Download AP Chemistry: Course Introduction Sheet

AP Chemistry Summer Assignment
About this assignment:
This assignment is meant to review the essential skills and knowledge that was covered in Regular/Honors
Chemistry. There are four units to this assignment:
•Unit 1:
•Unit 2:
•Unit 3:
•Unit 4:
Matter and measurement
Atomic Structure
Molecules and Compounds
Chemical Equations and Stoichiometry
Each unit is preceded by a list of learning targets. Pay attention to these targets! Not every learning target
has problems that go with it. Read the textbook! Make sure that after you complete a unit that you feel
confident that you have mastered each target.
It benefits you to take this assignment seriously and master these skills/concepts. We will not have time to
review them during AP Chemistry and if you do not understand them thoroughly then you will struggle all
year.
When This Assignment is due:
This assignment will be due during the first week of school (day 3 or 4) and you will have a test
covering the entire assignment the same week. Use the learning targets as your guide to prepare for the test!
Resources at Your Disposal
You have the following resources available to help you with this assignment
•Your Textbook: You had a textbook assigned to you before summer started. It is an excellent resource
and you should become accustomed to learning from it. We will use it a great deal throughout the year
•Videos: I have posted videos on select topics on my teacher webpage in the Online Resources section.
They go through several examples of the types of problems I have assigned. If you cannot find my webpage,
email me and I will send you the link.
•AP Chemistry Boot Camp: You are highly encouraged to sign up for AP Chemistry Boot Camp. We will
mostly be covering Units 3 and 4 during boot camp, but if time permits I can take questions from Units 1
and 2.
•Email Me: I typically check my school email daily throughout the summer (although I will be away from
the internet a few days here-and-there). If you have a specific question, please feel free to email me and I
will respond as soon as I can.
•Each Other: You are encouraged to study with a partner so you can help each other understand the
material. While studying with a friend will help you complete the assignment quicker, remember that in the
end it is you that is ultimately responsible for understanding the material. So make sure your friend helps
you, but doesn’t do it for you.
Unit 1 -Matter and Measurement
MATH SKILLS LEARNING TARGETS
Metric System
 You know the metric system.
 You know the meaning of the metric prefixes, kilo-, centi-, and milli-.
 You know that there are other metric prefixes and can look them up if needed (micro, mega, pico, etc.)
 You can convert one measurement into another (e.g., 0.532 cg = ______ mg).
 You can convert squared or cubed units (e.g., knowing that 2.54 cm = 1 inch, 38.5 in2 = _____ cm2).
Dimensional Analysis & Showing Your Work
 When you convert one unit to another, you can show your work using dimensional analysis or unit
analysis.
 You know that good examples of dimensional analysis are changing metric units, converting time units,
or using density to convert mass to volume or volume to mass.
 You know that you should always show enough work so that if your answer is incorrect, I can tell where
you went wrong.
Scientific Notation
 You can translate regular numbers into scientific notation and numbers written in scientific notation into
normal notation.
 You know the distinction between exponential notation and scientific notation.
Making Measurements
 You can use a ruler or other measuring device to make a measurement to the correct number of
significant figures, i.e. include all of the digits in the measurement that are a significant part of the
measurement.
 You can correctly assign a  value when making a given measurement.
 You always include a unit on a measurement.
 You know the distinction between a measurement and a defined number (e.g., 12 things in a dozen, pi).
 You can explain the difference between accuracy (how close a measurement is to a true or accepted
value) and precision (how close a set of measurements are to each other).
Significant Figures
 You can determine the number of significant figures in a given measurement (i.e., you know whether a
“0” in a measurement is significant or not.)
 You can determine the precision in a calculation involving measurements when the measurements are
written with the correct number of significant figures.
 You can determine the precision in a calculation involving measurements when the measurements are
written with  notation.
Unit 1 -Matter and Measurement
PROPERTIES OF MATTER
Define the following:
physical property –
chemical property –
physical change –
chemical change –
intensive property –
extensive property element –
compound –
mixture –
Differentiate between the three states of matter.
State
Gas
Liquid
Solid
Picture
Movement
Shape
Volume
Compression
List the commonly used metric prefixes and their meanings.
Unit
Symbol
Giga Mega
Kilo Deci
Centi Milli Micro Nano Pico Femto
Meaning
Perform the following calculations involving density.
Example 1: A rectangular solid has the following dimensions, find it’s density:
length:
2.00 cm
width:
75.0 mm
height:
4.00 cm
mass:
436 grams
Example 2: Assume you had a silver sphere with a mass of 1.50 kg. Calculate the diameter of the sphere
(in cm). The density of silver is 10.5 g/cm3. The formula for the volume of a sphere is . . . V = 4/3 π r3.
Convert temperatures between Celsius and Kelvin.
Formula: K = oC + 273
Example 1: convert 195 Kelvin to Celsius.
Example 2: convert 55C to Kelvin.
DIMENSIONAL ANALYSIS PRACTICE
You can rent a bucket of golf balls at the driving range to practice your swing.
Make the following conversion factor cards and use them to set up the problems below.
