Download AP Chemistry Summer Assignment

Name: ________________________
Oregon City High School AP Chemistry Summer Assignment
Welcome to AP chemistry! This summer homework is meant to be a review of the material covered in Chemistry
A and B as well as an introduction to the a few concepts in the first three chapters of the AP Chemistry Textbook that we
haven’t covered yet. Having the following skills will be essential to your success in AP Chemistry and I will expect that
you already have a firm grasp on these topics as we start the year. The following assignment is to be completed over the
summer and brought in COMPLETED on the first day of class. NO LATE PAPERS WILL BE ACCEPTED!!! Because
this information is part of the first unit of study you will be responsible for all of these concepts on your first test.
While you may need to reference materials to help remind you how to do some of these problems (your notes
from chemistry A and B, your AP Chemistry textbook, the internet, etc.) please make sure that your work is YOUR OWN
as you will be the one responsible for understanding this information.
Included is a copy of the periodic table used in AP Chemistry. Notice that this is not the table used in first year
chemistry. The AP table is the same that the College Board allows you to use on the AP Chemistry test. Notice that it has
the symbols of the elements but not the written names. You may want to spend some time familiarizing yourself with the
names and symbols of each element.
About the AP Chemistry Course:
Since this is a college level course taught in high school, it is very demanding, both in time and effort required. Much
of the work involves solving math-type story problems. Homework is assigned each day through all three trimesters. The
three weeks before the AP Exam in May will be used for review. The amount of work outside of class depends upon the
student and his/her background; however, students should be prepared to spend anywhere from 45 minutes to an hour
each night after school on just their Chemistry homework. Those students who are heavily involved in after school
activities and / or jobs will have to learn to budget their time very carefully.
Why Take AP Chemistry?
There are several reasons why a student might want to take AP Chemistry, including (but not limited to!) the following:
1. AP Chemistry will challenge you to the limits of your academic ability. In the past you may have found classes
"too easy", and therefore may not have stimulated you to do your very best. This will not be the case in AP Chemistry.
2. AP Chemistry should allow you to achieve college credit while still enrolled in high school. This will save time and
money. Some students who have passed the AP Exam elect to take first year college chemistry anyway, where they
find the material an easy review, and achieve top grades while others around them are frustrated and struggling in a
class which is too large and/or the instructor is unavailable for help.
3. AP Chemistry looks great on your transcript or on a letter of recommendation. More and more colleges and
universities are looking for ways a student has distinguished themselves in high school. Being a "straight A" student
no longer carries the weight it once did. Taking AP Chemistry is a way of distinguishing yourself in high school.
4. While AP chemistry will be challenging and time consuming (sorry… it is a college class… :/) we will get to
perform some fun labs, get to know each other really well, have an unbelievably deep understanding of chemistry, and
we will get to do some fun projects (including a field-trip near the end).
I look forward to seeing you all at the beginning of the next school year! If you need to contact me during the summer,
you can email me and I will get back to you as soon as possible.
Jenaya Hoffman
OCHS AP Chemistry Teacher
[email protected]
AP CHEMISTRY CHAPTER 1 REVIEW: MATTER AND MEASUREMENT
1. SI Units and the Metric System
Every measurement has two parts: a number and a unit.
SI System – 1960 an international agreement was reached to set up a system of units so scientists everywhere
could better communicate measurements. Le Système International (SI) was all based on the metric system.
Prefix + base unit
Exp: 1 kg = 1000g
1,000,000 micrograms in 1 gram
Prefix tells you the power of 10 to multiply by - decimal system -easy conversions
*Density of water = 1 g/mL
lL
Length
How many millimeters, mm, are in a meter? _________ Find an object that is 1 mm thick: _________
How many centimeters, cm, are in a meter? __________ Find an object that is 2.5 cm in diameter:_________
Figure out YOUR height: ___________ m ____________ cm ____________ mm
Mass
How many grams are in a kilogram? _________
A 150 lb person has a mass of __________ kg.
Hint: 1 kg = 2.2 lbs
Volume – derived from length see picture above.
