 73.94 x 67 x 780 = __________<br />
9) 0.62 x 0.097 x 40 = __________<br />
10) 600 x 10 x 5030 = __________<br />
11) 5200 ÷ 4.46 = __________<br />
12) 0.0052 x 0.4 x 107 = __________<br />
13) 0.099 x 8.8 = __________<br />
14) 0.0095 x 5.2 = __________<br />
15) 8000 ÷ 4.62 = __________<br />
16) 0.6 x 0.8 = __________<br />
17) 2.84 x 0.66 = __________<br />
18) 0.5 x 0.09 = __________<br />
19) 8100 ÷ 34.84 = __________<br />
20) 8.24 x 6.9 x 8100 = __________
Question Sig Figs Question Add & Subtract Question Multiple & Divide<br />
1 4 1 55.36 1 20,000<br />
2 4 2 84.2 2 94<br />
3 3 3 115.4 3 300<br />
4 3 4 0.8 4 7<br />
5 4 5 245.53 5 62<br />
6 3 6 34.5 6 0.005<br />
7 3 7 74.0 7 4,000<br />
8 2 8 53.287 8 3,900,000<br />
9 2 9 54.876 9 2<br />
10 2 10 40.19 10 30,000,000<br />
11 3 11 7.7 11 1,200<br />
12 2 12 67.170 12 0.2<br />
13 3 13 81.0 13 0.87<br />
14 4 14 73.290 14 0.049<br />
15 4 15 29.789 15 2,000<br />
16 3 16 39.53 16 0.5<br />
17 4 17 70.58 17 1.9<br />
18 2 18 86.6 18 0.05<br />
19 2 19 64.990 19 230<br />
20 1 20 36.0 20 460,000
To use the Stair-Step method, find the prefix the original measurement starts with. (ex. milligram)<br />
If there is no prefix, then you are starting with a base unit.<br />
Find the step which you wish to make the conversion to. (ex. decigram)<br />
Count the number of steps you moved, and determine in which direction you moved (left or right).<br />
The decimal in your original measurement moves the same number of places as steps you moved and in the<br />
same direction. (ex. milligram to decigram is 2 steps to the left, so 40 milligrams = .40 decigrams)<br />
If the number of steps you move is larger than the number you have, you will have to add zeros to hold the<br />
places. (ex. kilometers to meters is three steps to the right, so 10 kilometers would be equal to 10,000 m)<br />
That’s all there is to it! You need to be able to count to 6, and know your left from your right!<br />
1) Write the equivalent<br />
a) 5 dm =_______m b) 4 mL = ______L c) 8 g = _______mg<br />
d) 9 mg =_______g e) 2 mL = ______L f) 6 kg = _____g<br />
g) 4 cm =_______m h) 12 mg = ______ g i) 6.5 cm 3 = _______L<br />
j) 7.02 mL =_____cm 3 k) .03 hg = _______ dg l) 6035 mm _____cm<br />
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<br />
than 1 kg? Explain your answer.<br />
3. Wanda needs to move 110 kg of rocks. She can carry l0 hg each trip. How many trips must she<br />
make? Explain your answer.<br />
4. Dr. O is playing in her garden again She needs 1 kg of potting soil for her plants. She has 750 g.<br />
How much more does she need? Explain your answer.<br />
5. Weather satellites orbit Earth at an altitude of 1,400,000 meters. What is this altitude in kilometers?<br />
6. Which unit would you use to measure the capacity? Write milliliter or liter.<br />
a) a bucket __________<br />
b) a thimble __________<br />
c) a water storage tank__________<br />
d) a carton of juice__________<br />
7. Circle the more reasonable measure:<br />
a) length of an ant 5mm or 5cm<br />
b) length of an automobile 5 m or 50 m<br />
c) distance from NY to LA 450 km or 4,500 km<br />
d) height of a dining table 75 mm or 75 cm<br />
8. Will a tablecloth that is 155 cm long cover a table that is 1.6 m long? Explain your answer.<br />
9. A dollar bill is 15.6 cm long. If 200 dollar bills were laid end to end, how many meters long would<br />
the line be?<br />
10. The ceiling in Jan’s living room is 2.5 m high. She has a hanging lamp that hangs down 41 cm.<br />
Her husband is exactly 2 m tall. Will he hit his head on the hanging lamp? Why or why not?
Using SI Units<br />
Match the terms in Column II with the descriptions in Column I. Write the letters of the correct term in<br />
the blank on the left.<br />
Column I Column II<br />
_____ 1. distance between two points<br />
a. time<br />
_____ 2. SI unit of length<br />
_____ 3. tool used to measure length<br />
_____ 4. units obtained by combining other units<br />
_____ 5. amount of space occupied by an object<br />
_____ 6. unit used to express volume<br />
_____ 7. SI unit of mass<br />
_____ 8. amount of matter in an object<br />
_____ 9. mass per unit of volume<br />
_____ 10. temperature scale of most laboratory thermometers<br />
_____ 11. instrument used to measure mass<br />
_____ 12. interval between two events<br />
_____ 13. SI unit of temperature<br />
_____ 14. SI unit of time<br />
_____ 15. instrument used to measure temperature<br />
b. volume<br />
c. mass<br />
d. density<br />
e. meter<br />
f. kilogram<br />
g. derived<br />
h. liter<br />
i. second<br />
j. Kelvin<br />
k. length<br />
1. balance<br />
m. meterstick<br />
n. thermometer<br />
o. Celsius<br />
Circle the two terms in each group that are related. Explain how the terms are related.<br />
16. Celsius degree, mass, Kelvin _____________________________________________________<br />
________________________________________________________________________________<br />
17. balance, second, mass __________________________________________________________<br />
________________________________________________________________________________<br />
18. kilogram, liter, cubic centimeter __________________________________________________<br />
________________________________________________________________________________<br />
19. time, second, distance __________________________________________________________<br />
________________________________________________________________________________<br />
20. decimeter, kilometer, Kelvin _____________________________________________________<br />
________________________________________________________________________________
1. How many meters are in one kilometer? __________<br />
2. What part of a liter is one milliliter? __________<br />
3. How many grams are in two dekagrams? __________<br />
4. If one gram of water has a volume of one milliliter, what would the mass of one liter of water be in<br />
kilograms?__________<br />
5. What part of a meter is a decimeter? __________<br />
In the blank at the left, write the term that correctly completes each statement. Choose from the terms<br />
