Chemistry I Booklet

Fall Semester

Class: Ms. Manakul

Student Name: ____________________

Student Hour: _________

Periodic Table of the Elements

hydrogen

1

H 1.0079

helium

2

He 4.0026

lithium

3

Li 6.941

beryllium

4

Be 9.0122

boron

5

B 10.81

carbon

6

C 12.011

nitrogen

7

N 14.007

oxygen

8

O 15.999

fluorine

9

F 18.998

Neon

10

Ne 20.179

sodium

11

Na 22.990

magnesium

12

Mg 24.305

aluminium

13

Al 26.982

silicon

14

Si 28.086

phosphorus

15

P 30.974

sulphur

16

S 32.06

chlorine

17

Cl 35.453

argon

18

Ar 39.984

potassium

19

K 39.098

calcium

20

Ca 40.08

scandium

21

Sc 44.956

titanium

22

Ti 47.90

vanadium

23

V 50.941

chromium

24

Cr 51.996

manganese

25

Mn 54.938

iron

26

Fe 55.847

cobalt

27

Co 58.933

nickel

28

Ni 58.71

copper

29

Cu 63.546

zinc

30

Zn 65.38

gallium

31

Ga 69.72

germanium

32

Ge 72.59

arsenic

33

As 74.922

selenium

34

Se 78.96

bromine

35

Br 79.904

krypton

36

Kr 83.80

rubidium

37

Rb 85.468

strontium

38

Sr 87.62

yttrium

39

Y 88.906

zirconium

40

Zr 91.22

niobium

41

Nb 92.906

molybdenum

42

Mo 95.94

technetium

43

Tc [98]

ruthenium

44

Ru 101.07

rhodium

45

Rh 102.91

palladium

46

Pd 106.4

silver

47

Ag 107.87

cadmium

48

Cd 112.41

indium

49

In 114.82

tin

50

Sn 118.69

antimony

51

Sb 121.75

tellurium

52

Te 127.60

iodine

53

I 126.90

xenon

54

Xe 131.30

caesium

55

Cs 132.91

barium

56

Ba 137.33

lutetium

71

Lu 174.97

hafnium

72

Hf 178.49

tantalum

73

Ta 180.95

tungsten

74

W 183.85

rhenium

75

Re 186.21

osmium

76

Os 190.2

iridium

77

Ir 192.22

platinum

78

Pt 195.09

gold

79

Au 196.97

mercury

80

Hg 200.59

thallium

81

Tl 204.37

lead

82

Pb 207.2

bismuth

83

Bi 208.98

polonium

84

Po [209]

astatine

85

At [210]

radon

86

Rn [222]

francium

87

Fr [223]

radium

88

Ra [226]

lawrencium

103

Lr [262]

rutherfordium

104

Rf [261]

dubnium

105

Db [262]

seaborgium

106

Sg [263]

bohrium

107

Bh [264]

hassium

108

Hs [265]

meitnerium

109

Mt [268]

darmstadtium

110

Ds [269]

roentgenium

111

Rg [272]

copernicium

112

Cn [277]

ununtrium

113

Uut [284]

flerovium

114

Fl [289]

ununpetium

115

*Uup [288]

livermorium

116

Lv [293]

ununseptium

117

*Uus [294]

ununoctium

118

*Uuo [299]

*Discovery reported by not verified

lanthanum

57

La 138.91

cerium

58

Ce 140.12

praseodymium

59

Pr 140.91

neodymium

60

Nd 144.24

promethium

61

Pm [145]

samarium

62

Sm 150.4

europium

63

Eu 151.96

gadolinium

64

Gd 157.25

terbium

65

Tb 158.93

dysprosium

66

Dy 162.50

holmium

67

Ho 164.93

erbium

68

Er 167.26

thulium

69

Tm 168.93

ytterbium

70

Yb 173.04

actinium

89

Ac [227]

thorium

90

Th 232.04

protactinium

91

Pa 231.04

uranium

92

U 238.03

neptunium

93

Np [237]

plutonium

94

Pu [244]

americium

95

Am [243]

curium

96

Cm [247]

berkelium

97

Bk [247]

californium

98

Cf [251]

einsteinium

99

Es [252]

fermium

100

Fm [257]

mendelevium

101

Md [258]

nobelium

102

No [259]

H 2.2

Electron Affinity Table He ---

Li 1.0

Be 1.6

B 2.0

C 2.6

N 3.0

O 3.4

F 4.0

Ne ---

Na 0.9

Mg 1.3

Al 1.6

Si 1.9

P 2.2

S 2.6

Cl 3.2

Ar ---

K 0.8

Ca 1.0

Sc 1.4

Ti 1.5

V 1.6

Cr 1.7

Mn 1.6

Fe 1.8

Co 1.9

Ni 1.9

Cu 1.9

Zn 1.7

Ga 1.8

Ge 2.0

As 2.2

Se 2.6

Br 3.0

Kr ---

Rb 0.8

Sr 1.0

Y 1.2

Zr 1.3

Nb 1.6

Mo 2.2

Tc 2.1

Ru 2.2

Rh 2.3

Pd 2.2

Ag 1.9

Cd 1.7

In 1.8

Sn 2.0

Sb 2.1

Te 2.1

I 2.7

Xe ---

Cs 0.8

Ba 0.9

La 1.1

Hf 1.3

Ta 1.5

W 1.7

Re 1.9

Os 2.2

Ir 2.2

Pt 2.2

Au 2.4

Hg 1.9

Tl 1.8

Pb 1.8

Bi 1.9

Po 2.0

At 2.2

Rn ---

Fr 0.7

Ra 0.9

Ac 1.1

General Physical Constants

Atomic mass unit 1 amu = 1.6605 x 10-24 g

Avogadro’s Number N = 6.02 x 1023 particles/mol

Gas constant R = 8.31 L kPa K-1 mol-1

Ideal gas molar volume Vm = 22.4 L mol-1

Masses of subatomic particles

Electron (e-)

Proton (p+)

Neutron (n0)

me = 0.0005486 amu = 9.1096 x 10-28 g

mp = 1.007277 amu = 1.67261 x 10-24 g

mn = 1.008665 amu = 1.67492 x 10-24 g

Speed of Light (in a vacuum) c = 3.00 x 108 m s-1

Planck’s Constant h = 6.626 x 10-34 J s

Solubility Rules

Negative Ion Rule

NO3- All compounds formed with the negative ion are soluble

I-, Br-, Cl- All compounds formed with the negative ion are soluble except Ag+, Pb2+, Hg2+, and Cu+

SO42-

Most compounds formed with the negative ion are soluble; exceptions include SrSO4, BaSO4, CaSO4, RaSO4, Ag2SO4, and PbSO4

CO32-, PO4

3-, SO3

2-

All compounds formed with the negative ion are insoluble except those of the alkali metals and NH4

+

OH- All compounds formed with the negative ion are insoluble except those of the alkali metals, NH4

+, Sr2+, and Ba2+. (Ca(OH)2 is slightly soluble)

S2- All compounds formed with the negative ion are insoluble except those of the alkali metals, alkaline earth metals, and NH4

+.

Reactivity Series

Metals Halogens

Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Tin Lead Hydrogen Copper Mercury Silver Platinum Gold

Fluorine Chlorine Bromine Iodine

SI Units and Equivalents

Quantity SI Unit Common Equivalents

Length Meter (m) 1 meter = 1.0936 yards 1 centimeter = 0.39370 inch

1 inch = 2.54 centimeters 1 mile = 5280 feet = 1.6093 kilometers

Volume Cubic meter (m3)

1 liter = 10-3 m3 = 1.0567 quarts 1 gallon = 4 quarts = 8 pints = 3.7852 liters

1 quart = 32 fluid ounces = 0.94635 liter

Temperature

Kelvin (K) °C = 5/9 (F -32) K = °C + 273.15

Mass Kilogram (kg)

1 kilogram = 1000 grams = mass weighing 2.2046 pounds

1 amu = 1.66057 x 10-27 kilograms

Time Second (s) 1 hour = 60 minutes = 3600 seconds

Energy Joule (J) 1 joule = 1 kg m2 s-2 (exact) = 0.23901 calorie 1 calorie = 4.184 joules

Pressure Pascal (Pa) 1 atmosphere = 101.3 kilopascals = 760 mm Hg (torr) 1 atmosphere = 14.70 pounds per square inch

Table of Contents

I. Syllabus page i

II. Classroom Map page iv

III. School Safety Drills page v

IV. Lab Safety Contract page vi

V. Unit Objectives page viii

VI. Objective Graphs page xi

VII. Homework

Unit 1: Measurement page 1

Unit 2: Matter & the Atom page 4

Unit 3: Electromagnetic Spectrum page 9

Unit 4: Metals & Ionic Bonding page 14

Unit 6: Covalent Bonding page 19

Unit 7: Moles & Chemical Reactions page 22

VIII. Labs

Significant Figures Lab page 31

Density Lab page 33

Beanium Lab page 36

Spectroscopy Spectrum Lab page 38

Flame Test Lab page 40

Paper Chromatography of Food Dyes page 42

Precipitation Lab page 45

IX. Appendix

Flowchart for Naming Chemical Compounds page A

Flowchart for Writing Formulas for Chemical Compounds page B

VSEPR Sheet page C

Mole Map page D

i

Chemistry I Syllabus

Lincoln College Preparatory Academy

Teacher: Melissa Manakul E-mail:mmanakul@kcpublicschools.org

Room: 207 Google Classroom: xbzb8sl

*Any communication to me either through e-mail or Google Classroom will be responded to as soon as humanly possible.

MY WEEKLY SCHEDULE:

Day Monday Tuesday Wednesday Thursday Friday

Morning In classroom by

6:50 am

In classroom by

6:50 am

In classroom by

6:50 am

In classroom by

6:50 am

In classroom by

6:50 am

Afternoon Meeting Tutoring

2:30 – 3:30 Staff Meeting

Will leave by

3:00 pm

Will leave by

2:30 pm

COURSE DESCRIPTION:

Chemistry is an experimental science that combines academic study with the acquisition of practical and investigational

skills. It is called the central science, as chemical principles underpin both the physical environment in which we live and

all biological systems. Apart from being a subject worthy of study in its own right, chemistry is a prerequisite for many

other courses in higher education, such as medicine, biological science and environmental science, and serves as a useful

preparation for employment.