Consider a cube 1m on each edge ∴1.0m3
- a decimeter is 1/10 of a meter so (1m)3 = (10dm)3 = 1,000 dm3
- 1dm3 = 1 liter (L)
How many cm3 are in 1 m3? ______________
If 1 cm3 is equal to 1 mL how many cm3 are in a liter? _____________
Using a measuring cup at home find the volume of water (in mL) that will fit in a pop can: _________
If 1 mL of water has a mass of 1 g, how many grams of water are in your pop can? _________
2. Temperature Conversions
Temperature is a measure of the average kinetic energy.
Different temperature scales are all talking about the same height of mercury.
Notice 1 °C is equal to 1 K
K = °C + 273
°C = K – 273
Derivation for the equation
for converting ºF to ºC below
Normal body temperature is 98.6°F. Convert this temperature to the Celsius and Kelvin scales.
____________ °C
____________ K
Liquid nitrogen, which is often used as a coolant for low-temperature experiments, has a boiling point of 77 K.
What is this temperature in degrees Celsius?
____________ °C
One interesting feature of the Celsius and Fahrenheit scales is that -40°F and -40°C represent the same
temperature. Verify that this is true. Show your work.
3. Density
Determining Density
A chemist, trying to identify the main component of a compact disc cleaning fluid, finds that 25.00 cm3 of the
substance has a mass of 19.625 g at 20°C. The following are the names and densities of the compounds that
might be the main component.
Compound
Chloroform
Diethyl ether
Ethanol
Isopropyl alcohol
Toluene
1.492
0.714
0.789
0.785
0.867
Which of these compounds is the most likely to be the main component of the compact disc cleaner?
__________________
An empty container weighs 121.3 g. Filled with carbon tetrachloride (density 1.53 g/cm3 ) the container weighs
283.2 g. What is the volume of the container?
__________________
A student has a cube of aluminum that measures 4 cm wide on each side. What is the volume of this cube?
(Volume = length x width x height)
__________________
When the student massed the cube on a scale they found that the cube of aluminum had a mass of 165 g.
What is the density of this aluminum cube?
__________________
Percent error is a measure of how inaccurate a measurement is.
Percent Error = |Your Value – Accepted Value| x 100%
Accepted value
Using a computer, look up the accepted value for the density of aluminum and calculate your percent error for
your calculated density.
Percent error for Al ___________%
4. Accuracy, Precision, and Types of Error
Accuracy – correctness; agreement of a measurement with the true value.
Precision – reproducibility; degree of agreement among several measurements.
Random error – equal probability of a measurement being high or low.
Systematic error – occurs in the same direction each time. Usually due to equipment or experimental flaw.
The results of several dart throws show the difference between precise and accurate.
(a) Neither accurate nor precise (large random errors).
(b) Precise but not accurate (small random errors, large systematic error).
(c) Bull’s-eye! Both precise and accurate (small random errors, no systematic error).
To check the accuracy of a graduated cylinder, a student filled the cylinder to the 25-mL mark using water
delivered from a buret and then read the volume delivered. Following are the results of five trials:
Is the graduated cylinder accurate?
Is the graduated cylinder precise?
What type of error occurred, random or systematic?
5. States of Matter
Step 1.) In the circles below, draw what water molecules (H2O) would look like in the three different
states of matter.
Step 2.) Describe the speed at which the molecules move in each state of matter.
Solid
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
Liquid
____________________________
____________________________
____________________________
____________________________
____________________________
____________________________
____________________________
____________________________
____________________________
____________________________
____________________________
____________________________
Gas
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
___________________________
What is a vapor and how does it differ from a gas?
_________________________________________________________________________________________
Give three examples of vapors below.
___________________
___________________
___________________
Give three examples of gases below.
___________________
___________________
___________________
6. Dimensional Analysis
Use conversion factors to change the units.