listed below.<br />
Metric SI standard ten<br />
prefixes ten tenth<br />
6. An exact quantity that people agree to use for comparison is a ______________ .<br />
7. The system of measurement used worldwide in science is _______________ .<br />
8. SI is based on units of _______________ .<br />
9. The system of measurement that was based on units of ten was the _______________ system.<br />
10. In SI, _______________ are used with the names of the base unit to indicate the multiple of ten<br />
that is being used with the base unit.<br />
11. The prefix deci- means _______________ .
Standards of Measurement<br />
Fill in the missing information in the table below.<br />
S.I prefixes and their meanings<br />
Prefix<br />
Meaning<br />
0.001<br />
0.01<br />
deci- 0.1<br />
10<br />
hecto- 100<br />
1000<br />
Circle the larger unit in each pair of units.<br />
1. millimeter, kilometer 4. centimeter, millimeter<br />
2. decimeter, dekameter 5. hectogram, kilogram<br />
3. hectogram, decigram<br />
6. In SI, the base unit of length is the meter. Use this information to arrange the following units of<br />
measurement in the correct order from smallest to largest.<br />
Write the number 1 (smallest) through 7 - (largest) in the spaces provided.<br />
_____ a. kilometer<br />
_____ b. centimeter<br />
_____ c. meter<br />
_____ e. hectometer<br />
_____ f. millimeter<br />
_____ g. decimeter<br />
_____ d. dekameter<br />
Use your knowledge of the prefixes used in SI to answer the following questions in the spaces<br />
provided.<br />
7. One part of the Olympic games involves an activity called the decathlon. How many events do you<br />
think make up the decathlon?_____________________________________________________<br />
8. How many years make up a decade? _______________________________________________<br />
9. How many years make up a century? ______________________________________________<br />
10. What part of a second do you think a millisecond is? __________________________________
Dimensional Analysis<br />
This is a way to convert from one unit of a given substance to<br />
another unit using ratios or conversion units. What this video<br />
www.youtube.com/watch?v=aZ3J60GYo6U<br />
Let’ look at a couple of examples:<br />
1. Convert 2.6 qt to mL.<br />
First we need a ratio or conversion unit so that we can go from quarts to milliliters. 1.00 qt = 946 mL<br />
Next write down what you are starting with<br />
2.6 qt<br />
Then make you conversion tree<br />
2.6 qt<br />
Then fill in the units in your ratio so that you can cancel out the original unit and will be left with the<br />
unit you need for the answer. Cross out units, one at a time that are paired, and one on top one on<br />
the bottom.<br />
2.6 qt mL<br />
qt<br />
Now fill in the values from the ratio.<br />
2.6 qt 946 mL<br />
1.00 qt<br />
Now multiply all numbers on the top and multiply all numbers on the bottom and write them as a<br />
fraction.<br />
2.6 qt 946 mL = 2,459.6 mL<br />
1.00 qt 1.00<br />
Now divide the top number by the bottom number and write that number with the unit that was not<br />
crossed out.
1qt=32 oz 1gal = 4qts 1.00 qt = 946 mL 1L = 1000mL<br />
2. Convert 8135.6 mL to quarts<br />
=<br />
3. Convert 115.2 oz to mL<br />
=<br />
4. Convert 2.3 g to Liters<br />
=<br />
5. Convert 8.42 L to oz<br />
=<br />
Go to http://science.widener.edu/svb/tutorial/ chose #7 “Converting Volume” and do 5 more in the<br />
space provided.<br />
1. Convert _________ to _________<br />
=<br />
2. Convert _________ to _________<br />
=<br />
3. Convert _________ to _________<br />
=<br />
4. Convert _________ to _________<br />
=<br />
5. Convert _________ to _________<br />
=
Converting Real Things<br />
Table 1<br />
Using the scale, come up with a conversion ratio just by looking at the scale and prove that it works<br />
but converting 10 grams to ounces. (Hint: 7grams and 15 grams)<br />
Table 2<br />
Using a 50 mL graduated cylinder, fill a 600mL beaker 66.7% full.<br />
Table 3<br />
Using the measuring cup determine how many mL are in 4 ounces of water. Make a conversion ratio<br />
you could use to do other conversions.<br />
Table 4<br />
Using a meter stick, measure the back table in inches, feet and yards and convert them into<br />
centimeters and meters.<br />
Table 5<br />
Using the scale, find the mass of a text book on your table in ounces (round to the nearest ounce)<br />
and cover it to grams.<br />
Table 6<br />
Using the ruler, measure the length of a piece of paper in inches and then convert that into meters.<br />
Make a conversion ratio you could use to do other conversions.
Atoms Are Building Blocks<br />
Atoms are the basis of chemistry. They are the basis for everything in the Universe. You<br />
should start by remembering that matter is composed of atoms. Atoms and the study of<br />
atoms are a world unto themselves. We're going to cover basics like atomic structure<br />
and bonding between atoms.<br />
Smaller Than Atoms?<br />
Are there pieces of matter that are smaller than atoms?<br />
Sure there are. You'll soon be learning that atoms are<br />
composed of pieces like electrons, protons, and neutrons.<br />
But guess what? There are even smaller particles moving<br />
around in atoms. These super-small particles can be found<br />
inside the protons and neutrons. Scientists have many<br />
names for those pieces, but you may have heard of<br />
nucleons and quarks. Nuclear chemists and physicists<br />
work together at particle accelerators to discover the<br />
presence of these tiny, tiny, tiny pieces of matter.<br />
Even though super-tiny atomic particles exist, you only<br />
need to remember the three basic parts of an atom: electrons, protons, and neutrons.<br />
What are electrons, protons, and neutrons? A picture works best to show off the idea.<br />
You have a basic atom. There are three types of pieces in that atom: electrons, protons,<br />
and neutrons. That's all you have to remember. Three things! As you know, there are<br />
almost 120 known elements in the periodic table. Chemists and physicists haven't<br />
stopped there. They are trying to make new ones in labs every day. The thing that<br />
makes each of those elements different is the number of electrons, protons, and<br />
neutrons. The protons and neutrons are always in the center of the atom. Scientists call<br />
the center region of the atom the nucleus. The nucleus in<br />
a cell is a thing. The nucleus in an atom is a place where<br />
you find protons and neutrons. The electrons are always<br />
found whizzing around the center in areas called shells or<br />
orbitals.<br />
You can also see that each piece has either a "+", "-", or a<br />
"0." That symbol refers to the charge of the particle. Have<br />
you ever heard about getting a shock from a socket, static<br />
electricity, or lightning? Those are all different types of<br />
electric charges. Those charges are also found in tiny particles of matter. The electron<br />
always has a "-", or negative, charge. The proton always has a "+", or positive, charge. If<br />
the charge of an entire atom is "0", or neutral, there are equal numbers of positive and<br />
negative pieces. Neutral means there are equal numbers of electrons and protons. The<br />
third particle is the neutron. It has a neutral charge, also known as a charge of zero. All<br />
atoms have equal numbers of protons and electrons so that they are neutral. If there are<br />
more positive protons or negative electrons in an atom, you have a special atom called<br />
an ion.
http://www.learner.org/interactives/periodic/basics_interactive.html
Looking at Ions<br />
We haven’t talked about ions before, so let’s get down to basics. The<br />
atomic number of an element, also called a proton number, tells you the<br />
number of protons or positive particles in an atom. A normal atom has a<br />
neutral charge with equal numbers of positive and negative particles.<br />
That means an atom with a neutral charge is one where the number of<br />
electrons is equal to the atomic number. Ions are atoms with extra<br />
electrons or missing electrons. When you are missing an electron or<br />
two, you have a positive charge. When you have an extra electron<br />
or two, you have a negative charge.<br />
What do you do if you are a sodium (Na) atom? You have eleven<br />
electrons — one too many to have an entire shell filled. You need to<br />
find another element that will take that electron away from you. When you lose that<br />
electron, you will you’ll have full shells. Whenever an atom has full shells, we say it is<br />
"happy." Let's look at chlorine (Cl). Chlorine has seventeen electrons and only needs<br />
one more to fill its third shell and be "happy." Chlorine will take your extra sodium<br />
electron and leave you with 10 electrons inside of two filled shells. You are now a happy<br />
atom too. You are also an ion and missing one electron. That missing electron gives you<br />
a positive charge. You are still the element sodium, but you are now a sodium ion (Na + ).<br />
You have one less electron than your atomic number.<br />
Ion Characteristics<br />
So now you've become a sodium ion. You have ten electrons.<br />
That's the same number of electrons as neon (Ne). But you<br />
aren't neon. Since you're missing an electron, you aren't really<br />
a complete sodium atom either. As an ion you are now<br />
something completely new. Your whole goal as an atom was<br />
to become a "happy atom" with completely filled electron<br />
shells. Now you have those filled shells. You have a lower<br />
energy. You lost an electron and you are "happy." So what<br />
makes you interesting to other atoms? Now that you have<br />
given up the electron, you are quite electrically attractive.<br />
Other electrically charged atoms (ions) of the opposite charge<br />
(negative) are now looking at you and seeing a good partner to<br />
bond with. That's where the chlorine comes in. It's not only chlorine. Almost any ion with<br />
a negative charge will be interested in bonding with you.