Bucket Facts:
1 bucket
918 grams of g.b.
918 grams of g.b.
1 bucket
1 bucket
20 golf balls
20 golf balls
1 bucket
1 bucket
2 Liters of g.b.
2 Liters of g.b.
1 bucket
1000 grams
1 kg
1 kg
1000 grams
1. Calculate the mass of 3.05 buckets of golf balls.
Given: 3.05 buckets x _____________ = ___________________
2.
How many individual golf balls would be in 8.75 buckets?
Given: 8.75 buckets x _____________ = _____________________
3.
Calculate the volume (in Liters) of 300. golf balls.
Given: 300. golf balls x ______________ x ____________________ = ______________________
4.
Calculate the mass of a Liter of golf balls.
Given: 1.00 Liter of golf balls x _________________ x _________________ = _________________
5.
Calculate the mass (in kg) of one golf ball.
Given: 1.00 golf ball x ______________ x _______________ x _________________ = ____________
6.
How many golf balls would you need to have a mass of 575 grams?
7.
What is the volume, in Liters, of 1500 g of golf balls?
8.
How many buckets would you need to hold 375 kg of golf balls?
9.
What is the volume (in Liters) of 155 golf balls?
10.
Determine the volume of 13.2 kg of golf balls.
11.
How many golf balls would have a mass of 19,9 kg?
12.
What is the mass (in grams) of 50.0 L of golf balls?
SCIENTIFIC NOTATION & UNIT ANALYSIS
Change the following to Scientific Notation (maintain the number of significant figures):
1.
5.280 =
_______________
11.
2,560 =
_______________
2.
2,000 =
_______________
12.
.0009 =
_______________
3.
15 =
_______________
13.
8,900,000 =
_______________
4.
6,589,000 =
_______________
14.
.0920 =
_______________
5.
70,400,000,000 = _______________
15.
6,300 =
_______________
6.
.00263 =
_______________
16.
.90 =
_______________
7.
.00589 =
_______________
17.
250 =
_______________
8.
.006 =
_______________
18.
.006087 =
_______________
9.
.400 =
_______________
19.
500,000 =
10.
.08060 =
_______________
20.
.0000000105 =
_______________
_______________
Make the following Metric System conversions using “unit analysis” (you may use scientific notation):
1.
100 mg
_______________ =
_______________ g
2.
20 cm
_______________ =
_______________ m
3.
50 L
_______________ =
_______________ kL
4.
22 g
_______________ =
_______________ cg
5.
825 cm
_______________ =
_______________ km
6.
2,350 kg
_______________ =
_______________ g
7.
19 mL
_______________ =
_______________ cL
8.
52 km
_______________ =
_______________ m
9.
36 m
_______________ =
_______________ cm
10.
18 cm
_______________ =
_______________ mm
11.
6g
_______________ =
_______________ mg
12.
4,259 mg
_______________ =
_______________ g
QUICK NOTES- SIGNIFICANT FIGURES
A. When taking measurements all certain digits plus the first uncertain number are significant.
Example: Your bathroom scale weighs in 10 Newton increments and when you step onto it, the pointer stops
between 550 and 560. Your look at the scale and determine your weight to 557 N. You are certain of the first
two places, 55, but not the last place 7. The last place is a guess and if it is your best guess it also is significant.
B. When given measurements, the numbers that are significant are the digits 1-9 and the 0 when it is not merely a
place holder.
1. When 0’s are between sig. fig., 0’s are always significant.
Example: 101 has 3 sig. fig. and 34055 has 5 sig. fig.
2. When the measurement is a whole number ending with 0’s, the 0’s are never significant.
Example: 210 has 2 sig. fig. and 71,000,000 also has 2 sig. fig.
3. When the measurement is less than a whole number, the 0’s between the decimal and other significant
numbers are never significant (they are place holders).
Example: .0021 has 2 sig. fig. and .0000332 has 3 sig. fig.
4. When the measurement is less than a whole number and the 0’s fall after the other significant numbers, the 0’s
are always significant.
Example: .310 has 3 sig. fig. and .3400 has 4 sig. fig.
5. When the measurement is less than a whole and there is a 0 to the left of the decimal, the 0 is not significant.
Example: 0.02 has only 1 sig. fig. and 0.110 has 3 sig. fig.
6. When the measurement is a whole number but ends with 0’s to the right of the decimal, the 0’s are significant.
Example: 20.0 has 3 sig. fig., 18876.000 has 8 sig. fig.
In case 4 and 6 the 0’s have no effect on the value (size) of the measurement. Therefore, these 0’s must have
been included for another reason and that reason is to show precision of the measurement. Since these 0’s
show precision they must therefore be significant.
In cases 2 and 3 removal of the 0’s DO change the value (size) of the measurement, the 0’s are place holders
and are thus not significant.
In case 5 the 0 is completely unnecessary, it is neither a place holder nor adds to the accuracy of the
measurement.
QUICK NOTES- UNCERTAINTIES IN CALCULATIONS
1. When adding or subtracting numbers written with the  notation, always add the  uncertainties and then round
off the  value to the largest significant digit. Round off the answer to match.