Each conversion factor = 1
1 foot = 12 inches (equivalence statement)
12 in
= 1 ft
=
1 ft
1 ft
12 in
72 ft | 12 in = 864 in
| 1 ft
To convert a unit, multiply your quantity by a conversion factor that “cancels” the undesirable unit and puts the
desired unit in the numerator. (Multiply all the numbers on top) ÷ (multiply all numbers on the bottom)
Common Conversions
1 inch (in) = 2.54 centimeter (cm)
1 foot (ft) = 12 inches (in)
1 yard (yd) = 3 feet (ft)
1 mile (mi) = 1760 yards (yd)
1 meter (m) = 3.28 feet (ft)
1 mile = 1.61 kilometer (km)
1 pound (lb) = 453.6 grams (g)
2.205 pounds (lb) = 1 kilogram (kg)
1 ounce = 28 grams
60 seconds (s) = 1 minute (min)
60 minute (min) = 1 hour (hr)
24 hours (hr) = 1 day
365 days = 1 year
1 fortnight = 14 days
100 years = 1 century
1 cup (c) = 16 tablespoons (tbsp)
1 gallon = 4 quarts
1000 milliliters (mL) = 1 liter (L)
1.057 quarts = 1 liter (L)
1 kilometer (km) = 0.62 miles (mi)
1000 years = 1 millennium
1 tablespoon (tbsp) = 3 teaspoons (tsp)
3
1 cubic centimeter (cm ) = 1 milliliter (mL)
3.785 liters (L) = 1 gallon (gal)
A pencil is 7.00 in. long. What is its length in centimeters?
You want to order a bicycle with a 25.5-in. frame, but the sizes in the catalog are given only in centimeters.
What size of frame (in cm) should you order?
A student has entered a 10.0-km run. How long is the run in miles?
The speed limit on many highways in the United States is 55 mi/h. What number would be posted in kilometers
per hour?
A Japanese car is advertised as having a gas mileage of 15 km/L. Convert this rating to miles per gallon.
7. Classification of Matter
Place the following words in the flow chart below:
Element
Homogeneous Mixture (solution)
Homogeneous
Heterogeneous mixture (solution)
Atom
Matter
Neutron
Proton
Pure Substance
Compound
Nucleus
Electron
Mixtures can be separated by
methods involving only
physical changes
-
Compounds can be
decomposed to
elements only through
chemical changes.
List the six physical properties found in
Chapter 1 of your book:
•
•
•
•
•
•
What is a common chemical property?
____________________
Distillation
What physical property would allow you to separate a
mixture containing water, alcohol, cyclohexane (a liquid), and
salt using the above apparatus? _______________
What would be left over in the distilling flask? __________
Paper chromatograph of ink.
(a) A line of the mixture to be separate is placed
at one end of a sheet of porous paper.
(b) The paper acts as a wick to draw up the liquid.
(c) The component with the strongest attraction
for the liquid travels faster than those that cling to
the paper.
8. Significant Figures
There are two kinds of numbers in the world:
• exact:
o example: There are exactly 12 eggs in a dozen.
o example: Most people have exactly 10 fingers and 10 toes.
• inexact numbers:
o example: any measurement.
If I quickly measure the width of a piece of notebook paper, I might get 220 mm (2 significant
figures). If I am more precise, I might get 216 mm (3 significant figures). An even more precise
measurement would be 215.6 mm (4 significant figures).
Significant figures are critical when reporting scientific data because they give the reader an idea of how well
you could actually measure/report your data. In any measurement, the number of significant figures is critical.
The number of significant figures is the number of digits believed to be correct by the person doing the
measuring and always includes one estimated digit.
Beaker
The first digit is definitely a 4 because
the liquid line is between 40 and 50.
The second digit is an estimate… 46? 47? 48?
This measurement has 2 significant figures.
Graduated cylinder
Buret
Graduated cylinder has more gradations (lines)
so we can get more significant figures.
Now we know it is 36 and we can
estimate the last digit 36.4? 36.5? 36.6?
This measurement has 3 significant figures.
A. Determining number of significant figures
1.
All non zero numbers are significant (meaning they count as sig figs)
613 has __3__ sig figs
2.
123456 has ____ sig figs
Buret has even more
gradations so it will give
you the most accurate
measurement.
We know 20.3 and can
estimate the last digit.
20.37? 20.38? 20.39?
This measurement has
4 sig figs.
17865332 has ____ sig figs
Zeros located between non-zero digits are significant (they count)
5004 has __4__ sig figs
602 has ____ sig figs
6000000000000002 has ___ sig figs
3. Trailing zeros (those at the end) are significant only if the number contains a decimal point; otherwise
they are insignificant (they don’t count)
5.640 has __4__ sig figs
120000. has ____ sig figs
120000 has ____ sig figs
4. Zeros to left of the first nonzero digit are insignificant (they don’t count); they are only placeholders!
0.000456 has __3__ sig figs
0.052 has ____ sig figs
0.000000000000000052 has ____ sig figs
B. Rules for addition/subtraction problems
Your calculated value will have the same number of digits to the right of the decimal point as that of
the least precise quantity.