Electrovalence<br />
Don't get worried about the big word. Electrovalence is just another word for something<br />
that has given up or taken electrons and become an ion. If you look at the periodic table,<br />
you might notice that elements on the left side usually become positively charged ions<br />
(cations) and elements on the right side get a negative charge (anions). That trend<br />
means that the left side has a positive valence and the right side has a negative<br />
valence. Valence is a measure of how much an atom wants to bond with other atoms. It<br />
is also a measure of how many electrons are excited about bonding with other atoms.<br />
There are two main types of bonding, covalent and electrovalent. You may have heard<br />
of the term "ionic bonds." Ionic bonds are electrovalent bonds. They are just groups of<br />
charged ions held together by electric forces. When in the presence of other ions, the<br />
electrovalent bonds are weaker because of outside electrical forces and attractions.<br />
Sodium and chlorine ions alone have a very strong bond, but as soon as you put those<br />
ions in a solution with H + (Hydrogen ion), OH - (Hydroxide), F - (Fluorine ion) or Mg ++<br />
(Magnesium ion), there are charged distractions that break the Na-Cl bond.<br />
Look at sodium chloride (NaCl) one more time. Salt is a very strong bond when it is<br />
sitting on your table. It would be nearly impossible to break those ionic/electrovalent<br />
bonds. However, if you put that salt into some water (H 2 O), the bonds break very<br />
quickly. It happens easily because of the electrical attraction of the water. Now you have<br />
sodium (Na + ) and chlorine (Cl - ) ions floating around the solution. You should remember<br />
that ionic bonds are normally strong, but they are very weak in water.
Neutron Madness<br />
We have already learned that ions are atoms that are<br />
either missing or have extra electrons. Let's say an atom<br />
is missing a neutron or has an extra neutron. That type of<br />
atom is called an isotope. An atom is still the same<br />
element if it is missing an electron. The same goes for<br />
isotopes. They are still the same element. They are just a<br />
little different from every other atom of the same element.<br />
For example, there are a lot of carbon (C) atoms in the<br />
Universe. The normal ones are carbon-12. Those atoms have 6 neutrons. There are a<br />
few straggler atoms that don't have 6. Those odd ones may have 7 or even 8 neutrons.<br />
As you learn more about chemistry, you will probably hear about carbon-14. Carbon-14<br />
actually has 8 neutrons (2 extra). C-14 is considered an isotope of the element carbon.<br />
Messing with the Mass<br />
If you have looked at a periodic table, you may have noticed that the atomic mass of<br />
an element is rarely an even number. That happens because of the isotopes. If you are<br />
an atom with an extra electron, it's no big deal. Electrons don't have much of a mass<br />
when compared to a neutron or proton.<br />
Atomic masses are calculated by figuring out the<br />
amounts of each type of atom and isotope there are in<br />
the Universe. For carbon, there are a lot of C-12, a<br />
couple of C-13, and a few C-14 atoms. When you<br />
average out all of the masses, you get a number that is a<br />
little bit higher than 12 (the weight of a C-12 atom). The<br />
average atomic mass for the element is actually 12.011.<br />
Since you never really know which carbon atom you are<br />
using in calculations, you should use the average mass<br />
of an atom.<br />
Bromine (Br), at atomic number 35, has a greater variety of isotopes. The atomic mass<br />
of bromine (Br) is 79.90. There are two main isotopes at 79 and 81, which average out<br />
to the 79.90amu value. The 79 has 44 neutrons and the 81 has 46 neutrons. While it<br />
won't change the average atomic mass, scientists have made bromine isotopes with<br />
masses from 68 to 97. It's all about the number of neutrons. As you move to higher<br />
atomic numbers in the periodic table, you will probably find even more isotopes for<br />
each element.
I did not use publisher for this part because my computer does<br />
not have it to be used. I used a PDF file and used as much of it<br />
as I could.
I did not use publisher for this part because my computer does<br />
not have it to be used. I used a PDF file and used as much of it<br />
as I could.
I did not use publisher for this part because my computer does<br />
not have it to be used. I used a PDF file and used as much of it<br />
as I could.
I did not use publisher for this part because my computer does<br />
not have it to be used. I used a PDF file and used as much of it<br />
as I could.
P<br />
N<br />
P<br />
N
P<br />
N<br />
P<br />
N
Electron Configuration<br />
Color the sublevel:<br />
s = Red<br />
d = Green<br />
p = Blue<br />
f = Orange<br />
Write in sublevels<br />
Write period, sublevel and super scripts.<br />
Ctrl Shift =<br />
gives you super scripts
www.youtube.com/watch?v=jtYzEzykFdg<br />
www.youtube.com/watch?<br />
annotation_id=annotation_2076&feature=iv&src_vid=jtYzEzykFdg&v=cOlac8ruD_0<br />
www.youtube.com/watch?<br />
annotation_id=annotation_570977&feature=iv&src_vid=cOlac8ruD_0&v=lR2vqHZWb5A
Electron Configuration<br />
In order to write the electron configuration for an atom you must know the 3 rules of<br />
electron configurations.<br />
1. Aufbau<br />
Notation<br />
nO e<br />
where<br />
n is the energy level<br />
O is the orbital type (s, p, d, or f)<br />
e is the number of electrons in that orbital shell<br />
Principle<br />
electrons will first occupy orbitals of the lowest energy level<br />
2. Hund rule<br />
when electrons occupy orbitals of equal energy, one electron enters each orbital until<br />
all the orbitals contain one electron with the same spin.<br />
3. Pauli exclusion principle<br />
an orbital contains a maximum of 2 electrons and<br />
paired electrons will have opposite spin
In the space below, write the unabbreviated electron configurations of the following elements:<br />
1) sodium ________________________________________________<br />
2) iron ________________________________________________<br />
3) bromine ________________________________________________<br />
4) barium ________________________________________________<br />
5) neptunium ________________________________________________<br />
In the space below, write the abbreviated electron configurations of the following elements:<br />
6) cobalt ________________________________________________<br />
7) silver ________________________________________________<br />
8) tellurium ________________________________________________<br />
9) radium ________________________________________________<br />
10) lawrencium ________________________________________________<br />
Determine what elements are denoted by the following electron configurations:<br />
11) 1s²s²2p⁶3s²3p⁴ ____________________<br />
12) 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s¹ ____________________<br />
13) [Kr] 5s²4d¹⁰5p³ ____________________<br />
14) [Xe] 6s²4f¹⁴5d⁶ ____________________<br />
15) [Rn] 7s²5f¹¹ ____________________<br />
Identify the element or determine that it is not a valid electron configuration:<br />
16) 1s²2s²2p⁶3s²3p⁶4s²4d¹⁰4p⁵ ____________________<br />
17) 1s²2s²2p⁶3s³3d⁵ ____________________<br />
18) [Ra] 7s²5f⁸ ____________________<br />
19) [Kr] 5s²4d¹⁰5p⁵ ____________________<br />
20) [Xe] ____________________<br />
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<br />
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<br />
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<br />
7)silver [Kr] 5s 2 4d 9 8)tellurium[Kr] 5s 2 4d 10 5p 4<br />
9)radium [Rn] 7s 2 10)lawrencium[Rn] 7s 2 5f 14 6d 1<br />
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<br />
[Kr] 5s 2 4d 10 5p 3 antimony [Xe] 6s 2 4f 14 5d 6 osmium<br />
[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”)<br />
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)<br />
[Kr] 5s 2 4d 10 5p 5 valid iodine<br />
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<br />
page.<br />
Color your periodic table similar to the one on Pages 168—169.