Example: (22.4  .5) + (14.76  .25) = 37.16 .75 = 37.2  .8
The uncertainty begins in the tenths place… it is the last significant digit.
2. When adding or subtracting numbers written in significant figures, show the uncertainty by rounding the answer to
match the largest place with uncertainty.
Example: 267 + 11.8 = 278.8 = 279
The least accurate original measurement is only accurate to the ones place.
4. When multiplying or dividing measurements written in significant figures, show the uncertainty of your
calculations by rounding off your answer to match the same number of significant figures as your least
precise measurement (the measurement with the least number of significant figures).
Example: 477.85  32.6 = 14.657975 = 14.7
32.6 is the least accurate measurement with only 3 significant figures.
NOTE: There are two types of precision: “absolute precision” and “relative precision.”
Example: 322.45 x 12.75 x 3.92 = 16116.051 = 16100
All the measurements are accurate to the hundredth place (absolute precision) but the answer is rounded to 3
significant figures because 3.92 has only 3 significant figures (relative precision).
In Summary:
Adding and Subtracting
Multiplying and dividing
#’s with  notation
Rule 1
Don’t Do This Case
#’s with significant figures
Rule 2
Rule 4
SIGNIFICANT FIGURES & ROUNDING
Indicate the number of significant figures then round each to the number of significant figures indicated.
For example:
1.234
has ______4___ significant figures and, rounded to
2
significant figures, is ___1.2____
1.
0.6034
has __________ significant figures and, rounded to
2
significant figures, is __________
2.
12,700
has __________ significant figures and, rounded to
2
significant figures, is __________
3.
12,700.00 has __________ significant figures and, rounded to
1
significant figures, is __________
4.
0.000983
has __________ significant figures and, rounded to
2
significant figures, is __________
5.
123342.9
has __________ significant figures and, rounded to
5
significant figures, is __________
6.
6.023 x 1023
significant figures, is __________
7.
.005600
has __________ significant figures and, rounded to
1
significant figures, is __________
8.
10000.5006 has __________ significant figures and, rounded to
5
significant figures, is __________
9.
2.0 x 10-3
has __________ significant figures and, rounded to
1
significant figures, is __________
10.
3.456110
has __________ significant figures and, rounded to
3
significant figures, is __________
has __________ significant figures and, rounded to
2
Given calculations with the calculator answer, write the answers with the appropriate number of significant figures.
Example:
6.00 x 3.00
= 18
The answer should be
______18.0_____
1. 23 + 46
= 69
The answer should be
_______________
2. 23.0 + 46.0
= 69
The answer should be
_______________
3. 253 + 345.8
= 598.8
The answer should be
_______________
4. 56 – 35
= 21
The answer should be
_______________
5. 56.00 – 35.0
= 21
The answer should be
_______________
6. 46 x 12
= 552
The answer should be
_______________
7. 3.24 x 5.63
= 18.2412
The answer should be
_______________
8
(2.355 + 2.645) x 10.00 = 50
The answer should be
_______________
9
654  32
The answer should be
_______________
= 20.4375
Unit 2 -Atoms and Elements
LEARNING TARGETS
Parts of the Atom:
The Families of the Periodic Table:
 State the three particles that make up an atom,
their symbol, their charge, their mass, and
their location.
 State the number of protons, neutrons, and
electrons in any atom or ion.
 Explain that isotopes are two atoms with the
same atomic number (number of protons) but
different mass numbers (number of
nucleons—protons + neturons).
 Represent the nucleus with isotopic notation,
220
such as: 86 Rn
 Recognize when two nuclei are isotopes of
each other.
 List the common families of the periodic table
and recognize to which family any element
belongs.
 Recognize metals, non-metals, and metalloids
(semi-metals) on the periodic table.
 State and define the terms conductivity,
malleability, ductility, and sectility.
 State some element facts such as which
elements are too radioactive to exist, which is
the largest non-radioactive element, which
element has the greatest density, and which
element has the highest melting point.
 Explain how Dmitri Mendeleev put together
the periodic table and why we give him credit
for the table even though others were working
along the same lines.
 List the three elements that Mendeleev
predicted and where they are located on the
periodic table.
Molar Mass Calculations:
 Calculate the isotopic mass of an atom given
the resting mass of protons and neutrons.
 Explain that a mole of any element is actually
made up of various isotopes in constant
percentage abundance.
 Calculate the average atomic mass of an
element using the percent abundance and mass
of each isotope.
 Calculate the percent abundance of isotopes
given the average atomic mass and isotopic
masses of an element.
A Little Nuclear Chemistry:
 State that Henri Becquerel discovered
radioactivity and Marie Curie studied it.
 List the three “Becquerel rays” (alpha, beta,
and gamma) and state why alpha particles
were the perfect tool for Ernest Rutherford to
study the structure of atoms.
 State that the alpha particle is the same as a
helium nucleus, a beta particle is a high-speed
electron, and a gamma ray is a high-energy
form of light.