6.22
53.6
 limiting term has 1 decimal place
14.311
+ 45.09091
119.22191  round to 119.2 (1 decimal place)
5365.999
 limiting term has 3 decimal places
- 234.66706
5131.33194  round to 5131.332
Underline the limiting term and write the answer using the correct amount of significant figures:
17.12 + 30.123 = _________
1000.00 – 62.5 = ________
15.05 + 0.0044 + 12.34 = ________
C. Rules for multiplication/division problems
When multiplying/dividing, the answer should have the same number of significant figures as the
limiting term. The limiting term is the number with the least number of significant figures.
503.29 x 6.177 = 3108.82233 → round to 3109
↑
limiting term has 4 sig figs
1000.1 = 4.11563786 → round to 4.12
243
↑ limiting term has 3 sig figs
Underline the limiting term and write the answer using the correct amount of significant figures:
35.010 / 1.23 = __________
0.1700 x 1700. x 1700 = __________
D. Rules for conversions
When converting a number, the answer should have the same number of significant figures as the
number started with.
Some numbers are exact because they are known with complete certainty.
Most exact numbers are integers: exactly 12 inches are in a foot, there might be exactly 23
students in a class. Exact numbers are often found as conversion factors or as counts of objects.
*Exact numbers can be considered to have an infinite number of significant figures.
Thus, the number of apparent significant figures in any exact number can be ignored as a
limiting factor in determining the number of significant figures in the result of a calculation.
52.4 in | 1 ft
| 12 in
↑
3 sig figs
= 4.366666667 ft
→ round to 4.37 ft
E. Rules for rounding off numbers
1.) If the digit to be dropped is greater than 5, the last retained digit is increased by one.
For example: 12.6 is rounded to 13.
2.) If the digit to be dropped is less than 5, the last remaining digit is left as it is.
For example: 12.4 is rounded to 12.
3.) If the digit to be dropped is 5, and if any digit following it is not zero, the last remaining digit is
increased by one.
For example: 12.51 is rounded to 13.
Significant Figure Practice Problems
How many significant figures does each of the following contain?
1.) 54
_____
4.) 4.00
_____
7.) 0.041
_____
2.) 45678 _____
5.) 400
_____
8.) 0.00010
_____
3.) 4.03
6.) 400.
_____
9.) 190909090
_____
_____
Round the following numbers as indicated:
10.) To four figures:
3.682417
21.860051
375.6523
112.511
45.4673
________
_______
_______
_______
_______
2.473
5.687524
7.555
8.235
_______
_______
_______
_______
79.2588
0.03062
3.4125
41.86632
_______
_______
_______
_______
11.) To one decimal place: 1.3511
________
12.) To two decimal places: 22.494
_________
Which number in each of the additions/subtractions is the limiting term, and how many decimal places should
the answer of each addition/subtraction have? Write the answer with the correct amount of significant figures.
13.)
55.43 + 44.333 + 5.31 + 9.2
= _______________
# of sig figs _______
14.)
890.019 + 890.1234 + 890.88788
= _______________
# of sig figs _______
15.)
69.99999 – 45.44444444
= _______________
# of sig figs _______
16.)
3.461728 + 14.91 + 0.980001 + 5.2631
= _______________
# of sig figs _______
Which number in each of the multiplication/division problems is the limiting term, and how many sig figs
should the answer of each multiplication/division have? Write the answer with the correct amount of sig figs.
17.)
343.4 / 34.337
= _______________
# of sig figs _______
18.)
0.000000003 x 30.03030
= _______________
# of sig figs _______
19.)
(1.3) x (5.724)
= _______________
# of sig figs _______
20.)
(6305) / (0.010)
= _______________
# of sig figs _______
21.)
(6.78 x 10-4) x (1.4 x 102)
= _______________
# of sig figs _______
22.)