GOALS FOR SEMSTER:

1. You will be prepared for IB Chemistry II, AP Chemistry, or collegiate General Chemistry course.

2. You will be able to raise your ACT score 3 points over the course of the year.

REQUIRED MATERIALS:

Booklet: This is your bible in Chemistry. Each student receives a booklet at the beginning of the semester, which includes

all unit objectives, homework, labs, and supplemental material. This is to go home every night.

Lab Book: This will be a composition notebook (college ruled or grid paper preferred) which will house all your

completed labs and activities. This will be graded throughout the year, and left in the classroom.

Laptop: This will be used to access the textbook and Google Classroom where any announcements, calendar changes,

copies of activities, additional supplemental information, and extra copies of homework. We will also use it for a

variety of activities to be done at home or in the classroom through a variety of educational websites.

SUGGESTED MATERIALS:

Binder/Folder: This will hold any paper work, including but not limited to tests, quizzes, lab reports, and activities. You

must take this home every night. Any time Tyler or my record is incorrect, it is the student’s job to report this to

Ms. Manakul with a graded copy. This is the only way a grade can be changed for full credit.

Scientific Calculator: TI-30x is a preferred calculator, however any scientific calculator will do.

TOPICS COVERED:

(1) Scientific Method (2) Matter & Atomic Structure (3) Electromagnetic Spectrum (4) Ions & Ionic Bonding (5)

Covalent Bonding (6) Moles & Chemical Reactions (7) Stoichiometry (8) Gases (9) Solutions (10)Acids & Bases

(11)Thermochemistry

METHOD OF INSTRUCTION:

We will be moving through material at a steady pace. We will be working through an objective in about a week. There

will be at least 1 quiz per unit and 1 test per unit. There will be around 1 to 2 major labs per unit and activities/discussions

scattered throughout the unit. Participation in labs and activities is a requirement and are done to ensure you properly

comprehend material. You will have homework most nights. All work must be completed for the next class period

unless specified otherwise.

ii

DO NOWS/EXIT TICKETS:

Do Now’s & Exit Tickets are Ms. Manakul’s way to see who does or does not get what is happening in class. Ms.

Manakul also believes any work that you do should be worth something. Thus, you will turn your Do Now/Exit Ticket in

to Ms. Manakul BEFORE she goes over the answer. If you answer the question correctly your Do Now/Exit Ticket sheet

will go into the class piggy bank. If you would like to save paper, and show the class that “you got this,” you may go over

the answer in Ms. Manakul’s place. If you choose this route, your name will automatically go into the class piggy bank.

Before each quiz Ms. Manakul will pull out two random Do Now/Exit Ticket Sheets and those two people will gain

automatic extra credit towards the quiz. Before a test Ms. Manakul will pull out four sheets. If you are absent, you may

see Ms. Manakul before or after school (not during school) to make up your Do Now so you do not miss out on any of the

extra credit drawing entries.

HOMEWORK:

As mentioned above, homework will be assigned at the end of class and will be due in class the following school day.

Homework is a place to better comprehend material, gain proper practice, and solidify knowledge. However, sometimes

homework is confusing or you did not learn material the first time. Should this happen you should ALWAYS ASK

QUESTIONS either with a classmate or Ms. Manakul. Because we all struggle with homework, Ms. Manakul takes

homework as a completion grade. You must have tried the problem, i.e. you do not necessarily need to complete the

problem. Be aware an IDK is not trying and will thus your homework to be considered incomplete. After the class Do

Now, Ms. Manakul will check homework and go through the homework to ensure everyone comprehends the material.

Remember if you didn’t get it the first time ALWAYS ASK QUESTIONS!

LAB WORK:

Ms. Manakul will go over each lab the day before. You are to prepare your lab book by having your lab book filled out at

home based on the rubric. To prepare you for designing your own lab, each lab is designed to be very similar to your

typed lab report. We will discuss the specifics of the designed lab when we get there. For your safety you will be wearing

lab coats, goggles, and closed toed shoes in the lab room for the majority of labs. If you wear open toed shoes during a

lab, you will wear gloves over your shoes to protect your feet. If this is to change, Ms. Manakul will verbally state it. Ms.

Manakul has high expectations for lab, thus, you will gather, set up, and put away your own lab materials. Furthermore,

YOU MUST BE SAFE at all times during a lab. Ms. Manakul will give ONLY ONE verbal warning. If she must give

another, you will be ejected from lab, and receive an automatic zero for the lab.

LATE WORK:

Ms. Manakul dislikes late work and generally cannot spare the time to regrade items. Thus, any late work will receive

30% off. This means by the 3rd class day, your work will receive an automatic zero. This is to ensure work is turned in a

timely manner, and Ms. Manakul has adequate time to grade your work.

QUIZZES/TESTS:

For every quiz/test you will be given a clean copy of the periodic table. For unit tests only, you will be allowed to have

your own hand written 3” by 5” one sided notecard with any notes you need on them. At the end of the semester you may

take a cumulative test. This will be determined by Ms. Manakul and announced in class when necessary. If you are caught

cheating or talking during a test, Ms. Manakul will take your test and give you an automatic zero. All tests and quizzes are

graded as fast as humanly possible and will be placed in Tyler at the same rate. Any changes will be announced verbally.

ABSENT:

Lesson plans for the week are posted in the classroom. A large calendar is also posted which will label future events as

well. You can also check chemistrywithmanakul.weebly.com for any updates or changes. If you are absent, it is your

responsibility to make up the lecture notes on your own time. Should you need assistance, you may see Ms. Manakul

before or after school (reference My Weekly Schedule). Remember to turn in work as soon as possible. Any work

assigned during your absence will have a 2 school day grace period. If you are absent on the day an assignment is due, it

will be due the first day you return. If you miss a lab or activity, you are to collect the data from a friend, and finish the

rest of your lab report on your own time. *If Ms. Manakul is unexpectedly absent, you are to refer to the lesson plans

posted and watch online for specific details of any work to be completed in class. If an activity or lab is scheduled for that

day, it will be rescheduled. Ms. Manakul will recap all information and class activities/labs will be done with the class

when she returns.

iii

CLASS RULES:

1. Pay close attention to class, and be prepared to work hard.

2. 10/10 Rule. No passes will be given during the first 10 minutes and the last 10 minutes of class. Restroom passes

only in extreme emergencies. We have very little time and will be moving through material quickly. This means

that class time is IMPORTANT and Ms. Manakul will do her best to maximize the time we have together.

3. IF ABSENT: DON’T GET LEFT BEHIND. Remember chemistry is a class you must practice to understand,

which means missing a day can be essential to failing. Get help! See My Weekly Schedule.

4. Turn in work on time! Every day your work is late you will lose 30% that means that by the end of day 3 your

work will be an automatic zero. Ms. Manakul will give 2 extra days to make up work for every day you have an

excused absent.

5. No food or drinks will be allowed in the lab room. You may eat in the classroom as long as it does not become a

distraction. If you need to throw away large items or liquids, please toss them in trash cans in the hallways.

6. ALL electronic devices except your computer may not be used in the classroom. There will be occasions where

you may use them. Ms. Manakul will either say this or you may ask if the time is appropriate. If the time is

inappropriate the device will be taken on the second verbal warning and can be retrieved by the end of class.

7. Per district policy: Every 3 tardies = detention. Ms. Manakul considers a tardy as not being in the classroom at

the start of class.

GRADES:

A = 90.0 – 100.0 10% Homework

B = 80.0 – 89.9 20% Activities/Discussions

C = 70.0 – 79.9 20% Lab

D = 60.0 – 69.9 20% Quizzes

F = 0.0 – 59.9 30% Tests

Quarter Grades: Based on breakdown Semester Grade: Average of Quarters Year Grades: Average of Semesters

EXTRA CREDIT:

Extra credit is given throughout the school year and will be collected and graded each quarter. Students have the option to

read, analyze, and explain an approved scientific article following the extra credit rubric. Copies of the articles are located

in the Google Classroom under the Extra Credit Tab. Only one extra credit article can be done per quarter, and if one is

not done for the quarter students do not have an option to make it up later. Due dates for extra credit will be announced in

class, on Google Classroom, and will be posted on the calendar in the classroom.

TIPS TO SURVIVING CHEMISTRY:

1. Come to class prepared. If you bring you’re “A” game, you are more likely to get an “A” for the year.

2. Participate in class. The more you participate, the better you learn.

3. Pay attention in class. Ms. Manakul loves tell you exactly what will be on the test. If you’re listening you will

know what to expect and what to do.

4. Do the practice tests as if it was a real test. This will give you a better idea of what you need to study more and

how long you need for each problem so you can manage your time wisely.

5. Keep all your notecards. They will make it easier to study for the semester tests.

6. If you ever think something is wrong on a test, homework, notes, etc. TELL Ms. Manakul right away. We all get

confused; including Ms. Manakul and it is better to get it clarified then answer a problem wrong.

7. Keep up with any missed notes or work. Chemistry BUILDS! Missing one day of class can put you two to three

days behind your classmate, which will make you struggle more during the test.

8. Do your best on all homework and classwork. Practicing chemistry is like a mental workout. The best workouts

are the ones that require the most mental energy, so DON’T GIVE UP, even if it is tough.

9. Do your own work. Copying others’ homework or labs will only hurt you on the quizzes and tests.

10. Get your own your calculator!

iv

Classroom Map

v

School Safety Drills

On this page you will fill in Manakul’s Classroom procedures for all School Safety Drills.

1. Fire Safety

2. Tornado

3. Lock Down

4. Earthquake

5. Where are the maps located for emergency drills?

vi

vii

viii

Chemistry I Semester I Objectives

Unit 1: Measurement

Objective Accomplishment Goal Quiz

You will be able to do complete

mathematical problems in

chemistry

You will be able to set up

dimensional

analysis/conversions and

scientific notation problems to

ensure a change in units.

You will be able to report

quantitative results with the

correct number of significant

figures and in correct scientific

notation.

You know how to write numbers

in scientific notation and know

all the rules of significant

figures.