Unit 2 -Atoms and Elements
ATOMIC STRUCTURE
Summarize the sub-atomic particles
proton
neutron
symbol
charge
location
mass
electron
The mass of the atom is due to the _____________________________
The size of the atom is due to the __________________
How Many Particles in Each Atom?
The particle that defines the identity of an atom is the _____________
Every hydrogen atom has ___ proton.
Every magnesium atom has ___ protons.
Any atom that has 23 protons is _________________.
Any atom that has 92 protons is _________________.
The mass of an atom is mostly from the ___________ and ____________.
Find O on the periodic table. It’s mass is ______ amu. O has ___ protons. It must have ___ neutrons.
Electrically neutral atoms (as opposed to ions) have one electron for every proton.
Fill in this chart for these neutral atoms:
Atom
Mass
protons neutrons electrons
He
Si
Be
H
Rn
Ar
F
If the mass is not close to a whole number, it is because the atom has several _____________. These are
atoms with the same number of ___________ but different numbers of _____________.
Chlorine has two isotopes: Cl-35 ( ___ p+ & ___ n) and Cl-37 ( ___ p+ & ___ n).
THE NUCLEAR ATOM
ALL of the answers to this worksheet can be logically figured out by looking at the Schematic Diagrams for
Various Atoms , the Periodic Table, and discussing with your partners. All of the information you need is
here somewhere. Determine each answer and be able to give convincing reasons for each answer. Good luck.
C?
1.
How many protons are found in
12
13
13
2.
How many neutrons are found in
12
13
13
3.
How many electrons are found in
12
13
13
4.
Based on the model,
a)
what do all carbon atoms (and ions) have in common?
b)
C?
C?
C?
C?
C?
C?
C?
C?
what do all hydrogen atoms (and ions) have in common?
5.
What is the significance of the atomic number, Z, above each atomic symbol in the periodic chart?
6.
What do all nickel (Ni) atoms have in common?
7.
How is the mass number, A, determined?
8.
What structural feature is different in isotopes of a particular element?
9.
a)
What feature distinguishes a neutral atom from an ion?
b)
How is the charge on an ion determined?
10.
Where is most of the mass of an atom, within the nucleus or outside of the nucleus?
Explain your reasoning.
11.
Complete the following table:
Isotope
31
18
P
Atomic
Number
Z
O
Ni2+
Number of
electrons
15
8
19
58
Mass
Number
A
39
18
58
ISOTOPES
1.
Give the mass number of each of the following atoms:
(a) an iron atom with 30 neutrons
(b) an americium atom with 148 neutrons
(c) a tungsten atom with 110 neutrons
2.
Give the complete symbol ( AZ X ) for each of the following atoms:
(a) nitrogen with 8 neutrons
(b) zinc with 34 neutrons
(c) xenon with 75 neutrons
3.
How many electrons, protons, and neutrons are there in an atom of:
(a) carbon-13, 13 C
(b) copper-63, 63 Cu
(c) bismuth-205, 205 Bi
4.
Fill in the blanks in the table (one column per element).
65
86
Symbol
Cu
Kr
Number of protons
Number of neutrons
Number of electrons
in the neutral atom
Name of element
78
117
46
36
5.
Radioactive americium-241 is used in household smoke detectors and in bone mineral analysis. Give the
number of electrons, protons, and neutrons in an atom of americium-241.
6.
Verify that the atomic mass of magnesium is 24.31 amu, given the following information:
24
Mg , mass = 23.985042 amu; percent abundance = 78.99%
25
Mg , mass = 24.985837 amu; percent abundance = 10.00%
26
Mg , mass = 25.982593 amu; percent abundance = 11.01%
7.
Copper has two stable isotopes, 63 Cu and
8.
Strontium has four stable isotopes, Strontium-84 has a very low natural abundance, but
Cu , with masses of 62.939598 amu and 64.927793 amu,
respectively. Calculate the percent abundances of these isotopes of copper.
88
9.
65
86
Sr are all reasonably abundant. Which of these more abundant isotopes predominates?
There are three naturally occurring isotopes of neon:
Sr , 87 Sr , and
neon-20
neon-21
neon-22
mass 19.9924 amu
mass 20.9940 amu
mass 21.9914 amu
abundance 90.84%
abundance 0.260%
abundance 8.90%
a. Without calculation, what is the approximate atomic mass of neon?
b. Calculate the actual atomic mass.
10.
Uranium has an atomic mass equal to 238.0289. It consists of two isotopes: uranium-235 with an
isotopic mass of 235.044 amu and uranium-238 with an isotopic mass of 238.051. Calculate the % abundance
of the uranium-235 isotope.
11.
From amongst the elements sodium, chlorine, nickel, argon, calcium, uranium, and oxygen, select the
alkali metal, the alkaline earth metal, the transition metal, the actinide, the halogen, the noble gas, and the
chalcogen (Group 6A).
Unit 3 - Molecules & Compounds
LEARNING TARGETS
I can:
Formulas
 Look at a formula and state how many
elements and atoms are in that compound.
 Calculate the molecular mass or molar
mass of any compound.