12.5 x 75
= _______________
# of sig figs _______
AP CHEMISTRY CHAPTER 2 REVIEW: ATOMS, MOLECULES, & IONS
1. Using either your purple notes from Chemistry A and B or the internet define the following terms.
LAW OF CONSERVATION OF MASS:
THE LAW OF DEFINITE PROPORTIONS: Define and give an example.
THE LAW OF MULTIPLE PROPORTIONS: Define and give an example.
DALTON’S ATOMIC THEORY OF MATTER: Give the four parts of his theory.
I.
II.
III.
IV.
Using the knowledge gained in Chemistry A what are TWO MODIFICATIONS that have been made to
Dalton’s Atomic Theory?
In your own words describe how J.J. Thomson’s experiment with the cathode ray tube led to the discovery of
electrons.
Draw a picture of the cathode ray tube and what happened when a magnet was brought close to the ray.
What was the name of the “model” that was born from this experiment? Draw it.
In your own words describe how Rutherford’s Gold Foil experiment led to the discovery of the nucleus.
Draw a picture of the gold foil experiment.
What was the name of the “model” that was born from this experiment? Draw it.
2. NAMING COVALENT MOLECULES, IONIC COMPOUNDS, AND ACIDS
Rules for Naming Covalent Compounds (nonmetal + nonmetal)
A.
Use prefixes to indicate the number of each element in the molecule.
mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-.
B.
Drop the mono prefix if there is only one of first element in the molecule. Exp: CO2, carbon dioxide. (no mono)
C.
Ending of the last element changes to –ide.
Rules for Naming Ionic Compounds (metal + nonmetal)
A.
Balance Charges (charges should add up to zero).
B.
Cation (+ ion) is always written first (in name and in formula).
Cation has same name as on periodic table. You may need to indicate the charge of the cation in the name using
roman numerals if it is multivalent. Exp: FeCl3 is Iron(III) chloride whereas FeO is iron(II) oxide.
C.
Change the ending of the anion (-ion) to –ide (unless polyatomic ion, then named as given).
I. Name these binary compounds of two nonmetals.
IF7__________________________
N2O5____________________________
N2O4________________________
As4O10___________________________
PCl3_________________________
S2Cl2____________________________
XeF2 _____________________________
SF6_______________________________
II. Name these binary compounds with a fixed charge metal.
AlCl3 _______________________
MgO____________________________
KI__________________________
SrBr2 ___________________________
CaF2 ________________________
Al2O3___________________________
BaI2______________________________
Na2S _____________________________
III. Name these binary compounds of multivalent cations (use roman numerals).
CuCl2 ______________________
Fe2O3____________________________
PbCl4_______________________
Cu2S_____________________________
AuI3________________________
CoP_____________________________
SnO______________________________
HgS______________________________
IV. Name these compounds with polyatomic ions.
Fe(NO3)3_________________________ NaOH__________________________
Ca(ClO3)2________________________ KNO2__________________________
NH4NO2_________________________ Cu2Cr2O7 _______________________
Cu2SO4___________________________
NaHCO3__________________________
Acids- If the formula has hydrogen written first, then this usually indicates that the hydrogen is an H+ cation and that the compound is
an acid.
Rules for Naming an Acid
A. When the name of the anion ends in –ide, the acid name begins with the prefix hydro-, the stem of the anion has
the suffix –ic and it is followed by the word acid.
-ide becomes hydro _____ic Acid
Example: Cl- is the Chloride ion so HCl = hydrochloric acid
HCl ________________________
HI _________________________
H2S _______________________
HF _______________________
B. When the anion name ends in –ite, the acid name is the stem of the anion with the suffix –ous, followed by the
word acid.
-ite becomes ______ous Acid
Example: ClO2- is the Chlorite ion so HClO2. = Chlorous acid.
C. When the anion name ends in –ate, the acid name is the stem of the anion with the suffix –ic, followed by the
word acid.
-ate becomes ______ic Acid
Example: ClO3- is the Chlorate ion so HClO3 = Chloric acid.
**I like to remember this rule as “I ate something and it was icky.”
I.
HNO3, which contains the polyatomic ion nitrate, is called nitric acid.
HNO2, which contains the polyatomic ion nitrite, is called nitrous acid.
Name the following acids using the correct naming rule.