Unit 2: Matter & Atom

Objective Accomplishment Goal Quiz Test

You can distinguish between an

element, compound, and

chemical reaction.

You are able to distinguish

between elements, compound,

and list the 6 types of chemical

reactions.

You are able to describe the

progression of scientific work

in distinguishing components of

the atom.

You can describe the

experiments associated with the

scientists who discovered the

atom and its components

You are able to distinguish the

components of the atom, and

calculate the average atomic

mass of an element.

You are can to determine the

number of electrons, protons,

neutrons, mass number, charge,

and atomic number when given

a combination of these. You can

calculate the average atomic

mass and know which isotope is

more abundant given specific

data.

You will be able to explain how

radioactive decay works.

You can complete radioactive

decay equations.

Unit 3: Electromagnetic Spectrum

Objective Accomplishment Goal Quiz Test

You will be able to explain the

atomic emission spectra by

transition of electrons between

energy levels.

You can calculate frequency and

wavelength of emitted light

using Plank’s constant and

explain what photoelectric effect

is.

You can explain how electrons

are arranged the modern atom

structure.

You can organize electrons in

any element by their energy

order and orbital type.

ix

You will be able to explain

how elements are organized in

the periodic table.

You can identify each element

as either a metal, metalloid, or a

non-metal and determine the

element’s family name.

You will be able to explain

periodic trends.

You can explain ion formation,

atomic radius, ionization energy,

and electronegativity energy.

Unit 4: Metals & Ionic Bonding

Objective Accomplishment Goal Quiz Test

You are able to define valence

electrons and determine their

number for a particular group

of chemical elements, explain

octet rule and Lewis dot

diagrams, which atoms are

likely to lose or gain electrons,

how cations and anions form.

You can explain why

representative elements like to

be noble gases, and draw a

picture to depict what a

representative element looks like

as a cation or anion.

You can explain the formation

of Lattice Structures and its

determination of chemical and

physical properties for a

compound.

You can determine the

coordination number for a

compound and describe the

different types of lattice

structures.

You can explain how ionic and

metallic bonds are formed,

describe the properties of each,

and explain why alloys are

important.

You can list at least three

properties of an ionic and

metallic compound, and explain

the importance of alloys.

Unit 5: Covalent Bonding

Objective Accomplishment Goal Quiz Test

You can explain how electrons

are shared in covalent bonds,

what are coordinate covalent

bonds, compare and contrast

ionic and covalent bonds.

You can draw covalent

compounds and their resonance

structures.

You can explain the difference

between atomic and molecular

orbitals, predict the shapes of

molecules

You can determine a covalent

compounds shape using VSEPR

theory.

Molecules using VSEPR

theory, and explain how

bonding occurs between

molecules.

You can use VSEPR theory to

explain polarity within a

covalent compound and the

attraction between compounds.

x

Unit 6: Moles & Chemical Reactions

Objective Accomplishment Goal Quiz Test

You will be able to execute

conversion problems

You are able to determine the

givens, set up the train tracks,

and execute your solution.

You will be able to explain

what a mole is and use it in

conversion problems

You can explain what a mole is

and be able to write it as a

conversion factor.

You will be able to determine

the chemical formula of an

ionic or covalent compound.

You are able to determine if a

compound is going to be ionic or

covalent and then write its

chemical formula based on its

name.

You will be able to calculate

percent composition of a

chemical compound.

You can set up the steps to

determine percent composition

of a chemical compound.

xi

Objective Graphs

On this page you will create a bar graph for each unit of your goal, quiz and test scores per

objective (these are obtained from the trackers done after each quiz and test). Use test data for

best accuracy & remember to put units and labels. SUGGESTION: Use different colors.

Unit 1: Measurement Unit 2: Matter & Atom

Unit 3: Electromagnetic Spectrum Unit 4: Metals & Ionic Bonding

xii

Unit 5: Covalent Bonding Unit 6: Moles & Chemical Reactions

Spare Graphs

1

Unit 1: Measurement Homework

SECTION 1.1: MEASUREMENTS AND THEIR UNCERTAINTY

1. Explain why there are two units for density (g/L vs. g/m3)?

2. If the temperature in Europe is 23.0°C, what would this temperature be in Fahrenheit

(°F)?

3. If the temperature in the lab is 73.2°F, what is the temperature in Kelvin?

4. Determine the number of significant figures are in the following numbers:

a. 209 m

b. 4140 inches

c. 1.330 mi

d. 0.0034 cm

e. 3.22 x 1016 g

f. 214 dogs

5. Perform the following operations (try these without a scientific calculator):

a. 2.567 m x 0.00456 m

b. 103 L / 2.33 L

c. 6.23 x 102 g + 5.33 x 103 g

d. 7.34 x 10-2 cm – 1.33 x 10-1 cm

2

Use the following information to answer questions 6-11.

Using different rulers, Bruce and Pete each measured the length of the same object three times.

Bruce’s three measurements are 19 cm, 20 cm, and 22 cm. Pete’s three measurements are 20.9

cm, 21.0 cm, and 21.0 cm.

6. Calculate the average value of Bruce’s measurements and express the answer with the

correct number of significant figures.

7. Calculate the average value of Pete’s measurements and express the answer with the

correct number of significant figures.

8. Whose measurements are more precise?

9. The actual length of the object is 20 cm. whose measurements are more accurate?

10. What is the error of Pete’s average measurement?

11. What is the percent error of Pete’s average measurement?

3

SECTION 1.2: CONVERIONS

1. Convert the following quantities. See periodic table for necessary information.

a. 565,900 seconds into days

b. 17 years into minutes

c. 43 miles into feet

d. 165 pounds into kilograms

e. 27 milliohms to ohms

f. 49 micrometers to meters

g. 469 Joules to calories

4

Unit 2: Matter & Atom Homework

SECTION 2.1 PROPERTIES OF MATTER 1. Complete the following table.

Physical state Definite Shape? Definite

Volume?

Easily

Compressed?

gas

no no

yes

2. Which of the following is not a type of matter?

a. plasma b. aqueous c. solid d. gas

3. Which of the following is not a property of a gas?

a. has a definite shape c. assumes the shape of its container

b. has an indefinite volume d. is easily compressed

4. Which of the following is not a physical property of sucrose?

a. solid at room temperature c. dissolves in water

b. decomposes when heated d. tastes sweet

5. Which of the following is in a different physical state at room temperature than the other

three?

a. salt b. sugar c. flour d. water

6. Classify the following properties as extensive or intensive.

a. color b. volume c. mass d. boiling point

SECTION 2.2 MIXTURES

1. How might you separate a mixture of water and salt?

2. What is a homogeneous mixture?

3. Which of the following mixtures are homogeneous? Which are heterogeneous?

a. gasoline b. chunky peanut butter c. oil and vinegar salad dressing

4. Which of the following are substances? Which are mixtures?

a. ethanol b. motor oil c. vinegar d. neon

5. Classify the following as elements, compounds, or mixtures.

a. table salt b. water c. iron d. stainless steel

5

SECTION 2.3 ELEMENTS AND COMPOUNDS

1. What elements make up ammonia, chemical formula NH3?

2. Name the elements represented by the following chemical symbols.

a. Pb b. K c. Au d. Fe

3. Write the chemical symbol for each of the following elements.

a. tin b. sodium c. silver d. carbon

4. A liquid is allowed to evaporate and leaves no residue. Can you determine whether it was

an element, a compound, or a mixture? How?

5. Which of the following is not an element?

a. copper c. sulfur

b. sucrose d. helium

SECTION 2.4 CHEMICAL REACTIONS

1. Which one of the following is a chemical change?

a. Gasoline boils. c. Gasoline burns.

b. Oxygen is added to gasoline. d. Gasoline is poured into a tank.

2. Classify each of the following changes as physical or chemical.

a. A puddle is dried by the sun. c. Bread is toasted.

b. A dark cloth is faded by sunlight. d. Soap is mixed with water.

3. Which of the following is not a type of chemical reaction?

a. combustion c. triple replacement

b. double replacement d. decomposition

4. Carbon dioxide plus water yields carbonic acid.

a. Name the product(s) of this reaction.

a. Name the reactant(s) of this reaction.

6

5. List 3 key terms to determine a physical change.

6. List 3 key terms to determine a chemical change.

7. Which of the following is not a physical change?

a. dissolving sugar in water c. evaporating sea water to obtain salt

b. burning gasoline in an engine d. slicing a piece of bread

SECTION 2.5 STRUCTURE OF THE NUCLEAR ATOM

1. A sulfur-32 atom contains 16 protons, 16 neutrons, and 16 electrons. What is the mass (in

grams) of a sulfur-32 atom?

2. The mass of a neutron is 1.67 x 10-24 g; approximately what number of neutrons would

equal a mass of one gram?

3. Which statement is consistent with the results of Rutherford’s gold foil experiment?

a. All atoms have a positive charge.

b. Atoms are mostly empty space.

c. The nucleus of an atom contains protons and electrons.

d. Mass is spread uniformly throughout an atom.

SECTION 2.6 DISTINGUISHING BETWEEN ATOMS

1. How many protons are found in an atom of each of the following?

a. boron c. neon

b. sulfur d. lithium

2. How many neutrons are in each atom?

a. c.

b. d.

23 Na 11

81 Br 35

238 U 92

19 F 9

7

3. Complete the table for the following elements.

Element Number of

Protons

Number of

Electrons

Number of

Neutrons

Atomic

Number

Mass

Number

Manganese 25 30

Sodium 11 12

Bromine 35 45

Yttrium 39 89

Arsenic 33 75

Actinium 227

4. The two most abundant isotopes of carbon are carbon-12 (mass = 12.00 amu) and carbon-

13 (mass = 13.00 amu). Their relative abundances are 98.9% and 1.10%, respectively.

Calculate the atomic mass of carbon.

5. Element X has two isotopes: X-100 and X-104. If the atomic mass of X is 101 amu, what

is the relative abundance of each isotope in nature?

8

SECTION 2.7: Nuclear Chemistry

1. Identify the following as alpha, beta, gamma, or neutron.

a. 1

0 n

b. 0

−1 e

c. 4

2 He

d. 0

0 γ

e. Nuclear decay with no mass and no charge.

f. An electron.

g. Least penetrating nuclear decay.

h. Most damaging nuclear decay to the human body.

i. Nuclear day that can be stopped by skin or paper.

j. Nuclear decay that can be stopped by aluminum.