 State that the mass of a molecule is
measured in amu’s and the mass of a mole
is measured in grams.
 Give examples of empirical formulas,
molecular formulas, and structural
formulas.
 Identify a formula as empirical,
molecular, or structural.
 Calculate the percent composition (by
mass) for any compound.
 Calculate the empirical formula from
percent composition data.
 Determine the molecular formula of a
compound given its empirical formula
and molar mass.
Ionic Compounds
I can:
 Name and determine the charge of a
monoatimic ion from its place on the
periodic table
 List the names the common polyatomic
ions.
Hydrates
 Give examples of hydrates and anhydrous
compounds.
 Calculate the formula of a hydrate from
 State whether a compound is an ionic
compound or a nonmetal compound.
 Write the formula of an ionic compound
given the two ions or its name. Know
when to use parentheses.
 Name an ionic compound given the
dehydration data.
The Mole
 State the significance of the mole.
 State the three mole facts for any
substance (molar volume, molar mass,
Avogadro’s number)
formula.
 Determine the charge on an ion from
1 mole = 22.4 Liters @ STP (gases only)
1 mole = 6.02 x 1023 particles
(particles = molecules or atoms)
1 mole = gram molecular mass of chemical
information in an ionic formula.
Nonmetal Compounds
aka Molecular Compound
 Write the formula of a binary nonmetal
compound (molecular compound) given
its name.
 Name a binary nonmetal compound
(molecular compound) given its formula.
Percent Composition
 Use dimensional analysis to convert
between moles, mass, volume, and number
of particles for a chemical.
 Use density as a conversion factor in mole
problems.
 Use gas density to calculate molar mass.
Unit 3 -Molecules & Compounds
QUICK NOTES
THE PERIODIC TABLE
Groups & families
Periods
= vertical columns
= horizontal rows
Group IA
Group IIA
Group VIA
Group VIIA
Group VIIIA
Alkali metals
Alkaline Earth metals
Chalcogens
Halogens
Noble gases / rare gases / inert gases
Metals - elements found on the left side of the “stair case” on the periodic table as well as the Lanthanoids and
Actinides on the bottom, good conductors of heat & electricity, ductile, malleable, solids at room temperature
(except Hg)
Nonmetals - elements found on the right side of the staircase, gases, liquid, & solid; usually poor conductors
and are brittle
Metalloids - elements that lie along staircase which have properties of both metals and nonmetals (except Al,
which is usually considered a metal)
IONS
Cations = ions with a positive charge formed by a metal atom losing one or more electrons
Na (atom)
11 protons
11 electrons
Na+ (ion)
11 protons
10 electrons
Anions = ion with a negative charge formed by a nonmetallic atom gaining 1 or more electrons
Charges of common monatomic ions
Group 1A ions
+1 charge
Group 2A ions
+2 charge
Group 3A ions
+3 charge (usually just aluminum)
Group 5A ions
-3 charge (usually just nitrogen, sometimes P)
Group 6A ions
-2 charge
Group 7A ions
-1 charge
DIATOMICS
Elements found as diatomic molecules in nature include:
hydrogen, oxygen, fluorine, bromine, iodine, nitrogen, chlorine
H2
O2
F2
Br2
I2
N2
Cl2
These were discovered by Prof. HOFBrINCl or was his name BrINClHOF?
QUICK NOTES- POLYATOMIC NAMING PRACTICES
Polyatomic ions that don't appear on the above tables do NOT always follow these naming practices. If you can
remember the formula of the ion whose name ends with ate, you can usually work out the formulas of the
other family members as follows:
modify stem name with:
meaning
examples
-ate
a common form, containing oxygen
chlorate, ClO3nitrate, NO3sulfate, SO42-
-ite
one less oxygen than -ate form
chlorite, ClO2sulfite, SO32nitrite, NO2-
per-, -ate
same charge, but contains one more
oxygen than -ate form
perchlorate, ClO4perbromate, BrO4-
hypo-, -ite
same charge, but contains one less oxygen hypochlorite, ClOhypobromite, BrOthan the -ite form
thio-
replace an O with an S
thiosulfate, S2O32thiosulfite, S2O22-
Some anions can capture hydrogen ions. For example, carbonate (CO32- can capture an H+ to produce hydrogen
carbonate HCO3- (often called bicarbonate). Each captured hydrogen neutralizes one minus charge on the anion.
modify stem name with: meaning
examples
hydrogen
or bi-
(1) captured H+ ions hydrogen carbonate, HCO3- (a.k.a. bicarbonate)
hydrogen sulfate, HSO4- (a.k.a. bisulfate)
dihydrogen
(2) captured H+ ions dihydrogen phosphate, H2PO4-
WRITING FORMULAS AND NAMING COMPOUNDS
QUICK NOTES
Writing formulas and naming compounds can be confusing because there are different types of compounds that
follow different rules. Additionally, some compounds (H2O, NH3, CH4, etc.) simply have common names that
must be memorized.