Create groups for these Scientist and explain your groupings<br />
(use the information you got from your research)
Research these Scientist and summarize their contributions to chemistry<br />
Antoine Henri Becquerel<br />
Niels Bohr<br />
Louis de Barogilie<br />
Glenn Seaborg<br />
Hantaro Nagaoka<br />
Democritus<br />
Marie and Pierre Curie<br />
Eugene Goldstein<br />
Dmitri Mendeleev<br />
J.J. Thomson<br />
James Chadwick<br />
Erwin Shrodinger<br />
John Dalton<br />
Lothar Meyer<br />
Robert Millikan<br />
J.W. Dobereiner<br />
Ernest Rutherford
Chapter 25 Notes<br />
Vocab:<br />
-25.1 Nuclear Radiation-<br />
<br />
<br />
<br />
<br />
<br />
<br />
Radioactivity – reference to the spontaneous emission of rays or particles from certain<br />
elements<br />
Nuclear radiation – the name of the rays and particles emitted from a radioactive source<br />
Radiosiotopes – unstable isotopes<br />
Alpha particle – the helium nuclei emitted by some radioactive sources, has 2 protons<br />
and 2 neutrons<br />
Beta particles – an electron resulting from the breaking apart of an atom, has 1 electron<br />
and 1 proton, but no neutron<br />
Gamma Ray – a high-energy photon emitted by a radioisotope, electromagnetic<br />
radiation with high penetration through the human body<br />
Notes:<br />
<br />
<br />
<br />
<br />
<br />
Antoine Henri Bequerel and Marie and Pierre Curie accidentally discovered that uranium<br />
salt emitted radioactive rays because of a fog left on a photographic plate, and used<br />
radioactivity to explain why this happened.<br />
Atoms become more stable in nuclear or chemical reactions.<br />
Nuclear reactions begin with unstable radioisotopes, which later become much more<br />
stable after the reaction.<br />
Unlike chemical reactions, nuclear reactions are not affected by changes in<br />
temperature, pressure, or the presence of catalysts. Nuclear reactions of a given<br />
radioisotope can’t be slowed down, sped up, or stopped.<br />
Radioactive decay doesn’t require input energy. Radioactivity is continued until unstable<br />
isotopes are changed to become stable isotopes of a different element.<br />
Type Particles Symbol Charge Mass Source Penetration<br />
Alpha Radiation Helium nuclei α 2+ 4 Radium Low (0.05<br />
mm body<br />
tissue)<br />
Beta Radiation Electrons β 1- 1/1837 Carbon Moderate (4<br />
mm body<br />
tissue)
Gamma<br />
Radiation<br />
High energy<br />
electromagnetic<br />
radiation<br />
γ 0 0 Cobalt Very High<br />
(penetrates<br />
body easily)<br />
-Chapter 25.2 Nuclear Transformations-<br />
Vocab:<br />
<br />
<br />
<br />
<br />
<br />
Nuclear force – an attractive force that acts between all nuclear particles that are<br />
extremely close together<br />
Positron – a particle with the mass of an electron but a positive charge<br />
Half-life – the time required for one-half of the nuclei in a radioisotope sample to decay<br />
to products<br />
Transmutation – the conversion of one atom of one element into another atom of<br />
another element<br />
Transuranium elements – elements with atomic numbers above 92 (uranium)<br />
Notes:<br />
<br />
<br />
<br />
The neutron-to-proton ratio in a radioscope determines the type of decay that occurs.<br />
All nuclei that have an atomic number over 83 are radioactive, these contain too many<br />
neutrons and protons that make the atom unstable.<br />
During each half-life, half of the remaining radioactivity atoms decay into atoms of a<br />
new element.<br />
Half-Life Amount Remaining<br />
0 A^0 X (1/2)^0 = A^O<br />
1 A^0 X (1/2)^1 = A^0<br />
2 A^0 X (1/2)^2 = A^0<br />
<br />
<br />
Radiocarbon dating measures the amount of the isotope carbon-14 an object contains.<br />
Can only be used the measure the time line of a carbon-based material or life form.<br />
Transmutation can occur by radioactive decay or when particles attack the nucleus of<br />
an atom.
Learning Goal: The students will learn how molecular bonding is different than ionic bonding and<br />
electrons affect the shape of a molecule and its properties.<br />
Metallic or Ionic Bonds<br />
CCSTD HS <strong>Chemistry</strong> 2.a..<br />
A chemical bond is the “glue” that holds atoms together. Chemical bonds<br />
can be weak or strong, and everything in between!<br />
One thing that all chemical bonds share is that they are just an atom’s<br />
attempt to complete its outer (valence) shell of electrons. The closer an atom’s<br />
valence shell comes to being completely full, then the closer that atom comes to<br />
resemble an inert, or noble, gas.<br />
It turns out that all atoms strive to have the electronic configuration of an<br />
inert gas atom! Why is that such a desirable state? Inert gases are extremely stable.<br />
In fact, inert gases are so stable that they don’t bond, ever, to anything. They are<br />
perfectly content to remain just as they are.<br />
Ionic Bonds<br />
Another type of chemical bond, the ionic bond, involves a somewhat<br />
different arrangement of electrons within a molecule. Ionic bonds form when a<br />
“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<br />
takes an electron (to result in a negatively charged atom).<br />
You can just imagine what will happen if you put a positively charged atom<br />
(cation) next to a negatively charged atom (anion)! They are attracted to each other<br />
and we call this type of attraction an ionic bond.<br />
Think of an ionic bond as one valence shell cloud with a positive charge<br />
being attracted to another valence shell cloud with a negative charge. When those<br />
two clouds are next to each other, they “stick” together and the resultant molecule<br />
has no net charge.<br />
Table salt, or Sodium Chloride (NaCl), is a compound that has an ionic bond<br />
(see Figure 3.1.2). The Sodium within the molecule donates its electron and the<br />
Chloride accepts the electron. This condition gives the Sodium atom the electronic<br />
configuration of a Neon atom (an inert gas), and the Chloride atom the electronic<br />
configuration of an Argon atom (another inert gas!).<br />
` An ionic bond can involve the loss of more than one electron, with the other<br />
atom in the ionic compound gaining the same number of electrons. Magnesium<br />
Chloride (MgCl2) is an example of a doubly ionic compound. Once again,<br />
Magnesium Chloride has no net charge, and in this case, one Magnesium atom has<br />
lost two electrons (achieving the electronic configuration of Neon), with each of<br />
two Chloride atoms accepting one electron.<br />
Ionic bonds have distinctive properties of:<br />
When in solid form, all ionic compounds are crystalline solids at room temperature<br />
(the alternating positive and negative ions are arranged in a crystal lattice<br />
structure),
Strong ionic bonding forces making the solid structures hard with high melting and<br />
boiling points,<br />
Multiple charges within the ionic bond lead to stronger bonding forces (i.e. the<br />
bonding of MgCl2 is stronger than the bonding of NaCl), and<br />
Most ionic compounds are soluble in water (once they are fused (melted), or<br />
dissolved in water, ionic compounds conduct electricity because the anions are<br />
then free to move freely). Most often, ionic compounds are combinations of<br />
cations of atoms from the furthest left of the periodic (groups IA and IIA), with<br />
anions of atoms from the right of the periodic table (group VIIA), immediately<br />
preceding the inert gases.<br />
Metallic Bonds<br />
The metallic bond is actually a special case of an ionic bond. Think of metallic<br />
bonding as a more or less static arrangement of positive ions within a moving array<br />
of mobile electrons.<br />
This unique arrangement of cations and electrons give metals their characteristic<br />
properties. Metallic properties include.<br />
Heat conductivity: mobile electrons can carry the kinetic energy of heat.<br />
Shiny appearance: the rapidly moving electrons emit energy in the form of light.<br />
Electricity conductors: electricity is the flow of electrons.<br />
Malleability: ability to be easily shaped into flat sheets or drawn into wires.