2. Complete the following nuclear equations.

a. 42

19𝐾 →

0

−1𝑒 + ________ d.

239

94𝑃𝑢 →

4

2𝐻𝑒 + ________

b. 9

4𝐵𝑒 →

9

4𝐵𝑒 + ________ e.

235

92𝑈 → _______ +

231

90𝑇ℎ

c. 6

3𝐿𝑖 →

4

2𝐻𝑒 + ________ f. _______ →

142

56𝐵𝑎 +

91

36𝐾𝑟 + 3

1

0𝑛

9

Unit 3: Electromagnetic Spectrum

SECTION 3.1 MODELS OF THE ATOM

1. How many sublevels are in the following principal energy levels?

a. n = 1 c. n = 3 e. n = 5

b. n = 2 d. n = 4 f. n = 6

2. How many orbitals are in the following sublevels?

a. 1s sublevel d. 4f sublevel g. fifth principal energy level

b. 5s sublevel e. 7s sublevel h. 6d sublevel

c. 4d sublevel f. 3p sublevel

3. What are the types of sublevels and number of orbitals in the following energy levels?

a. n = 1 c. n = 3 e. n = 5

b. n = 2 d. n = 4

SECTION 3.2 ELECTRON ARRANGEMENTS IN ATOMS

1. Write a complete electron configuration of each atom.

a. hydrogen d. barium g. krypton

b. vanadium e. bromine h. arsenic

c. magnesium f. sulfur i. radon

10

SECTION 3.3 PHYSICS AND THE QUANTUM MECHANICAL MODEL

1. What is the wavelength of the radiation whose frequency is 5.00 × 1015 s–1? In what region

of the electromagnetic spectrum is this radiation?

2. An inexpensive laser that is available to the public emits light that has a wavelength of 670

nm. What are the color and frequency of the radiation?

3. What is the energy of a photon whose frequency is 2.22 × 1014 s–1?

4. What is the frequency of a photon whose energy is 6.00 × 10–15 J?

5. Arrange the following types of electromagnetic radiation in order of increasing frequency.

a. infrared c. visible light e. microwaves

b. gamma rays d. radio waves f. ultraviolet

6. Suppose that your favorite AM radio station broadcasts at a frequency of 1600 kHz. What

is the wavelength in meters of the radiation from the station?

11

SECTION 3.4 ORGANIZING THE ELEMENTS

1. Which element listed below should have chemical properties similar to fluorine (F)?

a. Li b. Si c. Br d. Ne

2. Identify each element as a metal, metalloid, or nonmetal.

a. fluorine

b. germanium

c. zinc

d. phosphorus

e. lithium

3. Which of the following is not a transition metal?

a. Magnesium b. titanium c. chromium d. mercury

4. Name two elements that have properties similar to those of the element potassium

5. Elements in the periodic table can be divided into three broad classes based on their

general characteristics. What are these classes and how do they differ?

SECTION 3.5 CLASSIFYING THE ELEMENTS

1. Use the periodic table to write the electron configuration for silicon. Explain why silicon

is considered a metalloid and representative element.

2. Use the periodic table to write the electron configuration for iodine. Explain why iodine

is classified as a halogen.

3. Which group of elements is characterized by an s2p3 configuration?

12

4. Name the element that matches the following description.

a. one that has 5 electrons in the third energy level

b. one with an electron configuration that ends in 4s24p5

c. the Group 6A element in period 4

5. Identify the elements that have electron configurations that end as follows.

a. 2s22p4

b. 4s2

c. 3d104s2

6. What is the common characteristic of the electron configurations of the elements Ne and

Ar? In which group would you find them?

7. Why would you expect lithium (Li) and sulfur (S) to have different chemical and physical

properties?

8. What characterizes the electron configurations of transition metals such as silver (Ag)

and iron (Fe)?

13

SECTION 3.6 PERIODIC TRENDS

1. Explain why a magnesium atom is smaller than atoms of both sodium and calcium.

2. Predict the size of the astatine (At) atom compared to that of tellurium (Te). Explain your

prediction.

3. Would you expect a Cl– ion to be larger or smaller than an Mg2+ ion? Explain.

4. Which effect on atomic size is more significant, an increase in nuclear charge across a

period or an increase in occupied energy levels within a group? Explain.

5. Explain why the sulfide ion (S2–) is larger than the chloride ion (Cl–).

6. Compare the first ionization energy of sodium to that of potassium.

7. Compare the first ionization energy lithium to that of beryllium.

8. Is the electro negativity of barium larger or smaller than that of strontium? Explain.

9. What is the most likely ion for magnesium to form? Explain.

10. Arrange oxygen, fluorine, and sulfur in order of increasing electro negativity.

14

Unit 4: Metals & Ionic Bonding

SECTION 4.1 IONS

1. For each element below, state (i) the number of valence electrons in the atom, (ii) the

electron dot structure, and (iii) the chemical symbol(s) for the most stable ion.

a. Ba b. I c. K

2. How many valence electrons does each of the following atoms have?

a. gallium b. fluorine c. selenium

3. Write the electron configuration for each of the following atoms and ions.

a. Ca c. Na+ e. O2–

b. chlorine atom d. phosphide ion

4. What is the relationship between the group number of the representative elements and the

number of valence electrons?

5. How many electrons will each element gain or lose in forming an ion? State whether the

resulting ion is a cation or an anion.

a. strontium c. tellurium e. bromine

b. aluminum d. rubidium f. phosphorus

6. Give the name and symbol of the ion formed when

a. a chlorine atom gains one electron.

b. a potassium atom loses one electron.

c. an oxygen atom gains two electrons.

d. a barium atom loses two electrons.

15

7. How many electrons are lost or gained in forming each of the following ions?

a. Mg2+ b. Br– c. Ag+ d. Fe3+

8. Classify each of the following as a cation or an anion.

a. Na+ c. I– e. Ca2+

b. Cu2+ d. O2– f. Cs+

SECTION 4.2 NAMING IONS

1. What is the charge on the ion typically formed by each element?

a. Oxygen c. sodium e. nickel, 2 electrons lost

b. iodine d. aluminum f. magnesium

2. How many electrons does the neutral atom gain or lose when each ion forms?

a. Cr3+ c. Li+ e. Cl–

b. P3– d. Ca2+ f. O2–

3. Name each ion. Identify each as a cation or an anion.

a. Sn2+ c. Br– e. H–

b. Co3+ d. K+ f. Mn2+

4. Write the formula (including charge) for each ion. Use Table 9.3 if necessary.

a. carbonate ion c. sulfate ion e. chromate ion

b. nitrite ion d. hydroxide ion f. ammonium ion

16

5. Name the following ions. Identify each as a cation or an anion.

a. CN– c. PO43– e. Ca2+

b. HCO3– d. Cl– f. SO32–

SECTION 4.3 IONIC BONDS AND IONIC COMPOUNDS

1. Use electron dot structures to predict the formula of the ionic compounds formed when

the following elements combine.

a. sodium and bromine d. aluminum and oxygen

b. sodium and sulfur e. barium and chlorine

c. calcium and iodine

2. Which of these combinations of elements are most likely to react to form ionic

compounds?

a. sodium and magnesium c. potassium and iodine

b. barium and sulfur d. oxygen and argon

3. What is the meaning of coordination number?

4. How is the coordination number determined?

17

SECTION 4.4 NAMING AND WRITING FORMULAS FOR IONIC COMPOUNDS

1. Write the formulas for these binary ionic compounds.

a. magnesium oxide c. potassium iodide e. sodium sulfide

b. tin(II) fluoride d. aluminum chloride f. ferric bromide

2. Write the formulas for the compounds formed from these pairs of ions.

a. Ba2+, Cl– c. Ca2+, S2– e. Al3+, O2–

b. Ag+, I– d. K+, Br– f. Fe2+, O2–

3. Name the following binary ionic compounds.

a. MnO2 c. CaCl2 e. NiCl2 g. CuCl2

b. Li3N d. SrBr2 f. K2S h. SnCl4

4. Write formulas for the following ionic compounds.

a. sodium phosphate c. sodium hydroxide e. ammonium chloride

b. magnesium sulfate d. potassium cyanide f. potassium dichromate

5. Write formulas for compounds formed from these pairs of ions.

a. NH4+, SO4

2– c. barium ion and hydroxide ion

b. K+, NO3– d. lithium ion and carbonate ion

18

6. Name the following compounds.

a. NaCN c. Na2SO4 e. Cu(OH)2

b. FeCl3 d. K2CO3 f. LiNO3

7. Name and give the charge of the metal cation in each of the following ionic compounds.

a. Na3PO4 c. CaS e. FeCl3

b. NiCl2 d. K2S f. CuI

SECTION 4.5 BONDING IN METALS

1. What is a metallic bond?

2. How is the electrical conductivity of a metal explained by metallic bonds?

3. Are metals crystalline? Explain.

4. Give three possible crystalline arrangements of metals. Describe each.

5. What is an alloy?

6. Name the principal elements present in each of the following alloys.

a. brass d. sterling silver

b. bronze e. cast iron

c. stainless steel f. spring steel

19

Unit 5: Covalent Bonding

SECTION 5.1 MOLECULAR COMPOUNDS

1. Classify each of the following as an atom or a molecule.

a. Be c. N2 e. Ne

b. CO2 d. H2O

2. Which of the following are diatomic molecules?

a. CO2 c. O2 e. CO

b. N2 d. H2O

3. What types of elements tend to combine to form molecular compounds?

4. What information does a molecule’s molecular structure give?

5. How do ionic compounds and molecular compounds differ in their relative melting and

boiling points?

SECTION 5.2 NAMING AND WRITING FORMULAS FOR MOLECULAR

COMPOUNDS

1. Name the following molecular compounds.

a. PCl5 c. NO2 e. P4O6 g. SiO2

b. CCl4 d. N2F2 f. XeF2 h. Cl2O7

2. Write the formulas for the following binary molecular compounds.

a. nitrogen tribromide c. sulfur dioxide

b. dichlorine monoxide d. dinitrogen tetrafluoride

20

SECTION 5.3 THE NATURE OF COVALENT BONDING

1. Draw the electron dot structure for hydrogen fluoride, HF.

2. Draw the electron dot structure for phosphorus trifluoride, PF3.

3. Draw the electron dot structure for nitrogen trichloride, NCl3.

4. How many resonance structures can be drawn for CO32–? Show the electron dot

structures for each.