The two types of compounds we will focus on first are ionic compounds (formed from positive and negative
ions) and binary nonmetal compounds (molecular compounds). Later we will add acids. So… you must
recognize the type of compound before you try to name it. [Note: + ion = “cation” and – ion = “anion”.]
Ionic
Formula
Naming
I.
+ ion before – ion
ex: NaCl (NH4)2SO4 Al2S3
Name of cation + name of anion
sodium chloride
ammonium sulfate
aluminum sulfide
Binary Nonmetal
usually the less electronegative atom is first
ex: CO
CO2
N2O
Indicate the number (mono, di, tri, and kind of atoms.
First element is simply name of element. Second
element name ends with “ide”
carbon monoxide
carbon dioxide
dinitrogen monoxide
Writing Ionic Formulas
Cl
NO3
S2
CO32
N3
PO43
Na+
NH4+
Sn2+
Hg22+
Al3+
Sn4+
II.
Naming Ionic Compounds
Cation
Anion
Cu2+
OH
Ba2+
SO42
NH4+
Cr2O72
Ag+
C2H3O2
Fe3+
S2
Formula
Name
OH
III. Write the appropriate number next to each prefix
mono
di
tri
tetra
penta
hexa
IV.
hepta
octa
deca
Writing Formulas of Binary Nonmetal Compounds
Name
Formula
Name
Formula
nitrogen trifluoride
phosphorus trichloride
nitrogen monoxide
phosphorus pentachloride
nitrogen dioxide
sulfur hexafluoride
dinitrogen tetroxide
disulfur decafluoride
dinitrogen monoxide
xenon tetrafluoride
V.
nona
Naming Binary Nonmetal Compounds
Name
VI.
Formula
Name
Formula
CCl4
HBr
P4O10
N2F4
ClF3
XeF3
BCl3
PI3
SF4
SCl2
Practice for Both Types of Compounds
Formula
Name
Formula
Name
HCl
carbon dioxide
PCl5
ammonium carbonate
K2S
sulfur dichloride
NiSO4
calcium iodide
ClF3
boron trifluoride
OF2
phosphorus triiodide
Al(OH)3
magnesium perchlorate
NCl3
potassium permanganate
(NH4)3PO4
aluminum phosphate
S2Cl2
dioxygen difluoride
M O L A R
M A S S
&
%
C O M P O S I T I O N
I. Molar Masses
Given a periodic table, you should be able to calculate the molecular mass (in u’s) or the molar mass (in grams)
for any element or compound.
Examples: (give answers to two decimal places)
H2SO4
Cl2
CO2
N2O
Ca(OH)2
HC2H3O2
NaOCl
Al2S3
II. Fraction and Percent Composition
It is useful to determine how much of a compound’s mass is made up of each element. Water, H2O, for
example has a molar mass of 18.02 g. The H’s mass is 2(1.0079) = 2.02 g. The O’s mass is 16.00 g.
2.02
16.00
We can set up fractions for each element: H =
= 0.112 = 11.2%.
O=
= 0.888 = 88.8%.
18.02
18.02
This is called the percent composition. The fraction composition is a good in-between step.
Determine the fraction and percent composition of each element below (answer to one decimal place):
1. H2SO4
2. Ca(OH)2
3. HC2H3O2
4. CO2
5. N2O
6. NaOCl
7. Al2S3
HYDRATES & COMPOSITION PROBLEMS
1.
Cupric chloride, CuCl2, when heated to 100C is dehydrated. If 0.235 g of CuCl2 · x H2O gives 0.185 g of
CuCl2 on heating, what is the value of x?
2.
The “alum” used in cooking is potassium aluminum sulfate hydrate, KAl(SO4)2 · x H2O . To find the value
of x, you can heat a sample of the compound to drive off all of the water and leave only KAl(SO4)2.
Assume you heat 4.74 g of the hydrated compound and that the sample loses 2.16 g of water. What is the
value of x?
3.
If “Epsom salt,” MgSO4 · x H2O is heated to 250C, all the water of hydration is lost. On heating a 1.687g sample of the hydrate, 0.824 g of MgSO4 remains. What is the formula of Epsom salt?
4.
When CaSO4 · x H2O is heated, all of the water is driven off. If 34.0 g of CaSO4 (molar mass = 136) is
formed from 43.0 g of CaSO4 · x H2O, what is the value of x?
Mole Calculations - Difficulty Level 1
1 mole = 6.02 x 1023 molecules = 22.4 L (@ STP)
1.
Calculate the mass of 1.58 moles CH4. [molar mass CH4 = 16.0 g/mol]
G: 1.58 moles CH4
D: ? g CH4
1.58 moles CH4
=
2.
What volume will 7.29 moles of CO2 gas occupy at STP?
G: 7.29 moles CO2
D: ? L CO2
7.29 moles CO2
=
3.
How many molecules are there in a 0.00583 mole sample of H2O?
G: 0.00583 moles H2O
D: ? molecules H2O
0.00583 moles H2O
=
Mole Calculations - Difficulty Level 2
4.
What mass of CO2 gas occupies a volume of 395 Liters at STP? [molar mass CO2 = 44.0 g/mol]
G:
D:
=
5.