More notes…<br />
Ionic Bonds<br />
Ionic bonds are the fusing together of metals and nonmetals<br />
Cross multiply the charges, and the charges become subscripts<br />
Names: Metal Nonmetal (NaCl), the nonmetal adds the suffix ‘ide’ for the full<br />
name (Sodium Chloride)<br />
The metal atoms give away their electrons to become a positively charged ion, and<br />
the nonmetal takes electrons to become a negatively charged ion.<br />
H 1 0<br />
H 2 0<br />
H 3 0<br />
H 4 0<br />
1-No pronunciation (Hydrogen Oxide)<br />
2-Di (Hydrogen Dioxide)<br />
3- Tri (Hydrogen Trioxide)<br />
4-Tetra (Hydrogen Tetraoxide)
Learning Goal: The students will learn how molecular bonding is different than ionic bonding and<br />
electrons affect the shape of a molecule and it's properties.<br />
Covalent Bond Notes<br />
When 2 nonmetals bond together, they share electrons<br />
Charges are not needed, because the mixtures are always stabilized.<br />
Sometimes they are called strong bonds because of the high difficulty of breaking these atoms<br />
apart aside from using special enzymes.<br />
Hydrogen and ionic bonds are easily broken by anything as small as temperature change.<br />
Covalent bonds form definite and easily predictable shapes.<br />
Another form of covalent bonding is polar bonding, where an atom shares atoms with<br />
hydrogen atoms. Water is an example of polar bonding.<br />
The Octet rule includes that all atoms excluding hydrogen and helium have to have 8 total<br />
electrons in their valence shell to become stable.<br />
This process can be achieved through many forms of bonding, with covalent being the most<br />
powerful.
A single bond is when two pair of electrons are shared between two atoms.<br />
A Double pair is when an atom shares two pairs of electrons to another atom<br />
A Triple pair is when an atom shares 3 pairs of electrons (total of 6) with another atom<br />
A polar covalent bond is a bond made, but is not equally shared because of the higher<br />
electronegativity of one of the atoms is higher than the other, which causes higher attraction<br />
with the electron<br />
A nonpolar covalent bond is bond made and is equally shared because of the equal<br />
electronegativity and therefore equally magnetic pull of the electron between the atoms
Name Formula Charge<br />
Dichromate Cr₂O₇ 2-<br />
Sulfate SO₄ 2-<br />
Hydrogen Carbonate HCO₃ 1-<br />
Hypochlorite ClO 1-<br />
Phosphate PO₄ 3-<br />
Nitrite NO₂ 1-<br />
Chlorite ClO₂ 1-<br />
Dihydrogen phosphate H₂PO₄ 1-<br />
Chromate CrO₄ 2-<br />
Carbonate CO₃ 2-<br />
Hydroxide OH 1-<br />
Hydrogen phosphate HPO₄ 2-<br />
Ammonium NH₄ 1+<br />
Acetate C₂H₃O₂ 1-<br />
Perchlorate ClO₄ 1-<br />
Permanganate MnO₄ 1-<br />
Chlorate ClO₃ 1-<br />
Hydrogen Sulfate HSO₄ 1-<br />
Phosphite PO₃ 3-<br />
Sulfite SO₃ 2-<br />
Silicate SiO₃ 2-<br />
Nitrate NO₃ 1-<br />
Hydrogen Sulfite HSO₃ 1-<br />
Oxalate C₂O₄ 2-<br />
Cyanide CN 1-<br />
Hydronium H₃O 1+<br />
Thiosulfate S₂O₃ 2-
Orbital Equation Lone Pairs Angle Name
Linear AX 2<br />
sp<br />
Bent AX 2 E<br />
Bent AX 2 E 2<br />
One pair of electrons Two pairs of electrons.<br />
sp 2 sp 3<br />
Tetrahedral AX 4<br />
Octahedral
Trig.plannar<br />
AX 3 sp 3 Trig. Pyramidal AX 3 E<br />
sp 3 d<br />
T-shaped AX 3 E 2<br />
sp 3 d<br />
Trig.Bipyramidal<br />
AX 5<br />
sp 3 d<br />
Square plannar<br />
AX 4 E 2 sp 3 d 2
www.youtube.com/watch?v=AsqEkF7hcII<br />
www.youtube.com/watch?v=tEn0N4R2dqA<br />
www.youtube.com/watch?v=Pft2CASl0M0<br />
www.youtube.com/watch?v=rwhJklbK8R0<br />
The Mole
www.youtube.com/watch?v=BTRm8PwcZ3U<br />
www.youtube.com/watch?v=F9NkYSKJifs<br />
www.youtube.com/watch?v=xPdqEX_WMjo<br />
Molar Mass
Mole Conversions
Steps for Mole Conversions
Answer the following questions:<br />
1) How many moles are in 25 grams of water?<br />
Mole Calculation Practice<br />
2) How many grams are in 4.5 moles of Li 2 O?<br />
3) How many molecules are in 23 moles of oxygen?<br />
4) How many moles are in 3.4 x 10 23 molecules of H 2 SO 4 ?<br />
5) How many molecules are in 25 grams of NH 3 ?<br />
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?<br />
8) How many molecules are there in 0.655 moles of C6H14?<br />
9) How many moles are there in 2.35 x 1024 molecules of water?<br />
10) How many grams does 5.60 x 1022 molecules of SiO2 weigh?<br />
11) How many molecules are there in 21.6 grams of CH4?<br />
1) 1.39 moles<br />
2) 134.1 grams<br />
3) 1.38 x 10 25 molecules<br />
4) 0.56 moles<br />
5) 8.85 x 10 23 molecules<br />
6) 106.7 grams<br />
7) How many grams does 0.500 moles of CuBr weigh? 31.8 grams<br />
8) How many molecules are there in 0.655 moles of C 6H 14? 3.94 x 10 23 molecules<br />
9) How many moles are there in 2.35 x 10 24 molecules of water? 3.90 moles<br />
10) How many grams does 5.60 x 10 22 molecules of SiO 2 weigh? 5.59 grams<br />
11) How many molecules are there in 21.6 grams of CH 4? 8.13 x 10 23
How to Balance Chemical Equations<br />
A chemical equation is a theoretical or written representation of what happens during a chemical<br />
reaction. The law of conservation of mass states that no atoms can be created or destroyed in a<br />
chemical reaction, so the number of atoms that are present in the reactants has to balance the<br />
number of atoms that are present in the products. Follow this guide to learn how to balance chemical<br />
equations.<br />
Step 1<br />
Write down your given equation. For this example, we will use:<br />
C 3 H 8 + O 2 --> H 2 O + CO 2<br />
Step 2<br />
Write down the number of atoms that you have on each side of the equation. Look at the subscripts<br />
next to each atom to find the number of atoms in the equation.<br />
Left side: 3 carbon, 8 hydrogen and 2 oxygen<br />
Right side: 1 carbon, 2 hydrogen and 3 oxygen
Step 3<br />
Always leave hydrogen and oxygen for last. This means that you will need to balance the carbon<br />
atoms first.<br />
Step 4<br />
Add a coefficient to the single carbon atom on the right of the equation to balance it with the 3 carbon<br />
atoms on the left of the equation.<br />
C 3 H 8 + O 2 --> H 2 O + 3CO 2<br />
The coefficient 3 in front of carbon on the right side indicates 3 carbon atoms just as the subscript 3<br />
on the left side indicates 3 carbon atoms.<br />
In a chemical equation, you can change coefficients, but you should never alter the subscripts.