SECTION 5.4 BONDING THEORIES

1. Predict the shape and bond angle for the compound carbon tetrafluoride, CF4.

2. Predict the shape and bond angle for phosphorus trifluoride, PF3.

3. Predict the shape and bond angle of xenon tetrachloride, XeCl4.

4. Predict the shape and bond angle of sulfur hexafluoride, SF6.

21

SECTION 5.5 POLAR BONDS AND MOLECULES

1. What type of bond—nonpolar covalent, polar covalent, or ionic—will form between each

pair of atoms?

a. Na and O b. O and O c. P and O

2. Explain why most chemical bonds would be classified as either polar covalent or ionic.

3. Would you expect carbon monoxide and carbon dioxide to be polar or nonpolar

molecules?

4. Draw the structural formulas for each molecule and identify polar covalent bonds by

assigning the slightly positive (δ+) and slightly negative (δ –) symbols to the appropriate

atoms.

a. NH3 b. CF3

5. Which would you expect to have the higher melting point, CaO or CS2?

22

Unit 6: Moles & Chemical Reactions Homework

SECTION 6.1 THE MOLE: A MEASUREMENT OF MATTER

1. What is the molar mass of sucrose (C12H22O11)?

2. What is the molar mass of each of the following compounds?

a. phosphorus pentachloride (PCl5)

b. uranium hexafluoride (UF6)

3. Calculate the molar mass of each of the following ionic compounds:

a. KMnO4

b. Ca3(PO4)2

4. How many moles is 3.52 × 1024 molecules of water?

5. How many atoms of zinc are in 0.60 mol of zinc?

6. What is the mass of 1.00 mol of oxygen (O2)?

23

SECTION 6.2 MOLE–MASS AND MOLE–VOLUME RELATIONSHIPS

1. What is the molar mass of each of the following compounds?

a. C6H12O6 b. NaHCO3 c. C7H12 d. KNH4SO4

2. Calculate the mass in grams of each of the following:

a. 8.0 mol lead oxide (PbO)

b. 0.75 mol hydrogen sulfide (H2S)

c. 0.00100 mol silicon tetrahydride (SiH4)

d. 1.50 × 10–2 mol molecular oxygen (O2)

e. 2.30 mol ethylene glycol (C2H6O2)

3. How many grams are in 1.73 mol of dinitrogen pentoxide (N2O5)?

24

4. How many grams are in 0.658 mol of calcium phosphate [Ca3(PO4)2]?

5. Calculate the number of moles in each of the following:

a. 0.50 g sodium bromide (NaBr)

b. 13.5 g magnesium nitrate [Mg(NO3)2]

c. 1.02 g magnesium chloride (MgCl2)

d. 0.00100 g monochloromethane (CH3Cl)

e. 1.50 × 10–3 g propylene glycol [C3H6(OH)2]

6. A chemist plans to use 435.0 grams of ammonium nitrate (NH4NO3) in a reaction. How

many moles of the compound is this?

7. A solution is to be prepared in a laboratory. The solution requires 0.0465 mol of quinine

(C20H24N2O2). What mass, in grams, should the laboratory technician obtain in order to

make the solution?

25

8. What is the volume at STP of 2.66 mol of methane (CH4) gas?

9. How many moles is 135 L of ammonia (NH3) gas at STP?

SECTION 6.3 PERCENT COMPOSITION AND CHEMICAL FORMULAS

1. A sample of a compound analyzed in a chemistry laboratory consists of 5.34 g of carbon,

0.42 g of hydrogen, and 47.08 g of chlorine. What is the percent composition of this

compound?

2. Find the percent composition of a compound containing tin and chlorine if 18.35 g of the

compound contains 5.74 g of tin.

3. If 3.907 g of carbon combines completely with 0.874 g of hydrogen to form a compound,

what is the percent composition of this compound?

26

4. From the formula for calcium acetate, Ca(C2H3O2)2, calculate the mass of carbon that can

be obtained from 65.3 g of the compound.

5. How many grams of aluminum are in 25.0 g of aluminum oxide (Al2O3)?

6. How many grams of iron are in 21.6 g of iron(III) oxide (Fe2O3)?

7. Determine the empirical formula of each of the following compounds from the percent

composition:

a. 7.8% carbon and 92.2% chlorine

b. 10.0% C, 0.80% H, 89.1% Cl

27

SECTION 6.4 DESCRIBING CHEMICAL REACTIONS

1. Write the skeleton equation for the reaction between hydrogen and oxygen that produces

water.

2. Write the skeleton equation for the reaction that produces iron(II) sulfide from iron and

sulfur.

3. Write the skeleton equation representing the heating of magnesium carbonate to produce

solid magnesium oxide and carbon dioxide gas.

4. Write a balanced equation for the production of HCl gas from its elements.

5. Write a sentence that completely describes the chemical reaction represented by this

balanced equation.

2HCl(aq) + CaCO3(s) CO2(g) + CaCl2(aq) + H2O(l)

6. Write the word equation for the following equation. Write a sentence fully describing the

reaction. Is the equation correctly balanced? Explain.

2Ag(s) + S(s) Ag2S(s)

7. Write a balanced equation representing the formation of aqueous sulfuric acid from water

and sulfur trioxide gas.

28

8. Write a balanced equation from this word equation.

aqueous silver nitrate + copper metal silver metal + aqueous copper nitrate

9. Write a balanced equation for the following word equation.

phosphorus + oxygen tetraphosphorous decoxide

SECTION 6.5 TYPES OF CHEMICAL REACTIONS

1. Write a balanced equation representing the reaction of magnesium with oxygen gas to

produce magnesium oxide.

2. Write the balanced equation for the reaction that occurs between aluminum and fluorine.

3. Write the balanced equation for the production of oxygen gas and potassium chloride

from the decomposition of potassium chlorate.

4. Write the balanced equation for the reaction between hydrochloric acid and calcium

metal. The products are hydrogen gas and calcium chloride.

5. Write the balanced equation for the combustion of propane (C3H8) to produce carbon

dioxide and water vapor.

29

6. Write the balanced equation for the reaction between iron(III) chloride and sodium

hydroxide. The products are iron(III) hydroxide and sodium chloride.

7. Classify each of the reactions in problems 1–6 as to type.

8. Use the activity series of metals (Table 11.2) and your knowledge of the relative

reactivity of the halogens to predict whether the following reactions will occur. Write

balanced equations for those reactions that do occur.

a. Br2(l) + NaCl(aq) →

b. Ca(s) + Mg(NO3)2(aq) →

c. K(s) + H2SO4(aq) →

d. Zn(s) + NaOH(aq) →

SECTION 6.6 REACTIONS IN AQUEOUS SOLUTION

1. Write the net ionic equation for the reaction between aqueous barium nitrate, Ba(NO3)2,

and sodium sulfate, Na2SO4.

2. Magnesium reacts with HCl to form hydrogen and magnesium chloride. Write the

balanced net ionic equation for this reaction.

30

3. The double-replacement reaction below results in the formation of the precipitate lead

chloride. Balance the equation and write the net ionic equation.

a. Pb(NO3)2(aq) + NH4Cl(aq) → PbCl2(s) + NH4NO3(aq)

4. Identify the precipitate formed when solutions of the following ionic compounds are

mixed. If no precipitate is formed, write no precipitate.

a. Zn(NO3)2 + SnCl2 →

b. KCl + AgNO3 →

c. Cu(NO3)2 + Na2S →

d. Al2(SO4)3 + 3Mg(OH)2 →

31

Significant Figures Lab

Background:

This lab will test your skills with a triple balance beam and a graduated cylinder. A

mechanical balance works on the principle that equal masses will be pulled downwards by

gravity with equal force. Thus, if two equal masses are suspended by a string around a pulley,

they balance each other out. Hence the name balance beam. A triple balance beam uses three

(triple) balances to create equivalent mass as what is being weighted. Calibrating a triple balance

beam requires accurate sight and motion control. By calibrating or zeroing the balance (placing

all weights onto the zero marker, then adjusting the left side of the beam so to create a perfect

line from the left to the right side of the beam) the user ensures more accurate results. A

calibrated triple balance beam will give users accurate measurements up to five significant

figures (see information about significant figures below).

A graduated cylinder is considered a measuring cylinder, and is used to measure the

volume of a liquid, generally in milliliters. There are different types of graduated cylinders: 10

ml, 25ml, 50ml, and 100ml will be available in the lab room. Each has its own grading divisions,

for instance 100 mL cylinders have 1 mL grading divisions while 10 ml cylinders have 0.1 ml

grading divisions.

Any measurement established from a scientific tool requires good eye sight and some

estimation like the triple balance beam. To do this, you will need to move the weights so that the

scale re-aligns. The value of the measurement will be the summation of all the weights.

Likewise, when liquid is poured into a graduated cylinder, it creates a meniscus at the top of the

liquid (looks like a small concave curve at the top of the liquid). Almost all liquids create a

meniscus, because of the capillary (attractive) force which occurs between most liquids and the

sides of the graduated cylinder. When measuring the value of a graduated cylinder you must

determine its value from the bottom of the meniscus.

Significant figures are the value of the measurement which is reported. When reading

scales, the amount of significant figures is determined by the grading divisions on the tool plus

one. For instance the 100 mL graduated cylinder has 1 mL grading divisions. Thus, a

measurement on a graduated cylinder would report a value with one decimal place. The plus one

value is included in measurements to ensure all variations that could occur. Many times a

meniscus will not be directly on a line, thus the scientist must estimate its final location, hence

the added digit.

When doing calculations with significant figures there are a couple rules to follow.

1. When adding and subtracting, the smallest amount of decimals must be the same

as your final number.

2. When multiplication and division, the smallest number of digits must be the same

as your final number.

Since your measurements are done through estimation, your final digit will contain

uncertainty. The measurement uncertainty is often taken as the standard deviation of a state-of-

knowledge probability distribution over the possible values that could attribute to a measured

quantity. For Chemistry I, you will be judging this uncertainty through error calculation, which

the equation is given below.

𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝐸𝑟𝑟𝑜𝑟 = |𝐴𝑐𝑡𝑢𝑎𝑙 𝑤𝑒𝑖𝑔ℎ𝑡 − 𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑤𝑒𝑖𝑔ℎ𝑡|

𝐴𝑐𝑡𝑢𝑎𝑙 𝑊𝑒𝑖𝑔ℎ𝑡 × 100

32

Objective:

1. Students gain experience with various scientific tools and learn best practices with each

tool.

2. Students will be able to determine the number of significant figures/digits for each

scientific tool.

Materials:

Triple Balance Beam

Nuts and Bolts Bag

100 mL Graduated cylinder

250 mL Beaker

250 mL Flask

Lab Book

Pencil or Black or Blue Pen

Methods:

In this lab you will be conducting two lab experiments. During each lab you must

determine the proper number of significant figures/digits for each scientific tool when recording

data. Then you will need to utilize the rules of significant figures to make and report the correct

end values for each experiment.

In the first lab you are working towards determining the mass of a single washer, bolt,

and nut. In your Nuts & Bolts Bag you will see several bolts, nuts, and washer glued together.

On some bolts you will have more than one nut, and on some bolts you will have more than one

washer. You will need to weigh each bolt combination individually. Then do some creative

calculations to determine the weight of one bolt, one nut, and one washer. The triple balance

beam used to weight each combination should be calibrated or zeroed before weighting each

combination. This is to ensure accurate results are recorded. All weights should be placed into

the Nuts & Bolts Lab Data Table provided in the data recording section. All calculations done

should be placed under the data processing section.

In the second lab you will be comparing the density of water in a 250 mL Beaker, 250

mL Flask, and a graduated cylinder. To do this you will need to record the mass of 50 mL of

water in each of the vessels. Thus, determine the mass of each vessel, the mass of each vessel

with 50 mL of water, the mass of water within each vessel, and the density of pure water. Record

all values in the Density of Water Comparison Data Table, and do all calculations under the data

processing section. (Hint: the actual density of water is 1.0 g/ml)

Safety:

For your safety it is important to proceed with caution. The bolt, nut, and washer combinations

are heavy and can cause severe damage to the lab room and to people if thrown or mishandled.

Likewise, you are handling glassware, which should always be inspected for cracks before use,

has the ability to shatter, especially when heated. Thus, handle glassware with extreme caution.

Should you shatter or break any glassware, Alert Ms. Manakul immediately, and ensure no one

touches the broken glassware. Ms. Manakul will handle the cleaning processes. Since we are not

heating glassware, you will not be required to wear goggles or lab coats for this lab. However,

closed toed shoes must be worn at all times. Hard toed shoes would be preferable since the bolt,

nut, and washer combinations are heavy and can cause damage to a person’s foot if dropped on

to them. Likewise stepping on broken glass or having glassware break on your feet is dangerous.

This lab has no environmental concerns, since no hazardous materials are being utilized.

32

Data Recording:

Nuts & Bolts Data Table

Weight (g)

Com

bin

ati

on

s

1 Bolt, 1 Nut, 1 Washer

1 Bolt, 3 Nuts, 0 Washer

1 Bolt, 1 Nut, 3 Washers

1 Bolt, 1 Nut, 2 Washers

1 Bolt, 2 Nuts, 2 Washer

Density of Water Comparison Data Table

Vessel Mass of Vessel Mass of Vessel +

Water

Mass of Water

250 mL Beaker

250 mL Flask

100 mL Graduated

Cylinder

Data Processing:

1. Calculate the weight of each individual bolt, nut, and washer.

(Weight of 1 Bolt, 1 Nut, 3 Washers) – (Weight of 1 Bolt, 1 Nut, 2 Washers) = Weight of 1

Washer

2. Calculate your percent error for a bolt, nut, and washer. You will need the actual weights

for the bolt, nut, and washer. Ask Ms. Manakul for this data when you reach this step.

Remember to show all work.

3. Calculate the density of water for each vessel. Show all work here. Calculate the percent

error for each vessel. You will need the actual density of water. Ask Ms. Manakul for this

data when you reach this step. Remember to show all work.

Challenge Questions:

1. Explain why having accurate significant figures, especially when doing calculations, is

important to the scientific community.

2. You were supposed to get large values of percent error for the bolt, nut, and washer.

Why?

3. Why was using tap water in experiment two not going to give us the best values?

Conclusion & Evaluation

1. Write your own conclusion and evaluation of this lab.

33

Density Lab

Background:

The density, or more precisely, the volumetric mass density, of a substance is its mass per unit

volume. The symbol most often used for density is ρ (the lower case Greek letter rho).

Mathematically, density is defined as mass divided by volume, where ρ is the density, m is the

mass, and V is the volume.

𝜌 = 𝑚𝑎𝑠𝑠

𝑣𝑜𝑙𝑢𝑚𝑒=

𝑚

𝑉

For a pure substance the density has the same numerical value as its mass concentration.

Different materials usually have different densities, and density may be relevant to buoyancy,

purity and packaging. The density of a material varies with temperature and pressure. This

variation is typically small for solids and liquids but much greater for gases. Below is a simple

list of densities of some common materials:

Solids and Liquids

Material Density at 20°C (kg/m3)

Gold 19.3

Mercury 13.6

Lead 11.3

Aluminum 2.70

Glass (flint) 2.9-5.9

Silver 10.5

Brass 8.49

Diamond 3-3.5

Epoxy glass fiber 1.5

Iron 7.87

Objectives:

1. You will be able to determine the substance of the compound based on its density.

2. You will gain practice calculating using conversions.

Materials:

1 white/clear marble

1 silver marble

Triple balance beam

Weigh boat

10 mL graduated cylinder

Tap water

50 mL Beaker

Methods:

In this lab you will be determining what the white/clear and silver marble are made of based on

their densities. You will use the graduated cylinder to find the volume of each marble. You will

first need to fill your 50 mL beaker with some tap water (DO NOT use the distilled water). Once

you have your water you can fill your 10 mL graduated cylinder with 5 mL of water. Once you

have the bottom meniscus on the 5 mL maker you can CAREFULLY place the white/clear

marble into the beaker. DO NOT drop the marble to the bottom of the graduated cylinder,

because it could break the graduated cylinder. You can then record the change in water level.

34

You can then CAREFULLY remove the marble from the graduated cylinder by pouring the

water and marble into your 50 mL beaker. To extract the marble you may use tweezers. You will

need to do the same procedure four more times (this will give you five trials). Make sure to dry

your marble between trials; otherwise your volumes will be incorrect, because you are adding

extra water with your marble. You will then need to repeat the procedure for the silver marble;

again making sure you do five trials. Once you have done that you will need to take the weight of

your marble. The weigh boat is to help ensure the marble doesn’t roll off the balance. Thus, you

will need to determine the weight of the weigh boat first. Then the weight of the marble and

weigh boat. Once you have those numbers you will need to calculate the weight of the marble by

itself. Once you have all your information, you can clean up and start your data processing

portion.

Safety:

In this lab you will need to wear, goggles, closed toed shoes, and lab coats. Since you will be

working with glass you will need to ensure the safety of your body and eyes. DO NOT lose any

materials especially the marbles. Throwing marbles or being carless with them is grounds for

expulsion from the lab. Any long hair will need to be tied back to ensure it does knock anything

over.

Environmental Safety:

There is very little environmental concerns for this lab. Any dangerous particulates that are being

transferred from the marbles to the water will be highly diluted within the water. Thus when the

water goes down the drain, the environment see little to no change.

Data Collection:

Marble Volume Change

Volume Change

Trial 1 Trial 2 Trial 3 Trial 4 Trial 5 Average

Marb

les

White/Clear

Silver

Marble Mass Values

Mass (g)

Weigh Boat

Weigh Boat and White/Clear Marble

White/Clear Marble

Weigh Boat and Silver Marble

Silver Marble

Data Processing:

1. Create a bar graph that shows the averages of the volume change for the white/clear and

silver marble.

35

2. Determine the density in g/mL of the white/clear marble by using the weight of the

white/clear marble and the average volume change for the white/clear marble.

3. Determine the density in g/mL of the silver marble by using the weight of the silver

marble and the average volume change for the silver marble.

Challenge Questions:

1. What type of composition was the white/clear marble?

2. What type of composition was the silver marble?

3. Most if not all densities determined in a lab are never exactly on the stated value given.

Why does this occur?

4. Explain another way determine the composition of the marble without destroying the

marble.

Conclusion/Evaluation:

1. Write your own conclusion and evaluation of this lab.

36

“Beanium” Lab

Background:

Most elements on the Periodic Table exist in at least two isotopic forms. Isotopes are

atoms with the same atomic number but with different mass numbers due to varying numbers of

neutrons. The atomic mass shown on the Periodic Table for each element, is actually an average

of all the isotopes of that element, weighted by the percentage of the abundance in which they

occur.

NEWS FLASH!!! A NEW ELEMENT HAS BEEN DISCOVERED.

Springfield USA—Nuclear Chemists, performing basic research on food products at Springfield

Power Plant, have discovered what is believed to be a new element. Mr. Burns, the plant’s

owner, says, “We have tentatively named this element Beanium.” Mr. Smithers, assistant to Mr.

Burns adds, “We derived this element from the protein nodules we put into our chili.”

Further research of the new element will be conducted in more

suitable surroundings, namely laboratories in a nearby school. Because

Springfield apparently only has an elementary school, research work

has been contracted to neighboring Clayton High School. “Student

excitement regarding this discovery is running at a fever pitch!” says

Lisa Simpson, student. Many chemistry students have generously

volunteered their time and expertise to help with the follow-up

experiments involving the new element.

Dr. Julius Hibbert says the first follow-up experiments

conducted at Clayton High School will determine how many isotopes

of this element exist. The second experiment will determine the mass

of each isotope. The third experiment will determine the percent

abundance of each isotope. The final calculations will discover the

average atomic mass of the new element.

“One unique property of Beanium should make these experiments particularly easy—

unlike normal atoms, Beanium atoms are very large.” says Mr. Smithers. “They can be easily

seen, and different isotopes can be sorted by hand.”