How many molecules are in a 0.250 gram sample of H2O? [molar mass H2O = 18.0 g/mol]
G:
D:
=
6.
What volume will 3.01 x 1022 molecules of CH4 occupy at STP?
G:
D:
=
Mole Calculations - Difficulty Level 3
1 mole = 6.02 x 1023 molecules = 22.4 L (@ STP)
1.
Calculate the mass of 7.23 moles CH4. [molar mass CH4 = 16.0 g/mol]
G:
D:
2.
What volume will 9.35 moles of CO2 gas occupy at STP?
G:
D:
3.
How many molecules are there in a 0.0752 mole sample of H2O?
G:
D:
4.
What mass of CO2 gas occupies a volume of 10.8 Liters at STP? [molar mass CO2 = 44.0 g/mol]
G:
D:
5.
How many molecules are in a 1.44 gram sample of H2O? [molar mass H2O = 18.0 g/mol]
G:
D:
6.
What volume will 1.21 x 1024 molecules of CH4 occupy at STP?
G:
D:
Unit 4 -Chemical Equations and Stoichiometry
LEARNING TARGETS
I can:
Chemical Equations
Limiting Reactant Problems



Give examples of products and reactants in a
chemical equation.
State that Antoine Lavoisier introduced the
law of conservation of matter.
Combustion



State that combustion is another name for
burning.
Write an equation for a combustion reaction
given only the fuel that is burned.
Correctly label substances in an equation as
solid (s) , gas (g), liquid (l), or aqueous (aq)
Balancing Equations



Balance equations by adding coefficients.
Recognize when an equation is balanced.
State that the formulas of reactants and
products should not be changed in order to
balance equations.
Stoichiometry Problems



Use the stoichiometric factor (mole ratio) to
convert from moles of one substance to moles
of a different substance. (i.e. In the equation:
N2 + 3H2  2NH3,
3 mol H2  2 mol NH3)
Convert between the quantities of mass,
volume, molecules and moles using
dimensional analysis
(i.e. use 1 mol = 22.4 L, 1 mol = 6.02 x 1023
molecules, and 1 mol = gram molecular mass)



Recognize that a problem with two “given
values” is a limiting reactant problem.
Determine the limiting reactant and excess
reactant in a problem.
Solve problems involving Limiting Reactants
Calculate how much excess chemical is left
over after a reaction.
Percent Yield Problems



Use stoichiometry to calculate the theoretical
yield (mass of a product) in a problem.
State that actual yields are usually given in a
problem.
Use the theoretical yield and actual yield to
calculate the percent yield.
Chemical Analysis Problems



Calculate the mass of each element in a given
compound given data such as the masses of
CO2 and H2O formed in a combustion
reaction.
Use mass and mole information to calculate
the empirical formula of an unknown
substance.
Use percent composition to equalize mass
and mole information derived from different
samples.
Show the units of molar mass as grams/mol or
g·mol-1.
Unit 4 -Chemical Equations and Stoichiometry
COMBUSTION EQUATIONS
For burning to occur, you need a fuel, an oxidizer, and heat. When hydrocarbons are the fuel and O2 in the air is
the oxidizer, then CO2 and H2O are the products.
Example:
Write the balanced equation for the complete combustion of propane, C3H8, in air.
Solution:
First, set up the basic equation. You memorize the “+ O2  CO2 + H2O” part.
C3H8 + O2  CO2 + H2O
Next, balance. 3 C’s in C3H8 result in 3CO2’s; 8 H’s in C3H8 result in 4 H2O’s;
C3H8 + __ O2  3 CO2 + 4 H2O
Total O’s on the product side = 10 [(3 x 2) + (4 x 1)] = total O’s on the reactant side.
This would mean that 5 O2’s were involved.
Tip: If an UNEVEN number of O’s need to be represented, a fraction should be used. 7 O’s = 7/2 O2
Tip: Take into account fuels that contain oxygen. Subtract the O’s from that represented as O2’s
Practice: Write the balanced combustion equations for the following substances.
1.
CH4
2.
C5H12
3.
C9H20
4.
C2H6
5.
C8H18
6.
C4H10
7.
C2H5OH
8.
C3H7OH
9.
HC2H3O2
10.
CH3COCH3
BALANCING EQUATIONS
Balance the following chemical equations:
1.
__ZnS + __HCl  __ZnCl2 + __H2S
2.
__HCl + __Cr  __CrCl2 + __H2
3.
__Al + __Fe3O4  __Al2O3 + __Fe
4.
__H2 + __Br2  __HBr
5.
__Na2S2O3 + __I2  __NaI + __Na2S4O6
6.
__LaCl3 + __Na2CO3  __La2(CO3)3 + __NaCl
7.
__NH4Cl + __Ba(OH)2  __BaCl2 + __NH3 + __H2O
8.
__Ca(OH)2 + __H3PO4  __Ca3(PO4)2 + __H2O
9.
__La2(CO3)3 + __H2SO4  __La2(SO4)3 + __H2O + __CO2
10.