HClO4__________________________
H2SO4__________________________ HC2H3O2_________________________
H3PO4__________________________
HNO2__________________________ H2CrO4__________________________
H2C2O4_________________________
H2CO3__________________________
II. Name these compounds appropriately.
CO____________________________ NH4CN _________________________
HIO3__________________________
NI3____________________________ AlP ____________________________
OF2___________________________
LiMnO4________________________ HClO __________________________
SO2___________________________
CuCr2O7_______________________
K2O____________________________
HF____________________________
FeF3__________________________
KC2H3O2________________________
MnS__________________________
III. Write the chemical formulas.
Tin (IV) phosphide
_____________
copper (II) cyanide
____________
Magnesium hydroxide
_____________
sodium peroxide
____________
Sulfurous acid
_____________
lithium silicate
____________
Potassium nitride
_____________
chromium (III) carbonate ____________
Gallium arsenide
_____________
cobalt (II) chromate
____________
Zinc fluoride
_____________
dichromic acid
____________
3. PERCENT COMPOSITION
Complete the following problems showing all work.
1. A 0.941 gram piece of magnesium metal is heated and reacts with oxygen. The resulting magnesium oxide product weighed
1.560 grams. Determine the percent composition of each element in the compound.
2. Determine the empirical formula given the following data for each compound:
a) Fe = 63.53%, S = 36.47%
b) Fe = 46.55%, S = 53.45%
3. A compound contains 21.6% sodium, 33.0% chlorine, 45.1% oxygen. Determine the empirical formula of the compound.
AP CHEMISTRY CHAPTER 3 REVIEW: CHEMICAL REACTIONS AND
STOICHIOMETRY
1. SOLUBILITY RULES
I. Review solubility rules and identify each of the following compounds as soluble or insoluble in water.
You must memorize these solubility rules: Salts of NH4+, Na+, K+ and NO3- are always soluble.
SOLUBILITY GUIDELINES
Compounds
Solubility
Exceptions
Salts of alkali metals (group 1A) and
Soluble
Some lithium compounds
ammonium (NH4+)
All nitrate, chlorate and acetate salts
Soluble
Sulfate salts
Soluble
Cation is Pb, Ag, Hg, Ba, Sr, or Ca
Halide (group 7A halogen ions) salts
Soluble
Cation is Ag, Hg or Pb
Acids (H in front)
Soluble
carbonates, phosphates, chromates,
Cation is alkali metal (group 1A) or
Insoluble
sulfides, hydroxides and oxides
ammonium
*salts = ionic compounds
Na2CO3
___________
CoCO3
_____________
Pb(NO3)2
_____________
K 2S
___________
BaSO4
_____________
(NH4)2S
_____________
AgI
___________
Ni(NO3)2
_____________
KI
_____________
FeS
___________
PbCl2
_____________
CuSO4
_____________
Li2O
___________
Mn(C2H3O2)2
_____________
Cr(OH)3_
_____________
AgClO3
___________
Sn(SO3)4
_____________
FeF2
_____________
II. Write out the balanced chemical equation for each of the following double replacement reactions. Predict whether each of
these double replacement reactions will give a precipitate or not based on the solubility of the products. If yes, identify the
precipitate and write it on the line to the right.