Scientists are expecting a complete, comprehensive summary of this new element within

two days, including diagrams and collected data tables. “This is the most exciting Chemistry

discovery this century!” exclaimed Mr. Burns.

Objective:

Be able to calculate the amu for beanium.

Materials:

Cup of beans

Triple balance beam

Calculator

Methods:

In this lab your group will be given a cup of beans, which is a representation of all the beanium

isotopes, which exist in nature. Careful, each group has been given different number of beans,

thus no two groups will have the same amu for beanium. Organize your beans, calculate the

37

percentage of each type of bean you have been given; then determine the amu for your beanium

cup.

Safety:

All items for this experiment are natural products. However, they have been exposed to a

chemical environment. Thus, consumption of the beans is prohibited. No lab coat or safety

googles are required for this lab. Closed toed shoes are required, in case a triple balance beam is

dropped. Moreover, throwing any beans around the lab room is grounds for expulsion from the

lab.

Environmental Safety:

There is no environmental safety for this lab, since we will not be disposing of any component of

this lab.

Data Recording:

Total # of beanium atoms in sample ____________

In the table below, draw a picture of the Location and Shape of Object:

Isotope # of atoms of

this isotope

present

Total mass of all

the atoms of this

isotope

Average mass of

this isotope

(Show

calculation in

data processing)

% abundance of

this isotope

(Show

calculation in

data processing)

Whitebeanium

Blackbeanium

Redbeanium

Pintobeanium

(has brown

spots)

Data Processing:

1. Calculate the average mass of each isotope. Remember to record your answers into the

table above and to show all your work.

2. Calculate the percent abundance of each isotope. Remember to record your answers into

the table above and to show all your work.

Challenge Questions

1. Explain why how you could have estimated the amu of your beanium cup before you did

any calculations but after sorting?

2. Explain how another isotope would effect your calculations if it had a percent abundance

of 25%?

Conclusion/Evaluation:

1. Create your conclusion and evaluation of this lab.

38

Spectrophotometry Lab

Background:

We see very differently than we hear. With sound, we are able to pick out many different

frequencies, i.e. different pitches. For example, if we listen to music, we can pick out the drums

and voice separately, even though they are happening at the same time. We don’t have that

capability with light. Instead, we end up seeing one individual color, which most likely is made

up of many different wavelengths of light. The electromagnetic spectrum, shown in Fig. 10.1,

covers a huge range of wavelengths, from gamma rays at 10−14 m to AM radio waves at 104 m.

In this lab we are going to be concerned with the narrow band of wavelengths, ∼ 400 − 750 nm

(a nm = 10−9 m), that make up visible light. In order to know very accurately what wavelengths

are being emitted by a source of light, we will use a digital spectrometer.

Figure 10.1: The electromagnetic spectrum with the visible light region blown up.

Objective:

You will be able to determine the atomic emission spectrum of different lights.

Materials:

Lab book

Calculator

Methods:

In this lab you will be using a spectrophotometer to see the different types of light in which

various light sources give off. Since we do not have an excess amount of light sources, this lab

has been set up as lab stations. Each lab station has space in which you can look at the various

light sources with its own spectrophotometer. DO NOT REMOVE SPECTROPHOTOMETERS

FROM ITS STATION. The light sources you will look at will be sun light, a light bulb, a heat

lamp, and a neon light. Your group will rotate around the four different stations to see and collect

your data. Be aware you will be sharing the stations with your classmates, so please do not move

39

anything in the stations. Furthermore, please move through the stations quickly so that everyone

has a chance to see each station.

Safety:

As you know many different light sources can be harmful to our retinas (a component of our

eyes which enables us to see). Hence, staring at any light source or an extended period of time

can be very harmful. The spectrophotometer is designed to help separate the light into its various

colors, and ensure we are not directly looking at the light source. However, you will still need to

ensure you stare no longer than 30 seconds at any of the light sources in the lab. Furthermore,

since you will need your eyes inhibited (vision blocked) during this lab, you will not need to

wear goggles. A lab coat will also not need to be worn, since you are not working with

chemicals. However, the heat lamp can become very hot. Do Not touch or play with the heat

lamp in any way. Clothing, paper, or spectrophotometer to close to the heat lamp, could catch on

fire, so be sure to stand at least an arm distance away from the lab. Furthermore, do not move,

touch or disrupt any station, since they have been previously prepared for you class, and will be

used in other classes.

Environmental Safety:

The components of light sources can be dangerous to the environment if not disposed of

properly. Luckily much of the components that make up lights can be recycled and reused. Thus,

all light sources will be recycled for the environment.

Data Collection:

Light Source Colors Seen Corresponding

Wavelengths (nm) Observations

Sun Light

Lamp

Heat Lamp

Neon Light

Data Processing:

1. Using Figure 10.1 from the background calculate the frequency at which each light

source gives off. Remember there can be more than one color seen, thus you will need to

calculate every color seen. Show all your work.

Challenge Questions:

1. Explain how wavelength and frequency are related?

2. Explain how wavelength/frequency and energy are related?

3. Which light source gave off more energy? Does this make sense? If so explain, if not

explain.

Conclusion/Evaluation:

1. Write your own conclusion and evaluation for this lab.

40

Flame Test Lab

Background:

A Bunsen burner, named after Robert Bunsen, is a common piece of laboratory

equipment that produces a single open gas flame, which is used for heating, sterilization, and

combustion. The gas can be natural gas (which is mainly methane) or a liquefied petroleum gas,

such as propane, butane, or a mixture of both. The hose barb is connected to a gas nozzle on the

laboratory bench with rubber tubing. The gas then flows up through the base through a small

hole at the bottom of the barrel and is directed upward. There are open slots in the side of the

tube bottom to admit air into the stream via the venturi effect, and the gas burns at the top of the

tube once ignited by a flame or spark. The most common methods of lighting the burner are

using a match or a spark lighter.

The amount of air mixed with the gas stream affects the completeness of the combustion

reaction. Less air yields an incomplete and thus cooler reaction (orange flame), while a gas

stream well mixed with air provides oxygen in an equimolar amount and thus a complete and

hotter reaction (blue flame). The air flow can be controlled by opening or closing the slot

openings at the base of the barrel, similar in function to the choke in a carburettor.

The hottest part of the flame is the tip of the inner flame, while the coolest is the whole

inner flame. Increasing the amount of fuel gas flow through the tube by opening the needle valve

will increase the size of the flame. However, unless the airflow is adjusted as well, the flame

temperature will decrease because an increased amount of gas is now mixed with the same

amount of air, starving the flame of oxygen.

Objective:

You will be able to determine the cation in an unknown solution.

Materials:

Two to three sticks of each metal solution

250 mL beaker

Methods:

Each solution and station has been prepared for all classes. DO NOT mix sticks around the stock

solution, as it will contaminate the solutions for other classes. Each stock solution will have

sticks soaking in them. You will need one stick from each stock solution. Make sure to keep

track of each stick, to ensure you know what metal is producing the particular color. You will

need a trash beaker for all your sticks. Using your 250 mL beaker place at least 50 to 100 mL of

Bunsen burner flames depend on air

flow in the throat holes (on the

burner side, not the needle valve for

gas flow): 1. air hole closed (safety

flame used for lighting or default), 2.

air hole slightly open, 3. air hole half

open, 4. air hole fully open (roaring

blue flame).

41

water to create a proper trash beaker for this lab. All Bunsen burners have been set up. When you

are prepared with all your sticks and trash beaker call Ms. Manakul over and she will light your

flame. Do not play around with the flame, as this will qualify as horseplay (leading to automatic

ejection from the lab). Once you have a medium heat flame, move the stick within the flame

until you see some color change. This may need to be done two to three times (with different

sticks). BE CAREFUL to not light the stick on fire. When you are done with the stick or in case

it catches fire, make sure to place the stick into your trash beaker – flame side into the water.

Once you have observed all metals, dry your sticks and thrown them away. To turn off the

Bunsen burner, turn off the gas first (at the source). Burner may stay out to cool for next class.

Clean trash beaker and place everything away.

Safety:

A Bunsen burner can produce a very hot flame. Thus, full lab gear (lab coat, goggles, and closed

toed shoes) should be worn at all times. Since flames will consume any food source it comes into

contact with, lab coat sleeves should be rolled (not pushed) up at all times. Pushed up sleeves can

fall down at any time, thus rolling sleeves is a better method of safety. Lastly, all used sticks

should be clearly no longer lit before being thrown away.

Environmental Safety:

Since the chemicals are burned clean from the test sticks, no chemicals are released to the

environment when the sticks are thrown away. However, the room will be well vented to ensure

that when the chemicals are vaporized (turned into a gas) are diluted, ensuring less harm to the

environment.

Data Collection:

Stick Main metal Color produced

Data Processing:

1. What is the metal in the unknown?

Challenge Questions:

1. Each solution produces a unique color. Would you expect this result based on the modern

view of the atom?

2. Some commercially available fireplace logs burn with a red and/or green flame. What

elements would be responsible for these colored flames?

3. Aerial fireworks contain gunpowder and chemicals that produce colors. What element

would you include to produce crimson red? Yellow?

Conclusion/Evaluation:

1. Write your own conclusion and evaluation for this lab.

42

Paper Chromatography of Food Dyes Lab

Background:

Paper chromatography is an analytical method used to separate colored chemicals or

substances. It is primarily used as a teaching tool, having been replaced by other chromatography

methods, such as thin-layer chromatography and gel electrophoresis. This is useful for separating

complex mixtures of compounds having similar polarity, for example, amino acids. The setup

has two components. The mobile phase is a solution that travels up the stationary phase, due to

capillary action (the ability of a liquid to flow through narrow spaces even against gravity). The

mobile phase is generally an alcohol solvent mixture, while the stationary phase is a strip of

chromatography paper, also called a chromatogram. The chromatogram is loaded with the

compound which is to be separated. The molecules which are smaller have the ability to move

more quickly through the chromatography paper. This can be quantified by the retardation factor

(Rf) value.

The retardation factor (Rƒ) may be defined as the ratio of the distance traveled by the

substance to the distance traveled by the solvent. Rƒ values are usually expressed as a fraction of

two decimal places.