__Na2O + __(NH4)2SO4  __Na2SO4 + __H2O + __NH3
11.
__C4H10 + __O2  __CO2 + __H2O
12.
__C7H6O2 + __O2  __CO2 + __H2O
13.
__P4O10 + __H2O  __H3PO4
14.
__FeS2 + __O2  __Fe2O3 + __SO2
15.
__NH3 + __O2  __NO + __H2O
16.
__Fe + __HCl  __H2 + __FeCl2
17.
__PbO2 + __HCl  __H2O + __PbCl2 + __Cl2
18.
__Fe2O3 + __H2SO4  __Fe2(SO4)3 + __H2O
19.
__NO2 + __H2O  __NO + __HNO3
20.
__C2H6S + __O2  __CO2 + __H2O + __SO2
STOICHIOMETRY
General Stoichiometry
1.
Several brands of antacid tablets use aluminum hydroxide to neutralize excess acid.
Al(OH)3(s) + 3 HCl(aq)  AlCl3(aq) + 3 H2O(l)
[Molar masses:
78.01
36.46
133.4
18.02]
What quantity of HCl, in grams, can a tablet with 0.750 g of Al(OH)3 consume? What quantity of water is
produced?
2.
The equation for one of the reactions in the process of reducing iron ore to the metal is
Fe2O3(s) + 3 CO(g)  2 Fe(s) + 3 CO2(g)
[Molar masses:
159.7
28.01
55.85
44.01]
(a) What is the maximum mass of iron, in grams, that can be obtained from 454 g (1.00 lb) of iron(III)
oxide?
(b) What mass of CO is required to reduce the iron(III) oxide to iron metal?
Limiting Reactants
3.
The reaction of methane and water is one way to prepare hydrogen:
CH4(g) + H2O(g)  CO(g) + 3 H2(g)
[Molar masses:
16.04
18.02
28.01
2.02]
If you begin with 995 g of CH4 and 2510 g of water, what is the maximum possible yield of H2?
4
Disulfur dichloride, S2Cl2, is used to vulcanize rubber. It can be made by treating molten sulfur with
gaseous chlorine:
S8(l) + 4 Cl2(g)  4 S2Cl2(l)
[Molar masses: 256.6
70.91
135.0]
Starting with a mixture of 32.0 g of sulfur and 71.0 g of Cl2, which is the limiting reactant? What mass of
S2Cl2 (in grams) can be produced? What mass of the excess reactant remains when the limiting reactant is
consumed?
Percent Yield
5.
Diborane, B2H6, is a valuable compound in the synthesis of new organic compounds. One of several ways
this born compound can be made is by the reaction
2 NaBH4(s) + I2(s)  B2H6(g) + 2 NaI(s) + H2(g)
[Molar masses:
37.84
253.8
27.67
149.9
2.02]
Suppose you use 1.203 g of NaBH4 with an excess of iodine and obtain 0.295 g of B2H6. What is the
percent yield of B2H6?
6.
Disulfur dichloride, which has a revolting smell, can be prepared by directly combining S8 and Cl2, but it
can also be made by the following reaction:
3 SCl2(l) + 4 NaF(s)  SF4(g) + S2Cl2(l) + 4 NaCl(s)
[Molar masses:
103.0
41.99
108.1
135.0
58.46]
Assume you begin with 5.23 g of SCl2 and excess NaF. What is the theoretical yield of S2Cl2? If only
1.19 g of S2Cl2 is obtained, what is the percent yield of the compound?
CHEMICAL ANALYSIS
Show all of your work (on a separate sheet of paper if necessary)
Chemical Analysis
32. A mixture of CuSO4 and CuSO4 • 5H2O has a mass of 1.245 g, but, after heating to drive off all the water,
the mass is only 0.832 g. What is the weight percent of CuSO4 • 5H2O in the mixture?
34. A 1.25-g sample contains some of the very reactive compound Al(C6H5)3. On treating the compound with
aqueous HCl, 0.951 g of C6H6 is obtained.
Al(C6H5)3(s) + 3HCl(aq)  AlCl3(aq) + 3C6H6(l)
Assuming that Al(C6H5)3 was converted completely to products, what is the weight percent of Al(C6H5)3 in
original 1.25-g sample?
Determination of Empirical Formulas
36. Styrene, the building block of polystyrene, is a hydrocarbon, a compound consisting only of C and H. If
0.438 g of styrene is burned in oxygen and produces 1.481 g of CO2 and 0.303 g of H2O, what is the
empirical formula of styrene?
38. Menthol, from the oil of mint, has a characteristic cool taste. The compound contains only C, H, and O. If
95.6 mg of menthol burns completely in O2, and gives 269 mg of CO2 and 110 mg of H2O, what is the
empirical formula of menthol?
40. Silicon and hydrogen form a series of compounds with the general formula SixHy. to find the formula of one
of them, a 6.22-g sample of the compound is burned in oxygen. On doing so, all of the Si is converted to
11.64 g of SiO2 and all of the H to 6.980 g of H2O. What is the empirical formula of the silicon compound?