Step 5<br />
Balance the hydrogen atoms next. You have 8 on the left side, so you'll need 8 on the right side.<br />
C 3 H 8 + O 2 --> 4H 2 O + 3CO 2<br />
On the right side, we added a 4 as the coefficient because the subscript showed that we already<br />
had 2 hydrogen atoms.<br />
When you multiply the coefficient 4 times the subscript 2, you end up with 8.<br />
Step 6<br />
Finish by balancing the oxygen atoms.<br />
Because we've added coefficients to the molecules on the right side of the equation, the number of<br />
oxygen atoms has changed. We now have 4 oxygen atoms in the water molecule and 6 oxygen<br />
atoms in the carbon dioxide molecule. That makes a total of 10 oxygen atoms.<br />
Add a coefficient of 5 to the oxygen molecule on the left side of the equation. You now have 10<br />
oxygen molecules on each side.<br />
C 3 H 8 + 5O 2 --> 4H 2 O + 3CO 2.<br />
The carbon, hydrogen and oxygen atoms are balanced. Your equation is complete.
1) ___ NaNO 3 + ___ PbO ___ Pb(NO 3 ) 2 + ___ Na 2 O<br />
2) ___ AgI + ___ Fe 2 (CO 3 ) 3 ___ FeI 3 + ___ Ag 2 CO 3<br />
3) ___ C 2 H 4 O 2 + ___ O 2 ___ CO 2 + ___ H 2 O<br />
4) ___ ZnSO 4 + ___ Li 2 CO 3 ___ ZnCO 3 + ___ Li 2 SO 4<br />
5) ___ V 2 O 5 + ___ CaS ___ CaO + ___ V 2 S 5
6) ___ Mn(NO 2 ) 2 + ___ BeCl 2 ___ Be(NO 2 ) 2 + ___ MnCl 2<br />
7) ___ AgBr + ___ GaPO 4 ___ Ag 3 PO 4 + ___ GaBr 3<br />
8) ___ H 2 SO 4 + ___ B(OH) 3 __ B 2 (SO 4 ) 3 + ___ H 2 O<br />
9) ___ S 8 + ___ O 2 ___ SO 2<br />
10) ___ Fe + ___ AgNO 3 ___ Fe(NO 3 ) 2 + ___ Ag
1) 2 NaNO 3 + PbO Pb(NO 3 ) 2 + Na 2 O<br />
2) 6 AgI + Fe 2 (CO 3 ) 3 2 FeI 3 + 3 Ag 2 CO 3<br />
3) C 2 H 4 O 2 + 2 O 2 2 CO 2 + 2 H 2 O<br />
4) ZnSO 4 + Li 2 CO 3 ZnCO 3 + Li 2 SO 4<br />
5) V 2 O 5 + 5 CaS 5 CaO + V 2 S 5<br />
6) Mn(NO 2 ) 2 + BeCl 2 Be(NO 2 ) 2 + MnCl 2<br />
7) 3 AgBr + GaPO 4 Ag 3 PO 4 + GaBr 3<br />
8) 3 H 2 SO 4 + 2 B(OH) 3 B 2 (SO 4 ) 3 + 6 H 2 O<br />
9) S 8 + 8 O 2 8 SO 2<br />
10) Fe + 2 AgNO 3 Fe(NO 3 ) 2 + 2 Ag<br />
Additional Notes:
Categories of Reactions<br />
All chemical reactions can be placed into one of six categories. Here they are, in no<br />
particular order:<br />
1) Synthesis: A synthesis reaction is when two or more simple compounds combine to form a<br />
more complicated one. These reactions come in the general form of:<br />
A + B ---> AB<br />
One example of a synthesis reaction is the combination of iron and sulfur to form iron (II) sulfide:<br />
8 Fe + S 8 ---> 8 FeS<br />
If two elements or very simple molecules combine with each other, it’s probably a synthesis reaction.<br />
The products will probably be predictable using the octet rule to find charges.<br />
2) Decomposition: A decomposition reaction is the opposite of a synthesis reaction - a<br />
complex molecule breaks down to make simpler ones. These reactions come in the general form:<br />
AB ---> A + B<br />
One example of a decomposition reaction is the electrolysis of water to make oxygen and hydrogen<br />
gas:<br />
2 H 2O ---> 2 H 2 + O 2<br />
If one compound has an arrow coming off of it, it’s probably a decomposition reaction. The products<br />
will either be a couple of very simple molecules, or some elements, or both.<br />
3) Single displacement: This is when one element trades places with another element in a<br />
compound. These reactions come in the general form of:<br />
A + BC ---> AC + B<br />
One example of a single displacement reaction is when magnesium replaces hydrogen in water to<br />
make magnesium hydroxide and hydrogen gas:<br />
Mg + 2 H 2O ---> Mg(OH) 2 + H 2<br />
If a pure element reacts with another compound (usually, but not always, ionic), it’s probably a single<br />
displacement reaction. The products will be the compounds formed when the pure element switches<br />
places with another element in the other compound.<br />
Important note: these reactions will only occur if the pure element on the reactant side of the equation<br />
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<br />
switch places, forming two entirely different compounds. These reactions are in the general form:<br />
AB + CD ---> AD + CB<br />
One example of a double displacement reaction is the reaction of lead (II) nitrate with potassium<br />
iodide to form lead (II) iodide and potassium nitrate:<br />
Pb(NO 3) 2 + 2 KI ---> PbI 2 + 2 KNO 3<br />
If two ionic compounds combine, it’s probably a double displacement reaction. Switch the cations<br />
and balance out the charges to figure out what will be made.<br />
Important note: These reactions will only occur if both reactants are soluble in water and only one<br />
product is soluble in water.<br />
5) Acid-base: This is a special kind of double displacement reaction that takes place when an<br />
acid and base react with each other. The H + ion in the acid reacts with the OH - ion in the base,<br />
causing the formation of water. Generally, the product of this reaction is some ionic salt and water:<br />
HA + BOH ---> H 2O + BA<br />
One example of an acid-base reaction is the reaction of hydrobromic acid (HBr) with sodium<br />
hydroxide:<br />
HBr + NaOH ---> NaBr + H 2O<br />
If an acid and a base combine, it’s an acid-base reaction. The products will be an ionic compound<br />
and water.<br />
6) Combustion: A combustion reaction is when oxygen combines with another compound to<br />
form water and carbon dioxide. These reactions are exothermic, meaning they produce heat. An<br />
example of this kind of reaction is the burning of napthalene:<br />
C 10H 8 + 12 O 2 ---> 10 CO 2 + 4 H 2O<br />
If something that has carbon and hydrogen reacts with oxygen, it’s probably a combustion reaction.<br />
The products will be CO 2 and H 2 O.<br />
Follow this series of questions. When you can answer "yes" to a question, then<br />
stop!<br />
1) Does your reaction have two (or more) chemicals combining to form one chemical? If yes, then it's<br />
a synthesis reaction<br />
2) Does your reaction have one large molecule falling apart to make several small ones? If yes, then<br />
it's a decomposition reaction<br />
3) Does your reaction have any molecules that contain only one element? If yes, then it's a single<br />
displacement reaction<br />
4) Does your reaction have water as one of the products? If yes, then it's an acid-base reaction<br />
5) Does your reaction have oxygen as one of it's reactants and carbon dioxide and water as<br />
products? If yes, then it's a combustion reaction<br />
6) If you haven't answered "yes" to any of the questions above, then you've got a double<br />
displacement reaction.