𝑅𝑓 = 𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑝𝑖𝑔𝑚𝑒𝑛𝑡 𝑚𝑜𝑣𝑒𝑑

𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 𝑚𝑜𝑣𝑒𝑑

If Rƒ value of a solution is zero, the solute remains in the stationary phase and thus it is

immobile. If Rƒ value = 1 then the solute has no affinity for the stationary phase and travels with

the solvent front. To calculate the Rƒ value, take the distance traveled by the substance divided

by the distance traveled by the solvent (as mentioned earlier in terms of ratios). For example, if a

compound travels 9.9 cm and the solvent front travels 12.7 cm, (9.9/12.7) the Rƒ value = 0.779 or

0.78. Rƒ value depends on temperature and the solvent used in experiment, so several solvents

offer several Rƒ values for the same mixture of compound.

Food coloring, or color additive, is any dye, pigment or substance that imparts color

when it is added to food or drink. They come in many forms consisting of liquids, powders, gels,

and pastes.

Objective:

You will be able to infer differences in polarity by observing the separation of food dyes.

Materials:

Chromatogram paper

Sharpies – Black, Blue, and Red

250 mL beaker

Stirring rod

100 mL of Ethanol

Methods:

Obtain a precut chromatogram paper strip. Using a PENCIL

DO NOT USE A PEN (pens consist of similar dyes and

could affect your results, while pencils are made of graphite,

a non-polar molecule, and will not affect your results) draw

a line 1 mm from the bottom of the paper, as you see

represented by the dashed line in the picture to the right.

X X X Black Blue Red

43

Again with a pencil place the labels for the colors we will be using – as shown in the picture.

Where the X’s are (about 2 mm from the bottom of the strip), draw a small dot for each color.

Make sure the dots do not touch – this will make the colors blead and you will not get good

separation. Allow the spots to dry for a minute before continuing. Place the glass stirring rod

through the hole in the chromatogram paper strip; this will allow the strip to stand upright in the

250 mL beaker as seen in the picture below.

Gently pour enough ethanol into the 250 mL so that the top of the ethanol touches the first pencil

line (1 mm from the bottom of the strip); WHILE MAKING SURE NONE OF THE LIQUID

TOUCHES THE PAPER ANYWHERE ELSE!. Wait for the solvent to travel through the

chromatography paper (this will take several minutes). Once the solvent has reached the glass

stirring rod, take the chromatogram out and with a pencil mark a line indicating the distance

traveled by solvent. If you do not mark your end line, you will not know your measurements

after the chromatogram dries, which only takes a few minutes. Once you have collected your

data, you may throw the chromatography paper away, and the solvent may go down the drain

with lots of water.

Safety:

Sharpie dyes are designed to stain organic substances. Since people and clothes are made up of

natural/organic substances, you will need to wear full lab gear (lab coat, goggles, and closed toed

shoes). This is to ensure you don’t walk away very colored.

Environmental Safety:

Most sharpie dyes are made mostly from natural colors. Thus, the use and disposal of these dyes

will show little to no effect to the environment.

Glass stirring rod

44

Data Collection:

Black

Color Produced Distance solvent

traveled

Distance color

traveled

Rf (Remember to

show work in the

data processing

section)

Blue

Color Produced Distance solvent

traveled

Distance color

traveled

Rf (Remember to

show work in the

data processing

section)

Red

Color Produced Distance solvent

traveled

Distance color

traveled

Rf (Remember to

show work in the

data processing

section)

Data Processing:

1. Calculate the Rf values for each color for each pen. Remember to show all work and

place your answers in the correct columns in the data collection section.

2. Based on the Rf values, which dye is the most polar? Which dye is the least polar?

Challenge Questions:

1. Describe what would happen if we used a different solvent like water, vinegar, or

ammonia.

2. Describe what would happen if we used a different chromatogram substance like paper

towels and note-book paper.

Conclusion/Evaluation:

1. Write your own conclusion and evaluation for this lab.

45

Precipitation Lab

Background:

When doing these reactions involving ionic compounds remember they are made of

positive and negative ions held together by the attractive, electrostatic forces that occur between

oppositely charged particles. In water, soluble ionic compounds break apart completely into their

respective ions. Example: NaCl (s) when put into water yields Na+ (aq) and Cl- (aq), and AgNO3

also dissociates in water to form these respective ions Ag+(aq) and NO3-(aq). It turns out that

when these two solutions of sodium chloride and silver nitrate are mixed a solid falls out – a

process called precipitation. This is visually seen in small samples, when two clear solutions

mixed to form a cloudy or opaque looking solution.

Objective:

You will be able to write a set of rules for which compounds produce solid precipitates.

Methods:

On a separate piece of paper copy the grid shown below. Make sure to make the sizes match that

of your well plate. Place your well plate over the grid. Add one to two drops of each chemical in

their respective wells. Record your results in your data collection table/grid. When done

carefully rinse your well plate with lots of water.

AgNO3 (Ag+) Pb(NO3 (Pb+) CaCl2 (Ca2+)

Na2CO3 (CO32-)

Na3PO4 (PO43-)

NaOH (OH)

Na2SO4 (SO42-)

NaCl (Cl-)

Safety:

Since you are working with chemicals which can cause irritation to skin and lungs, full lab gear

(goggles, lab coat, and closed toed shoes) will be worn for the duration of the lab. Should you

come in contact with any of the chemicals, rise with lots of water and see Ms. Manakul for

further instructions.

Solution containing Solution containing Solution containing aqueous NaCl aqueous AgNO3 aqueous NaNO3 and

solid AgCl

46

Environmental Safety:

Since we are using small amounts for this lab, the reactants used and the products produced will

have little to no effect on the environment.

Data Collection:

Fill in the table with a description of what is produced when the two substances are

mixed.

AgNO3 (Ag+) Pb(NO3 (Pb+) CaCl2 (Ca2+)

Na2CO3 (CO32-)

Na3PO4 (PO43-)

NaOH (OH)

Na2SO4 (SO42-)

NaCl (Cl-)

Data Processing:

1. Write complete balanced chemical equations for each reaction. Remember to include

states of matter.

2. Based on this lab, how would you know which ions always created a solid and which

ones always didn’t?

Challenge Questions:

1. Of the balanced chemical equations you wrote in data processing, which one(s) are not

proper chemical reactions? Why would you consider these not proper chemical reactions?

Conclusion/Evaluation:

1. Write your own conclusion and evaluation for this lab.

Appendix

A

Flowchart for Naming Chemical Compounds

How many elements are in the

compound?

Is the first element a metal? Does the formula begin with

NH4?

Two Three or more

The compound is ionic.

Does it contain a transition

metal?

The compound is

covalent. Use the

prefixes to name the first

element (except don’t use

mono- for the first one),

then use the prefixes to

name the second element as

well.

The compound name begins

with ammonium.

Name the metal, then

name the polyatomic

anion using your list. Don’t

change any parts of the

names.

Yes

No

Yes No

How many elements are left

after NH4?

Name the non-metal element, but

change its ending to –ide. The

compound name is ammonium

something-ide.

Name the polyatomic anion. Don’t

change any parts of the name. The compound

name is ammonium something.

One Two or more Name the

metal, then

name the non-

metal, but

change its

ending to

–ide.

Name the metal,

give the charge as a

Roman

numeral in

brackets, then name

the non-metal,

but change its

ending to –ide.

No

Yes

END

END

END

END

END

START

END

B

Flowchart for Writing Formulas for Chemical Compounds

Yes

Does the name contain

prefixes?

Is the first name

ammonium?

The compound is

covalent. Write the

symbol for each element

and use the prefixes to find

the subscripts.

No

Yes

The compound is ionic.

Does the name contain a

Roman numeral?

The compound is ionic and

contains NH41+. Is the

anion name on the list of

polyatomic anions?

No Yes

Determine the charge of the

cation from the periodic table (the

group number).

Use the charge of the polyatomic anion

with the charge of the cation to

balance the charges.

No Yes

The Roman numeral

gives the charge on the

metal ion. Use this

when balancing the

charges.

Is the anion name

on the list of

polyatomic ions?

No

Yes

END

END

No

Write the formula using

subscripts to show how many of

each ion are needed to balance

the charges.

Find the charge of the anion using the periodic

table and use this with the cation charge to

balance the charges. You can draw Lewis dot

(electron transfer) diagrams if you like.

START

C

VESPR Theory

Species Type Geometry Shape Bond Angle Example

A2 Linear Linear 180° H2

AX2 Linear Linear 180° CO2

AX3 Planar

Triangular

Planar

Triangular 120° BF3

AX2E1 Planar

Triangular V-Shape 104.5° SO2

AX4 Tetrahedral Tetrahedral 109.5° CH4

AX3E1 Tetrahedral Pyramidal 109.5° NH3

AX2E2 Tetrahedral V-Shape 104.5 H2O

AX5 Triangular

Bipyramidal

Triangular

Bipyramidal

90°, 120°,

180° PCl5

AX4E1 Triangular

Bipyramidal See Saw

90°,

120°¸180° SF4

AX3E2 Triangular

Bipyramidal T-Shape 90°, 180° ClF3

A3E3 or

AX2E3

Triangular

Bypiramidal Linear 180° I3

-

AX6 Octahedral Octahedral 90°, 180° SF6

AX5E1 Octahedral Square

Pyramidal 90°, 180° BrF5

AX4E2 Octahedral Square Planar 90°, 180° XeF4

D

MOLE MAP

Representative particles (Atoms, molecules, ions use Avogadro’s

number)

MASS (Use average atomic mass off of the

periodic table or GFM by adding all

atomic masses in formula together to

calculate the molar mass)

Volume of a GAS Used only for gases

@ STP

MOLE ALL MOLE RATIOS ARE DONE

HERE

6.022 x 1023 rep part 1 mole

1 mole 6.022 x 1023 rep part

Molar mass g 1 mole

1 mole molar mass g

22.4 L 1 mole

1 mole 22.4 L

Volume of a solid, liquid or gas (Use density)

Density= MASS VOLUME

1 = VOLUME Density MASS

Volume of a Solution (Use Molarity = M)

X mole solute 1 L of solution (aq)

1 L of solution (aq) X mole solute