silver nitrate and potassium chloride
____________
magnesium nitrate and sodium carbonate
____________
strontium bromide and potassium sulfate
____________
cobalt (III) bromide and potassium sulfide
____________
ammonium hydroxide and copper (II) acetate
____________
lithium chlorate and chromium (III) fluoride
____________
2. BALANCING EQUATIONS AND TYPES OF REACTIONS
I. Balance the following equations with the lowest whole number coefficients.
____S8
 ____ SO3
+ ____ O2
____C10H16 + ____ Cl2
 ____ C
____Fe
 ____ Fe2O3
+ ____ O2
+ ____ HCl
 ____CO2
____C7H6O2 + ____ O2
____KClO3  ____ KCl
+ ____ H2O
+ ____ O2
____H3AsO4 ____As2O5 + ____ H2O
____V2O5
+ ____ HCl
 ____VOCl3
+ ____ H2O
____Hg(OH)2 + ____ H3PO4  ____ Hg3(PO4)2 + ____ H2O
II. Balance the following equations, indicate the states of matter for each compound, and indicate the type of reaction
taking place:
1)
____ NaBr (
)
+
)

____ Al2(SO4)3 (
) 
____ H3PO4 (
____ Na3PO4 (
)
+
____ HBr (
)
Type of reaction: ____________________
2)
____ Ca(OH)2 (
)+
____ CaSO4 (
) +
____ Al(OH)3 (
)
Type of reaction: ____________________
3)
____ Mg (
)
+
____ Fe2O3 (
)

____ Fe (

____ CO2 (
)
+
____ MgO (
+
____ H2O (
+
____H2 (
)
Type of reaction: ____________________
4)
____ C2H4 (
)
+
____ O2 (
)
)
)
Type of reaction: ____________________
5)
____ PbSO4 (
)

____ PbSO3 (
)
+
____ O2 (

____N2I6 (
)
Type of reaction: ____________________
6)
____NH3 (
)
+
____I2 (
)
Type of reaction: ____________________
)
)
3. WRITING CHEMICAL EQUATIONS
Mole Ratio
from
coefficients
4. STOICHIOMETRY AND LIMITING REACTANTS
1. Given the equation below, what mass of water would be needed to react with 10.0g of sodium oxide?
Na2O + H2O  2NaOH
2.
2NaClO3  2NaCl + 3O2
What mass of sodium chloride is formed along with 45.0g of oxygen gas?
3.
4NH3 + 5O2  4NO + 6 H2O
What mass of water will be produced when 100.0g of ammonia is reacted with excess oxygen?
4.
If the reaction in #3 is done with 25.0g of each reactant, what is the maximum amount of product that could be made?
Which reactant would be the limiting reactant?
5.
Na2S + 2AgNO3  Ag2S + 2NaNO3
If the above reaction is carried out with 50.0g of sodium sulfide and 35.0g of silver nitrate what is the maximum amount of
silver sulfide that could be made?
What is your limiting reactant?
What mass of the excess reactant remains?
6.
6NaOH + 2Al  2Na3AlO3 + 3H2
What volume of hydrogen gas (measured at STP) would result from reacting 75.0g of sodium hydroxide with 50.0g of
aluminum?
SOLUBILITY GUIDELINES
Activity Series of Metals
Name
Symbol
Lithium
Li
Potassium
K
Calcium
Ca
Sodium
Na
Magnesium
Mg
Aluminum
Al
Zinc
Zn
Iron
Fe
Lead
Pb
Hydrogen
H
Copper
Cu
Mercury
Hg
Silver
Ag
Platinum
Pt
Gold
Au
Compounds
Solubility
Exceptions
Soluble
Some lithium compounds
Soluble
-
Soluble
Cation is Pb, Ag, Hg, Ba, Sr, or
Ca
Soluble
Cation is Ag, Hg or Pb
acids (have an ___ in front)
carbonates, phosphates, chromates,
Soluble
-
sulfides, hydroxides, and oxides
Insoluble
Cation is alkali metal or
ammonium (NH4+)
Salts of alkali metals and ammonium
All nitrate, chlorate and acetate salts
Water
line
sulfate salts
halide salts (halogens)
Acid
line
*salt is an ionic compound
1 J = 0.2390 cal
4.184J = 1 cal
ΔH = m x C x ΔT
M = mol / L
%v / v =
M1V1 = M2V2
volume of solute
× 100%
volume of solvent
Pressures: 1 atm = 101.3 kPa = 760 mmHg = 760 torr
Boyle’s Law
€
P1V1 = P2V2
Combined Gas Law
Charles’ Law
V1 V2
=
T1 T2
Ideal Gas Law
%m / v =
€
K = °C + 273
P1 V1 P2 V2
=
T1
T2
PV=nRT
R = 8.31
P1 P2
=
T1 T2
mass of solute (g)
× 100%
volume of solvent (mL)
STP = 1 atm and 273K
€
Gay-Lussac’s Law
€
1000 cal = 1 Cal
Partial Pressures
L × kPa
mol × K
Ptotal = P1 + P2 + P3
€
€
€
or
R = 0.0821
L × atm
mol × K
moles gasX
x Ptotal = PX
€ total moles