1) double displacement<br />
2) combustion<br />
3) single displacement<br />
4) double displacement<br />
5) acid-base<br />
6) synthesis<br />
7) decomposition<br />
List what type the following reactions are: (answers)<br />
Predicting Reaction Products – Answers<br />
Predict the products of each of the following chemical reactions. If a reaction will not occur, explain why not:<br />
1) ____ Ag 2 SO 4 + ____ NaNO 3 → no reaction!<br />
Examining this reaction, it appears that a double displacement reaction will occur. This would lead to the conclusion that<br />
the products would be AgNO3 and Na2SO4. However, for this reaction to occur, both reactants and only one of the<br />
products must be soluble in water. If you look up the solubilities on a chart, you’ll find that Ag2SO3 is partly soluble in<br />
water, and all of the other compounds are totally soluble in water. This tells us that this reaction will not occur.<br />
2) ____ NaI + ____ CaSO 4 → no reaction!<br />
Another double displacement reaction, this time with Na2SO4 and CaI2 as products. Because both products are soluble<br />
in water and CaSO4 is only partially soluble in water, the conditions for a successful double displacement reaction are not<br />
met.<br />
3) 2 HNO 3 + 1 Ca(OH) 2 → 1 Ca(NO 3 ) 2 + 2 H 2 O<br />
It’s an acid-base reaction, and acid-base reactions occur readily whether or not the reactants are both soluble in water.<br />
4) 1 CaCO 3 → 1 CaO + 1 CO 2<br />
It’s a decomposition reaction. If you didn’t guess that these were the products, you should have at least known that it was<br />
a decomposition reaction and predicted that this would have broken into its constituent elements, Ca, C, and O2.<br />
5) 1 AlCl 3 (aq) + 1 (NH 4 ) 3 PO4(aq) → AlPO 4 (s) + 3 NH 4 Cl(aq)<br />
This is a double displacement reaction, except in this case both of the reactants and only one product are soluble in<br />
water. Because the conditions for a successful reaction are met, the reaction does occur!<br />
6) ____ Pb + ____ Fe(NO 3 ) 3 →<br />
Though this is a single displacement reaction, lead is lower on the activity series than the iron it would replace. As a<br />
result, this reaction does not occur.<br />
7) 2 C 3 H 6 + 9 O 2 → 6 CO 2 + 6 H 2 O<br />
The reactants suggest that this is a combustion reaction, meaning that the products must be carbon dioxide and water.<br />
Once you figure this out, the only thing left to do is balance it, as shown.<br />
8) 2 Na + 1 CaSO 4 → 1 Na 2 SO 4 + 1 Ca<br />
This should clearly be a single displacement reaction. Because sodium is higher on the activity series than calcium, this<br />
reaction does occur.
Balancing Equations Practice Problems<br />
From Widener <strong>Chemistry</strong> Practice<br />
1 page
Stoichiometry Practice Problems<br />
From Widener <strong>Chemistry</strong> Practice<br />
2 pages
Percent Yield Practice Problems<br />
from Widener <strong>Chemistry</strong> Practice<br />
1 pages
Cody Nichols<br />
Between the particles in gas, it is empty; this allows for more<br />
movement and kinetic energy that is caused by the particles hitting<br />
each other<br />
Because of no attractive and repulsive force between them, it<br />
allows for a balance state of gas. Due to this, the particles stay in a<br />
random walk .<br />
The collisions that occur during a random walk are perfectly<br />
elastic and the kinetic energy is transferred without any lose.<br />
Gases are very small particles, but humans<br />
have learned to harness the thousands<br />
of types of gases that exists,<br />
which aid them in everyday life, like<br />
propane, nitrogen, co 2, and other gases.
The particles in liquid have kinetic<br />
energy , this makes the particles<br />
flow past each other. The flow of theses particles is<br />
referred to as a fluid.<br />
Unlike gas though, liquid particles are atattracted<br />
to each other, depending on weather the<br />
attraction is disrupted determines the physical<br />
Liquid evaporates and then turns into a vapor<br />
or a gas, then it condensates and turns back<br />
into a liquid. The vapor pressure or the force exerted<br />
by a gas above a liquid is caused by the collision<br />
of particles in a container that contains evaporated<br />
liquid.<br />
Water is the most abundant<br />
source of liquid on<br />
earth. H 2 O can take the<br />
form of all three states of<br />
matter( solid, liquid, and<br />
gases). Examples of this<br />
are Ice as a solid, regular<br />
water as a liquid , and a<br />
rain cloud.
The particles in a solid are packed tightly together, they are dense<br />
and not easy to compress.<br />
When heating a solid, the particles start to break down and at<br />
freezing point, the solid is stuck in a fixed position.<br />
There are crystal, allotropes, and non-crystalline solids, which either<br />
have simple cubic unit cells, body centered unit cells or face<br />
centered unit cells.
Changes from a solid to a gas without going through liquid is called<br />
sublimation. Solid goes through sublimation to a vapor, then the vapor goes<br />
to deposition, back to a solid. A phase diagram shows the three states and<br />
what points that they will change at, the only point where they are equal is<br />
the triple point.<br />
Learning goal<br />
The students will learn what are the factors that determine and characteristics that distinguish<br />
gases liquids and solids and how substances change from one state to another.
Boyles law: V1 + P2 = P 2 + V2<br />
rearrange the equation to isolate V2:<br />
substitute the known values for P1 V1 and P2 into the equation and solve<br />
The finished equation is<br />
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<br />
directly proportional to its Kelvin<br />
temperature if the pressure is<br />
kept constant.<br />
As temperature increases, so does<br />
the volume of the substance being<br />
exposed to the heat.
V1<br />
T1 = V2<br />
T2<br />
Cross multiply the<br />
numbers<br />
11.<br />
6.80 L<br />
325 C = X<br />
25 C = 6.80<br />
L* 25 C/ 325 C = 170<br />
L/325=0.52 L
http://www.youtube.com/watch?v=7JKVtbe-hV8
Pressure of a gas<br />
corresponds to the kelvin<br />
temperature if the volume<br />
remains constant.
you put the numbers<br />
where they need to be<br />
108/41 = / 22Cross<br />
multiple 2376/41 then<br />
divide=57.9512195,<br />
then you use the least<br />
significant figures, so<br />
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<br />
Direct: 6/1=6 12/2=6 P1 / T1 = P2 / T2<br />
Combined Gas Law<br />
P1 x V1/ T1 = P2 x V2/ T2<br />
V1 = 3110 mL T1 = x V2 = 651 mL T2 = 80.60 K Indirect equation<br />
3110 / x = 651 / 80.6 3110 * 80.6 / 651 = x 250, 666 / 651 x=385 K<br />
V1 = 2600 mL P1 = x V2 = 6570 mL P2 = 1500 torr Direct equation<br />
2600x = 6570 * 1500 2600x = 985500 x = 3790 torr 3790 torr = 4.965 atm x = 5.0 atm<br />
P1 = 1320 torr V1 = x T1 = 615 C P2 = 1.28 atm V1 = 7.11 L T2 = 905 K CGL equation<br />
Conversions: T1615 C = 888 K P11320 torr = 1.7292 atm<br />
1.7292*x/888 = 1.28*7.11/9051.7292x/888=0.010051.7292x=8.9244x=5.16 L<br />
P=Pressure V=Volume T=Tempurature n=# in mol r=8.31kpq OR 0.082 atm<br />
Ideal Gas Law PxV=nrT P=nrT/V<br />
P = 1470 torr V = 7.87 L T = 712 K n = x r = 0.082 atm IGL equation<br />
Conversions: P1470 torr = 195.51<br />
195.51 = 58.384x/7.87 1538.6637=58.384x 26.354=x x = 26.4<br />
P = 1.36 atm V = 8.53 L T = 499 K n = x r = 0.082 atm<br />
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<br />
V = Volume must be in same units<br />
T = Temperature must be in Kelvin<br />
Direct relationship = as the temperature increases, the volume increases<br />
V1/T1 = V2/T2<br />
P1/T1 = P2/T2<br />
Indirect relationship = as volume increases, pressure increases<br />
V1 x P1 = V2 x P2<br />
Combined Gas Law<br />
P = Pressure in atm or kPa<br />
V = Volume must be in Liters<br />
N = number of moles<br />
R = Constant either 0.082 atm OR 8.31 kPa, depending on pressure<br />
T = Tempurature in Kelvin<br />
PV = NRT<br />
Ideal Gas Law<br />
P1 = 2.27 atm V1 = 7.94 L P2 =? V2 = 10.8 L<br />
2.27 x 7.94/ 10.8 = x x = 1.67 atm 1.67 atm = 1270 torr P2 = 1270 torr<br />
V1 = 12.6 L T1 = 383 K V2 = ? T2 = 45.3 K<br />
12.6/383=x/45.3 1.49 = x V2 = 1.5 L<br />
V1 = 2.83L T1 = 832K P1 = ? V2 = 3.85L T2 = 411K P2 = 0.396 atm<br />
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<br />
P = 1.6 atm N = 0.433 mol T = 346 K R = 0.082 atm V = ?<br />
1.6*x = 0.433*346*0.082 x= 7.678 7.678 L = 7678 mL V = 7680 mL<br />
P1 = 0.494 atm V1 = 2230 mL T1 = ? P2 = 0.150 atm V2 = 3590 mL T2 = 279 K<br />
0.494*2230/ x = 0.15*3590 / 279 x=0.00175 T1=0.002 K
Unit 5 & 6 Test Review<br />
Heat of Fusion<br />
q = m ∙ΔH f<br />
Specific Heat<br />
c = ____q____<br />
m ∙ ΔT ᵒC<br />
Heat of Fusion of Water<br />
334 J/g = 80 cal/g<br />
Specific Heat of Water<br />
4.18 J/g ∙ᵒC<br />
Heat of fusion is the amount of heat energy required to change the state of a substance from solid to<br />
liquid. This example problem demonstrates how to calculate the amount of energy required to melt a<br />
sample of water ice.<br />
With Graham's Law, you can find the effusion rates for two gases or the molecular mass of a gas.<br />
This ratio of effusion rates follows the pattern that the gas with the lesser molecular mass has a<br />
greater rate of effusion.<br />
Ideal Gas Law<br />
Combined Gas Law<br />
p∙v=n∙R∙T V 1 ∙P 1 = V 2 ∙P 2<br />
T 1 T 2<br />
R = .082 atm or 8.31kpa<br />
Boyle’s Law<br />
P 1 ∙V 1 = P 2 ∙V 2<br />
Charles’ Law<br />
Gay – Lussac’s Law<br />
V 1 = V 2 P 1 = P 2<br />
T 1 T 2 T 1 T 2<br />
Molarity<br />
M = moles<br />
L<br />
Gas Solubility<br />
S 1 = S 2<br />
P 1 P 2
Unit 5<br />
Chapter 13 States of Matter<br />
Chapter 14 The Behavior of Gases<br />
Chapter 15 Water and Aqueous Systems<br />
Unit 6<br />
Chapter 16 Solutions<br />
Chapter 17 Thermochemistry<br />
Chapter 18 Reaction Rates and Equilibrium<br />
Chapter 19 Acid and Bases<br />
Learning Goals<br />
The students will learn what are the factors that determine and characteristics that distinguish gases<br />
liquids and solids and how substances change from one state to another.<br />
The students will learn how gases respond to changes in pressure, volume, and temperature and why<br />
the ideal gas law is useful even though ideal gases do not exist.<br />
The students will learn how the interactions between water molecules account for the unique<br />
properties of water and how aqueous solutions form.<br />
The student will learn how energy is converted in a chemical or physical process and how to<br />
determine the amount of energy is absorbed or released in that process.
#1<br />
What is the heat in Joules required to melt 25 grams of ice? What is the heat in calories?<br />
#2<br />
How much energy is released when 20 g of water is frozen at 0ᵒ C?
#3<br />
A sample of water of unknown mass undergoes a temperature change from 58.85ᵒC to 16.06 ᵒC with<br />
a concomitant change in heat content of -123200 Joules. What is the mass of the water?<br />
#4<br />
A sample of water with a mass of 597.9g and an unknown temperature loses 18890 Joules. If the<br />
final temperature is found to be 35.76ᵒC .What was the initial temperature?
#5<br />
A gas system has volume, moles and temperature of 9.76L, 0.264moles and 308.0C, respectively.<br />
What is the pressure in atm?<br />
#6<br />
A gas system has pressure, volume and temperature of 1.43atm, 6.86L and 394K, respectively. How<br />
many moles of gas are present?
#7<br />
A gas system has pressure, volume and moles of 1.44atm, 8.16L and 0.427moles, respectively. What<br />
is the temperature in K?<br />
#8<br />
What is the molarity of a solution of sodium bromide which contains 64.1 grams of solute in a total<br />
volume of 0.436L?
#9<br />
What volume in mL of a 0.942M solution of sodium bromide would be required if you wanted<br />
0.692moles of solute?<br />
#10<br />
What is the percent by mass of water in CuSO 4 ∙ 5H 2 0?